WBCHSE Class 11 Chemistry S Block Elements Notes

Class 11 Chemistry S Block Elements Group 1 Elements (Alkali Metals) Introduction

Class 11 S-Block Notes All the alkali metals have one valence electron ( ns¹ ) outside the noble gas core. The loosely held s -electronin the outermost valence shell makes them the most electropositive metals.

• To get the stable electronic configurations of noble gases, they readily lose the valence electron to generate the monovalent (M⁺) ions. Hence, they are never found in a free state but in the combined state of nature.
• Since the last electron enters ns -orbital, these are called s -block elements.
• Since all these elements have similar valence shells or outer electronic configurations, all the alkali metals exhibit a striking resemblance in their physical and chemical properties and they are placed in a definite group (Gr-1).
• Lithium shows some abnormal behavior as its electronic configuration is slightly different from the rest of the members of Gr-1 and also because of its extremely small atomic and ionic radii.

Again, lithium shows some similarities with magnesium present in the group- 2 of the third period

Electronic configuration of alkali metals: 

Occurrence Of Alkali Metals

• Since the alkali metals are highly reactive, they do not exist in a free state. In nature, they mostly occur as compounds like halides, oxides, silicates, borates, and nitrates.
• According to abundance, lithium is placed at the 35th position. It mainly occurs in nature in the tire form of silicates,
• For example: Spodumene: LiAl(SiO₃)₂ and Lepidolite: Li₂Al₂(SiO₃)₃(F, OH)₂
• Sodium and potassium are respectively placed at 7th and 8th position in order of their abundance. Sea water is a major source of NaCl and KCl.
• Sodium is abundantly present in the form of rock salt (NaCl). Other important minerals are Chile salt petre: NaNO₃, borax: Na₂B₄O₇.10H₂O, mirabilite: Na₂SO₄, and trona: Na₂CO₃-NaHCO₃-2H₂O.
• Important ores of potassium are sylvite: KCl, carnallite: KCl-MgCl₂-6H₂O, and feldspar: (K₂O-Al₂O₃-6SiO₂).
• Rubidium and cesium are much less abundant than lithium. Radioactive francium does not occur appreciably in nature. It is obtained from the radioactive decay of actinium
• ²²⁷Ac ₈₉ → ²²³Fr₈₇  +⁴He₂  Its longest-lived isotope ²²³Fr₈₇  has a half¬life period of only 21 minutes.
• Since most of the compounds of alkali metals are water soluble, they are found in adequate amounts in seawater.

General Trends In Atomic And Physical Flv Properties Of Alkali Metals

The alkali metals show regular trends in their physical and chemical properties with an increase in atomic number. Some important atomic and physical properties of alkali metals are given in the following table:

Atomic and physical properties of alkali metals:

Class 11 S-Block Notes

General trends in different atomic and physical properties of alkali metals and their explanations:

1. Atomic and ionic radii 

The atomic and ionic radii of alkali metals are the largest in their respective periods and these values further increase on moving down the group from Li to Cs.

Atomic and ionic radii  Explanation:

On moving from left to right in a period, the number of electronic shells remains the same but the nuclear charge increases with each succeeding element Thus, the valence shell electrons experience a greater pull towards the nucleus and this results in successive decreases in atomic and ionic radii with an increase in atomic number

• Thus, the atomic and ionic radii of alkali metals are the largest in their respective periods.
• On moving down the group, a new electronic shell is ) added to each element and the nuclear charge increases with an increase in atomic number.
• The addition of an electronic shell tends to increase the size of the atom but the increase in nuclear charge tends to decrease the atomic radii by attracting the electron cloud inward. Thus, the two factors oppose each other.
• However, the increase in the number of shells increases the screening effect of the inner electrons on the outermost s -electron and as the screening effect is, quite large, it overcomes the contractive effect of the increased nuclear charge.
• The net result is an increase in atomic and ionic radii down the group from Li to Cs.

2. Ionization enthalpy

The first ionization enthalpies of alkali metals are the lowest in their respective periods. Explanation: Since the alkali metal atoms are largest in their respective periods, their outermost electrons being far away from the nucleus experience less force of attraction and hence, can be removed easily.

1. Ionization enthalpy of group-1 alkali metals decreases down the group.

Explanation:

Since the alkali metal atoms are largest in their respective periods, their outermost electrons being far away from the nucleus experience less force of attraction and hence, can be removed easily.

2. Ionisation enthalpy7 of group-1 alkali metals decreases down the group

Explanation:

On moving down the group from Li to Cs, the distance of the valence s-electron from the nucleus progressively increases due to the addition of a new shell with each succeeding element With an increase in the number of inner shells, the screening effects progressively increase and as a result, the effective nuclear charge experienced by the valence electron progressively decreases and hence, the ionization enthalpies decrease down the group.

3. The second ionization enthalpies of alkali metals are very high.

Explanation:

The monovalent cation formed by the removal of an electron from the alkali metal atom has a very stable noble gas configuration,

For example – 

1. Li⁺: 1s² or [He], Na⁺: 1s²2s²2p⁶ 
2. [Ne], k⁺: 1s²2s²2p⁶3s²3p⁶ or [Ar] etc.

Removal of another electron from the monovalent ion having a stable noble gas configuration is very difficult and requires a huge amount of energy. For this reason, the second ionization enthalpies of alkali metals are very high.

3. Hydration of ions, hydrated radii, and hydration enthalpy

The salts of alkali metals are generally ionic and are soluble in water because the cations get hydrated in water to form hydrated cations: M⁺ + aq —> [M(aq)]⁺.

1. The degree of hydration of ions and the hydrated radii decrease as we move down the group

Explanation:

The smaller the cation, the greater its degree of hydration. Since ionic radii increase down the group, the degree of hydration decreases, and consequently, the radii of die-hydrated ions decrease from Li⁺ to Cs⁺.

2. The order of mobilities of die alkali metal ions in aqueous solution is: Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺

Explanation:

Smaller ions are more easily hydrated. As Li+ is the smallest ion among the given ions, it is most easily hydrated and has the least ionic mobility in an aqueous solution whereas Cs+ is the largest and is least hydrated. So its mobility is die highest.

3. Ionic conductance of the hydrated ions increases from [Li(m7)]+ to [Cs(aq)]+.

Explanation:

The ionic conductance of these hydrated ions increases from [Li(aq)]⁺ to [Cs(ag)]⁺ because die size decreases and mobility increases in this order. Hydration of ions is an exothermic process. The energy released when 1 gram-mol of an ion undergoes hydration is called hydration energy or hydration enthalpy,

4. Hydration enthalpy of alkali metal ions decreases from Li⁺ to Cs⁺.

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Class 11 S-Block Notes

Explanation:

The hydration enthalpy of an ion depends upon the ratio of charge to radius (q: r). Since the radii of alkali metal ions increase down the group, the hydration enthalpies decrease from Li⁺ to Cs⁺. Li⁺ ion has the maximum degree of hydration and for this reason, most of the lithium salts are found to be hydrated

For example: LiCl-2H₂O, LiClO₄-3H₂O etc.

4. Oxidation State

Alkali metals exhibit a +1 oxidation state in their compounds and it remains restricted in a +1 state only.

Oxidation State Explanation: 

Alkali metals have low ionization enthalpies and by losing their valence s -electrons they acquire the stable electronic configurations of the nearest noble gases. Thus, they have a strong tendency to form M⁺ ions and exhibit a +1 oxidation state in their compounds.

The second ionization enthalpies required to pull out another electron from M⁺ ions having very unstable noble gas electronic configuration are very high indeed and are not available under the conditions of chemical bond formation. Hence,v the alkali metals do not form M²⁺ ions, i.e. their oxidation state remains restricted to +1 state.

5. Metallic character

The elements of this group are typical metals that are soft (can be easily cut with a knife) and light. When freshly cut, they are silvery white but on exposure to air, they turn tarnished. The metallic character, which refers to the level of reactivity of a metal, increases on moving down the group.

Metallic character Explanation:

As the ionization enthalpy decreases down the group,  the tendency to lose the valence electron increases, and consequently, the metallic character increases.

6. Photoelectric effect

Alkali metals (except Li) exhibit a photoelectric effect. The emission of electrons from the surface of a metal exposed to electromagnetic radiations of suitable wavelength is called the photoelectric effect.

Photoelectric effect Explanation:

Due to low ionization enthalpies, the alkali metals exhibit a photoelectric effect. It is to be noted that lithium having the highest ionisation enthalpy does not exhibit a photoelectric effect. Cesium having the lowest ionisation enthalpy possesses the highest tendency to exhibit a photoelectric effect.

Potassium anti-cesium, rather than lithium is used in photoelectric cells:

The ionization enthalpies of potassium and cesium are much lower than that of lithium. For this reason, these two metals on exposure to light easily emit electrons from their surface but lithium does not. Hence, potassium and cesium rather than lithium are used in photoelectric cells.

7. Electronegativity

The alkali metals have low electronegativity which further decreases down the group.

Electronegativity Explanation:

The alkali metals having ns¹ electronic configuration preferably show electron releasing tendency rather than electron accepting. Thus, they have low electronegativities. Since the atomic sizes increase down the group, the tendency of atoms to hold their valence electrons decreases down the group, and consequently, electronegativity decreases down the group.

Class 11 S-Block Notes

8. Conductivity

Alkali metals are good conductors of heat and electricity.

Conductivity Explanation:

Due to the presence of loosely bound valence electrons (ns¹) which are free to move throughout the metal structure, the alkali metals are good conductors of heat and electricity.

9. Melting and boiling points 

Melting and boiling points: Melting and boiling points of alkali metals are low and decrease down tire group.

Melting and boiling point Explanation:

The cohesive energy that binds the atoms in the crystal lattices of these metals is relatively low (weak metallic bonding) due to the presence of only one valence electron (ns¹) which can take part in bonding. Hence, their melting and boiling points are low. These further decrease down the group as the strength of the metallic bonds and cohesive energy decrease with increasing atomic size.

10. Nature of bonds formed

Alkali metals form ionic compounds and the ionic character of compounds increases down the group from Li to Cs.

Nature of bonds formed Explanation:

For low ionization enthalpies, alkali metals readily form monovalent cations by losing their valence electrons. As ionization enthalpies decrease down the group, the ionic character of the compounds increases down the group.

11. Density

The densities of alkali metals are quite low and increase down the group from Li to Cs.

Density Explanation:

Due to their large atomic size and weak metallic bond, alkali metals have low density. Both the atomic volume and the atomic mass increase down the group but the corresponding increase in atomic mass is not balanced by the increase in atomic volume. As a result, the densities of alkali metals increase down the group. However, the density of K is less than that of Na because the atomic size and atomic volume of potassium are quite higher than that of sodium. As a result, the ratio of mass/ volume decreases.

Li is the lightest metal having a density of 0.53 g. cm^-3. It cannot be preserved in kerosene because it floats over it. Generally kept wrapped in paraffin wax.

12. Flame coloration

Alkali metals or their salts on heating in the flame of the bunsen burner, impart characteristic colors and they can be easily identified from the color of the flame

Class 11 S-Block Notes

Flame coloration Explanation:

Ionization enthalpies of alkali metals are not much higher. Thus, when an alkali metal or its salt (especially chloride due to its more volatile nature) is heated in a Bunsen burner flame, the electrons in the valence shell get excited and jump to higher energy levels by absorbing energy. When the excited electron drops back to its ground state, the emitted radiation falls in the visible region and as a result, alkali metals or their salts impart color to the flame.

Alkali metals can be detected by flame tests and can be estimated by flame photometry or atomic absorption spectroscopy.

13. Softness

Alkali metals are soft (can be cut easily with die help of a knife) and their softness increases down the group.

Softness Explanation:

The softness of alkali metals is due to their low cohesive energy and weak metallic bonding. Further, on moving down the group, the strength of metallic bonding decreases due to an increase in atomic size and as a result, the softness of the metals increases down the group.

Chemical Properties Of Alkali Metals

Alkali metals are highly reactive. Such reactivity may be attributed to their large atomic size, low ionization enthalpies, and low heats of atomization. –

Action of air and moisture

Alkali metals, being highly reactive, react readily with atmospheric oxygen to form oxides. These oxides further react with moisture to form hydroxides which in turn produce carbonates by reacting with atmospheric CO₂

Metal oxides also react with CO₂ to form carbonates.

These metals lose their glossiness and become tarnished due to the formation of carbonate layers on their surface. To protect from atmospheric oxygen and moisture, these metals are always stored in inert hydrocarbon solvents such as kerosene, petroleum ether, etc.

Reaction with oxygen

When the alkali metals are heated with oxygen or excess air, they form different types of oxides depending upon the nature of the metal involved. Lithium mainly forms monoxide (Li₂O), sodium forms peroxide (Na₂O₂), and the other alkali metals (K, Rb, and Cs) mainly form superoxides having the general formula MO₂. The temperature required for the reaction decreases down the group from Li to Cs.

4Li + O₂ → 2Li₂ O₂ (Lithium monoxide)

2Na + O₂ → Na₂O₂ (Sodium peroxide)

M +O₂ → MO₂ (Superoxide) [here, M = K, Rb, Cs]

 Oxygen Explanation

A smaller cation can stabilize a smaller anion while a larger cation can stabilize a larger anion. If both the ions are similar in size, the coordination number will be high and this results in higher lattice energy.

A cation having a weak positive electric field can stabilize an anion having a weak negative electric field. Li⁺  ion and oxide ion (O^2-) have small ionic radii and high charge densities. Hence, these small ions combine to form a very’ stable lattice of Li₂O.

Sodium forms peroxide (Na₂O₂) but potassium forms superoxide (KO₂), even though the peroxide ion,  is larger in size than the superoxide ion, This can be explained in terms of charge densities. Due to the bigger size

Na+ ion has a weaker positive field around It and therefore it can stabilise peroxide Ion which also has a weaker negative field around it. Thus, Na⁺ forms peroxide. K⁺ ion is still bigger in size and the magnitude of the positive field around it is much weaker so it can’t stabilize superoxide ion which also has a much weaker negative field around it. Hence, K forms superoxide.

According to valence bond theory, two O-atoms in superoxide ion (O₂^–) are attached by a 2- 2-electron bond (common covalent bond) and a 2-electron bond. As the unpaired electron is present in the 2-electron bond, the superoxide ion is paramagnetic and all tire superoxides are colored (LiO₂ and NaO₂ are yellow, KO₂ & CsO₂ are orange, RbO₂ is brown).

According to MO theory, there is an unpaired electron in one of the n antibonding orbitals and because of this, the superoxide ion is paramagnetic. The electronic configuration of O₂^–

Class 11 S-Block Notes

O₂^–  ion present in common oxides [For example,  Li₂O, NaO₂etc.) and the 02~ ion in peroxides (For example,  Na₂O₂), contain no unpaired electron and for this reason, these are diamagnetic and colorless.

Reaction with water

The alkali metals having high negative reduction potentials (E°) can act as a better-reducing agent than hydrogen. Hence, they react with water to form water-soluble hydroxides and liberate hydrogen gas.

2M + 2H₂O → 2MOH + H₂T [M = alkali metal]

Reactivity with water increases down the group as the electropositive character of the metals increases clown the group. Lithium decomposes water slowly. Sodium reacts with water vigorously. K, Rb, and Cs react with water explosively and the evolved hydrogen gas catches fire.

Alkali metals also react with compounds containing acidic H-atoms

For example:  halogen halides (HX), alcohols (ROH), acetylene HC=CH, etc., to form their corresponding salts and H₂

2Na + 2HX → 2NaX  (Sodium halide) + H₂ ↑

Li + 2C₂H₅OH → 2 C₂ H₅ OLi (Lithium ethoxide) + H₂ ↑

2Na + 2HC=CH →  2NaC = CH (Sodium acetylide)+ H₂ ↑

The standard electrode potential of Li is most negative while that of sodium is least negative i.e., in the reaction with water, Li releases a greater amount of heat than sodium. Despite that, Li reacts less vigorously with water than Na. j This can be explained concerning chemical kinetics. Na has a low melting point and the heat of the reaction is sufficient j to melt it. Molten metal spreads out and exposes a relatively large surface to water and as a result, it reacts with water 1 readily and violently. On the other hand, the melting point of Li is much higher and the heat of the reaction is not sufficient to melt it. Hence, its surface area does not increase and it reacts slowly with water.

Reaction with dihydrogen

All alkali metals react with dihydrogen at about 673 K (Li at 1073 K) to form colorless, crystalline hydrides (MH). These hydrides have a high melting point.

1. The reactivity of the alkali metals towards dihydrogen decreases down the group.

Explanation:

As the size of the metal cation increases down the group, the lattice energy of the hydrides decreases down the group. Consequently, the reactivity of the alkali metals towards hydrogen decreases down the group,

2. The ionic character of the alkali metal hydrides increases from Li to Cs.

Explanation:

Since the ionization enthalpy of alkali metals decreases down the group, the tendency to form cations as well as the ionic character of the hydrides increases.

Reaction with halogens

All alkali metals react vigorously with halogens to form crystalline halide compounds having the general formula MX. Lithium halides are covalent due to the very high polarising power of the Li⁺ ion. Halides of other alkali metals are ionic in nature’

⇒ \(2 \mathrm{M}+\mathrm{X}_2 \rightarrow 2 \mathrm{M}^{+} \mathrm{X}^{-}\) ( X = F, Cl, Br or I)

The reactivity of the alkali metals towards a particular halogen increases down the group.  Due to a decrease in ionization enthalpies or an increase in the electropositive character of the metals down the group, the reactivity increases down the group.

The reactivity of halogens towards a particular alkali metal decreases in the order:

F₂> Cl₂ > Br₂ > I₂

Reducing nature

The alkali metals act as strong reducing agents because of their low ionization enthalpies. Since the ionization enthalpies decrease on moving down the group, therefore, in the free state the reducing power also increases in the same order, i.e., Li  < Na < K < Rb < Cs.

Class 11 S-Block Notes

The tendency of a metal to lose an electron in solution is measured by its standard electrode potential (E°). The alkali metals have low values (higher negative values) of E° and so they have a strong tendency to lose electrons and can act as strong reducing agents. Lithium, although, has the highest ionization enthalpy, is the strongest reducing agent in solution (E° = -3.04 V). On the other hand, Na is the weakest reducing agent and the reducing character increases from Na to Cs, i.e., Na < K < Rb < Cs.

Explanation of anomalous behavior of lithium

The anomalous behavior of lithium can be explained because the ionization enthalpy is the property of an isolated atom in the gaseous state while the standard electrode potential is concerned when the metal atom goes into solution.

The ionization enthalpy involves the change:

M(g) → M⁺+(g) + e, while the standard electrode potential involves the change: M(s) → M⁺(aq) + e.

The latter change occurs in three steps as follows:

1. M(s)→ M(g) – sublimation enthalpy
2. M(g)→ M+(g) + e – ionisation enthalpy
3. M⁺ (g) + H₉O → M⁺(aq) + hydration enthalpy

The overall tendency for the change depends on the net effect of these three steps. Among the alkali metal cations, Li⁺ ion has the maximum tendency to get hydrated due to its very small size. The high hydration enthalpy compensates the energy required in the first two steps to a large extent and the overall energy required to convert M(s) to M⁺ (aq) is minimum for lithium.

Thus, small size and high hydration enthalpy are responsible for the strong reducing character of lithium.

The solution in liquid ammonia

Alkali metals dissolve in liquid ammonia to give highly conducting deep blue solutions which are highly reducing and paramagnetic. As the concentration increases (> 3M), the color of the solution changes to copper-bronze. These concentrated solutions are diamagnetic.

Solution in liquid ammonia Explanation:

1. When an alkali metal is dissolved in liquid v ammonia, ammoniated cations, and ammoniated electrons are formed as shown below:

M + (x+ y)NH₃ →  [M(NH₃)_x]⁺ (Ammoniated cation) + [e(NH₃)_y]^–(Ammoniated electron)

2. The blue color of these solutions is due to the excitation of the free ammoniated electrons to higher energy levels by absorbing energy corresponding to the red region of visible light.  The transmitted light is blue which imparts a blue color to the solutions.

3. With the increase in the concentration of the alkali metal, the formation of clusters of metal ions starts and because of this, at a much higher concentration (> 3M) the solutions possess metallic luster and attain the color of copper-bronze.

4. These blue solutions are highly conducting because of the presence of ammoniated electrons and ammoniated cations but the conductivity decreases with increasing concentration as the ammoniated cations get attached to the free unpaired electrons.

5. These blue solutions are paramagnetic due to the presence of unpaired electrons. However, tire paramagnetism decreases with increasing concentration due to the association of ammoniated electrons to yield diamagnetic species.

6. The free ammoniated electrons make these solutions very powerful reducing agents.

7. These solutions when kept, form metal amides and release H₂. However, these solutions can be stored in anhydrous conditions in the absence of impurities like Fe, Pt, Zn, etc.

M⁺(am) + e(am) + NH₃(l)→ MNH₂(am) + ½H₂(g)

Where ‘am’ stands for ‘solution in ammonia

2M + 2NH₃→ 2MNH₂ (metal amide)+ H₂

Class 11 S-Block Notes

Extraction of alkali metals

Alkali metals cannot be extracted by applying common processes used for the extraction of other metals.

Alkali metals Explanation:

• The alkali metals are strong reducing agents. Hence, they cannot be extracted by reduction of their oxides or other compounds.
• Since they are highly electropositive, the method of displacing them from their salts by any other element is not possible.
• The aqueous solution of their salts cannot be used for extraction by electrolytic method because hydrogen, instead of the alkali metal is discharged at the cathode (discharge potentials of alkali metals are much higher).
• However, by using Hg as a cathode, the alkali metals can be deposited but in that case, the alkali metals readily combine with mercury to form amalgams from which the recovery of metals becomes quite difficult.
• The electrolysis of their fused salts (usually chlorides) is the only successful method for their extraction, Another metal . ‘ chloride is generally added to lower its fusion temperature

General Characteristics Of The Compounds Of Alkali Metals

The compounds of alkali metals are predominantly ionic. Some of the general characteristics of these compounds are described below.

1. Oxides and hydroxides

1. Typical oxides or monoxides of alkali metals

For example: Li₂O and Na₂O are white ionic solids and basic. These oxides react with water to form strong alkalis (MOH).

Example:  Na₂O + HO → 2NaOH

2. All peroxides are strong oxidizing agents. They react with water or acid to give hydrogen peroxide (H₂O₂) and the corresponding metal hydroxide. Na₂O₂ is widely used as an oxidizing agent in inorganic chemistry.

M₂O₂ + 2H₂O →  2MOH + H₂O₂

Example: Na₂O₂ + 2H₂O₂→ 2NaOH + H₂O₂

3. Superoxides are stronger oxidizing agents than peroxides and react with water or acid to give both H₂O₂ and O₂ along with metal hydroxide.

2MO₂ + 2H₂O→  2MOIH + H₂O₂+ O₂

Example: 2KO₂ + 2H₂O₂ →2KOH + H₂O₂ + O₂

The alkali metal hydroxides (MOH) are all white crystalline solids and corrosive. They are the strongest of all bases and readily dissolve in water. Due to excess hydration, a large amount of heat is released. These hydroxides are thermally stable except Li OH. The basic strength of alkali metal hydroxides increases on moving down the group from Li to Cs.

Explanation:

The ionization enthalpies of alkali metals decrease on moving down the group and this causes a weakening of the bond between the alkali metal and the hydroxyl group (M —OH). This results in an increase in the concentration of hydroxyl ions in the solution, i.e.,  the basic character of the solution increases on moving down the group.

Thus, the basic strength of the hydroxides follows the order:

CsOH > RbOH > KOH > NaOH > LiOH

2. Halides

The alkali metal halides can be prepared by combining metals directly with halogens or by reacting appropriate oxides, hydroxides or carbonates with aqueous halogen acids (HX).

2M + X₂  →  2MX ; M₂ O + 2HX↓ 2MX + H₂ O

MOH + HX→ MX + H₂ O

M₂ CO₃+ 2HX→ 2MX + HO₂ + CO₂

The enthalpy of formation (ΔH°_f) of alkali metal halides is highly negative. For a given metal, AHj values decrease from fluoride to iodide. These halides are colorless crystalline solids having high melting and boiling points.

1. The melting point of halides of a particular alkali metal decreases as:

Fluoride > Chloride > Bromide > Iodide.

Explanation:

For a particular alkali metal ion, the lattice enthalpies decrease as the size of the halide ion increases.

Lattice enthalpies of NaP, NaCl, Nalir, and Nal are 893, 770,730, and 685 kJ. mol^-1 respectively, A.s the lattice enthalpy decreases, the energy required to break the crystal lattice decreases, and consequently, the melting points decrease. Thus, the melting points of NaF, NaCl, NaBr, and Nal are found to be 1201K, 10IMK, 1028 K, and 944 K respectively.

2. For a particular halide ion, the melting point of IJX is less than that of

NaX and thereafter the melting points decrease on moving down the group from Na to Cs.

Class 11 S-Block Notes

Explanation:

The melting point of LiCl (887K) is less than that of NaCl (I084K), because LiCl is covalent (for smaller atomic size of Li compared to that of Na), but NaCl is ionic. Thereafter, the order of melting point is:

NaCl(1084K)>KCl(1039K)>RbCI(988K)>CsCl(925K) This is observed because the lattice enthalpies decrease as the size of the alkali metal atom increases.

3. Solubilities of the alkali metal halides (except fluorides) decrease on moving down the group since the decrease in hydration enthalpy is more than the corresponding decrease in the lattice enthalpy.

For example, the difference in lattice enthalpy between NaCl and KCl is 67kJ. mol^-1 whereas the difference in hydration enthalpy between Na⁺ and K⁺ ion is 76 kj -mol^-1 Thus, KCl is relatively less soluble in water compared to NaCl.

Explanation:

The solubility of a salt in water depends on its lattice enthalpy as well as its hydration enthalpy. In general, if hydration enthalpy > lattice enthalpy, the salt dissolves in water but if the hydration enthalpy < lattice enthalpy, the salt does not dissolve.

Further, the extent of hydration depends on the ionic size. The smaller the size of the ion, the more it will get hydrated and the greater will be its hydration enthalpy. LiF, for example, is almost insoluble in water because of its higher lattice enthalpy (-1005 kJ . mol^-1 ).

On the other hand, the low solubility of Csl in water is due to smaller hydration enthalpies of the two large ions [-276(Cs⁺)-305(I^–) = -581 kJ.-mol^-1]. o Due to the smaller size and relatively higher electronegativity of Li, lithium halides except LiF are predo¬minantly covalent and hence, are soluble in organic solvents such as acetone, alcohol, ethyl acetate etc.

In contrast, sodium chloride, being ionic, is insoluble in organic solvents.

3. Soils of oxoacids

Alkali metals react with c to acids such as carbonic acid (H₂CO₃), nitric acid (HNO₃), sulphuric acid (H₈SO₄), etc., to form corresponding salts and release H₂. Due to the high polarising power and lattice energy of small Li ions, lithium salts behave abnormally.

4. Nature of carbonates and bicarbonates

All alkali metals form carbonates of the type M₂CO₃. Since the alkali metals are highly electropositive, their carbonates are remarkably stable up to l000°C above which they first melt and then decompose to form oxides. These salts are readily soluble in water. As electropositive character increases down the group, the stability of carbonates increases in the same order:

Cs₂CO₃ > Rb₂CO₃ > K₂CO₃ > Na₂CO₃ > Li₂CO₃

Li₂CO₃ is insoluble in water and unstable towards heat. It decomposes readily to give Li₂O and CO₂.

Explanation:

1. The very small L ion exerts a strong polarising power on the large carbonate (CO₃^2-) ion and distorts the electron cloud of its nearby oxygen atom.

This results in the weakening of the C—O bond and the strengthening of the Li — O bond. This eventually facilitates the decomposition of Li₂CO₃ leading to the formation of Li₂O and CO₂. %

2. The crystal lattice formed by a smaller Li⁺ ion with a smaller O₂ ion is more stable than that 2  formed by a larger CO₃ ion and a smaller Li+ ion. This also favors the decomposition of Li₂CO₃

Class 11 S-Block Notes

3. The aqueous solution of carbonates is alkaline. This is because carbonates being the salts of strong bases and weak acids (H₂CO₃) undergo hydrolysis.

M₂CO₃ + 2H₂O  ⇌  2MOH (strong base) + H₂CO₃ (weak acid)

4. Bicarbonates or hydrogen carbonates (MHCO₃) of the alkali metals except LiHCO₃ are obtained in the solid state. These bicarbonates are soluble in water and stable towards heat. On strong heating, all the bicarbonates undergo decomposition to yield carbonates with the evolution of carbon dioxide.

2MHCO₃ (heat)→ M₂CO₃ + CO₂ + H₂O

As the electropositive character of the metals increases down the group from Li to Cs, the stability of the bicarbonates increases in the same order.

5. Nature of nitrates

The alkali metal nitrates (MNO₃) are prepared by the action of HNO₃ on the corresponding carbonates or hydroxides. They are ionic crystalline solids having low melting points and are highly soluble in water. On strong heating, they (except LiNO₃ ) decompose into nitrites and at higher temperatures oxides.

For example:

LiNO₃ decomposes readily on heating to give

6. Nature of sulphates

The alkali metals form sulfates of the type M₂SO₄. All the sulfates except Li₂SO₄ are soluble in water. The sulfates when fused with charcoal, form sulphides.

M₂SO₄.+ 4C→  M₂ S + 4CO

Sulfates of alkali metals form double salts with the sulfates of trivalent metals like Fe, Al, Cr, etc. These double salts crystallize with a large number of water molecules to form alum. A typical example is potassium aluminum

[K₂SO₄→ Al₂(SO₄)₃ -24H₂O].

Lithium sulfate (Li₂SO₄) is not known to form alum.

Anomalous Behaviour Of Lithium (Li) And Similarity Between Li And Mg

Although lithium, the first element of group 1, exhibits most of the characteristic properties of this group, yet it differs from other members of this group in several respects.

Reasons for anomalous behavior of lithium

1. Both Li- atom and Li⁺ ion have very small sizes.
2. Much higher polarising power of very small Li⁺ ion results in increased Points of difference between lithium and other alkali metals covalent character of its compounds.
3. Lithium has the lowest electropositive character, the highest ionization enthalpy, and the highest electronegativity compared to the rest of the members.
4. Non-availability of d -d-orbital in its valence (outermost) shell.
5. Strong intermetallic bonding (cohesive force) due to its small size. On the other hand, lithium shows a diagonal relationship with magnesium

Class 11 S-Block Notes

Points of difference between lithium and other alkali metals:

Reasons for the similarities between lithium & magnesium:

Lithium exhibits a diagonal relationship with the 3rd-period group-2 element, magnesium. Reasons for the similarities between lithium and magnesium

• The atomic as well as ionic radii of Li and Mg are almost the same (Li⁺ = 76 pm and Mg²⁺ = 72 pm).
• Both lithium and magnesium have almost similar electronegativities (Li = 0.98 and Mg = 1.2).

Similarities between lithium and magnesium:

Uses Of Alkali Metals

Class 11 S-Block Notes

Preparation, Properties, And Uses Of Some Important Compounds Of Sodium

1. Sodium carbonate (washing soda), (Na₂CO₃-10H₂O)

1. Manufacture: Ammonia-soda or Solvay process

Sodium carbonate is commonly known as washing soda. It is generally manufactured by the Solvay process or ammonia-soda process.

Principle: When carbon dioxide is passed through an aqueous solution of NaCl (brine, 28% NaCl solution) saturated with ammonia, sodium bicarbonate is formed.

NH₃ + CO₂ + H₂O > NH₄HCO₃

NH₄HCO₃ + NaCl  ⇌ NaHCO₃ + NH₄Cl

Due to the common ion effect of Na⁺ ion, sodium bicarbonate so formed gets precipitated. Such removal of solid NaHCO₃ shifts the reaction more and more towards the right. This results in a greater yield of NaHCO₃.  In this way, a nearly two-thirds portion of NaCl is converted into NaHCO₃. The precipitated NaHCO₃ is then filtered off, dried, and heated at 150°C to get sodium carbonate.

Evolved CO₂ is reused to saturate the ammoniated brine.

Class 11 S-Block Notes

Raw materials:

1. Brine solution (28% aqueous solution of NaCl),
2. Limestone or calcium carbonate (CaCO₃) it is the source of CO₂ and
3. Ammonia.

2. Description of the process

Preparation of ammoniated brine: 

1. This process is carried out in the absorption tower made of iron

2.  From an overhead tank, brine is allowed to trickle down slowly along the tower and ammonia gas from the ammonia recovery tower which is mixed with a small amount of CO₂ is allowed to pass through a tube situated near the bottom of the tower. As a result, the brine solution gets saturated with ammonia while calcium chloride and magnesium chloride are present as impurities in commercial.

3. Sodium chloride gets precipitated as their corresponding insoluble carbonates.

2NH₃ + CO₃ + H₂O → (NH₄)₂CO₃

CaCl₂ + (NH₄)CO₃ → 2NH₄Cl + CaCO₃↓

MgCl₂ + (NH₄)2CO₃ → 2NH₄Cl + MgCO₃↓

4. The ammoniated brine is then filtered to remove the precipitated calcium and magnesium carbonates and the filtrate thus obtained is passed into the carbonation tower.

Class 11 S-Block Notes

Carbonation of ammoniated brine:

1. This operation is carried out in a long cast iron tower (carbonation or Solvay tower). The tower is fitted with several horizontal plates.

2. The ammoniated brine solution is trickled down from the top of the tower while CO₂ gas from the lime kiln is introduced into the tower under high pressure through a pipe fitted at the base of the tower.

3. In this way, CO₂ comes in contact with the descending stream of ammoniated brine and they react with each other to form ammonium bicarbonate which subsequently combines with NaCl to produce sodium bicarbonate and ammonium chloride.

NH₃ + CO₂ + H₂O →  NH₄HCO₃

NaCl + NH₄HCO₃ →  NaHCO₃↓ + NH₄Cl

Separation of sodium bicarbonate:

1. The solution coming out of the carbonation tower contains crystals of NaHCO₃. These are separated by passing the solution through vacuum filters.

2. The separated sodium bicar¬bonate is washed with water to remove any sodium or ammonium chloride that may adhere to it and then dried.

3. The filtrate containing NH₄Cl and a small amount of NH₄HCO₃ is taken to the ammonia recovery tower where it comes in contact with Ca(OH)₂.

Calcination:

When the dry NaHCO₃ is heated strongly in a furnace at 180°C, it decomposes to form anhydrous Na₂CO₃. It is called soda ash. It is nearly 99.5 % pure.

Evolved CO₂ is reused in the carbonation tower or absorption tower

Recovery of ammonia:

The filtrate from the carbonation tower which contains ammonium chloride and a little ammonium bicarbonate is made to flow down the ammonia recovery tower. NH₄HCO₃ is decomposed by the heat of steam and NH₄Cl reacts with calcium hydroxide to form ammonia, carbon dioxide, and CaCl₂. The mixture of NH₃ and CO₂ is used for tire saturation of brine while calcium chloride is obtained as a by-product.

1.

2.

Potassium carbonate (K₂CO₃) cannot be prepared by the Soh’ay process. This is because unlike sodium bicarbonate, potassium bicarbonate (KHCO₃) which is fairly soluble in water does not get precipitated when CO₂ is passed through the ammoniated solution of KCl.

 Properties of Sodium carbonate: 

1. State:

Sodium carbonate is available either as anhydrous salt or as hydrated salt. The hydrated salts are white crystalline substances and are mainly of two types

1. Decahydrate (Na₂CO₃-10H₂O) and
2. Monohydrate (Na₂CO₂-H₂O).

The decahydrate is also called washing soda. The anhydrous salt commonly known as soda ash, is a white powder. When sodium carbonate is crystallized from water, the decahydrate is obtained as white transparent crystals. These crystals are efflorescent. When exposed to air for a long time, crystals of decahydrate partially lose their water of crystallization and is converted to monohydrate, a powdery substance which is known as crystal carbonate.

Class 11 S-Block Notes

⇒ \(\mathrm{Na}_2 \mathrm{CO}_3 \cdot 10 \mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{Na}_2 \mathrm{CO}_3 \cdot \mathrm{H}_2 \mathrm{O}+9 \mathrm{H}_2 \mathrm{O}\)

2. Action of heat:

When the decahydrate is heated up to 100°C, it slowly loses nine molecules of water of crystallization and gets converted into monohydrate. When the monohydrate is heated above 100°C, the anhydrous salt (Na₂CO₃) is produced as a white powder which melts at high temperatures but never undergoes decomposition.

1.

2.

Anhydrous Na₂CO₃ or soda ash melts at higher temperatures (melting point 852°C) but does not decompose. It turns to monohydrate when kept in the air.

3. Hydrolysis:

It dissolves in water with the evolution of a considerable amount of heat. Being a salt of a weak acid (HCO₃) and a strong base (NaOH), it undergoes hydrolysis in water to give an alkaline solution.

⇒\(\mathrm{Na}_2 \mathrm{CO}_3+2 \mathrm{H}_2 \mathrm{O} \rightleftharpoons 2\left[\mathrm{Na}^{+}+\mathrm{OH}^{-}\right]+\mathrm{H}_2 \mathrm{CO}_3\)

4. Reaction with acid:

At ordinary temperature, Na₂CO₃ reacts with dilute mineral acids to form the corresponding sodium salts and water along with the evolution of CO₂

Na₂CO₃ + 2HCl→2NaCl + CO₂↑ + H₂O

Na₂CO₃ + 2CH₃COOH→2CH₃COONa + CO₂↑ + H₂O

Reaction with slaked lime: When a solution of Na₂CO₃ is heated with slaked lime (milk of lime) at 80°C, sodium hydroxide with insoluble calcium carbonate is obtained.

Na₂CO₃  + Ca(OH)₂ → CaCO₃↓+ 2NaOH

Class 11 S-Block Notes

Uses of sodium carbonate:

• Sodium carbonate is mainly used for softening hard water and for washing clothes.
• It is used in fire extinguishers.
• It is largely used in the manufacture of soap, glass, borax, and caustic soda.
• It is used in the paper, paint, and textile industries.
• A mixture of Na₂CO₃ and K₂CO₃ is used as a fusion mixture.
• It is used as an important laboratory reagent both in qualitative and quantitative analysis.

2. Sodium bicarbonate or sodium hydrogen carbonate (baking soda), NaHCO₃

Preparation of sodium bicarbonate:

Sodium hydrogen carbonate is obtained as the intermediate product in the Solvay process of manufacturing sodium carbonate.

It can also be prepared by passing CO₂ through a saturated solution of sodium carbonate. Being less soluble, the white crystals of sodium hydrogen carbonate can be filtered out and dried at room temperature.

Na₂CO₃ + H₂O+ CO₂ ⇌ 2NaHCO₃↓

Properties of sodium bicarbonate:

1. State: It is a white crystalline solid and is sparingly soluble in cold water. It is also stable in air.

2. Hydrolysis: Being a salt of weak acid (H₂CO₃) and strong base (NaOH), it hydrolyses to give a faintly alkaline solution.

NaHCO₃ + H₂O ⇌  [Na⁺ + OH^–] + H₂CO₃

3. Action of heat: On heating, it decomposes to form CO₂, water, and sodium carbonate.

4. Reaction with acids: At ordinary temperature, it reacts with mineral acids to form CO₂, water, and the sodium salt of the acid:

NaHCO₃ + HCl→ NaCl + CO₂ ↑ + H₂O

Uses of sodium bicarbonate:

• It is used as an antacid (known as soda bi-carb). It is also used as a mild antiseptic for skin infections.
• It is the chief ingredient of ‘baking powder’ which is used in preparing breads, biscuits, cakes etc.
• It is used in the preparation of soft drinks like soda- water, lemonades, etc.
• It is also used in fire extinguishers.

Class 11 S-Block Notes

3. Sodium hydroxide (caustic soda), NaOH

Manufacture of sodium hydroxide:

Sodium hydroxide is industrially prepared by the electrolysis of an aqueous solution of NaCl (brine) in a specially designed cell called the Castner-Kellner cell or mercury cathode cell.

Sodium hydroxide Principle:

When a brine solution is electrolyzed in a cell using a mercury cathode and graphite anode, metallic sodium discharged at the cathode combines with mercury to form sodium amalgam. Now, electrolysis of slightly alkaline water in the cell using sodium amalgam as anode and iron rod as cathode produces NaOH. The reaction between sodium amalgam and water also produces NaOH.

Sodium hydroxide  Procedure:

1.  The cell consists of a large rectangular iron tank divided into three compartments by two slate partitions which do not touch the bottom of the tank but remain suspended in mercury placed in the grooves

2. The graphite anodes are fixed in the two outer compartments and the cathode which consists of several iron rods is fitted in the central compartment.

3. The layer of mercury at the bottom serves as an intermediate electrode as a cathode in the outer compartment and as an anode in the central compartment by induction.

4. The brine solution is taken in the two outer compartments and a very dilute NaOH solution is taken in the central compartment.

5. The mercury layer is made to flow from one compartment to another by rocking the cell with the help of an eccentric wheel.  On passing electric current, the following reactions take place in the outer and central compartments.

6. In the outer compartment, NaCl undergoes electrolysis. Cl₂ gas formed at the anode comes out from the outlet tube while sodium liberated at the cathode combines with mercury to form sodium amalgam.

NaCl → Na⁺+ Cl^– ; H₂ O → H⁺ + OH^–

• At cathode: Na⁺ + e →Na; Na + Hg→ Na/Hg
• At anode,: Cl^–→ Cl + e; Cl + Cl→ Cl₂ ↑

7. In the central compartment, sodium amalgam (Na/Hg) acts as an anode by induction.

⇒ \(\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{H}^{+}+\mathrm{OH}^{-}, \mathrm{NaOH} \rightarrow \mathrm{Na}^{+}+\mathrm{OH}^{-}\)

• At cathode: \(\mathrm{H}^{+}+e \longrightarrow \mathrm{H} ; \mathrm{H}+\mathrm{H} \longrightarrow \mathrm{H}_2 \uparrow\)
• At anode: \(\mathrm{Na} / \mathrm{Hg} \longrightarrow \mathrm{Na}^{+}+e+\mathrm{Hg}\)

Net reaction: \(2 \mathrm{Na} / \mathrm{Hg}+2 \mathrm{H}_2 \mathrm{O} \rightarrow 2\left(\mathrm{Na}^{+}+\mathrm{OH}^{+}\right)+2 \mathrm{Hg}+\mathrm{H}_2 \uparrow\)

The concentration of NaOH in the central compartment gradually increases with the progress of electrolysis and when it becomes 20%, the solution is withdrawn, evaporated and converted into pellets or flakes of NaOH.

Class 11 S-Block Notes

Properties of sodium hydroxide:

1. State: It is a white, crystalline hygroscopic solid having a melting point of 318°C.

2. Solubility: It dissolves in water with the evolution of heat, giving a strong alkaline solution. It also dissolves in alcohol.

3. Hygroscopic and corrosive nature:

The crystals of NaOH are deliquescent (hygroscopic). When exposed to air they absorb moisture from air and dissolve in the absorbed water. Moist caustic soda generally absorbs CO₂ from air to form sodium carbonate which forms a coating over the surface of the material. As Na₂CO₃ is non-hygroscopic, wet sodium hydroxide becomes dry again.

⇒ \(2 \mathrm{NaOH}+\mathrm{CO}_2 \rightarrow \mathrm{Na}_2 \mathrm{CO}_3+\mathrm{H}_2 \mathrm{O}\)

It is corrosive. When its concentrated solution comes in contact with the skin it produces a burning sensation. It breaks down the proteins of the skin and because of this property, it is commonly called caustic soda,

4. Reaction with acids, acidic oxides, and amphoteric oxides: Being a strong alkali, it reacts with acids, acidic oxides, and amphoteric oxides to form corresponding salts.

NaOH + HCl→ NaCl +H₂O

2NaOH + SO₂→ Na₂SO₃ +H₂O

Al₂O₃ + 2NaOH →2NaAlO₂ (Sodium aluminate) + H₂O

ZnO + 2NaOH → Na₂ZnO₂ (Sodium zincate) + H₂O

Uses of sodium hydroxide:

It is used

• In the manufacture of soap, paper, artificial silk, dyes, and several chemicals
• In the refining of [etroleum and vegetable oil,
• In the purification of bauxite,
• As a cleaning agent for greasy machines and metal NItoots,
• As a laboratory reagent etc.

4. Sodium chloride (common salt), NaCI

Preparation of sodium chloride:

• The main source of sodium chloride Is seawater which contains 2.7-2,9% of the salt by mass. In tropical countries like India, common salt is generally obtained by the evaporation of seawater.
• Crude sodium chloride obtained by this process contains calcium sulfate (CaSO.), sodium sulfate (Na₂SO₄), calcium chloride (CaCl₂), magnesium chloride (MgCl₂), etc as impurities.
• Since MgCI₂ and CaCl₂ are deliquescent (absorb moisture from the air), impure common salt gets wet in the rainy season.
• For purification, a saturated solution of crude NaCI is prepared and the insoluble impurities are removed by filtration.
• The filtrate is then saturated with hydrogen chloride gas and crystals of pure NaCI separate out due to the common ion effect.
• Chlorides of Ca and Hg being more soluble remain dissolved in the solution. NaCI can also be prepared from rock salt.

Properties of sodium chloride:

• NaCI is a white crystalline solid that melts at 1081K.
• 36 g of NaCI is soluble in 100g of water at 373K. However, solubility does not increase much with an increase in temperature.

Uses of sodium chloride:

• It is used as common salt or table salt for domestic purposes.
• It is used in the manufacture of sodium, caustic soda (NaOH), chlorine, washing soda, sodium peroxide, sodium sulfate, etc.
• It is used in soap industry, in softening hard water, in freezing mixtures, and for regenerating ion exchange resins.

Biological Importance Of Sodium And Potassium

Sodium and potassium ions are the most common cations present in biological fluids. A person weighing 70kg contains about 90g of Na and 170g of K along with 5g of Fe and 0.06 g of Cu.

The daily requirement of Na and K for the human body is about 2 g each.

1. The Na⁺ ions are mainly found outside the cells, in blood plasma, and in the interstitial fluid that surrounds the cells. These ions take part in the transmission of nerve signals, in regulating the flow of water across the cell membranes, and in the transportation of various amino acids and sugars into the cells.
2.  K⁺ ions are the most abundant cations in the cell fluids, where they activate a variety of enzymes, and promote the oxidation of glucose into ATP (adenosine triphosphate), and Na⁺ ions promote the transmission of nerve signals.
3. The Na⁺ and K⁺ ions differ considerably in concentration on the opposite sides of the cell membrane. In blood plasma, for example, the concentrations of Na+ and K+ ions are 143 million-L^-1 and 5 million-L^-1 respectively.
4. Within the blood cells, however, the concentrations of these ions are 10 millimol-L^-1 and 105 millimol-L^-1 respectively.
5. The activities in a nerve cell depend upon the sodium-potassium ion gradient. These ionic gradients are maintained by an ion transport mechanism that operates the active inclusion of K⁺ ions and active exclusion of Na⁺ ions across the cell membrane.
6. The transportation of ions requires energy which is obtained by hydrolysis of ATP. The hydrolysis of one ATP molecule to ADP provides enough energy to move three Na⁺ ions out of the cell two K⁺ ions and one H+ ion back into the cell.

Class 11 Chemistry S Block Elements Group-2 Elements (Alkaline Earth Metals) Introduction

The outermost shell of these elements contains two electrons and the penultimate shell contains eight electrons except forthe first member beryllium which contains two electrons.0 Since the last electron enters the ns orbital, these are also called s-block elements.

Their outermost electronic configuration may be represented as ns², where n- 2 to 7. Due to their similarity in electronic configuration, they are placed in the same group (Group- 2) of the periodic table and closely resemble each other in physical and chemical properties. Two valence electrons are always involved together giving rise to uniform bivalency of the elements.

Class 11 S-Block Notes

Beryllium shows some abnormal properties as its electronic configuration is slightly different from the rest of the members. The main reason is that both the beryllium atom and Be²⁺ ion are extremely small. Beryllium also shows some similarities with aluminum of group 13. Like alkali metals, the compounds of these metals are also predominantly ionic. The electronic configurations of alkaline earth metals are given in the following table

Electronic configuration!) of alkaline earth metals

Occurrence Of Alkaline Earth Metals

Due to low ionization enthalpies and high electropositive character, the alkaline earth metals are chemically very reactive and hence, do not occur in the free state but are widely distributed in nature as silicates, carbonates, sulfates, and phosphates.

• Relative abundance of Be, Mg, Ca, Sr, Ba, and Ra in the earth’s crust is 2, 27640, 46600, 384, 390, and 1.3 x 10-6 ppm respectively. 0 Beryllium, the fifty-first most abundant element by mass in the earth’s crust, is found as silicate minerals like beryl (Be₃Al₂Si₆O₁₈) and phenacite (Be₂SiO₄).
• Magnesium, the sixth most abundant element is found as carbonate, sulphate, and silicate. Its two important minerals are magnesite (MgCO₃) and dolomite [MgCO₃.CaCO₃].
• It is also found in seawater at 0.13% as MgCl₂ and MgSO₄.
• Calcium, the fifth most abundant element by mass found in the earth’s crust, occurs mainly as CaCO₃ in the form of limestone, marble and chalk. Its other important minerals are fluorspar (CaF2), fluorapatite, [3Ca₃(PO₄)2-CaF₃], gypsum (CaSO₄-2H_2O) anhydride, (CaSO₄).
• Strontium and barium are respectively the fifteenth and sixteenth most abundant element. Strontium occurs principally as the mineral celestite (SrSO₄) and strontianite (SrCO₃) while barium occurs mainly as the mineral barytes (BaSO₄).
• Radium is radioactive and extremely scarce. It occurs in very small amounts (1 gin 7 ton) in pitchblende as the decay product of uranium.

General Trends In Atomic And Physical Properties Of Alkaline Earth Metals

The alkaline earth metals show regular trends in their physical and chemical properties with an increase in atomic number. Some important atomic and physical properties of alkaline earth metals are given in the

Atomic and physical properties of alkaline earth metals:

General trends in different atomic and physical properties of alkali metals and their explanations

1. Atomic and ionic radii

The atomic and ionic radii of alkaline earth metals are fairly large but smaller than those of the corresponding alkali metals and these increase on moving down a group.

Atomic and ionic radii Explanation:

The electrons of alkaline earth metals having a higher nuclear charge are more strongly attracted towards the nucleus.  On moving down the group, the atomic as well as ionic radii increase. The addition of new shells and the increasing screening effect jointly overcome the effect of increasing nuclear charge down the group

2. Ionization enthalpy

1. The first and second ionization enthalpies of alkaline earth metals are quite low and decrease down the group from Be to Ra

Explanation:

The low ionization enthalpies of alkaline earth metals are due to their smaller nuclear charge and larger atomic size (compared to the other succeeding elements of the same period) which result in weaker forces of attraction between valence electrons (ns2) and nucleus. On moving down the group, atomic size increases and the screening effect of the inner shell electrons also increases

Class 11 S-Block Notes

Since the alkali metal atoms are largest in their respective periods, their outermost electrons being far away from the nucleus experience less force of attraction and hence, can be removed easily. These two effects jointly overcome the effect of increasing nuclear charge down the group. Thus, first and second ionization enthalpies decrease down the group.

2. The first ionization enthalpies of alkaline earth metals are higher than those of the corresponding alkali metals but their second ionization enthalpies are lower than those of the corresponding alkali metals.

Explanation:

The alkaline earth metals have higher values of first ionization enthalpy than those of the corresponding alkali metals because they have smaller size and higher nuclear charge which result in stronger forces of attraction between the valence electrons and the nucleus.

The second ionization enthalpy values of alkaline earth metals are much lower than those of the corresponding alkali metals because the loss of the second electron from an alkaline earth metal cation (M⁺) leads to the attainment of a stable noble gas configuration (ns²np⁶) while the loss of the second electron from an alkali metal cation (M⁺) causes loss ofits stable noble gas configuration

Example:

3. Electropositive or metallic character

The alkaline earth metals are highly electropositive and possess high metallic character. However, they are less electropositive than the alkali metals. Their electropositive or metallic character increases on moving down the group.

Electropositive Explanation:

• Due to their relatively low ionization enthalpies, alkaline earth metals have a strong tendency to lose both valence electrons to form dipositive ions. Thus, they exhibit high electropositive or metallic character.
• As their atoms have smaller sizes and higher ionization enthalpies compared to those of the corresponding alkali metals, their tendency to lose valence electrons is less than that of alkali metals. Hence, alkaline earth metals have less electropositive or metallic character as compared to tine alkali metals.
• On moving down the group from Be to Ra, ionization enthalpies decrease due to an increase in atomic radii. Therefore, the tendency to lose electrons increases and so does the electropositive character

Class 11 S-Block Notes

4. Hydration enthalpy

Hydration enthalpies of alkaline earth metal ions are much greater than that of the alkali metal ions & decrease down the group from Be²⁺ to Ba²⁺

Be²⁺ > Mg²⁺ > Ca²⁺ > Sr2+ > Ba²⁺

Hydration enthalpy Explanation:

Due to the smaller size of alkaline earth metal ions, their hydration enthalpies are much greater than those of the alkali metal ions. Therefore, the compounds of alkaline earth metals are found to be more extensively hydrated than those of alkali metals. Magnesium chloride and calcium chloride, for example, exist as hexahydrates (MgCl₂-6H₂O and CaCl₂-6H₂O) while sodium chloride and potassium chloride do not form such hydrates.

The ionic conductance of hydrated alkaline earth metal ions increases from [Be(H₂O)_x]²⁺ to [Ba(HO)_2x]²⁺ due to a decrease in the extent of hydration. The hydration enthalpy of an ion is directly proportional to its charge/radius ratio {q/r). On moving down a group, the radii of the alkaline earth metals increase. As a result, the hydration enthalpies of these metals decrease.

5. Oxidation State

Alkaline earth metals exhibit an oxidation in their compounds. Although the second date of +2 in their compounds. Although the second ionization enthalpy of these elements is nearly double that of the first ionization enthalpy, yet they exist as divalent ions (M²⁺) in most of their compounds.

Oxidation State Explanation:

1. The divalent ions (M²⁺) of alkaline earth metals have stable noble gas configurations. Thus, M²⁺ ion is more stable than M+ ion.

M ([Noble gas] ns²) → M²⁺ [Noble gas] + 2e

2. Due to greater charge and smaller size, the divalent cations lead to the formation of very stable lattices, and hence, a huge amount of energy is released. The high lattice enthalpy easily compensates for the high second ionization enthalpy.

3. Divalent cations for their smaller size get hydrated in water to a greater extent and the energy thus released (hydration enthalpy) is large enough to compensate for the second ionization enthalpy

The ΔH_i(3) values of alkaline earth metals are very high because the electron now has to be removed from the stable noble gas configuration. For this reason, the alkaline earth metals do not exhibit an oxidation state of more than +2.

6. Melting & boiling points

Alkaline earth metals have higher melting & boiling points than that of alkali metals. However, on moving down the group, no regular trend is observed.

Melting & boiling points Explanation: 

• Due to their smaller size, the atoms of alkaline earth metals form a more close-packed crystal lattice. Moreover, alkaline earth metals have two electrons in their valence shell whereas alkali metals have only one.
• The larger number of valence electrons leads to the formation of stronger metallic bonds.
• No regular trend in melting and boiling point is observed down the group because the atoms adopt different crystal structures.

7. Nature of bonds formed:

Like alkali metals, alkaline earth metals predominantly form ionic compounds. However, these are less ionic than the corresponding alkali metal compounds. Beryllium, the first member of this group, is an exception as its compounds are covalent. Magnesium also tends to form covalent compounds to some extent. On moving down the group, the tendency to form ionic compounds increases.

Nature of bonds formed Explanation:

Alkaline earth metals form ionic compounds because they have low ionization enthalpies. Their compounds, however, are less ionic because their ionization enthalpies are higher than those of the corres¬ ponding alkali metals. Due to its much smaller size and much higher ionization enthalpy, beryllium forms compounds that are predominantly covalent. Down the group, the tendency to form ionic compounds increases because ionization enthalpy decreases.

Class 11 S-Block Notes

8. Density and hardness

The alkaline earth metals are denser and harder than the corresponding alkali metals. However, on moving down the group, no regular trend is observed. It initially decreases from Be to Ca and then increases from Ca to Ba.

Density and hardness Explanation:

The extent of cohesive energy determines the density and hardness of metals and this depends on the number of electrons involved in metallic bonding and the size of the atom. In alkali metals, one electron per atom (the valence electron) is involved in metallic bonding while in alkaline earth metals, two electrons per atom (the valence electrons) are involved. Moreover, the atoms of alkaline earth metals are heavier and smaller in size.

Therefore, the extent of cohesive energy is relatively higher in the case of alkaline earth metals & consequently, the atoms in alkaline earth metals are packed more closely in their lattices. Cohesive energy decreases from Be to Ca due to a gradual increase in size while it is found to increase from Ca to Ba due to the formation of different crystal lattices.

9. Conductivity

The Gr-2 metals are good conductors of heat and electricity.

Conductivity Explanation:

Due to the presence of two loosely bound valence electrons (per atom) which can move freely throughout the crystal lattice, the alkaline earth metals are good conductors of heat and electricity.

10. Flame coloration

When the alkaline earth metals and their salts, except beryllium and magnesium, are heated in the flame of a bunsen burner, they impart characteristic color to the flame.

These colors are as follows:

• Ca: Brickred
• Ba: Apple green
• Sr & Ra: Crimson red

Flame coloration Explanation:

• When the alkaline earth metals or their salts are put into a flame, the electrons of their valence shell absorb energy and get excited to higher energy levels.
• When they drop back to the ground state, the absorbed energy is emitted in the form of visible light having characteristic wavelengths.
• Depending upon the wavelength of light emitted, different colors are imparted to the burner flame.
• Due to their smaller size, the valence electrons in Be and Mg are too strongly bound to get excited by the energy available from the flame. Therefore, they do not impart any color to the flame.

Alkaline earth metals (except Be and Mg) can easily be identified by flame test in qualitative analysis. Further, they can be estimated by flame photometry or atomic absorption spectroscopy.

11. Magnetic property

The alkaline earth metals and their salts are diamagnetic.

Magnetic property Explanation:

Since the divalent ions (M²⁺) of alkaline earth metals have noble gas configurations with no unpaired electrons, their salts are diamagnetic. The metals are also diamagnetic as all the orbitals are filled up with paired electrons.

Chemical Properties Of Alkaline Earth Metals (Group-2 Metals)

Due to their low ionization enthalpies and high electropositive character, alkaline earth metals have a strong tendency to lose their valence electrons. Therefore, they are highly reactive and do not exist in the free state in nature.

1. Reducing nature

The alkaline earth metals are strong reducing agents. However, they are weaker reducing agents than alkali metals. Again, like alkali metals, their reducing strength increases down the group.

Class 11 S-Block Notes

Reducing nature Explanation:

The alkaline earth metals except Be, have a fairly strong tendency to lose two valence electrons to form dipositive ions (M→ M²+  2e) i.e. they possess low ionization enthalpies and hence, they are strong reducing agents.

This is indicated by their high negative values of reduction potentials (E°). Their reducing strength, however, is less than the alkali metals as their atomization enthalpies and ionization enthalpies are relatively higher. Reducing strength increases on moving down the group as their ionization enthalpies decrease & electrode potentials become progressively more negative from Be to Ba.

2. Action of air

• Being fairly reactive, the alkaline earth metals are oxidized by the oxygen of the air and get tarnished due to the formation of a fine layer of oxide on their surface. With increasing atomic numbers, the effect of air on the metals gradually increases.
• Be and Mg being less reactive are not much affected by air. Ca and Sr get easily tarnished in air while Ba readily burns when exposed to air. Hence, Ca, Ba, and Sr are usually stored in paraffin.

3. Reaction with oxygen

Alkaline earth metals burn in oxygen to form oxides. Be, Mg, and Ca form monoxides while Sr and Ba form peroxides when they react with oxygen. This is because larger cation stabilizes a larger anion and hence the tendency to form peroxide increases as the size of the metal ion increases

(M = Be, Mg or Ca)

(M = Ba, Sr)

4. Reaction with water

• Alkaline earth metals except beryllium react with water to form the corresponding hydroxides along with the liberation of H₂ gas.
• Beryllium having the lowest negative standard electrodepotential (E° of Be²⁺/Be = -1.97V) among all

The group-2 metals is the least electropositive and hence, do not react with water or steam even at red hot conditions.

Ca, Sr, and Ba have relatively higher negative standard electrode potentials similar to those of the corresponding Gr-1 metals and hence, react even with cold water.

Mg has an intermediate value of E° and does not react with cold water but decomposes in boiling water.

M + 2H₂O→M(OH)₂ + H₂↑ (M = Mg, Ca; Sr or Ba)

Thus, the reactivity of the alkaline earth metals towards water increases on moving down the group. However, they are less reactive towards water as compared to the corresponding alkali metals.

5. Reaction with nitrogen

1. All alkaline earth metals burn in nitrogen to form nitrides of the type M3N2. However, Li forms Li3N.

3M + N₂ →M₂N₂ (M = Be, Mg, Ca, Sr and Ba)

2.  The ease of formation of nitrides decreases from Be to Ba. Since N₂ molecule is very stable, it requires very high energy to form nitride ions (N^3-). This large amount of energy is supplied from the lattice enthalpy evolved when crystalline solids containing ions with high charges (M²⁺ and N^3-) are formed.

Be₃N₂ is volatile because it is covalent. Other nitrides of this group are not volatile as they are ionic crystalline solids.

Class 11 S-Block Notes

6. Reaction with halogens

The alkaline earth metals directly combine with halogens at higher temperatures to form halides having the general formula, MX₂

Halides can also be obtained by the action of halogen acids on metals, their oxides, hydroxides, or carbonates

M + 2HX→ MX₂ + H₂; MO + 2HX→MX₂ + H₂O

M(OH)₂ + 2HX→MX₂ + 2H₂O

MCO₃ + 2HX→MX₂ + CO₂ + H₂O

BeCl₂ is, however, conveniently prepared by heating BeO with Cl₂ in the presence of charcoal at 1073K

7. Reaction with hydrogen:

All the elements of group-2 except Be, form metal hydrides of the general formula MH₂ when heated with hydrogen

Beryllium hydride can be prepared indirectly by reducing beryllium chloride with lithium aluminum hydride.

2BeCl₂ + LiAlH₄→ 2BeH₂ + LiCl + AlCl₃

Both beryllium hydride (BeH₂) and magnesium hydride (MgH₂) are covalent compounds. In these molecules, both Be and Mg have four electrons in their valence shell. Therefore, these molecules are electron deficient. To make up for their electron deficiency, these two compounds exist as polymers, (BeH₂)_n and (MgH₂)_n in which each Be or Mg -atom forms four three-centre two-electron (3c-2e) bonds or hydrogen bridge bonds or banana bonds.

Class 11 S-Block Notes

The structure of polymeric beryllium hydride is shown below:

CaH₂, SrH_2, and BaH₂ are ionic compounds in which a hydride ion (H^–) exists as an anion. Calcium hydride (CaH₂ ) which is also called hydrolith is used for the production of H₂ by the action of HaO on it.

8. Reaction with carbon

When the alkaline earth metals except for Be, are heated with carbon in an electric furnace or when their oxides are heated with carbon, carbides of the type MC₂ are obtained. These carbides are also called acetylides (containing discrete C₂ ions) as on hydrolysis they form acetylene.

(M = Mg, Ca, Sr or Ba)

At much higher temperatures ( ~ 1700°C), beryllium reacts with carbon to form Be₂C. This carbide is called methanide (containing discrete C^4- ion) as on hydrolysis it produces methane. On heating, MgC₂ forms Mg₂C₃, which is called allylide (containing discrete C₃^4- ion) as hydrolysis yields allylene (methyl acetylene).

CaC₂ + 2H₂O → HC≡ CH + Ca(OH)₂

Be₂C + 4H₂O →  2Be(OH)₂ + CH₄

Mg₂C₃+ 4H₂O →  CH₃C ≡ CH + 2Mg(OH)₂

When calcium carbide (CaC₂), an important chemical intermediate, is heated in an electric furnace with atmospheric nitrogen at 1375K, it produces calcium cyanamide (CaNCN).

The mixture of CaNCN and carbon is called nitrolim. It is used as a slow-acting nitrogen fertilizer as it undergoes very slow hydrolysis and evolves NH3 gas for a long period.

CaC₂ + 3H₂O → CaCO₃ + 2NH₃

9. Reaction with acids

The alkaline earth metals react with dilute acids to form the corresponding salt with the liberation of H₂ gas.

M + H₂SO₄ → MSO₄ + H₂↑ (M = Be, Mg, Ca, Sr or Ba)

Beryllium is the only group-2 metal which reacts with alkali to form H₂ and beryllate salt.

Be + 2NaOH + 2H₂O → Na₂ [Be(OH)₄] (Sodium beryllate)+ H₂ ↑

This is observed due to the diagonal relationship between aluminum and beryllium.

10. Solutions in liquid ammonia

Like alkali metals, alkaline earth metals dissolve in ammonia to give deep blue-colored solutions containing ammoniated cations and ammoniated electrons.

M + (x+ 2y)NH₃ → [M(NH₃)_x]²⁺ + 2[e(NH₃)_y ]^–

Class 11 S-Block Notes

Evaporation of ammonia from these solutions leads to the formation of hexammoniates M(NH₃)₆ which slowly decompose to yield the corresponding metal amides, M(NH₂)₂ and H₂

11. Tendency to form complexes

The group-2 elements tend to form stable complexes and it is found to be greater than that of alkali metals because their ions have smaller size and higher charge. The tendency to form complexes decreases down the group and this is due to the decrease in ion-dipole interaction with increasing size of the metal ion. Be and Mg have the maximum tendency to form complexes.

Examples of two stable complexes of Be and Mg are \(\left[\mathrm{BeF}_4\right]^{2-}\&\left[\mathrm{Mg}\left(\mathrm{NH}_3\right)_6\right]^{2+}\) respectively

1. Complexation of Ca²⁺ by EDTA and polyphosphates plays an important role in the removal of the metal in water softening.
2.  In chlorophyll, the complex formed by the combination of Mg and the tetrapyrrole system (porphyrin) is very crucial in photosynthesis.

12. Extraction of alkaline earth metals

Like alkali metals, alkaline earth metals are also very reactive and strong reducing agents. So they cannot be extracted by ordinary chemical reduction methods. These metals also cannot be prepared by electrolysis of aqueous solutions of their salts because in that case, hydrogen is discharged at the cathode instead of the metal which has a much higher discharge potential. However, electrolysis can be carried out using a Hg-cathode, but in that case, recovery of the metal from amalgam becomes difficult. These metals are best isolated by electrolysis of their fused salts, usually chlorides.

General Characteristics Of The Compounds Of Alkaline Earth Metals

Compounds of group-2 elements are predominantly ionic but are less ionic than the corresponding compounds of group 1 elements and this is due to their increased nuclear charge and smaller size. The general characteristics of some of the compounds of alkaline earth metals are discussed below.

1. Oxides of alkaline earth metals

1. Crystal structure:

Except for BeO (covalent solid), the oxides of the remaining alkaline earth metals are crystalline ionic solids and possess a rock-salt (NaCl) structure with coordination number 6. BeO though covalent, is an extremely hard solid because of its polymeric nature. BeO possesses a covalent lattice with coordination number 4. Both BeO and MgO have several properties that make them useful as refractory materials (for lining furnaces).

These properties are:

• They have high melting points (BeO-2500°C and MgO – 2800°C),
• They have very low vapor pressures,
• They are good conductors of heat, O they are chemically inert and
• They can act as electrical insulators.

2. Stability:

Due to much higher lattice enthalpies, the oxides are very stable towards heat. The lattice enthalpies decrease with an increase in the size of the metal ion.

3. Basic character:

Beryllium oxide, BeO reacts with both acids and alkalis, i.e., it is amphoteric while the oxides of other group-2 metals are basic.

BeO + 2HCl→BeCl₂ + H₂O

BeO + 2NaOH → Na₂ BeO₂ (Sodium beryllate) + H₂O

The basic strength of the oxides increases on moving down the group.

BeO (Amphoteric) < MgO (Weakly basic) < CaO (Basic ) < SrO, BaO( Strongly basic)

4. Reaction with water:

All these oxides except BeO and MgO, react with water to form sparingly soluble hydroxides. These reactions are exothermic.

MO + H₂O→M(OH)₂ + heat, M = Ca, Sr or Ba

2. Hydroxides

1. Basic character:

All the alkaline earth metal hydroxides are basic except Be(OH)₂ which is amphotericin nature. Their basic strength increases on moving from Be(OH)₂ to Ba(OH)₂

The alkaline earth metal hydroxides are, however, less basic than the alkali metal hydroxides.

Basic character Explanation:

Due to low ionization enthalpies of the alkaline earth metals, the M — O bond present in their hydroxides is weak and breaks up easily to give OH^– ions. For this reason, their hydroxides exhibit basic character. On moving down the group, the tendency of the M — OH bond to break heterolytically increases because ionisation enthalpies decrease and consequently, the basic character of the hydroxides increases.

Class 11 S-Block Notes

Due to larger ionic sizes and lower ionization enthalpies of alkali metals, the M — OH bonds in their hydroxides are still weaker than those in alkaline earth metal hydroxides. Thus, the alkali metal hydroxides are more basic than the alkaline earth metal hydroxides.

2. Solubility in water:

Hydroxides of alkaline earth metals are less soluble in water than the hydroxides of alkali metals. Again, the solubility of these hydroxides increases markedly on moving down the group

Solubility in water Explanation:

On moving down the group, both the lattice enthalpy and the hydration enthalpy decrease with increasing ionic size. However, the lattice enthalpy decreases more rapidly than the hydration enthalpy, and consequently, their solubility increases down the group.

3. Thermal stability:

The alkaline earth metal hydroxides decompose on heating to give the metal oxide and water.

Thermal stability Explanation:

Thermal stability of these hydroxides increases down the group as the polarising power of the M2+ ion and the lattice enthalpy of the oxide formed decreases with increasing ionic size down the group.

3. Halides

• Due to the high polarising power of the Be²⁺ ion, beryllium halides have a covalent nature having low melting points.
• All other alkaline earth metal halides are ionic and their ionic character increases as the size of the metal ion (M²⁺) increases down the group. These ionic halides are non-volatile solids having high melting points.
• Due to its covalent nature, beryllium halides are sparingly soluble in water but readily soluble in organic solvents. The halides of other group-2 alkaline earth metals are readily soluble in water.
• Except for BeCl₂, all other anhydrous halides of the alkaline earth metals are hygroscopic in nature and form hydrates.
• For example: MgCl₂-6H₂O, CaCl₂-6H₂O, SrCl₂-2H₂O and BaCl₂-2H₂O
• The tendency to form hydrate decreases down the group. Thus, anhydrous calcium chloride is used as a dehydrating agent in the laboratory.
• The dehydration of the hydrated chlorides, bromides, and iodides of Ca, Sr, and Ba can be achieved by heating. However, the corresponding hydrated halides of Be and Mg on heating suffer hydrolysis.
• BeF₂ is highly soluble in water due to the much higher hydration enthalpy of the very small Be²⁺ ion.
• All other fluorides (MgF₂ > CaF₂, SrF_2, and BaF₂ ) are almost insoluble in water because their lattice enthalpies are higher than their hydration enthalpies.
• Except for BeCl₂ and MgCl₂, all other alkaline earth metal chlorides impart characteristic color to a flame.
• For example: CaCl₂: is brick red, SrCl₂:  Is crimson red, BaCl₂: Is grassy green, etc.

Structure of BeCl₂ In the solid state, beryllium chloride has a polymeric chain structure with chlorine bridges as given below:

Which are bonded by two covalent bonds while the other two by coordinate bonds. The Be -atoms in (BeCl₂)n sp³ -hybridized.

In the vapor phase, BeCl₂ exists as a chlorine-bridged dimer which dissociates into linear triatomic monomer at about 1200K. In the dimer, Be is sp² -hybridized while in the monomeric is sp² -hybridized

• CaF₂ an industrially important compound, is the main source of both F₂ and HF

CaF₂ + H₂SO₄4-→ 2HF + CaSO₄

• CaF₂ is also used for making prisms and cell windows for spectrophotometers, an important instrument used in the spectroscopic analysis of compounds.
• In cold countries, CaCl₂ is Used for treating ice on roads, because 30% eutectic mixture of CaCl₂ +H₂O freezes at  – 55%C.

4. Salts of Oxoacids

Trends in the properties of some salts of group-2 elements are discussed below

1. Carbonates:

Solubility in water:

The carbonates of the alkaline earth metals are practically insoluble in water. Their solubilities decrease on moving down the group. BeCO₃ is sparingly soluble in water while BaCO₃ is insoluble in water.

Solubility in water Explanation:

On moving down the group, lattice enthalpy of carbonates remains almost unchanged” (the size of the metal ion is much smaller compared to CO₃^2-  ion) but hydration enthalpies of cations (M²⁺) decrease. Consequently, the solubilities of carbonates decrease down the group.

The extremely low solubility of alkaline earth metal carbonates in water is very important in the precipitation of Ba²⁺, Sr^2+, and Ca²⁺ ions as their carbonates, in Gr-IV qualitative analysis of basic radicals.

Thermal stability: Carbonates of gr.-2 metals decompose on heating to give metal oxide and carbon dioxide

Thermal stability of the carbonates increases down the group with increasing cationic size

Thermal stability Explanation:

The trend can be explained in terms of the stability of the monomeric is sp –hybridized. of the resulting metal oxides. With increasing stability of  Cl the metal oxide, the carbonate becomes more unstable BeCl₂ [monomer] towards heat. The stability of metal oxides decreases down the group due to a decrease in lattice enthalpy with increasing cationic (M²⁺) size. Hence, the stability of the carbonates towards heat increases down the group.

2. Bicarbonates:

The bicarbonates of alkaline earth metals do not exist in the solid state but are found in solutions. When such solutions are heated, bicarbonates decompose to form carbonates with the evolution of CO₂

3. Sulfates

1. Solubilityin water:

The sulfates of alkaline earth metals are relatively less soluble in water than the corresponding sulfates of alkali metals. Further, their solubilities decrease down the group. For example, BeSO₄ and MgSO₄ are more soluble in water, CaSO₄ is less soluble in water and SrSO₄, BaSO_4, and RaSO₄ are practically insoluble in water.

Solubilityin water Explanation:

On moving down the group, the hydration enthalpies decrease with increasing cationic size but the lattice enthalpy remains almost unchanged because the anion, SO₄ is much larger as compared to the metal ions (M²⁺). For this reason, the solubilities of sulfates decrease down the group.

2. Thermal stability: The alkaline earth metal sulfates white solids which dissociate on heating to give the metal oxides and sulfur trioxide.

The thermal stability of the sulfates increases down the group due to an increase in ionic character. This is due to a decrease in the polarising power of the metal ions with increasing ionic size. It becomes evident from the dissociation temperatures of the sulfates given below

4. Nitrates

Nitrates of alkaline earth metals are prepared by heating corresponding metal carbonate with dilute HNO₃

MCO₃  + 2HNO₃→ M(NO₃) + H₂ O + CO₂ (M = Be, Mg, Ca, Sr or Ba)

Magnesium nitrate crystallizes as Mg(NO₃)₂. 6H₂O whereas barium nitrate crystallizes as an anhydrous salt. This again shows that the tendency to form hydrates decreases with increasing ionic size and decreasing hydration enthalpy as we move down the group. Upon heating, all nitrates decompose to give the corresponding oxides with the evolution of NO₂ and O₂.

The nitrates of all these metals are soluble in water. Beryllium is unusual for the fact that it forms a basic nitrate in addition to the normal salt.

Anomalous Behaviour Of Be And Similarities Between Be And Al

Beryllium, the first member of group 2, shows some anomalous behavior, i.e.,it differs from the rest of the members of its family.

Reasons for anomalous behavior of beryllium:

• The extremely small size of the Be atom and Be²⁺ ion,
• Much higher polarising power of Be²⁺ ion,
• Higher ionization enthalpy and electronegativity as compared to the other members, absence of vacant d -d-orbitals in its valence shell.
• Again, beryllium resembles its diagonally placed element aluminum, the second typical element of group 13 and period 3, in several properties

Class 11 S-Block Notes

Difference between beryllium and other alkaline earth metals

Reasons for the similarities between beryllium and aluminum:

• They have approximately the same polarising power. The polarising power of Be²⁺ is 0.064, while that of Al³⁺is 0.060.0
• The standard electrode potentials of Be and A1 are much closer i.e. Be²⁺/Be = -1.85V and Al³⁺+/Al = -1.66V.0
• The electronegativity of both the elements i.e., beryllium and aluminum are the same (1.5in the Pauling scale).

Similarities between beryllium and aluminium:

Uses Of Alkaline Earth Metals

Preparation, Properties, And Uses of Some Important Compounds Of Calcium

1. Calcium oxide (quicklime), CaO

Preparation of calcium oxide:

Calcium oxide or quicklime is prepared commercially by heating limestone at about 1250K temperature in lime kilns.

CaCO₃ ⇌ CaO + CO₂ ↑- 425 kcal

The reaction is endothermic and reversible. To get a good yield of quicklime, the forward reaction is facilitated by removing CO₂ as soon as it is formed. Again, the temperature in the kiln is not allowed to rise above 1270K because above that temperature, silica (SiO₂ ) present as

impurity in limestone combines with CaO to form calcium silicate (CaSiO₃).

CaO + SiO₂ >127°K>CaSiO₃

Properties of calcium oxide:

1. State: It is a white amorphous solid having a melting point of of2870 K.

2. The action of heat:

Calcium oxide does not melt even when heated in an oxy-hydrogen flame (2270K). On strong heating, it becomes incandescent and emits bright white light (known as limelight). It melts only when heated in an electric furnace at 2850K.

3. The action of air:

When dry lumps of CaO are exposed to moist air, they absorb moisture and CO₂ from the air to form calcium hydroxide and calcium carbonate respectively. As a result, heat is evolved and the lumps of CaO are converted into powder. Calcium hydroxide thus produced also reacts with CO₂ of air to form calcium carbonate. Therefore, calcium carbonate is the final product we get. However, once the outer surfaces of the lumps of CaO become fully covered with CaCO₃, the core material is not further acted upon by moist air.

CaO + H₂O-+Ca(OH)₂; CaO + CO₂→CaCO₃

Ca(OH)₂ + CO₂→ CaCO₃ + H₂O

4. Reaction with water:

Calcium oxide possesses a high affinity towards water. Many organic liquids and moist gases are dried with CaO. It reacts vigorously with water to form Calcium hydroxide:

CaO + H₂O→ Ca(OH)₂

When a limited amount of water is sprinkled on the lumps of calcium oxide, a vigorous reaction starts. For its highly exothermic nature, the water added gets immediately transformed into steam with a hissing sound. As a result, the lumps of CaO swell up, crack, and finally crumble to a fine, dry, white powder of calcium hydroxide. Such powdered calcium hydroxide is known as slaked lime and the above process is called the slaking of lime

Quicklime when slaked with caustic soda (NaOH), produces a solid called soda lime (CaO + NaOH).

5. Reaction with acids and acidic oxides:

Calcium oxide is a basic oxide and hence reacts with acids to form the corresponding calcium salts and water.

Examples:

CaO + 2HCl→CaCl₂+ H₂O

CaO + H₂SO₄→CaSO₂ + H₂O

CaO + 2HNO₂→ Ca(NO₃)₂ + H₂O

In the reaction with sulphuric acid, insoluble calcium sulfate is produced and it forms a protective coating on the lumps of CaO. Consequently, the reaction proceeds only to a small extent and then stops.

CaO reacts with acidic oxides to form calcium salts.

Class 11 S-Block Notes

Examples:

CaO + CO₂ → CaCO₃ ; CaO + SO₂→ CaSO₃

6CaO + P₄O₁₀ →\(\rightarrow{\Delta}\) 2Ca₃(PO₄)₂

It reacts with silica at a much higher temperature to form calcium silicate: CaO + SiO₂→ CaSiO₃

6. Reaction with ammonium salts:

Being a strong base, CaO displaces ammonia forming ammonium salts. The reaction occurs rapidly on gentle heating. This reaction may be used for preparing NH3 in the laboratory.

2NH₄Cl + CaO→ 2NH₃↑ + CaCl₂ + H₂O

7. Reaction with chlorine gas:

When calcium oxide is heated in the presence of dry Cl₂ gas above 573K, calcium chloride is obtained with the evolution of oxygen gas.

8. Reaction with carbon:

When calcium oxide is heated with coke in an electric furnace at about 2273K, it forms calcium carbide and carbon monoxide. This reaction is used in the industrial preparation of calcium carbide.

Uses of calcium oxide:

• Quicldime is an important primary material for manufacturing cement and glass.
• It is used to prepare caustic soda from sodium carbonate.
• It is used in the purification of sugar.
• It finds application in the manufacture of dyes.
• It is largely used in the preparation of slaked lime which has many industrial uses.
•  It is used as a flux in metallurgy to remove siliceous impurities.
• It is used to dry several gases and alcohols.
• It is used in the preparation of calcium carbide and soda lime.
• It is used in softening of hard water and in tanning industries.

2. Calcium hydroxide or (slaked lime) Ca(OH₂)

Preparation of calcium hydroxide:

Calcium hydroxide or slaked lime is prepared by sprinkling a limited amount of water on the lumps of calcium oxide. The process is known as slaking of lime. The reaction is highly exothermic.

CaO + H₂O→Ca(OH)₂

It can also be prepared by treating a concentrated aqueous calcium chloride solution with a solution of caustic soda.

CaCl₂ + 2NaOH→Ca(OH)₂↓+ 2NaCl

Calcium hydroxide Physical properties:

Calcium hydroxide or slaked lime is available as a white amorphous powder. It is soluble in water to a very small extent. The clear dilute aqueous solution of calcium hydroxide is known as lime water. When a large amount of calcium hydroxide is added to water, a white suspension (like milk) of the substance in water is obtained. This is known as milk of lime.

Calcium hydroxide Chemical properties:

1. The action of air: When calcium hydroxide is exposed to air, it slowly absorbs CO₂ from the air and is converted into water-insoluble calcium carbonate.

Ca(OH)₂ + CO₂→ CaCO₃ ↓+ H₂O

It is to be noted that the formation of a white scum on the surface of clear lime water, exposed to air, is due to the formation of insoluble CaCO₃.

2. The action of heat: When calcium hydroxide is heated above 723K, it undergoes complete dehydration to yield CaO.

3. Reaction with carbon dioxide:

When carbon dioxide is passed through clear lime water, calcium carbonate is formed. The resultant insoluble calcium carbonate remains suspended in water as fine particles. As a result, clear lime water becomes milky (turbid) in appearance.

Ca(OH)₂ + CO₂ →CaCO₃↓ + H₂O

When excess of CO₂ gas is passed through this milky suspension, the water-insoluble particles of CaCO₃ further react with CO₂  in the presence of water to form water-soluble colorless calcium bicarbonate, Ca(HCO₃)₂. As a result, the turbidity of the solution disappears and it becomes transparent (clear) again.

CaCO₃ + CO₂ + H₂O→ Ca(HCO₃)₂(soluble)

When tiie clear solution of calcium bicarbonate is heated, the solution again becomes turbid due to the decomposition of Ca(HCO₃) into insoluble CnCO₃.

4. Reaction with sulfur dioxide:

When SO₂ gas is passed through clear lime water, water-insoluble, white calcium sulfite (CaSO₃) is formed and the clear solution becomes turbid. When an excess of SO₂ gas is passed through this turbid solution, SO₂ reacts with CaSO₃ in die presence of water to form water-soluble, colourless calcium bisulfite, Ca(HSO₃)₂. Thus, the turbid solution becomes clear again. When the resultant clear solution is heated, calcium bisulfite decomposes to give back insoluble calcium sulfite along with SO₂ and water. So, the clear solution becomes turbid again.

Ca(OH)2 + SO₂→ CaSO₃↓ + H₂O

CaSO₃ + SO₂ + H₂O→Ca(HSO₃)₂ (soluble)

Class 11 S-Block Notes

5. Reaction with acid:

Calcium hydroxide being a quite strong base reacts with acids and acidic oxides to form the corresponding salts and water.

Examples:

Ca(OH)₂ + 2HCl→CaCl₂ + 2H₂O

Ca(OH)₂ + H₂SO₄→ CaSO₄ ↓+ 2H₂O

Ca(OH)₂ + CO₂ →CaCO₃↓- + H₂O

Ca(OH)₂ + SO₂→ CaSO₃ + H₂O

3Ca(OH)₂+ P₂O₅→ Ca₃(PO₄)₂ + 3H₂O

In its reaction with sulphuric acid, insoluble white calcium sulfate is formed and it forms a protective coating on solid Ca(OH)₂ and stops the reaction.

Reaction with ammonium salts: Being a stronger base than ammonia, calcium hydroxide displaces ammonia from its salts when heated with an ammonium salt.

Examples: 2NH₄Cl + Ca(OH)₂→ 2NH₃T↑ + CaCl₂ + 2H₂O

This reaction is generally used for the preparation of ammonia in the laboratory.

6. Reaction with chlorine:

At about 313K, chlorine gas reacts with slightly moist slaked lime to form a dry, white, and powdery substance with a pungent smell. This powder having bleaching and disinfecting properties is commonly called bleaching powder. Regarding the formation and . composition of bleaching powder, the following two views have been proposed.

1. According to Odling (1861), bleaching powder Is a mixture of calcium hypochlorite, Ca(OCl)₂, and calcium chloride, CaCl_2, and Is called calcium chlorohypochlorite, Ca(OCl). Its formation can be shown as follows

2Ca(OH)₂+ 2Cl₂ → Ca(OCl)₂ + CaCl(Bleaching powder) +2 H₂O

Or, 2Ca(OH)₂ + 2CI₂ → 2Ca(OCl)Cl + 2H₂O

Or, Ca(OH)₂ + Cl₂ → Ca(OCl)Cl(Bleaching powder) + H₂O

2. According to Bunn, Clork, and Ghifford (1935), bleaching powder is supposed to be a mixture of

Ca(OCl)₂, CaCl₂ & Ca(OH)₂. Its formation can be represented as

2Cl₂ + 3Ca(OH)₂→ Ca(OCl)₂-Ca(OH)₂-CaCl₂-2H₂O (Bleaching powder)

Wlien Cl₂ gas is passed through excess of cold lime water, calcium chloride, and calcium hypochlorite are formed.

2Ca(OH)₂ + 2CI₂→ CaCl₂ + Ca(OCl)₂ + 2H₂O

When excess chlorine is passed through hot lime water calcium chloride and calcium chlorate are formed.

6Ca(OH)₂ + 6Cl₂→ 5CaCl₂ + Ca(ClO₃)₂ + 6H₂O

Uses of calcium hydroxide

• Calcium hydroxide is used in the manufacture of caustic soda, bleaching powder, superphosphate of lime (a chemical fertilizer), etc.
• It is used in the preparation of mortar, a building material.
• Slaked lime is mixed with three to four times of the weight of sand.
• The mixture is made into a thick paste with a gradual addition of water.
• The paste is called mortar. As the paste becomes dry, it hardens due to the formation of

CaCO₃:  Ca(OH)₂ + CO₂→ CaCO₃ + H₂O

• It is used in the manufacture of glass and in the manufacture of calcium hydrogen sulfate, Ca(HSO₄)₂ which is used in the paper industry.
• It is used in the recovery of ammonia from NH₄Cl (a by-product of the Solvay process), in coal gas purification, in tanneries for removing hair from hides, and for softening hard water. 0 Milk of lime is used for white washing due to its disinfectant properties.
• Lime water is a laboratory reagent for the detection of carbon dioxide.

3. Calcium carbonate (limestone), CaCO₃

Natural occurrence: In nature, calcium carbonate occurs in large quantities as limestone, marble, calcite, chalk, etc. It occurs also in the mineral, dolomite which is the double carbonate of calcium and magnesium (CaCO₃-MgCO₃). Besides these minerals, CaCO₃ occurs in abundance in corals, eggshells, outer covering of oysters, snail’s conch, and in teeth and bones of man and animals.

Preparation of calcium carbonate: Calcium carbonate may

Be prepared by passing CO₂ through lime water or by adding a solution of Na₂CO₃ to a solution of CaCl.

1 Ca(0H)₂ + CO₂ → CaCO₃↓+ H₂O

CaCl₂ + Na₂ CO₃→ CaCO₃↓+ 2NaCl

The precipitate of CaCO₃ is known as precipitated chalk.

Calcium carbonate Physical properties:

1. It is a white solid.
2. It is a stable compound which is almost insoluble in water.

Calcium carbonate Chemical properties:

1. The action of heat:

When heated at a higher temperature ( ~ 1270K), calcium carbonate decomposes to give calcium oxide (quicklime) & carbon dioxide. The reaction is reversible and endothermic. So, it proceeds towards completion when the reaction is carried out in an open vessel.

CaCO₃  ⇌  CaO + CO₂ ↑-heat

2. Reaction with dilute acids: It reacts with dilute acids to form the corresponding calcium salts.

CaCO₃  + 2HCl → CaCl₂ + CO₂↑ + H₂O

CaCO₃ + H₂SO₄ → CaSO₄ + CO₂↑ + H₂O

CaCO₃ + 2HNO₃ → Ca(NO₃)₂ + CO₂↑+ H₂O

3. Reaction with carbon dioxide:

When CO₂  gas is passed through a fine suspension of calcium carbonate in water, the latter slowly dissolves to produce calcium bicarbonate and thus, a clear solution is obtained.

CaCO₃ + CO₂ + H₂O→ Ca(HCO₃)₂

When the resultant clear solution is heated, calcium bicarbonate decomposes back to insoluble CaCO₃. Thus, the clear solution becomes turbid again. ‘

Uses of calcium carbonate:

• Calcium carbonate is largely used in the manufacture of quicklime, cement, and glass.
• Along with MgCO_3, it is used as a flux in the extraction of metals such as iron.
• It is used as a building material in the form of marble and the construction of statues.
• Specially precipitated>calcium carbonate is used in the manufacture of high-quality paper.
• Precipitated chalk is used in toothpowder and toothpaste, cosmetics, and also in some medicines (antacids).

Class 11 S-Block Notes

4. Plaster of Paris (hemihydrate of calcium sulfate), (CaSO₃)₂-H₂O

Preparation of Plaster of Paris:

Plaster of Paris is prepared by heating gypsum, (CaSO₄-2H₂O) at about 383-393K in a rotating burner.

Properties of Plaster of Paris:

• It is a white powder.
• On mixing with an adequate amount of water, it forms a concentrated mixture which solidifies in 5 to 15 minutes due to rehydration. This is called the setting of Plaster of Paris.

• When Plaster of Paris is heated at about 473K, it forms anhydrous CaSO₄. It is called dead burnt plaster and it does not solidify with water. For this reason, during the preparation of Plaster of Paris from gypsum, the temperature should not be allowed to rise above 393K.

Uses of Plaster of Paris:

• Plaster of Paris is largely used in the building industry.
• It is used in surgical bandages for plastering fractured bones.
• It is used for making casts of statues, molds in pottery work, ornamental castings, and blackboard chalks.
• It is also used in dentistry.

5. Cement

Cement is a mixture of finely powdered calcium silicates and aluminates along with small quantities of gypsum.

The raw materials used for cement are:

1. Limestone (CaCO₃) 
2. Clay containing silica (SiO₂) and
3. Alumina (Al₂O₃) and
4. Gypsum (CaSO₄ – 2H₂O).

Cement Composition:

Different types of cement have different compositions.  The composition of Portland cement is given below:

For cement to have good quality, the ratio of silica to alumina should be between 2.5 to 4 and the ratio of CaO to the total oxides of silicon, aluminum, and iron should be close to 2.

Cement Preparation:

For the manufacture of cement, limestone, and clay are fused by strong heating to form cement clinker. This is mixed with the gypsum and ground to a very thin powder.

Cement Settings:

When cement is mixed with water, it forms a plastic mass. After some time it becomes solid. This change is due to the three-dimensional linking between — Si—O — Si — and —Si—O—Al — chains. This transition from plastic to solid is called setting.

Fly ash is a waste product of the steel industry produced mainly due to the burning of coal and carbon compounds. It has similar properties to that of cement. Sometimes fly ash is used with the cement to reduce the cost without compromising on the quality.

Class 11 S-Block Notes

Cement Uses: 

• Cement is the most important construction material.
• It is used in the construction of tunnels, roads, bridges, etc.
• It is used in concrete and reinforced concrete. These are made by mixing cement with sand, pebbles, and water.
• Mixing with sand it is used for plastering.

Biological Importance Of Magnesium (Mg) And Calcium (Ca)

• Magnesium and calcium ions found in biological fluids play an important role in biological processes. Mg²⁺ ions are concentrated in cells while C²⁺ ions are concentrated in body fluids, outside the cell.
• It is known that the energy is stored in the form of ATP.  The formations of phosphate linkages are catalyzed by Mg²⁺ ions. Also, the hydrolysis of phosphate linkages, (Which is accompanied by the release of energy, is also catalyzed by Mg²⁺ ions.
• Mg²⁺ ions are present in chlorophyll-a, the green pigment of plants, which absorbs light and is essential for photosynthesis.
• Both these ions are also essential for the transmission of impulses along nerve fibers. About 99% of calcium in the body is present in bones and teeth as apatite, Ca₃(PO₄)₂. In the enamel of teeth, it is present as fluorapatite [3Ca₃(PO₄)₂.CaF₂].
• Ca²⁺ ions also play an important role in blood clotting and are necessary to trigger the contraction of muscles and to maintain regular heartbeats.
• The concentration of Ca²⁺ ions in blood plasma (about 100mg L^-1) is maintained by two hormones namely calcitonin and parathyroid.
• The calcium ions in bones exchange readily with those in blood plasma. About 400 mg of Ca²⁺ enters and leaves our bones every day.
• In normal adults, there is a balance between this exchange. However, in aged people, especially women, sometimes there occurs a net loss of calcium in the bone leading to a disease called osteoporosis.
• An adult human body contains 1200 g of Ca and 25 g of Mg compared to only 59 g of Fe and 0.06 g of Cu. The daily requirement of calcium and magnesium in the human body has been estimated to be about 200-300 mg. The sources of Mg in our food are nuts, green vegetables, wheat, coffee, etc. while that of Ca are milk, paneer, and different milk products.

Class 11 Chemistry S Block Elements Long Questions And Answers

Question 1. How can anhydrous magnesium chloride be prepared from magnesium chloride hexahydrate?
Answer:

Anhydrous magnesium chloride cannot be prepared by simply heating MgCl₂-6H₂O because it gets hydrolyzed by its water of crystallization.

However, when hydrated MgCl₂ is heated at 650K in the presence of HCl, its hydrolysis is prevented; and it loses its water of crystallisation to form anhydrous MgCl₂

However, when hydrated MgCl₂ is heated at 650K in the presence of HCI, its hydrolysis is prevented, and it loses its water of crystallization to form anhydrous MgCl₂.

Question 2. Which out of BeCl₂ and CaCl₂ would give an acidic solution when dissolved in water?
Answer: 

Being a covalent compound and a good Lewis acid, BeCl₂ forms a hydrated salt, Be(H₂O₂)4Cl₂. The hydrated ion undergoes hydrolysis in solution producing H₃O⁺. This occurs because the Be — O bond is very strong and so in the hydrated ion this weakens the O —H bond. Hence, there is a strong tendency to lose protons. For this reason, the aqueous solution of BeCl₂ is acidicin nature.

[Be(H₂O)₄]²⁺ + H₂O ⇌ [Be(H₂O)₃OH]⁺ + H₃O⁺

On the other hand, CaCl₂ is a strongly ionic compound and does not behave as a Lewis acid (the size of Ca is relatively large and its octet is filled up). Moreover, since it is a salt of strong acid and strong base, it does not undergo hydrolysis and therefore, its aqueous solution is neutral.

Question 3 Explain The below Observation

1.

2.

Answer:

1. The smaller Li⁺ ion exerts strong polarising power and distorts the electron cloud of the nearby oxygen atom of the OH^– ion. This results in the formation of a strong Li —O bond and the weakening of the O —H bond. This ultimately facilitates the decomposition of LiOH into Li₂O and H₂O. The polarising power of the large Na+ ion is much lower and thus, NaOH remains unaffected by heating.

Class 11 S-Block Notes

2. The smaller Li⁺ ion exerts a strong polarising power on highly polarisable H^– ion and as a result, the two atoms remain strongly attached by a covalent bond. On the other hand, due to the low polarising power of Na⁺ ion, NaH is essentially ionic & so it dissociates on heating to yield metallic sodium and dihydrogen.

Question 4. Discuss the roles of Na₂O₂, KO_2, and LiOH in the purification of air.
Answer:

Sodium peroxide (Na₂O₂) is used to purify the air in submarines and confined spaces as it removes carbon dioxide (CO₂) and produces O₂.

Potassium superoxide (KO₂) is used to purify air in space capsules, submarines, and breathing masks because it can absorb carbon dioxide (CO₂) thereby removing it and producing O₂ Both functions are important life support systems.

Lithium hydroxide (LiOH) is used for the absorption of carbon dioxide (CO₂) in space capsules and submarines

Question 5. Lithium forms monoxide while sodium forms) peroxide in the presence of excess oxygen why
Answer:

Larger cations can be stabilised by larger anions because if both the ions are comparable in size, the coordination number will be high and this gives rise to a high lattice enthalpy. Lithium-ion, Li⁺ as well as oxide ion, O^2- , have small ionic radii and high charge densities.

Hence these small ions combine and form a very stable lattice of lithium monoxide (Li₂O). Similarly, the formation of sodium peroxide (Na₂O₂) can be explained based on the stable lattice formed by the packing of relatively large Na⁺ ion and peroxide ion O₂

Question 6. A freshly cut piece of sodium metal appears shiny but its metallic lustre soon gets tarnished when exposed to air. Give reason
Answer:

When a freshly cut piece of metallic sodium is exposed to moist air, it readily reacts with oxygen to form sodium monoxide (Na₂O). The resultant sodium monoxide and also the metal itself readily react with the moisture of the air to form sodium hydroxide (NaOH).

In the subsequent step, both Na₂O and NaOH combine with CO₂ of air to form sodium carbonate (Na₂CO₃). Thus, a coating of sodium carbonate is formed on the surface of the metal and as a result of this, the metallic luster is tarnished.

4Na + O₂→ 2Na₂O; Na₂O + H₂O →2NaOH

2Na + 2H₂O→2NaOH + H₂↑; Na₂O+ CO₂→Na₂CO₃

2NaOH + CO₂ →Na₂CO₃ + H₂O

Question 7. What happens when each of the following compounds is heated? 

1. Li₂CO₃ 
2. Na₂CO₃
3. LiNO₃
4. KNO₃

Answer:

1. Lithium carbonate (Li₂CO₃ ) decomposes readily on heating to give lithium monoxide (Li₂O) and carbon dioxide (CO₂).

2. Na₂CO₃ does not decompose on heating.

(3) Lithium nitrate (LiNO₃) decomposes on heating to form lithium monoxide (Li₂O), nitrogen dioxide (NO₂), and dioxygen (O₂).

4.  Potassium nitrate (KNO₃) decomposes on heating to give potassium nitrite (KNO₂) and dioxygen

Question 8. What happens when

1. HCl gas is passed through a concentrated solution of common NaCl-containing impurities like Na₂SO₄, CaSO₄, CaCl₂, MgCl_2, etc.
2. caustic soda beads are exposed to air for a long time.
3. How will you convert Na₂CO₃ into NaHCO₃ and vice versa?

Answer:

1. When HCl gas is passed through a concentrated solution of common NaCl with impurities, crystals of pure NaCl separate out because of the common ion effect,  Caustic soda (NaOH) is a deliquescent substance and becomes wet on exposure to air.

On long exposure, the solid beads dissolve in the absorbed water Moist caustic soda then absorbs carbon dioxide (CO₂) from air to form sodium carbonate (Na₂CO₃) which forms a coating over the surface of the material. As sodium carbonate (Na₂CO₃) is not a deliquescent substance the wet sodium hydroxide becomes dry again.

2NaOH + CO₃→ Na₂CO₃+ H₂O

2.  Sodium bicarbonate (NaHCO₃) can be obtained by passing carbon dioxide (CO₂) through a saturated solution of sodium carbonate. Sodium bicarbonate, being less soluble gets separated from the solution as a white crystalline substance.

Na₂ CO₃ + CO₂+ H₂O→2NaHCO₃

When sodium bicarbonate (NaHCO₃) is heated, it decomposes to give sodium carbonate (Na₂CO₃) and carbon dioxide (CO₂)

Question 9. Why are the group-2 metals harder and have higher melting and boiling points than group-1 metals?
Answer:

The magnitude of the cohesive energy determines the hardness as well as melting and boiling points of the metals and it depends on the number of electrons involved in metallic bonding. In the case of group 1 metals, one electron per atom (valence electron, ns² ) is involved in metallic bonding while in group 2 metals, two electrons per atom (valence electrons, ns²) are involved in metallic bonding. Moreover, atoms of group 2 metals are smaller in size than those of group 1 metals.

Consequently, stronger metallic bonding exists in group 2 metals which results in higher cohesive energy and close packing of the atoms. This accounts for the greater hardness and higher melting and boiling points of group 2 metals as compared to group 1 metals.

Class 11 S-Block Notes

Question 10.

1. Give some common tests used for the detection of calcium compounds. 

2. A white solid When heated liberates a colorless gas that does not support combustion. The residue is dissolved in water to form (B) which can be used for whitewashing. When excess CO₂ gas is passed through the solution of (B), it gives a compound (C) which on heating forms (A). Identify (A), (B) and (C). Give the reactions
Answer:

The following tests may be performed for the detection of Ca -compounds:

1. Calcium salts give a brick-red color in the flame test,
2. When ammonium oxalate solution is added to a solution of a calcium salt, a white precipitate of crystalline calcium oxalate is obtained which is insoluble in acetic acid but soluble in mineral acids
3. In addition of a solution of sodium carbonate to a neutral (or ammoniacal) solution of a Ca-salt, white calcium carbonate is precipitated, which is soluble in acids.
4. The observations suggest that the compound (A) is limestone, i.e., CaCO₃, the compound (B) is calcium hydroxide & the compound (C) is calcium bicarbonate.

The corresponding reactions are as:

Question 11. Compare the alkali metals and alkaline earth metals concerning their

1. Basicity of oxides
2. Solubility of hydroxides and
3. Solubility of nitrates.

Answer:

1. Basicity of oxides:

The ionization enthalpy of alkali metals is less than the corresponding alkaline earth metals. So the alkali metal oxides are more basic than the corresponding alkaline earth metal oxides.

2. Solubility of hydroxides:

Due to of small size and higher ionic charge, the lattice enthalpies of alkaline earth metal hydroxides are much higher than those of alkali metal hydroxides and hence, the solubility of alkali metal hydroxides is much higher than that of alkaline earth metal hydroxides.

3. Solubility of nitrates:

Nitrates of alkali and alkaline earth metals are soluble in water. However, the solubility of alkali metal nitrates increases down the group because their lattice enthalpies decrease more rapidly than their hydration enthalpies. Nitrates of alkaline earth metals follow the reverse trend i.e., their solubility decreases down the group and this is because their hydration enthalpies decrease more rapidly than their lattice enthalpies.

Question 12. What are the properties that make oxides of MgO and  BeO useful for lining furnaces?
Answer:

The given properties make MgO and BeO useful for lining furnaces

• They have high melting points [melting point of Beo is 2500°C (approx) and MgO is 2800°C (approx)].
• They are very good conductors of heat.
• They have very low vapor pressure.
• They are chemically inert.
• They are insulators

Question 13. The halides of alkali metals are soluble in water except for LiF. Why?
Answer:

The solubility of a salt in water depends on its lattice enthalpy as well as its hydration enthalpy. A salt dissolves in water when its hydration enthalpy exceeds its lattice enthalpy value. The lattice enthalpy of LiF is very high and its hydration enthalpy value does not exceed this value. As a result, LiF is insoluble in water

However, the other halides of alkali metals possess higher hydration enthalpy values compared to their corresponding lattice enthalpies. Hence, they are soluble in water

Question 14. Why Is LiCO₃ decomposed at a lower temperature whereas Na₃ CO₃ at a higher temperature
Answer: 

Li+ ion being smaller in size forms a more stable lattice with the smaller oxide ion (O^2-) than the larger carbonate ion (CO₃ ) and consequently, Li₂CO₃  decomposes into Li₂ O at a much lower temperature. The high polarising power of very small Li⁺ ions also facilitates the decomposition of Li₂CO₃.

On the other hand, the larger Na+ ion forms a more stable lattice with the larger CO₃ ion than with the smaller O^2- ion. Therefore, Na₂ CO₃ is quite stable and decomposes only at very high temperatures.

Question 15. Compare the solubility and thermal stability of the given compounds of the alkali metals with those of the alkaline earth metals. 

1. Nitrates
2. Carbonates Sulphates.

Answer:

1. Nitrates:

The nitrates of alkali metals as well as alkaline earth metals are highly soluble in wOn heating, nitrates of both alkali and alkaline earth metals undergo decomposition.

Nitrates of alkali metals decompose to form metallic nitrite and oxygen. On the other hand, nitrates of alkaline earth metals decompose to form the corresponding oxides with the evolution of NO₂ and O₂

Due to the diagonal relationship, lithium nitrates behave similarly to magnesium nitrate

2. Carbonates:

1. Except for Li₂CO₃, other carbonates of alkali metals readily dissolve in water. However, carbonates of alkaline earth metals are practically insoluble in water. Their solubilities decrease on moving down the group. So, BeCO₃ is sparingly soluble in water while BaCO₃ is insoluble in water.

2. Carbonates of alkali metals except Li₂CO₃ are stable and do not decompose on heating, but carbonates of alkaline earth metals decompose on heating to give metal oxide and carbon dioxide.

The thermal stability of the carbonates increases down the group. Like MgCOg, Li₂CO₃ decomposes on heating.

3. Sulphates:

1. Except Li₂SO₄, the remaining sulfates of alkali metals are water-soluble. The sulfates of alkaline earth metals are relatively less soluble in water than the corresponding sulfates of alkali metals. Further, their solubilities decrease down the group,

2. The sulfates of alkali metals are stable compounds and do not decompose on heating. On the other hand, alkaline earth metals dissociate on heating to give metal oxides and sulfur trioxide.

The thermal stability of the sulfates increases down the group. Li₂SO₄ dissociates on heating just like MgSO₄

Class 11 S-Block Notes

Question 16. Starting with sodium chloride how would you proceed to prepare

1. Sodium metal
2. Sodium peroxide

Answer:.

1. Metallic sodium can be obtained by the electrolysis of a mixture of sodium chloride (40%) and calcium chloride (60 %) in a fused state. The function of calcium chloride is to lower the reaction temperature from 807°C (m.p. of NaCl) to about 577°C. The molten sodium metal thus obtained is liberated at the cathode

Overall reaction: Nacl → Na⁺ + Cl^–

At cathode: Na⁺ + e→ Na;

At anode: Cl^– → Cl + e; Cl + Cl → Cl₂ ↑

2.  Sodium chloride is first converted to sodium by electrolytic reduction. The metal is then heated more than air. The initially formed sodium oxide reacts with excess O₂ to form Na₂O₂.

Question 17. What happens when  

1. Magnesium is burnt in the air 
2. Calcium nitrate is heated?

Answer:

1. When magnesium bums in the air, magnesium oxide (MgO) and magnesium nitride (Mg₃N₂) are obtained as products

2. On heating calcium nitrate, it decomposes to form CaO, NO_2, and O₂ reaction

Question 18. The hydroxides and carbonates of sodium and potassium are easily soluble in water while the corresponding salts of magnesium and calcium are sparingly soluble in water. Explain.
Answer:

Alkali metals form monovalent cations (such as Na⁺, and K⁺) while alkaline earth metals form divalent cations (such as Mg²⁺, and Ca²⁺). Due to the increase in charge of the cations, the lattice energy of the corresponding salt increases. For this reason, hydroxides and carbonates of sodium and potassium have lower lattice enthalpy values than the hydroxides and carbonates of magnesium and calcium.

As hydration enthalpies of the hydroxides and carbonates of sodium and potassium are greater than their lattice enthalpies, these salts readily dissolve in water. However, in the case of the hydroxides and carbonates of calcium and magnesium, the lattice enthalpy values is greater than that of hydration enthalpy and consequently, these salts are less soluble in water.

Question 19. Why are lithium salts commonly hydrated and those of the other alkali ions usually anhydrous?
Answer:

As Li⁺ is the smallest ion among all the alkali metal ions, it can polarise water molecules more easily than the other alkali metal ions. So, numerous water molecules get attached to lithium salts as water of crystallization. However, this is not observed in the case of other alkali metal ions. Thus, lithium salts are commonly hydrated

Example: LiCl-2H₂O  and those of the other alkali ions are usually anhydrous.

Class 11 S-Block Notes

Question 20. Why is LiF almost insoluble in water whereas LiCl is soluble not only in water but also in acetone?
Answer:

The lattice enthalpy of ionic LiF (formed by small Li⁺ion and F^– ion) is higher than its hydration enthalpy. On the other hand, the lattice enthalpy of LiCl containing small Li⁺ ion and large Cl^– ion is considerably lower than its hydration enthalpy.

Thus, LiF is almost insoluble in water while LiCl is soluble. Furthermore, Li⁺ ions can polarise bigger Cl^– ions more easily than smaller F^– ions. As a result, LiCl has more covalent character than LiF and so, it is also soluble in the organic solvent acetone.

Question 21. What happens when

1. Sodium metal is dropped in water?
2. Sodium metal is heated in a free supply of air
3. Sodium peroxide dissolves in water?

Answer:

Sodium hydroxide is formed. H₂ gas is evolved which catches fire due to the exothermicity of the reaction.

2Na(s) + 2H₂O(l)→ 2NaOH(aq) + H₂(g)

2. Sodium peroxide is formed by heating sodium metal in a free supply of air.

2Na(s) + O₂ (g)→ 2Na₂O₂(s)

3. H₂O₂ is formed when sodium peroxide dissolves in water

Na₂O₂(s) + 2H₂O(l)→2NaOH(aq) + H₂O₂(Z)

Question 22. Comment on each of the given observations:

1. Lithium is the only alkali metal to form a nitride directly.
2.  M²⁺+(aq) + 2e → M(s) (where M = Ca, Sr, or Ba) is nearly constant.

Answer:

1. The lattice energy of lithium nitride (Li₃N) which consists of a small cation (Li⁺) and a small anion (N₃-) is much higher and this energy compensates for the high bond dissociation energy of the N=N bond and the energy to form N₃– ion. Larger alkali metal ions cannot compensate for these energy requirements. Hence, Li⁺ is the only alkali metal that forms nitride directly.

2. M²⁺ (aq) + 2e → M(s) Where M = Ca, Sr, Ba

Standard electrode potential, \(E_{\mathrm{M}^{2+} / \mathrm{M}}^0\) depends on 3

1. Enthalpy of vaporisation
2. Ionization enthalpy and
3. Enthalpy of hydration.

As the combined effect of these factors is almost the same for Ca, Sr and Ba, their E° values are nearly constant.

Class 11 S-Block Notes

Question 23. State as to why

1.  A solution of Na₂CO₃ is alkaline?
2. Alkali metals are prepared by electrolysis of their fused chlorides.
3. Sodium is found to be more useful than potassium.?

Answer:

1. Na₂CO₃ is a salt of weak acid H₂CO₃ and strong base NaOH. In an aqueous solution, it ionises to form Na⁺ and CO^2-₃ – ions.

Na₂CO₃ ⇌  2Na⁺  (aq) + CO^2-₃ (aq)

The formed CO^2-₃ ions hydrolyse in an aqueous solution to produce acetic acid and OH^– ion.

CO^2-₃ (aq) + 2H₂O(Z) ⇌ H₂CO₃ ⇌+ 2OH^– ( aq)

As H₂CO₃ is a weak acid, it remains mostly unionized. Consequently, the concentration of OH^– ions increases in the solution thereby making the solution alkaline.

2.

• As alkali metals are strong reducing agents, they cannot be extracted by chemical reduction from their oxides or other compounds,
• As alkali metals are highly electropositive, these metals cannot be displaced from their salts with the help of other elements,
• Alkali metals cannotbe obtained even by the electrolysis of aqueous solutions of their salts.

In this case, H₂, instead of the alkali metal is liberated at the cathode because the discharge potential of alkali metals is higher than that of hydrogen.

Thus, to prepare alkali metals, electrolysis of their fused chlorides is carried out. For example,

NaCl →Na⁺ + Cl^–

During electrolysis, at the cathode,

2Na⁺ + 2e →2Na; and at the anode, 2Cl→ Cl₂ + 2e

3. Sodium is found to be more useful because Na is not as reactive as K. For this reason, reactions of sodium with different substances can be controlled and usage of sodium is far more safe than potassium. Thus, sodium is more useful than potassium

Question 24. Write balanced equations for reactions between 

1. Na₂O₂ & Water,
2. KO₂ & water,
3. Na₂ O& CO₂

Answer: The balanced equations of the given reactions are.

1.

2. 2K O₂ (S) + 2H₂ O(l) → 2KOH(aq) + H₂ O₂ (aq) + O₂ (g)

3. Na₂ O + CO₂→ Na₂ CO₂

Question 25. How would you explain the given observations? 

1. BeO is almost insoluble but BeSO₄ is soluble in water.
2. BaO is soluble but BaSO4 is insoluble in water.
3. Lil is more soluble than KI in ethanol.

Answer:

1. O^2- is smaller in size than SO₄^2-. Consequently, a small Be²⁺ ion is tightly packed with a small O^2- ion, and thus, the lattice enthalpy of BeO is greater than its hydration enthalpy. So BeO is insoluble in water. On the other hand, a small Be²⁺ ion is loosely packed with a large SO^2-. ion and thus, the lattice enthalpy of BeSO₄ is less than its hydration enthalpy. So, BeSO₄ is soluble in water

2. Large Be²⁺ ion is tightly packed with large SO₄^2-.ion and thus, the lattice enthalpy of BaSO₄ is greater than its hydration enthalpy. So, BaSO₄ is insoluble in water On the contrary, a large Ba²⁺ ion is loosely packed with a small Oion, and thus lattice enthalpy of BaO is less than its hydration enthalpy. So, BaO is water soluble.

KI is predominantly ionic. On the other hand, due to the high polarising power of very small Li+ ions, Lil is predominantly covalent. For this reason, Lil is more soluble than KI in the organic solvent ethanol.

Question 27. How will you distinguish between:

1. Mg and Ca
2. Na₂SO₄ and BaSO₄
3. Na₂CO₃ and NaHCO₃,
4. LiNO₃ and KNO₃

Answer:

1. Calcium, when heated, imparts brick red color to the flame but magnesium does not.

2. Sodium sulfate (Na₂SO₄) is soluble in water but barium sulfate (BaSO₄) is insoluble.

3. Na₂CO₃ is stable to heat but NaHCO₃ decomposes on heating to produce CO₂ gas which turns limewater milky

4. When LiNO₃ is heated, it decomposes to yield reddish-brown vapors of NO₂. However, when KNO₃ is heated, it decomposes to yield colorless O₂ gas

Class 11 Chemistry S Block Elements Short Question And Answers

Question 1. Li+ ion is the smallest one among the ions of group- 1 elements. It would, therefore, be expected to have much higher ionic mobility and hence the solutions of lithium salts would be expected to have higher conductivity than the solutions of cesium salts. However, in reality, the reverse is observed. Explain
Answer: 

Due to high charge density, very small Li⁺ ion gets much more hydrated compared to the large Cs⁺ ion. Thus, the size of the hydrated lithium-ion is much larger than that of the hydrated cesium ion. For this reason, the mobility of Li⁺ ion is much lower than that of Cs⁺ ion and consequently, the solutions of lithium salts have much lower conductivity than the solutions of cesium salts.

Question 2. The E° value for Cl₂/Cl^– is +1.36, for I₂/I^– is + 0.53V, for Ag⁺/Ag is + 0.79 V, for Na⁺/Na is -2.71V and for Li⁺/Li is -3.04V. Arrange the following atoms and ions in decreasing order of their reducing strength: I^–, Ag, Cl^–, Li, Na
Answer:

The lower the value of standard reduction potential, the greater the tendency of the reduced form to be oxidized, i.e., the reduced form will serve as a stronger reductant. Therefore, the decreasing order of reducing strength of the given atoms and ions is

Li > Na > I^–> Ag > Cl^–

Class 11 S-Block Notes

Question 3. The alkali metals are obtained not by the electrolysis of the aqueous solutions of their salts but by the electrolysis of their molten salts. Explain.
Answer:

The solutions of alkali metal salts contain metal cations, anions, H⁺ ions, and OH^– ions. The discharge potential of H⁺ ions is lower than that of metal cations. So, on electrolysis of tire solutions of alkali metal salts, hydrogen is discharged at the cathode rather than the metal. However, when the molten salts of alkali metals are electrolysed, the metal cation being the only cation present, gets discharged at the cathode

Question 4. The alkali metals are paramagnetic but their salts are diamagnetic—why?
Answer:

Due to the presence of anupaired valence electron the alkali metals are paramagnetic. During salt formation, this unpaired electron of the outermost shell leaves the metal atom and becomes attached to a non-metal atom. As a result, the cation and the anion thus obtained contain no unpaired electron. Hence, the alkali metal salts are diamagnetic.

Question 5. Beryllium & magnesium do not give color to flame whereas other alkaline earth metals do so. Why?
Answer: 

Due to their smaller size, valence electrons of Be and Mg are more tightly held by the nucleus. Therefore, they need a large amount of energy for the excitation of their valence electrons to higher energy levels. Since such a large amount of energy is not available from Hansen flame, these two metals do not impart any color to the flame.

Question 6. E° for M²⁺(aq) + e →M(s) (where, M = Ca, Sr or Ba) Is nearly constant. Comment.
Answer:

The value of standard electrode potential (E°) of any M²⁺/M electrode depends upon three factors

1. Enthalpy of vaporisation
2. Ionization enthalpy, and
3. Enthalpy of hydration.

Since the combined effect of these factors is approximately the same for Ca, Sr, and Ba, their standard electrode potential (E°) values are nearly constant.

Question 7. Both alkaline earth metals and their salts are diamagnetic. Explain.
Answer:

The alkaline earth metals are diamagnetic as all the orbitals are filled with paired electrons. The ions of alkaline earth metals, M²⁺ have stable noble gas configurations in which all the orbitals are doubly occupied. Also, there are no unpaired electrons in the anions. Hence, the salts of alkaline earth metals are also diamagnetic.

Question 8. Beryllium salts can never have more than 4 molecules of water of crystallization, i.e., it can never achieve a coordination number > 4 while other metal ions tend to have a coordination number of 6, for example: [Ca(H₂O₆)²⁺. Explain.
Answer: 

Beryllium does not exhibit coordination numbers more than 4 because, in its valence shell of Be²⁺ ion, there are only four available orbitals (one s and three p) present. The remaining members of the group can have a coordination number of six by using their d -d-orbitals along with s -and p -orbitals.

Question 9. Anhydrous is used as a drying agent — why?
Answer:

Anhydrone or magnesium perchlorate, Mg(ClO₄)₂ is  used as a drying agent because, due to its smaller size and higher charge, Mg has a greater tendency to complexation with water molecules by forming Mg(ClO₄)₂ -6H₂O

Question 10. Li salts are commonly hydrated while other alkali metal salts are usually anhydrous. Explain.
Answer:

Due to the smallest size among all alkali metal ions, the Li⁺ ion can interact with water molecules more easily than the other alkali metal ions. Hence the salts of lithium are commonly hydrated. On the other hand, other alkali metal ions being larger have little tendency to get hydrated. Therefore, their salts are generally anhydrous. For example, lithium chloride crystallises as LiCl. 2H₂O₂ but sodium chloride crystallizes as NaCl.

Question 11. Although the abundance of Na and K in the earth’s crust are comparable, sodium is nearly 30 times more abundant than potassium in seawater—why?
Answer:

These metals were leached from the aluminosilicate rocks by weathering. The potassium salts having large anions are less soluble than the sodium salts because of higher lattice energy. Moreover, potassium is preferentially absorbed by the plants. For this reason, sodium is more abundant than potassium in seawater.

Question 12. When caustic soda solution is kept in a glass bottle, the inner surface of the bottle becomes opaque. Explain
Answer:

Caustic soda (NaOH) being strongly basic reacts with acidic silica (SiO₂) present in glass to form sodium silicate (Na₂SiO₃)

SiO₂ + 2NaOH → Na₂SiO₃ + H₂O

As a result, the inner surface of the bottle becomes opaque.

Question 13. Why are alkali metals stored in kerosene?
Answer:

When the highly reactive alkali metals are exposed to air,  they readily react with oxygen, moisture, and carbon dioxide of air to form oxides, hydroxides and carbonates respectively. To prevent these reactions, alkali metals are normally stored in kerosene, an inert liquid.

Question 14. The second ionization enthalpies of alkaline earth metals are much lower than those of the corresponding alkali metals. Explain.
Answer:

The loss of a second electron from an alkali metal cation (M⁺) causes a loss of stable noble gas configuration while the loss of a second electron from an alkaline earth metal cation leads to the attainment of a very stable noble gas configuration. This explains why the second ionization enthalpies of alkaline earth metals are much lower than those of the corresponding alkali metals.

Class 11 S-Block Notes

Question 15. MgO is used as a refractory material— Explain why
Answer:

Due to greater charge on both the cation (Mg²⁺) and die anion (O^2-), MgO possesses higher lattice energy and for tills, it has very high melting point and does not decompose on heating. For this reason, it is used as a refractory material

Question 16. BaSO₄ is insoluble in water whereas BeSO₄ is soluble in water—1-explain with reason._
Answer: 

The lattice enthalpy of BaSO₄ is much higher than its hydration enthalpy and hence it is insoluble in water. On the other hand, the hydration enthalpy of BeSO₄ is much higher than its lattice enthalpy because of the loose packing of the small Be²⁺ ion with the relatively large sulfate ion. Hence, it becomes soluble in water.

Question 17. BeCl₂ fumes in moist air but BaCl₂ does not. Explain
Answer:

BeCl₂ being a salt of a weak base, Be(OH)_2, and a strong acid, HCl undergoes hydrolysis in moist air to form HCl; which fumes in air. BaCl₂ on the other hand, being a salt of a strong base, Ba(OH)₂ and a strong acid, HCl does not undergo hydrolysis to form HCl and hence does not fume in moist air

BeCl₂ + 2H₂O→Be(OH)₂ + 2HCl↑

BaCl₂ +H₂O→ Ba(0H)₂ + 2HCl

Question 18. Mg₃N₂ when reacts with water, gives off NH₂ but HCl is not evolved when MgCl₂ reacts with water at room temperature. Give reasons.
Answer:

Mg₃N₂ is a salt of the strong base, Mg(OH)₂ and the weak acid, NH₃

Hence it gets hydrolyzed to give NH₃.

Mg₃N₂ + 6H₂O→ 3Mg(OH)₂ + 2NH₃T↑

MgCl₂,  On the other hand, is a salt of the strong base, Mg(OH)_2, and the strong acid, HCl. Hence, it does not undergo hydrolysis

Question 19. A piece of burning magnesium ribbon continues to bum in sulphur dioxide.
Answer:

A piece of magnesium ribbon continues to bum in SO₂ since it reacts to form MgO and S. This reaction is such exothermic that heat evolved keeps the magnesium ribbon burning.

This reaction is such exothermic that heat evolved keeps the magnesium ribbon burning

Question 20. Ba²⁺ ions are poisonous, still, they are provided to patients before taking stomach X-rays. Explain
Answer:

A barium meal (suspension of BaSO₄ in water) is generally given to patients when X-ray photographs of the alimentary canal are required. The salt provides a coating of the alimentary canal and hence X-ray photograph can be taken since it is quite transparent to the X-ray otherwise. BaSO₄ is almost insoluble in water and hence it does not pass from the digestive system to the circulatory system and can therefore be safely used for the purpose.

Question 21. Can sodium hydride be dissolved in water? Justify.
Answer:

Sodium hydride cannot be dissolved in water because it gets hydrolyzed with brisk effervescence of hydrogen gas.

NaH + H₂ O→  H₂ + NaOH

Question 22. Why does sodium impart a yellow color in the flame?
Answer:

The ionization enthalpy of Na is relatively low. Therefore when this metal or its salt is heated in Bunsen flame, its valence shell electron is excited to higher energies by absorption of energy. When the excited electron returns to its initial position in the ground state, it liberates energy in the form of light in the yellow region of the electromagnetic spectrum. That’s why sodium imparts yellow color to the flame.

Class 11 S-Block Notes

Question 23. Wind makes lithium exhibit uncommon properties compared to the rest of the alkali metals.
Answer:

The unusual properties of lithium as compared to other alkali metals is since

• Li – atom and L⁺  ion are exceptionally small in size and
• Li⁺ ion has the highest polarising power (i.e„ charge/size ratio).

Question 24. What is the common oxidation state exhibited by the alkali metals and why?
Answer:

The alkali metals easily lose their valence electrons (ns¹) to acquire a stable octet, (i.e., the stable electronic configuration of the nearest noble gas) and because of this, the common oxidation state exhibited by the alkali metals is +1.

Question 25. What is the difference between baking soda and baking powder?
Answer:

Both baking soda and baking powder are leavening agents. Baking soda is pure sodium bicarbonate. When baking soda is combined with moisture and an acidic ingredient (for example,  Yoghurt, buttermilk) the resulting chemical reaction produces CO₂ gas bubbles that cause baked goods to rise. Baking powder contains NaHCO₃, but it includes an acidifying agent (cream of tartar) already, and also a drying agent (usually starch).

Baking powder is available as single-acting baking powder and as double-acting baking powder. Single-acting baking powders are activated by moisture, so we must bake recipes that include this product immediately after mixing. Double-acting powder reacts in two phases and can stand for a while before baking.

Question 26. Though table salt is not deliquescent it gets wet ! in the rainy season— Explain.
Answer:

Pure NaCl is not deliquescent but table salt contains impurities like MgCl₂ and CaCl₂ These impurities being deliquescent absorb moisture from air in the rainy season. As a result, table salt gets wet

Question 27. What precautions should be taken while handling beryllium compounds and why?
Answer:

Contact of Be compounds with the skin dermatitis, and inhaling dust or smoke of Be-compounds causes a disease called berylliosis which is rather similar to success- Therefore, beryllium compounds should be handled with care

Question 28. Explain why the elements of group 2 form M²⁺ Ions, but not M3+ ions.
Answer:

Loss of third electron from group-2 metal atoms causes loss of stable noble gas configuration and for this reason group- 2 elements form M²⁺ ions but not M³⁺ ions. fifl Arrange Be(OH)₂, Ba(OH)₂ & Ca(OH)₂in order of increasing solubility in water and explain the order. Answer: Among the alkaline earth metal hydroxides having a common anion, the cationic radius influences the lattice enthalpy. Since the lattice enthalpy decreases much more than the hydration enthalpy with increasing cationic size, the solubility increases on moving down the group.

Question 29. The reaction between marble and dilute H₂SO₄ is not used to prepare CO₂ gas—why?
Answer:

Marble (CaCO₃) reacts with dilute H₂SO₄ to form insoluble CaSO₄ which deposits on the surface of marble and prevents further reaction. So, the evolution of CO₂ ceases after some time. Thus the reaction between marble and dilute H₂SO₄ prepares CO₂ gas

Question 30. Name the important compound of Li used in organic synthesis. How the compound is prepared?
Answer:

The compound is lithium aluminum hydride (LiAlH₄). It is a useful reducing agent and is used in organic synthesis. It is prepared by the reaction of lithium hydride and aluminum chloride in a dry ether solution.

4LiH + AlCl₃ → LiAlH₄ + 3LiCl

Question 31. What is the oxidation state of K in KO₂ and why is this compound paramagnetic?
Answer:

The superoxide ion is represented as O^–₂ . It has one unit of negative charge. Since the compound is neutral, therefore, the oxidation state of K is +1. The structure of Or is 6J0 O^–₂    is  Since it has one unpaired electron in π∗_2p MO, therefore, the compound is paramagnetic.

Question 32. The crystalline salts of alkaline earth metals contain more water of crystallization than the corresponding alkali metals. Explain.
Answer:

In the salts of alkaline earth metals, the metal ions have a smaller size and higher charge compared to the corresponding metal ions of the alkali metals of the same period. Thus, alkaline earth metals have a greater tendency to get hydrated & form crystalline salts compared to alkali metals. Thus, NaCl is completely anhydrous whereas MgCl₂ exists as MgCl₂-6H₂O

Question 33. Lithium salts are more stable if the anion present in the salt is small. Explain.
Answer:

Small anions have more ionic character and hence the salts of lithium containing those ions have more lattice enthalpy. Large anions, on the other hand, are highly polarisable and hence they impart covalent character to the salt. Thus, lithium salts are more stable with small anions than that with large anions

Question 34. Alkali metals become opaque when they are kept open in the air Why?
Answer:

As the alkali metals are highly reactive, they readily react 20; with oxygen to form oxides. These oxides undergo a reaction with the water vapor present in the air to produce hydroxides. The formed hydroxides immediately react with CO₂ of air to produce carbonate compounds. These carbonate compounds form layers on the surface of alkali metals. Consequently, they become opaque

Class 11 S-Block Notes

Question 35. BaSO₄ is insoluble in water, but BcSO₄ is soluble in water-Explain.
Answer:

The lattice enthalpy of BaSO₄ is much higher than its hydration enthalpy and hence, it is insoluble in water. On the other hand, the hydration enthalpy of BeSO₄ is much higher than its lattice enthalpy because of the loose packing of the small Be²⁺ ion with the relatively large sulfate ion. Hence, it becomes soluble in water.

Question 36. An aqueous solution of Be(NO₃)₂ is strongly acidic. Explain.
Answer:

In the hydrated ion, [Be(H₂O)₄]²⁺, water molecules are extensively polarised, ultimately leading to the weakening of the O —H bond.

Hydrolysis takes place and the solution becomes distinctly acidic:

(H₂O)₃Be₂ +—OH₂ + H₂O→(H₂O)₃ 3 Be⁺ —OH + H₃O⁺

Question 37. The hydroxides and carbonates of Na and K are readily soluble in water while the corresponding salts of Mg and Ca are sparingly soluble. Explain.
Answer:

Due to smaller size and higher ionic charge, the lattice enthalpies of alkaline earth metals are much higher than those of alkali metals and hence the solubility of alkali metal hydroxides and carbonates is much higher than those of alkaline earth metal hydroxides and carbonates;

Question 38. BeO is insoluble but BeSO₄ is soluble in water. Explain.
Answer:

The higher lattice enthalpy of BeO formed by the combination of a small cation and small anion is more than its hydration enthalpy but the lattice enthalpy of BeSO₄ formed by the combination of a small cation and large anion is less than its hydration enthalpy;

Question 39. BaO is soluble but BaSO₄ is insoluble in water — why?
Answer:

The lattice enthalpy of BaO formed by the combination of a large cation and a small anion is less than its hydration enthalpy but the lattice enthalpy of BaSO₄ formed by the combination of a large cation and a large anion is more than its hydration enthalpy;

Question 40. Lil is more soluble than KIin ethanol. Explain.
Answer:

Due to the much higher polarising power of very small Li⁺ ion, Lil is predominantly covalent but due to the low polarising power of relatively large K⁺ ion, KI is predominantly ionic and for this reason, Lil is more soluble in the organic solvent ethanol;

Question 41. How can fused calcium chloride be prepared? Give two important uses of it.
Answer:

When CaCl₂ 2H₂O is heated above 533K, anhydrous CaCl₂ forms. This melts at 1046K. When the molten salt is cooled, it solidifies as white lumps of crystalline mass which is known as fused calcium chloride;

Question 42. Hydrated magnesium chloride is heated in the presence of ammonium chloride (NH₄Cl).
Answer:

Magnesium chloride when heated with NH₄Cl forms an additional compound (MgCl₂-NH₄Cl₆H₂O) which on heating forms anhydrous MgCl₂

Question 43. Why are alkali metals not found in nature?
Answer:

Alkali metals are highly electropositive and extremely reactive elements. Thus, they easily react with atmospheric oxygen and carbon dioxide. These metals have a high tendency to lose electrons to form cations because of their low ionization potential. For this reason, they readily react with highly electronegative elements or radicals to form compounds. So, alkali metals are not found in a free state in nature

Question 44. Explain why is sodium less reactive than potassium.
Answer:

Due to its small size, the ionization enthalpy of sodium (496kL-mol^-1) is greater than potassium (419kJ. mol^-1) and the standard electrode potential value of potassium (-2.925V) is more negative than that of sodium (-2.714V). Thus, sodium is less reactive than potassium.

Question 45. Basicity of oxides
Answer:.

Basicity of oxides: The ionization enthalpy of alkali metals is less than that of the corresponding alkaline earth metals, i.e., the electropositive character of alkali metals is greater than that of alkaline earth metals. Thus, the basicity of oxides of alkali metals is more than the oxides of alkaline earth

Class 11 Chemistry S Block Elements Multiple Choice Questions

Question 1. NO₂ is not obtained on heating—

1. AgNO₃
2. KNO₃
3. Cu(NO₃)₂
4. Pb(NO₃)₂

Answer: 2. KNO₃

KNO₃ on heating decomposes to form potassium nitrite (KNO₂) and oxygen gas

Class 11 S-Block Notes

Question 2. Which one of the following has the lowest ionization

1. ls²2s²2p⁶
2. ls²s²2p⁶3s¹
3. ls²2s²2p⁵
4. ls²2s²2p³

Answer: 2. ls²s²2p⁶3s¹

Since the electronic configuration ls²s²2p⁶3s¹ is that of an alkali metal, its ionization potential value is the lowest.

Question 3. Which of the following represents the composition of carnallite mineral—

1. K₃O-Al₂O₃.6SiO₂
2. KNO₃
3. K₂SO4.MgSO₄-MgCl₂-6H₂O
4. KCl-MgCl₂-6H₂O

Answer: 4. KCl-MgCl₂-6H₂O

Carnallite:  KCl-MgCl₂-6H₂O

Question 4. Chlorine gas reacts with red-hot calcium oxide to give—

1. Bleaching powder and dichlorine monoxide
2. Bleaching powder and water
3. Calcium chloride and chlorine dioxide
4. Calcium chloride and oxygen

Answer: 4. Calcium chloride and oxygen

2CaO + 2Cl₂ → 2CaCl₂ + O₂

Question 5. The decreasing order of the basic characters of K₂O, BaO, CaO, and MgO is

1. K₂O > BaO > CaO > MgO
2. K₂O > CaO > BaO > MgO
3. MgO > BaO > CaO > K₂O
4. MgO>CaO>BaO>K₂O

Answer:   4. MgO>CaO>BaO>K₂O

The oxides of alkali metals are highly basic. The basic character of the oxides of alkaline earth metals increases on moving down the group.

Question 6. Match the flame colors of the alkaline earth metal salts in the Bunsen burner.

Answer: 4. 1 -A, 2-B,3-C

1. 1 -A, 2-C,3-B
2. 1 -C, 2-A,3-B
3. 1 -B, 2-C,3-A
4. 1 -A, 2-B,3-C

Question 7. The correct order of solubility in water is

1. CaSO₄>BaSO₄>BeSO₄>MgSO₄>SrSO₄
2. BeSO₄ > MgSO₄> CaSO₄> SrSO₄ > BaSO₄
3. BaSO₄>SrSO₄>CaSO₄>MgSO₄>BeSO₄
4. BeSO₄> CaSO₄ > MgSO₄ > SrSO₄> BaSO₄

Answer: 2. BeSO₄ > MgSO₄> CaSO₄> SrSO₄ > BaSO₄

Question 8. When BaCl₂ is added to an aqueous solution, a white precipitate is obtained. The anion among CO^2-₃ SO^2-₃ and SO^2-₄ that was present in the solution can be

1. CO^2-₃ But any of the other two
2. SO^2-₃ But not any of the other two
3. SO^2-₄ But not any of the other two
4. Any of them

Answer: 4. Any of them

BaSO₃, BaCO₃ and BaSO₄ are insoluble in water, thus they are precipitated out in an aqueous solution.

Question 9. Which of the following is least thermally stable

1. MgCO₃
2. CaCO₃
3. SrCO₃
4. BeCO₃

Answer: 4. BeCO₃

Question 10. Which one of the following orders presents the correct sequence of the increasing basic nature of the given oxides

1. MgO < K₂O < Al₂O₃ < Na₂O
2. Na₂O<K₂O<MgO<Al₂O₃
3. K₂O<Na₂O<Al₂O₃<MgO
4. Al₂O₃ < MgO < Na₂O < K₂ O

Answer: 4.Al₂O₃ < MgO < Na₂O < K₂ O

The basicity of the metallic oxides increases with the increase in electronegativity values of the metals

Class 11 S-Block Notes

Question 11. Which of the following on thermal decomposition yields a basic as well as an acidic oxide

1. KClO
2. CaCO₃
3. NH₄NO₃
4. NaNO₃

Answer: 2.CaCO₃

CaCO₃ at high temperature decomposes to form CaO and CO₂

Question 12. Which one of the following alkaline earth metal sulfates has its hydration enthalpy greater than its lattice

1. BaSO₄
2. SrSO₄
3. CaSO₄
4. BeSO₄

Answer: 4. BeSO₄

The hydration enthalpy of BeSO₄ is greater than Its lattice energy because of the very small size of Be ²⁺

Question 13. The main oxides formed on combustion of, Na and Kin excess of air respectively are

1. LiO₂, Na₂O and KOH
2. LiO₂, Na₂O₂ and K₂O
3. Li₂O₂, Na₂O₂ and KO
4. Li₂O₂, Na₂O₂ and KO₂

Answer: 4. Li₂O₂, Na₂O₂ and KO₂

4Li + O₂ →2Li₂O

2Na + O→Na₂O₂

K + O→KO

Question 14. Although lithium and magnesium resemble each other in properties due to a diagonal relationship, the following statement is not correct

1. Both of them form a nitride compound
2. When nitrates of both Li and Mg are heated, NO₂ and O₂ are obtained
3. Both of them form basic carbonate salt
4. Both of them form soluble bicarbonate salt

Answer: 3. Both of them form basic carbonate salt

Magnesium forms basic bicarbonate salt [3MgCO₃.Mg(OH)₂ .3H₂O] whereas lithium forms carbonate salt [Li₂CO₃]. I cannot form any basic bicarbonates

Question 15. Match Column – 1 Column – 2 for the compositions of substances and select the correct answer using the code given below

1. 1 -C, 2-D, 3-  A, 4- B
2. 1 -B, 2-C, 3-  D, 4- A
3. 1 -A, 2-B, 3-  C, 4- E
4. 1 -D, 2-C, 3-  A, 4- B

Answer: 2. 1 -B, 2-C, 3-  D, 4- A

Plaster of Paris: CaSO₄– ½H₂O

Epsomite: MgSO₄-7H₂O

Kieserite: MgSO₄-H₂O

Gypsum: CaSO₄-2H₂O

Question 16. Which one of the following is present as an active ingredient in bleaching powder for bleaching action

1. CaOCl₂
2. Ca(OCl)₂
3. CaO₂Cl
4. CaCl₂

Answer: 2.Ca(OCl)₂

Bleaching powder is – a mixture of calcium chlorohypochlorite [Ca(OCl)₂] and basic calcium chloride [CaCl₂-Ca(OH)₂-2H₂O]. The compound present as an active ingredient in bleaching powder for bleaching action is Ca(OCl)₂

Question 17. Which of the following has the lowest melting point

1. CaCl₂
2. CaBr₂
3. Cal₂
4. CaF₂

Answer: 3.Cal₂

With the increase in the size of the halogens (from toI), the covalent character of the corresponding compounds (from CaF₂ to Cal₂) increases. Therefore, the melting points of the compounds decrease

Class 11 S-Block Notes

Question 18. In the case of replacement reaction, the reaction will be most favorable if M happens to be

1. K
2. Rb
3. Li
4. Na

Answer: 2. Rb

For a particular alkyl halide, the reactivity of the alkali metal fluorides increases gradually from Li to Cs because their ionization enthalpies decrease and their electronegativities increase. Therefore, for the given reaction, if M = Rb, then the reaction becomes most favorable.

Question 19. Which one of the alkali metals forms only the normal oxide M₂O on heating in air

1. Li
2. Na
3. Rb
4. K

Answer: 1. Li

Question 20. The ease of adsorption of the hydrated alkali metal ions on the ion-exchange resins follows the order

1. K⁺ < Na+ ⁺ Rb⁺ < Li⁺
2. Na⁺ < Li⁺ < K⁺ < Rb⁺
3. Li⁺ < K⁺ < Na⁺ < Rb⁺
4. Rb⁺ < K+< Na+< Li⁺

Answer: 4. Li⁺ < K⁺ < Na⁺ < Rb⁺ 

The ease of adsorption of the hydrated alkali metal ions on the ion-exchange resins decreases with the increase in the size of the alkali metal. The correct order of the sizes of the alkali metals is: Li⁺ < Na⁺ < K⁺ < Rb⁺ .4.

The ease of adsorption of the hydrated alkali metal ions on the ion-exchange resins decreases with the increase in the size of the alkali metal. The correct order of the sizes of the alkali metals is: Li⁺ < Na⁺ < K⁺ < Rb⁺

Question 21. Be²⁺is isoelectronic with which of the following ions

1. Na⁺
2. Mg²⁺
3. H⁺
4. Li⁺

Answer: 4. Li⁺

No. ofelectrons in Be = 4

.-. No. of electrons in Be²⁺ = 2

Again, no. of electrons in Li = 3

No. ofelectrons in Li+ = 2

Therefore, Be²⁺ is isoelectronic with Li⁺

Question 22. On heating which of the following releases CO₂ most easily

1. K₂CO₃
2. Na₂CO₃
3. MgCO₃
4. CaCO₃

Answer: 3.

The thermal stability of the given compounds follows the order: of K₂CO₃ > Na₂CO₂ > CaCO₂ > MgCO₂.

Thus, MgCO₃ Can release CO₂ most easily on heating

Question 23. In contrast with beryllium, which one of the following statements is incorrect

1. It forms Be₂C
2. Its salts rarely hydrolyzed
3. Its hydride is electron deficient and polymeric
4. It is rendered passive by nitric acid

Answer: 2. Its salts rarely hydrolyzed

Beryllium forms covalent compounds, thus hydrolysis of beryllium compounds occurs readily.

Question 24. The suspension of slaked lime in water is

1. Quicklime
2. Milk of lime
3. Aqueous solution of slaked lime
4. Lime water

Answer: 2. Milk of lime

Suspension of slaked lime Is known as ‘milk ofIimei

Class 11 S-Block Notes

Question 25. Which of the following statements is false

1. Mg ²⁺ ions form a complex with ATP
2. Ca²⁺ ions are important in blood clotting
3. Ca²⁺ ions are not important in maintaining the regular beating of the heart
4. Mg²⁺ ions are important in the green parts of plants

Answer: 3. Ca²⁺ ions are not important in maintaining the regular beating of the heart

Ca²⁺ plays an important role in the regular beating of the heart

Question 26. The product obtained as a result of a reaction of nitrogen

1. Ca(CN)₂
2. CaCN₃
3. Ca₂CN
4. CaCN

Answer: The given options are not correct

Question 27. Which one of the following takes up CO₂ and releases

1. K₂O
2. CaO
3. KO₂
4. KOH

Answer: 3.KO₂

2KO₂ + CO₂ → K₂ CO₃ + \(\frac{3}{2}\) O₂

Question 28. Ionic mobility of which of the following alkali metal ions is lowest when the aqueous solution of these salts is put under an electric field

1. K
2. Rb
3. Li
4. Na

Answer: 3. Li

Question 29. Magnesium reacts with an element (X) to form an ionic compound. If the ground state electronic configuration of (X) is ls²2s²2p³ the simplest formula for this compound is

1. Mg₃X₂
2. Mg₂X₃
3. Mg₂X
4. Mg X₂

Answer: 1. Mg₃X₂

Electronic configuration of X: ls²2s²2p³

Valency of X = 3

Here, the valency of Mg = 2

Thus the compound will be Mg₃X₂

Question 30. Which of the following oxides is most acidic nature

1. CaO
2. MgO
3. BaO
4. BeO

Answer: 4. BeO

On moving down the group basic character of the gr-2 On moving down the group basic character of the gr-2

BeO < MgO < CaO < SrO < BaO

Thus, BeO is acidic among all the given oxides.

Class 11 S-Block Notes

Question 31. Among CaH₂, BeH₂, BaH₂ the order of ionic character

1. BaH₂< BeH₂ < CaH₂
2. BeH₂ < CaH₂ < BaH₂
3. BeH₂<BaH₂<CaH₂
4.  CaH₂ < BeH₂, < BaH₂

Answer: 2.BeH₂ < CaH₂ < BaH₂

Question 32. Which of the following is not hygroscopic

1. CsCl
2. MgCl₂
3. CaCl₂
4. LiCI

Answer: 1. CsCl is not hygroscopic in nature while MgCl,  CaCI_4, and LiCl are hygroscopic in nature

Question 33. Which is the correct order of solubility in water—

1. Ba(OH)₂ <Mg(OH)₂
2. BaCO₃ >CaCO₃
3. CaSO₄ <MgSO₄
4. Ca(OH)2^Mg(OH)₂

Answer: 3. CaSO₄ <MgSO₄

Question 34. Bleaching powder does not contain—

1. CaCl₂
2. Ca(OH)₂
3. Ca(OCl)₂
4. Ca(ClO₃ )₂

Answer: 4.Ca(ClO₃)₂

Bleaching powder is a mixture of calcium hypochlorite, Ca(OCI)_2, and the basic chloride CaCl₂, H₂O with some slaked lime, Ca(OH)₂.

Question 35. Which of the following statements is incorrect—

1. Li⁺ has a minimum degree of hydration
2. The oxidation state of K in KO₂ is +1
3. Na is used to make a Na/Pb alloy
4. MgSO₄ is readily soluble in water

Answer: 1.  Li⁺ has a minimum degree of hydration

The hydration enthalpies of alkali metal ions decrease with an increase in ionic sizes. Hence, the orderis Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺. Therefore Li⁺ has a maximum degree of hydration

Question 36. Which of the following has the highest hydration energy

1. MgCl₂
2. CaCl₂
3. BaCl₂
4. SrCl₂

Answer: 1. MgCl₂

Smaller-sized and highly charged metal ions have higher hydration energy. Therefore, the order of hydration energy will be

Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺

Question 37. Which is the correct order of solubility in water

1. Ba(OH)₂< Mg(OH)₂
2. BaCO₃ >CaCO₃
3. CaSO₄ < MgSO₄
4. Ca(OH)₂ ≅ Mg(OH)₂

Answer: 3.CaSO₄ < MgSO₄

Class 11 S-Block Notes

Question 38. What is the product of the reaction between the dilute solution of Ba(OH)₂ and H₂O₂ + ClO₂ 

1. HOCl
2. Ba(OCl)₂
3. Ba(ClO₃)₂
4. Ba(ClO₂)₂

Answer: 4.Ba(ClO₂)₂ 

Ba(OH)₂  +H₂ O₂  + 2ClO₂  → Ba(ClO₂ )₂  + 2H₂ O + O₂ 

Question 39. Which of the following compounds transforms baking soda into baking powder

1. KHCO₃
2. NaHCO₃
3. KHC₄H₄O₆
4. KCl

Answer: 3.KHC₄H₄O₃

Question 40. Which of the following pair of compounds does not undergo any chemical change on heating

1. MgCO₃,KHCO₃
2. CS₂CO₃,KNO₂
3. Na₂CO₃,NaNO₃
4. Li₂CO₃, KNO₂

Answer: 2.  CS₂CO₃, KNO₂

Question 41. Ca(OH)₂ (anhydrous)  CaCl₂ +B In this reaction, B is

1. CaOCl₂
2. Ca(ClO₃)₂
3. Ca(OH)₂
4. Ca(ClO₂)₂

Answer: 2. Ca(ClO₃)₂

Question 42. The blue color of potassium solution in liquid ammonia is due to the presence of

1. Solvated electron
2. Potassium amide
3. Impurities present in potassium
4. Potassium oxide

Answer: 1. Solvated electron

Question 43. The alkali metal that does not participate in the reaction MI + I₂ → MI₃ is

1. Na
2. K
3. Rb
4. Cs

Answer: 1. Na

Question 44. The lithium compound which is soluble in water is

1. Li₂CO₃
2. LiNO₃
3. LiF
4. Li₃PO₄

Answer: 2. LiNO₃

Question 45. An alkaline earth metal when heated along with nitrogen forms X. On hydrolysis, X forms an insoluble basic compound and a gas which turns CuSO₄ solution deep blue. The metal is

1. Be
2. Ca
3. Mg
4. Ba

Answer: 3. Mg

Class 11 S-Block Notes

Question 46. Which of the following compounds exists in only an aqueous solution

1. Li₃PO₄
2. NaHCO₃
3. [Li(NH₃)₄]I
4. LiHCO₃

Answer: 4.LiHCO₃

Question 47. Which of the following acids does not liberate CO₂ on reacting with sodium carbonate

1. Dilute HCl
2. Dilute H₃BO₃
3. Dilute H₃PO₄
4. Dilute H₂PO₄

Answer: 2. Dilute H₃BO₃

Question 48. Hydrated BeCl₂ acts as a/an

1. Lewis base
2. Arrhenius base
3. Arrhenius acid
4. Lewis acid

Answer: 4. Lewis acid

Question 49. Which of the following compounds or pair of compounds is responsible for turning yellow sodium peroxide white in the presence of air

1. Na
2. NaOH and Na₂CO
3. H₂O₂
4. NaOH and H₂O₂

Answer: 2.NaOH and Na₂CO

Question 50. The chemical formula ofmicrocosmic saltis

1. NaHPO₄.2H₂O
2. NH₄₃HPO₄.2H₂O
3. Na(NH₄)HPO₄.4H₂O
4. (NH₄)3HPO₄.4H₂O

Answer: 3.Na(NH₄)HPO₄.4H₂O

Question 51. The hydride which does not form as a result of a direct reaction between the metal and hydrogen is

1. CaH₂
2. MgH₂
3. BeH₂
4. NaH

Answer: 3.BeH₂

Question 52. An amphoteric oxide dissolves in HCl to form a salt. The salt does not impart any characteristic color to the flame but fumes in moist air. The oxide is

1. BaO₂
2. MgO
3. BeO
4. CaO

Answer: 3. BeO

Question 53. The difference in the number of water molecules in gypsum and Plaster of Paris is—

1. 5/2
2. 2
3. 1/2
4. 3/2

Answer: 4. 3/2

Class 11 S-Block Notes

Question 54. Compound A imparts a brick-red color to the flame and decomposes on heating to produce oxygen and a brown gas. A is

1. Mg(NO₃)₂
2. Ba(NO₃)2
3. Ca(NO₃)₂
4. Sr(NO₃)₂

Answer: 2. Ba(NO₃)₂

Question 55. The sodium salt of an unknown anion when heated with MgCl₂ forms a white precipitate. The anion is

1. SO^2-₄
2. CO^2-₃
3. HCO^–₃
4. NO^–₃

Answer: 3.HCO^–₃

Question 56. The salt which is added to table salt to keep it dry and free-flowing

1. KCl
2. Ca(PO₄)₂
3. KI
4. Na₃PO₄

Answer: 2.Ca(PO₄)₂

Question 57. The compound which acts as an oxidizing as well as a reducing agent

1. NaNO₃
2. Na₂O
3. Na₂O₂
4. KNO₃

Answer: 3.

Question 58. The compound which is used to extinguish fire caused by combustion of alkali metals is—

1. CC1₄
2. Sand
3. Water
4. Kerosene

Answer: 1.  CC1₄

Question 59. An aqueous solution of sodium sulphate is electrolyzed using inert electrodes. The products formed at the cathode and anode respectively are

1. O₂,H₂
2. H₂,O₂
3. O₂, Na
4. O₂,SO₂

Answer: 1.O₂, H₂

Question 60. Sodium when heated at 300°C in air forms X which absorbs CO₂ to form Na₂CO₃ and a compound Y. The compound Y is 

1. H₂
2. O₂
3. H₂O₂
4. O₃

Answer: 2.O₂

Question 61. Excess Na⁺ ions in the human body cause

1. Diabetes
2. Anaemia
3. Low blood pressure
4. High blood pressure

Answer: 4. High blood pressure

Class 11 S-Block Notes

Question 62. The metal carbide which on hydrolysis produces allylene or propyne is

1. Be
2. Ca
3. Al
4. Mg

Answer: 4. Mg

Question 63. A metal M readily forms water-soluble sulfate MSO₄, water-insoluble hydroxide M(OH)_2, and oxide MO. The oxide remains inert on heating. The hydroxide is soluble in NaOH. M is

1. Be
2. Mg
3. Ca
4. Sr

Answer: 1. Be

Question 64. When compound A is heated, it produces a colourless gas, and the residue obtained is dissolved in water to form compound B. The compound C is formed when excess CO₂ is passed through the aqueous solution of B. The compound C can be separated in a solid state from its solution. C in its solid state when heated forms the compound A. Ais

1. CaCO₃
2. Na₂CO₃
3. K₂CO₃
4. CaSO₄.2H₂O

Answer: 1.CaCO₃

Question 65. When metals A and B are heated in air, A only forms oxide but B forms both oxide and nitride. A and B are

1. Cs, K
2. Mg, Ca
3. Li, Na
4. K, Mg

Answer: 4. K, Mg

Question 66. The alkali metal which emits a light of the longest wavelength in the flame testis

1. Na
2. K
3. Cs
4. Li

Answer: 2. K

Question 67. The compound which does not form a double salt is

1. Li₂SO₄
2. Na₂SO₄
3. K₂SO₄
4. Rb₂SO₄

Answer: 1. Li₂SO₄

Question 68. The decreasing order of stability of the chloride salts of alkali metals is

1. LiCl > KCl > NaCl > CsCl
2. CsCl > KCl > NaCl > LiCl
3. NaCl > KCl > LiCl > CsCl
4. KCl > CsCl > NaCl > LiCl

Answer: 4. KCl > CsCl > NaCl > LiCl

Question 69. The affinity of sodium for water is used for drying

1. Alcohol
2. Ammonia
3. Benzene
4. Phenol

Answer: 3. Benzene

Class 11 S-Block Notes

Question 70. The ions present in an anhydrous mixture of potassium fluoride and hydrofluoric acid are

1. K⁺,H⁺ F^–
2. (KF)⁺+ (HF)^–
3. KH⁺, F^–
4. K⁺,HF₂^–

Answer: 4. K⁺,HF₂^–

Question 71. The correct order of covalent character in the following compounds is

1. LiCl < NaCl < BeCl₂
2. BeCl₂ > LiCl > NaCl
3. NaCl < LiCl < BeCl₂
4. BeCl₂ < NaCl < LiCl

Answer: 3. NaCl < LiCl < BeCl₂

Question 72. A chemical compound A is used for the recovery of ammonia during the preparation of washing soda. When CO₂ is passed through the aqueous solution of A, the solution turns turbid. A is used for whitewashing because of its disinfecting properties. The chemical formula of A is

1. Ca(HCO₃)₂
2. CaO
3. Ca(OH)₂
4. CaCO₃

Answer: 3. Ca(OH)₂

Question 73. The compound used for drying neutral or basic gases is—

1. Calcium carbonate
2. Sodium carbonate
3. Sodium bicarbonate
4. Calcium oxide

Answer: 4. Calcium oxide

Question 74. Which of the following compounds does not contain calcium carbonate

1. Dolomite
2. Marble statue
3. Burnt gypsum
4. Snail shell

Answer: 3. Burnt gypsum

Question 75. The compound which on thermal decomposition produces a basic and an acidic oxide is

1. KClO₃
2. Na₂CO₃
3. NaNO₃
4. CaCO₃

Answer: 4.CaCO₃

Question 76. Which of the following oxides does not react with water

1. BeO
2. CaO
3. MgO
4. SrO

Answer: 1. BeO

Question 77. Which of the following carbonates is soluble in water

1. SrCO₃
2. BaCO₃
3. Al₂(CO₃)₃
4. Rb₃CO₃

Answer: 3.Al₂(CO₃)₃

Class 11 S-Block Notes

Question 78. The compound whose aqueous solution is called ‘baryta water’ is

1. BaSO₄
2. BaO
3. BaCO₃
4. Ba(OH)₂

Answer: 4. Ba(OH)₂

Question 79. The alkali metal that emits a light of the shortest wavelength in the flame test is

1. Na
2. K
3. Cs
4. Li

Answer: 3. Cs

Question 80. The correct order of ionic mobility of the following ions in their aqueous solution is

1. Na⁺> K⁺> Rb⁺> Cs⁺
2. K⁺> Na⁺ > Rb⁺> Cs⁺
3. Cs⁺ > Rb⁺ > K⁺ > Na⁺
4. Rb⁺ > K⁺ > Cs⁺ > Na⁺

Answer: 3. Cs⁺ > Rb⁺ > K⁺ > Na⁺

Question 81. Which of the following compounds is not used for storing or immersing metallic sodium

1. Benzene
2. Kerosene
3. Ethanol
4. Toluene

Answer: 3. Ethanol

Question 82. Which of the following compounds is paramagnetic

1. KO₂
2. SiO₂
3. TiO₂
4. BaO₂

Answer: 1.KO₂

Question 83. KO₂ is used in oxygen cylinders that are used for submarines and spacecraft because—

1. It increases the amount of oxygen by absorbing CO₂
2. It eliminates water vapour
3. It absorbs CO₂
4. It forms O₃

Answer: 1.It increases the amount of oxygen by absorbing CO₂

Question 84. Which of the following compounds is the most stable

1. LiF
2. LiCl
3. LiBr
4. Lil

Answer: 1. LiF

Question 85. The alkali metal for which the photoelectric effect is maximum is

1. Cs
2. Na
3. K
4. Li

Answer: 1. Cs

Question 86. The melting points of alkali metals are low. Which of the alkali metals melts when room temperature becomes more than 30 °C 

1. K
2. Na
3. Cs
4. Rb

Answer: 3. Cs

Class 11 S-Block Notes

Question 87. Which of the following alkali metal hydroxides is the most basic in nature

1. CsOH
2. KOH
3. Li OH
4. RbOH

Answer: 1.CsOH

Question 88. The metal whose carbonate salt is the most stable is

1. Na
2. Mg
3. A1
4. Si

Answer: 1. Na

Question 89. Which of the following compounds is used for manufacturing soft soaps

1. KOH
2. NaOH
3. LiOH
4. Mg(OH)₂

Answer: 1.

Question 90. The atomic number of a radioactive alkali metal is

1. 55
2. 87
3. 19
4. 37

Answer: 2. 87

Question 91. Which of the following is used for the manufacture of high? temperature thermometers

1. An alloy of Li and Na
2. An alloy of Na and Cs
3. An alloy of Na and K
4. An alloy of K and Rb

Answer: 3. An alloy of Na and K

Question 92. The mixture of MgCl₂ and MgO is known as

1. Sorel’s cement
2. Portland cement
3. Alum
4. Magnesium oxychloride

Answer: 1. Sorel’s cement

Question 93. Which of the following alkaline earth metals does not form its corresponding hydride by directly reacting with hydrogen

1. Mg
2. Sr
3. Be
4. Ba

Answer: 3. Be

Question 94. Which of the following alkaline earth metal hydroxides is soluble in NaOH solution

1. Ba(OH)₂
2. Ca(OH)₂
3. Mg(OH)₂
4. Be(OH)₂

Answer: 4. Be(OH)₂

Question 95. The chloride salt which is soluble in ethanol is

1. BeCl₂
2. CaCl₂
3. SrCl₂
4. MgCl₂

Answer: 1.BeCl₂

Class 11 S-Block Notes

Question 96. Which of the following is produced when one mole of magnesium nitride reacts with an excess of water

1. One mole of ammonia
2. One mole of nitric acid
3. Two moles of ammonia
4. Two moles of nitric acid

Answer: 3. Two moles of ammonia

Question 97. Which of the following methods is used for the preparation of calcium

1. Reduction of CaO by carbon
2. Reduction of CaO by hydrogen
3. Electrolysis of a mixture of anhydrous CaCl₂ and KCl
4. Electrolysis of molten Ca(OH)₂

Answer: 3. Electrolysis of a mixture of anhydrous CaCl₂ and KCl

Question 98. Which alkaline earth metal ion plays a vital role in the contraction of muscles

1. Ba²⁺
2. Sr²⁺
3. Mg²⁺
4. Ca²⁺

Answer: 4. Ca²⁺

Question 99. Which of the following correctly indicates the formula of halides of alkaline earth metals

1. BeCl₂.2H₂O
2. BeCl₂.4H₂O
3. CaCl₂.6H₂O
4. SrCl₂.4H₂O

Answer: 1.BeCl₂.2H₂O

Question 100. The compounds of sodium which are used in the textile Industry are

1. Na₂CO₃
2. NaHCO₃
3. NaOH
4. NaCl

Answer: 1. Na₂CO₃

Question 101. Which of die following pairs of elements have similar properties

1. Be, Cs
2. K, Cs
3. Sr, Rb
4. Be, Al

Answer: 2. K, Cs

Question 102. The chlorides which are soluble in pyridine are

1. LiCl
2. CsCl
3. NaCl
4. BeCl₂

Answer: 1 and 4

Question 103. The gases in which magnesium burns are

1. CO₂
2. N₂O
3. N₂
4. SO₂

Answer: 1,2,3 and 4

Question 104. Which of the following oxides have rock salt structure with coordination number 6:6 

1. MgO
2. CaO
3. SrO
4. B₂O₃

Answer: 1,2 and 3

Question 105. Which of the following pairs of compounds cannot exist in aqueous solution

1. NaH₂PO₄ and Na₂HCO₃
2. Na₂CO₃ and NaHCO₃
3. NaOH and NaH₂ PO₄
4. NaHCO₃ and NaOH

Answer: 3 and 4

Question 106. The compounds which on heating do not form oxides are

1. NaNO₃
2. CsOH
3. LiOH
4. SrCO₃

Answer: 1 and 2

Question 107. Which of the following pairs of elements will give superoxides and peroxides respectively when heated with excess air

1. K, Br
2. Nay Rb
3. K, Rb d
4. Na, Ba

Answer: 3 and 4

Class 11 S-Block Notes

Question 108. Which of the following do not respond to flame test

1. Be
2. Mg
3. KO₂
4. Sr

Answer: 1 and 2

Question 109. Which of the following compounds are not paramagnetic nature

1. K₂O₂
2. NO₂
3. KO₂
4. K₂O

Answer: 1 and 4

Question 110. Which of the following is incorrect

1. Soda ash: Na₂CO₃
2. Pearl ash: Cu₂CO₃
3. Bone ash: K₂CO₃
4. Baking soda: NaHCO₃

Answer: 2 and 3

Question 111. Which ions of water are replaced by sodium ions when hard water is passed through zeolite (hydrated sodium aluminum silicate)

1. H⁺
2. Mg⁺
3. Ca²⁺
4. SOl₄^2-

Answer: 2 and 3

Question 112. The compounds which are soluble in organic solvents are

1. CaCl₂
2. BaCl₂
3. BeCl₂
4. AlCl₃

Answer: 3 and 4

Question 113. The compounds formed when potassium superoxide reacts with water are

1. KOH
2. H₂O₂
3. K₂O₂
4. O₂

Answer: 1, 2 and 4

Question 114. Which of the following hydrated salts undergo hydrolysis on heating

1. BaCl₂.2H₂O
2. MgCl₂.6H₂O
3. SrCl₂.2H₂O
4. CaCl₂.6H₂O

Answer: 2 and 4

Class 11 S-Block Notes

Question 115. Which of the following compounds are extensively used as drying agents

1. Anhydrous CaCl₂
2. Mg(ClO₄)2
3. BeC
4. Ca(OH)₂

Answer: 1 and 2

Question 116. Each of the following compounds reacts with water but which of these liberates the same gas

1. Na
2. Na₂O₂
3. KO₂
4. NaH

Answer: 1 and 4

Question 117. The polarisability of LiCl is higher than that of NaCl. Concerning this, which of the following statements are true

1. The melting point of LiCl is less than that of NaCl
2. LiCl is sparingly soluble in organic solvents
3. LiCl dissociates to a greater extent in water than NaCl
4. The conductivity of molten LiCl is less than that of NaCl

Answer: 1 and 4

Question 118. Which of the following statements are correct

1. Electronegativity of alkali metals decreases with an increase in atomic number
2. Lithium is the lightest metal
3. Alkali metals are strong reducing agents –
4. The electronegativity of alkali metals ranges from 1.0 to 0.7

Answer: 1, 3, 4

Question 119. Which of the following statements are incorrect about the hydrates of alkali metals

1. Conduct electricity in their molten states
2. These compounds act as oxidizing agents
3. These compounds dissolve in water to liberate
4. These compounds are covalent

Answer: 2 and 4

Question 120. Which of the following statements are correct about the ionic solids KI and CaO lattice enthalpy of CaO is greater than that of K₃

1. I am soluble in benzene
2. The melting point of CaO is high
3. The melting point of KI is high

Answer: 1, 2, 3

Question 121. The compounds which do not form NO₂ on undergoing thermal decomposition are

1. LiNO₃
2. NaNO₃
3. KNO₃
4. RbNO₃

Answer: 2, 3, and 4

Question 122. Metals are identified by their standard reduction potential, enthalpy of fusion, and atomic size. The alkali metals are identified by their

1. High boiling point
2. High negative standard reduction potential
3. High density
4. Greater atomic size

Answer: 2 and 4

Class 11 S-Block Notes

Question 123. Which of the following sulfates easily dissolve in water

1. BeSO₄
2. MgSO₄
3. BaSO₄
4. SrSO₄

Answer: 1 and 2

Question 124. The properties of beryllium nitride which are different from the nitrides of other alkaline earth metals are

1. Its volatility
2. Its covalent nature
3. Unable to undergo hydrolysis
4. Its ionic nature

Answer: 1 and 2

Question 125. Which of the following options is correct for RbO₂

1. It is a peroxide
2. It is diamagnetic
3. It is a superoxide
4. It is paramagnetic

Answer: 3 and 4

Question 126. Which of the following statements is correct for the alkaline earth metals

1. Hydration enthalpy of Sr²⁺is less than that of Ba²⁺
2. CaCO₄ undergoes decomposition at a higher temperature than BaCO₃
3. Ba(OH)₂ is a stronger base than Mg(OH)₂
4. SrSO₄  is more soluble in water than CaSO₄

Answer: 2 and 3

Question 127.  Which of the following statements are correct about the alloy formed by sodium and potassium

1. It is used in Lassaigne’s test
2. It is liquid at ordinary temperature
3. It is used in specially designed thermometers
4. It is used as a coolant in nuclear reactors

Answer: 2 and 3

Question 128. Which of the following statements is incorrect

1. BeCl₂ molecule is linear in the gaseous state
2. Calcium hydride is known as hydrolith
3. Carbides of both beryllium and calcium react with water to form acetylene
4. Oxides of both Be and Ca are amphoteric

Answer: 3 and 4

Question 129. Which of the following is less stable thermally

1. LiF
2. KCl
3. RbF
4. CsF

Answer: 2. KCl

The thermal stability of a compound increases with the increasing value of enthalpy of formation. Among the given compounds the value of enthalpy of formation is minimum (-428kJ.mol^-1) for KCl. Hence, it has the lowest thermal stability among the given compounds

Question 130. Which of the following pairs is responsible for developing an electric potential across the membrane of living cells

1. Ca²⁺and Na⁺
2. Na⁺ and K⁺
3. K⁺ and Ba²⁺
4. Mg²⁺ and Ca²⁺

Answer: 3. K⁺ and Ba²⁺

Class 11 S-Block Notes

Question 131. Which one of the following chlorides is a soluble organic solvent

1. CaCl₂
2. NaCl
3. MgCl
4. BeCl₂

Answer: 4.MgCl

Question 132. Which of the following alkaline earth metal carbonates is thermally least stable

1. BaCO₃
2. CaCO₃
3. SrCO₃
4. BeCO₃

Answer: 4.SrCO₃

Question 133. The cause of the different colors of the flame in the flame test is

1. Lowionisation potential
2. Low melting point
3. Malleability
4. The presence of one electron in the outermost orbit

Answer: 1. Lowionisation potential

Question 134. Which of the following alkaline earth metal sulfate is most soluble in water

1. CaSO₄
2. SrSO₄
3. BaSO₄
4. MgSO₄

Answer: 4. BaSO₄

Question 135. What will be the order of reducing powder of the following elements Li,Na,  K, Rb, Cs

1. Cs > Rb > K > Na > Li
2. Rb > Cs > K > Na > Li
3. K > Rb > Cs > Na > Li
4. Na > Li> K >Cs > Rb

Answer: 1. Cs > Rb > K > Na > Li

The order of reducing free states is:

Cs > Rb > K > Na > Li

Question 136. Which one of the following elements shows a diagonal relationship with magnesium

1. Na
2. Li
3. Be
4. Ca

Answer: 2. Li

Question 137. Sodium is preserved in which of the following liquids

1. Water
2. Ethanol
3. Kerosene oil
4. Methanol

Answer: 3. Kerosene oil

Question 138. Which of the alkali metals has having least melting point?

1. Na
2. K
3. Rb
4. Cs

Answer:  4. Cs

As the atomic size of the metal increases, the strength of metallic bonding decreases and consequently, the melting point decreases. Since the size of the cesium atom is the largest, it has the least melting point

Question 139. Which one of the following alkali metals gives hydrated salts? 

1. Li
2. Na
3. K
4. Cs

Answer: 1.  Li

 Li; Explanation:

Among all the alkali metal ions, Li+ is the smallest and thus it has the highest charge density. As a result, Li⁺ attracts water molecules more strongly than any other alkali metal cation and forms hydrated salts.

Class 11 S-Block Notes

Question 140. Which one of the alkaline earth metal carbonates is thermally the most stable? 

1. MgCO₃
2. CaCO₃
3. SrCO₃
4. BaCO₃
   

Answer:  4. BaCO₃

BaCO₃; Explanation:

Among all the alkaline earth metals, the Ba²⁺ ion is the largest. Again, CO₃^2- ion is also quite large. Thus, among the given carbonates, Ba²⁺ and CO₃^2-  ions are most tightly packed in the crystal lattice of BaCO₃. Consequently, the lattice enthalpy of BaCO₃ is the highest and so, is most thermally stable.

Class 11 Chemistry S Block Elements Very Short Answer Type

Question 1. Name the lightest and the heaviest metal.
Answer:   Lithium is the lightest (density: 0.53g-cm^-3) and osmium is the heaviest (density: 22.6g-cm^-3 ) metal.

Question 2. Which one between water and pyrene (CCl₄), can be used to extinguish fire caused by metallic sodium?
Answer:  Pyrene (CCl₄) can be used.

Question 3. Which among the alkali metal ions has the lowest mobility In aqueous solution?
Answer: Li⁺ ions have the lowest mobility because they remain highly hydrated in aqueous solution.

Question 4. Which of the alkali metal cations has the highest polarising power?
Answer: Due to its very small size, Li⁺ ion has the highest polarising power among the alkali metal ions

Question 5. The density of alkali metals is very low Why?
Answer: Due to large atomic size and weak metallic bonding, the densities of alkali metals are very low.

Question 6. Which out of LIF of Lil is more covalent?
Answer:  Lil is more covalent.

Question 7. Which alkali metal carbonate easily liberates CO₂ on heating?
Answer: Li₂CO₃ easily liberates CO₂ on heating.

Question 8. Which alkali metal chloride imparts violet color to the bunsen burner flame?
Answer: Potassium chloride (KCl)

Question 9. Which alkali metal hydride is used as a source of hydrogen for filling up meteorological balloons?
Answer: Lithium hydride (LiH).

Question 10. Name the products formed in the following reaction. Explain Lil + KF→? + ?
Answer: LiF and KI. The larger K⁺ ion stabilizes the larger I^– ion while the smaller Li⁺ ion stabilizes the smaller F- ion.

Question 11.  What is water glass?
Answer: Sodium silicate (NaOSiO₃).

Question 12. Which alkali metal combines with nitrogen to form the corresponding nitride?
Answer: Lithium (Li).

Question 13. What are the raw materials used for the manufacture of sodium carbonate by the Solvay process?
Answer: The raw materials are NaCl, CaCO and NH₃.

Question 14. Which out of the Na₂CO₃ solution & NaHCO₃ solution, changes the color of phenolphthalein into pink?
Answer: Na₂CO₃ solution.

Question 15. What is the main ingredient of baking powder?
Answer: Sodium bicarbonate (NaHCO₃).

Question 16. Which compound is used to treat the flue gases from coal and oil-fired power stations and to remove SO₂ and H₂SO₄ responsible for acid rain?
Answer: Na₂CO₃.

Question 17. What are the ingredients of the fusion mixture that is used in dry tests in inorganic analysis?
Answer: K₂CO₃ and Na₂CO₃

Question 18. Which alkaline earth metal is the most abundant in the earth’s crust?
Answer: Calcium (Ca).

Question 19. Which alkaline earth metal is radioactive?
Answer: Radium (Ra).

Question 20. Name the metal of group-2 which is used to prepare Grignard reagent.
Answer: Magnesium (Mg).

Question 21. Name the element of group 2 that resembles lithium in characteristics.
Answer: Magnesium (Mg).

Question 22. Which alkaline earth metals do not impart any color to the flame of a Bunsen burner?
Answer: Be anil Mg.

Question 23. Name an alkaline earth metal compound that can be used as a portable source of hydrogen.
Answer: Calcium hydride (CaH₂).

Question 24. Which Gr-2 metal forms covalent compounds?
Answer: Beryllium (Be).

Question 25. Which Gr-2 metal burns readily when exposed to air?
Answer: Barium (B a).

Question 26. Which alkaline earth metal reacts with alkali to form hydrogen gas?
Answer: Beryllium (Be).

Question 27. What is the composition of the alloy, electron?
Answer: 95% of Mg+ 5% of Zn.

Question 28. Which reagent is used to analyze Ca²⁺ and Mg²⁺ ions quantitatively?
Answer: EDTA (Ethylenediaminetetraacetic acid).

Question 29. What is the medicinal name of the aqueous solution of Mg(OH)₂?
Answer: Milk of magnesia.

Class 11 S-Block Notes

Question 30. What is anhydrous?
Answer: Magnesium perchlorate, Mg(ClO₄)₂ is known as anhydrous.

Question 31. How will you distinguish between BeSO₄ & BaSO₄?
Answer: BeSO₄ is water soluble but BaSO₄ is insoluble in water.

Question 32. Which out of MgCO₃, SrCO_3, and BaCO₄, possesses the highest thermal stability?
Answer: BaCO₄ has the highest thermal stability.

Question 33. Distinguish between Be(OH)₂ and Ba(OH)₂
Answer: Be(OH)₂ is soluble in caustic soda solution while Ba(OH)2 is insoluble in it.

Question 34. Why is BeCl₂ soluble in organic solvents?
Answer: BeCl₂ is covalent.

Question 35. Which Gr-2 metal carbonate is unstable in the air?
Answer: BeCO₃.

Question 36. Which alkaline earth metal sulfate is useful in diagnosing stomach ulcers by X-ray?
Answer: BaSO₄ (in ‘barium meal’ X-ray).

Question 37. BeO is covalent but still, it has a much higher melting point— mention the reason.
Answer: This is due to its polymeric structure.

Question 38. What is the difference between lime water and milk of lime?
Answer: Mg²⁺

Question 39. Write a reaction by which BeCl₂ can be prepared.
Answer:

Question 40. Explain why the compounds of beryllium are much more covalent than the other Gr-2 metal compounds.
Answer:

Due to the very small size and high charge of Be²⁺ ion, it has much higher polarising power and because of this, the compounds of Be are much more covalent than the other group-2 metal compounds.

Question 41. Beryllium compounds are extremely toxic why?
Answer:

Beryllium compounds are extremely toxic because of their very high solubility and their ability to form complexes with enzymes in the body.  Also, beryllium displaces magnesium from some enzymes

Question 42. What is the composition of soda lime used for the preparation of hydrocarbons in the laboratory?
Answer: NaOH and CEO
.
Question 43. ED Name two acidic oxides which react similarly with calcium hydroxide [Ca(OH)₂].
Answer: CO₂ and SO₂.

Question 44. Which alkaline earth metal oxide is used as a flux in metallurgy to remove siliceous impurities?
Answer: Calcium oxide (CaO)

Question 45. What is the commercial name of the disinfectant powder obtained when Cl2 reacts with slightly moist slaked lime at 40°C?
Answer: Bleaching powder.

Question 46. What is Plaster of Paris?
Answer: Hemihydrate of calcium sulphate, (CaSO₄)₂-H₂O is called Plaster of Paris

Question 47. What type of impurities in gypsum should be avoided in preparing Plaster of Paris from it?
Answer: Carbonaceous impurities are to be avoided.

Question 45. Why KNO₃ is used instead of NaNO₃ in gunpowder?
Answer: NaNO₃ is deliquescent. Hence, KNO₃ is preferred over NaNO₃ for the preparation of gunpowder.

Question 46. What do you mean by ‘black ash
Answer: A mixture of sodium carbonate (Na₂CO₃) and calcium sulfide (CaS) is called black

Question 47. What happens when magnesium is heated with | acetylene at 875K?
Answer: Magnesium on heating with acetylene at 875K forms magnesium carbide (Mg₃C₂)

Question 48. Which calcium salt causes the formation of the kidney? stones?
Answer:

Calcium oxalate, \(\mathrm{Ca}^2+\stackrel{\ominus}{\mathrm{O}}_2 \mathrm{C}-\mathrm{C} \stackrel{\ominus}{\mathrm{O}}_2\) causes formation of kidney stones

Question 49. Alkali metals are good reducing agents—Why?
Answer: 

The smaller the ionization enthalpy, the greater the reducing strength. Since alkali metals have lower ionization enthalpies, they are good reducing agents.

Question 50. Explain why the alkali metals cannot be obtained by the reduction method.
Answer: Alkali metals are strong reducing agents and it is difficult to reduce their oxide by any other reducing agent

Question 51. Which alkali metal ion has the maximum polarising j power and why?
Answer:  Li⁺ ion has the highest polarising power among all the alkali metal ions as the value of charge to size ratio of the smallest Li+ ion is the highest.

Question 52. The fire caused by sodium in the laboratory cannot be extinguished by spraying water Why?
Answer:

Sodium reacts vigorously with water producing H₂ gas which catches fire by the heat evolved in the reaction. So water cannot be used for extinguishing sodium-fire

Class 11 S-Block Notes

Question 53. Why does Li not exist with Na or K in their minerals?
Answer:

Lithium forms independent minerals and does not exist with Na or K because the Li⁺ ion is too small to replace the more abundant Na+ or K⁺ ions in their minerals

Question 54. How can you prepare propyne from magnesium carbide?
Answer:

When magnesium carbide (MgC₂) is heated, Mg₂C₃ (magnesium allylide) is obtained. This on hydrolysis, produces propyne (CH₃C=CH)

Question 55. Why do halides of Be dissolve in organic solvents while those of Ba do not?
Answer:

Halides of Be are covalent. Hence, they dissolve in organic solvents while of Ba are ionic. Hence, they do not dissolve in organic solvents.

Question 56. Write the composition of gunpowder.
Answer: Gunpowder is an explosive mixture containing KNO₃ along with charcoal and sulfur

Question 57. Why is Na₂S₂O₃ used in photography?
Answer: In photography, Na₂S₂O₃ is used to dissolve the unexposed AgBr.

Question 58. The affinity of sodium towards water is used in drying benzene. Explain.
Answer: Sodium does not react with benzene. Hence it can be used effectively for drying benzene.

Question 59. Write with a balanced equation the reaction for the manufacture of sodium bicarbonate from sodium carbonate.
Answer:

Sodium bicarbonate (NaHCO₃) Is manufactured by passing CO2 through a saturated solution of sodium carbonate (Na₂CO₃).

Na₂CO₃ + CO₂ + H₂O ⇌ 2NaHCO₃↓

Question 60.  Name a pair of elements that exhibits a diagonal relationship.
Answer:  Lithium (Li) and magnesium (Mg) CO₂

Question 62. Name an alkaline earth metal.
Answer: Calcium (ca)

Question 63. Write the balanced equation for the reaction when water is added to calcium carbide.
Answer: CaC₂ + 2M₂O →MC=CH + Ca(OH)₂

Question 64. Write the balanced equation(s) for the reaction when excess carbon dioxide is passed through brine saturated with ammonia
Answer:

When carbon dioxide is passed through an aqueous solution of NaCI (brine, 28% NaCl solution) saturated with ammonia, sodium bicarbonate is formed.

Question 65. Which alkali metal ion has the highest polarising power?
Answer: Li⁺

Question 66. What are the common oxidation states exhibited by group-1 and group-2- 2 metals respectively?
Answer: +2 And +1 respectively

Question 67. Name the alkali metals which form superoxides when heated with excess oxygen.
Answer: K, Rb and cs

Question 68. Which one is the lightest and which one is the heaviest of all the metals?
Answer: Lithium (Li) and Osmium(os) respectively

Question 69. What is the composition of the white powder obtained when metallic magnesium is burnt in the air?
Answer: A mixture of MgO and Mg₃N₂

Question 70. Name two metals of group 2 which do not impart any color to the flame.
Answer: Mg and Be

Question 71. Which alkali metal cannot be stored in kerosene?
Answer: Li

Question 72. Which alkali metal is used as a scavenger in metallurgy to remove O₂ and N₂ gases?
Answer: Li

Question 73. Which Gr-2 metal carbide reacts with water to produce methane?
Answer: Be₂ C

Question 74. In an aqueous solution, which alkali metal ion has the lowest mobility?
Answer: Li⁺

Question 76. Which is the most abundant alkaline earth metal in the earth’s crust?
Answer: Calcium

Question 77. Name one group-2 metal.
Answer: Chlorophyll

Question 78. Which alkali metal ion forms a stable complex with 18- crown-6
Answer: Potassium(K)

Question 79. Which alkaline earth metal is largely used as a lightweight construction metal?
Answer: Mg

Question 80. Which alkaline earth metal forms an organometallic compound known as Grignard reagent?
Answer: Mg

Class 11 S-Block Notes

Question 81. Which alkali and alkaline earth metals are radioactive?
Answer: Fr and Ra

Question 82.  The salts of which alkali metals are commonly hydrated?
Answer: Li

Question 83. Which alkali metal acts as the strongest reducing agent in aqueous solution?
Answer: Li

Question 84. Which is the least stable alkali metal carbonate?
Answer: Li₂CO₃

Question 85. What is baryta water?
Answer: Ba(OH ₂ solution

Question 86. Which alkaline earth metal hydroxide is most soluble in water?
Answer: Ba(OH)₂

Question 87. Which alkaline earth metal imparts crimson color to flame?
Answer: Srcl₂

Question 88. BeO is covalent and still has high melting points — why?
Answer: It is polymeric

Question 89. What is anhydrous?
Answer: Magnesium perchlorate Mg(ClO₄)₂

Question 90. Which alkaline earth metal carbonate can be kept only in an atmosphere of CO₂?
Answer: BeCO₃

Question 91. Which alkaline earth metal chloride is used as a desiccant in the laboratory?
Answer: CaCl₂

Question 92. The basic strength of which alkali metal hydroxide is the highest?
Answer: CsOH

Question 93. Which alkaline earth metal hydroxide is amphoteric?
Answer: Be(OH)₂

Question 94. How will you distinguish between Ba(OH)₂ and Be(OH)₂?
Answer: Be(OH)₂ dissolves in alkali but Ba(OH)₂ does not

Question 95. What are the raw materials used for the manufacture of washing soda by the Solvay process?
Answer: NaCl, CaCO₃ and NH₃

Question 96. What is soda ash?
Answer: . Anhydrous Na₂CO₃

Question 97. Which alkaline earth metal hydroxide and alkali metal carbonate are used for softening of hard water?
Answer: Ca(OH)₂ and Na₂CO₃;

Question 98. What is the formula of Plaster of Paris?
Answer: 2CaSO₄ .H₂O

Question 99. What is dead burnt plaster?
Answer: Anhydrous CaSO₄;

Class 11 S-Block Notes

Question 100. Which alkaline earth metal carbonate is used as an ingredient of chewing gum?
Answer: CaCO₃

Question 101. What is a fusion mixture?
Answer: A mixture of Na₂CO₃ and K₂CO₃

Question 102. What is the most abundant source of sodium chloride?
Answer: Seawater

Question 103. For which of its chief properties is Plaster of Pariswide used?
Answer: On mixing with water, it forms a plastic mass which sets into a hard mass within 5 to 15 minutes;

Question 104. What is used for making blackboard chalks?
Answer: CaCO₃

Question 105. Which compound is generally used for the detection of CO₂ in the laboratory?
Answer: Lime water [Ca(OH)₂]

Question 106. What is the diagonal relationship? Give one example
Answer:

See ‘General Discussion on s -block elements

Question 107. Why is LiCl soluble in organic solvents?
Answer:

As the polarising power of Li⁺ is very high, LiCl is covalent. Hence, it is soluble in an organic solvent

Question 108. Why are fumes seen when barium halides are kepi In open air?
Answer:

In the open air, barium halide undergoes hydrolysis by water vapor (moisture), and fumes of halogen acid (except HF) evolve

Question 109. Which of the alkaline earth metal hydroxides are amphotericin in nature?
Answer:

Beryllium hydroxide [Be(OH)₂] is amphoteric in nature

Question 110. What is hydrolysis?
Answer:

Calcium hydride [CaH₂] is known as hydrolytic

Question 111. Find out the oxidation state of sodium in Na₂O₂.
Answer:

Let the oxidation state of Na be x. The oxidation state of oxygen in the peroxides is -1

2x + 2(-1) = 0; 2x = 2,x= +1.

Question 112. Explain why can alkali and alkaline earth metals not be obtained by chemical reduction methods.
Answer:

Alkali and alkaline earth metals are strong, reducing agents and it is difficult to reduce their oxides or chlorides by any other reducing agent.

Question 113. Which two alkaline earth metals cannot be identified by flame test?
Answer:

Beryllium (Be) and magnesium (Mg) cannot be identified by flame test

Question 114. Beryllium chloride hydrate loses no water over P₄O₁₀ — why?
Answer:

Due to very small size and stronger hydrating tendency of Be²⁺ ion, it is not possible for P₄O₁₀ to abstract water molecules from beryllium chloride hydrate, [Be(H₂O)4]Cl₂;

Question 115. Magnesium occurs in nature largely as MgCO₃but beryllium never occurs as BeCO₃. Explain.
Answer:

The strong polarising power of Be²⁺ ion makes BeCO₃ unstable;

Class 11 S-Block Notes

Question 116. The crystalline salts of alkaline earth metals contain more water of crystallization than the corresponding alkali metal salts—why?
Answer:

Due to their smaller size and higher charge, the alkaline earth metal ions have a greater tendency to coordinate with water molecules as compared to alkali metal ions

Question 117. Out of Na and Mg which one has a higher second ionization enthalpy? Why?
Answer:

The second ionization enthalpy of Na is higher than that of Mg;

Question 118. The chloride of a metal is soluble in an organic solvent. The chloride can be CaCl₂, NaCl, MgCl₂, BeCl₂
Answer: BeCl₂

Question 119. Why does common salt become wet in the rainy season?
Answer:

Due to the absorption of aerial moisture by deliquescent impurities like MgCl₂, CaCl_2, etc;

Question 120. With the help of a drop of an indicator solution, how would you know whether a solution consists of Na₂CO₃ or NaHCO₃?
Answer:

A drop of phenolphthalein will change the colorless solution of Na₂CO₃ to purple but the colorless solution of NaHCO₃ will remain unchanged;

Class 11 Chemistry S Block Elements Fill In The Blanks

Question 1. _____________ has no d-orbital in its valence shell.
Answer: Lithium

Question 2. _____________ reacts with nitrogen to give nitrides.
Answer: Lithium

Question 3. _____________ imparts golden-yellow color to the flame.
Answer: Sodium

Question 4. Potassium, in reaction with dioxygen, produces _____________
Answer: KO₂

Question 5. The outermost electronic configuration of the radioactive alkali metal is _____________
Answer: 7s¹

Question 6. The bicarbonate salt_____________ of does not exist in solid state.
Answer: Lithium

Question 7. Ionic conductance of Li+ ion in aqueous solution is lowest _____________ is highest
Answer: Hydration

Question 8. _____________ion has maximum polarising power
Answer: Li⁺

Question 9. _____________ is the most abundant alkali metal in Earth’s
Answer: Sodium

Question 10. The basicity of the alkali metal hydroxides _____________ down the group.
Answer: Increases

Question 11._____________ is used as a source of oxygen in submarines, space shuttles, and oxygen masks.
Answer: KO₂

Class 11 S-Block Notes

Question 12. The alkali metals combine with mercury to give_____________
Answer: Amalgams

Question 13. The alkali metals exist as_____________ lattices having cordination number 8.
Answer: Body Center cubic

Question 14. Lil _____________ is more soluble than KI in ethanol.
Answer: More

Question 15. K₂CO₃ cannot be produced by the Solvay process because _____________ does not get precipitated in the aqueous solution.
Answer: KHCO₃

Question 16. _____________ and _____________ do not respond to flame test.
Answer: Be,  Mg

Question 17. _____________ ion exhibits maximum does not respond tendency to complexes.
Answer: Be²⁺

Question 18. Only _____________ can displace hydrogen from dilute HNO₃
Answer: Mg

Question 19. Common salt gets wet due to the presence of _____________ as impurity.
Answer: MgCl₂

Question 20. Hydration enthalpy of Mg²⁺ is _____________ Ca²⁺. as than that of
Answer: Greater

Question 21. Hydrolysis of calcium carbide produces _____________
Answer: Acetylene

Question 22. The commercial name of _____________is hydrolith
Answer: CaH₂,

Class 11 S-Block Notes

Question 23. BeCO₃ is stable only in an atmosphere of _____________
Answer: CO₂

Question 24. Lime water is a transparent aqueous solution of _____________
Answer: Ca(OH)₂

Question 25. The second ionisation enthalpy of the alkaline earth metals is _____________ than their first ionization enthalpy.
Answer: Greater

Question 26. _____________ is used to prepare Grignard reagents
Answer: Greater

Question 27. The melting point of the alkaline earth metals is _____________ than that of alkali metals
Answer: Mg

Question 28. Between Ca and Na, _____________ is used to dehydrat alcohols
Answer: Ca

Question 29. Among the alkaline earth metals, _____________ is abundant in the earth’s crust.
Answer: Ca

Question 30. Temperature of the mixture of _____________  -54°C. is most and ice is about
Answer: CaCl²⁺

Class 11 S-Block Notes

Class 11 Chemistry S Block Elements Warm-Up Exercise Question And Answers

Question 1. Name a radioactive alkali metal and write its atomic number.
Answer:  Francium, atomic number =87

Question 2. Mention the similarity shown in the electronic configurations of the alkali metals.
Answer: All alkali metals have similar valence shell electronic configuration of ns¹

Question 3. Why alkali metals are called s -s-block elements?
Answer: Alkali metals are called s -s-block elements because the last electron enters the ns-orbital

Class 11 S-Block Notes

Question 4. Which of the alkali metals exhibits abnormal behavior?
Answer: Lithium

Question 5. Which alkali metal is most abundant in the earth’s crust?
Answer: Sodium (Na)

Question 6. What is trona?
Answer:

Trona ( Na₂ CO₃ , NaHCO₃  .  2H₂O) is an important mineral of sodium

Question 7. Arrange lithium, sodium, and potassium according to their abundance in the earth’s crust.
Answer:

According to their abundance in nature, the elements are arranged as lithium < potassium < sodium.

Question 8. Give an example of a double salt formed by an alkali metal and alkaline earth metal.
Answer:

The double salt formed by an alkali metal and an alkaline earth metal is carnallite (KCl-MgCl₂-6H₂O)

Question 9. What type of crystal lattice is formed by the alkali metals?
Answer:

Alkali metals form a body-centered cubic lattice

Question 10. Arrange LiF, NaF, KF, RbF, and CsF in increasing order of their lattice energies.
Answer:

The increasing order of lattice energies of the given compounds is: CsF < RbF < KF < NaF < LiF

Question 11. Why are alkali metals paramagnetic?
Answer:

Due to the presence of unpaired electrons in the valence shell of alkali metals, they are paramagnetic.

Question 12. Which alkali metal is generally used in photoelectric cells?
Answer:

Cesium is generally used in photoelectric cells

Question 13. Which alkali metals form superoxides when heated in excess air?
Answer:

The alkali metals that form superoxides when they are heated in excess air are potassium (K), rubidium (Rb), and cesium (Cs).

Class 11 S-Block Notes

Question 14. Differentiate between Na₂CO₃ and NaHCO₃
Answer:

When NaHCO₃ is heated, it liberates CO₂ which turns lime-water milky. On the other hand, when Na₂CO₃ is heated, it does not undergo decomposition

Question 15. Arrange MCI, MBr, MF, and MI (where M= alkali metal) according to increasing covalent character.
Answer:

The covalent character of metallic chlorides increases with an increase in the size of the anion. Therefore, the order of the metallic chlorides according to increasing covalent character is MF < MCI < MBr < MI

Question 16. Explain why the peroxides and superoxides of the alkali metals act as strong oxidizing agents.
Answer:

In reaction with water, peroxides produce MOH along with H₂O₂ and superoxides produce MOH and O, along with H₂O₂. H₂O₂ is a strong oxidizing agent. Thus, peroxides and superoxides of the alkali metals act as strong oxidizing agents

M₂O₂ + 2H₂O → 2MOH + H₂O₂

2MO₂ + 2H₂O→ 2MOH + H₂O₂ + O

Question 17. Give a simple test to distinguish between KNO₃ & LiNO₃.
Answer:

Colourless O₂ gas evolves on heating KNO₃. But heating, LiNO₃ dissociates into colorless gas and brown NO₂ gas

Question 18. Explain why a solution of Na₂CO₃ is alkaline in nature whereas a solution of Na₂SO₄ is neutral.
Answer:

For Na₂CO₃, Na₂SO₄ is a salt of strong base (NaOH) and strong acid (H₂SO₄). So, the nature of the solution of Na₂SO₄ is neutral.

Question 19. Among the sulfate salts of lithium, sodium, potassium, and rubidium, which salt does not form double salt?
Answer:

Lithium sulfate (Li₂, SO₄) does not form any double salt

Question 20. What happens when sodium sulfate is fused with charcoal? Give equation.
Answer:

When sodium sulfate is fused with charcoal, it reduces to sodium sulfide, and carbonÿis oxidized to carbon monoxide:

Na₂SO₄ + 4→ Na₂S + 4CO↑

Class 11 S-Block Notes

Question 21. Which alkali metal bicarbonate has no existence in the solid state?
Answer:

The alkali metal bicarbonate which has no existence in a solid state is lithium bicarbonate (LiHCO₃)

Question 22. Mention the property for which lithium is used to separate N2 gas from a gas mixture.
Answer:

The alkali metal bicarbonate which has no existence in a solid state is lithium bicarbonate (LiHCO₃)

Question 23. Give a simple test to distinguish between Li₂CO₃ & Na₂CO₃. 
Answer:

On heating, Li₂CO₃ decomposes to CO₂ which turns lime water milky. On the other hand, Na₂ CO₃ does not decompose on heating

Class 11 S-Block Notes

Question 24. Write down the name of the alkali metal compound which i

1. Effective in the treatment of manic depressive psychosis
2. Used in baking powder

Answer:

1. Li₂CO₃
2. NaHCO₃

Question 25. Name an alkali metal that is used in Lassaigne s test for detection of nitrogen, sulphur and halogen in the organic compounds.
Answer: Sodium

Question 26. Explain why sodium carbonate is used in fire extinguishers.
Answer:

In fire extinguishers, Na₂CO₃ is used as it reacts with dil. H₂SO₄ to produce CO₂

Question 27. Why docs sodium hydrogen carbonate called baking soda?
Answer:

NaHCO₃ is an important ingredient in baking powder. Thus, it is also called baking soda.

Question 28. Why a standard solution of sodium hydroxide (NaOH) cannot be prepared?
Answer:

Sodium hydroxide, being a hygroscopic substance absorbs moisture from the atmosphere. It also absorbs CO₂ from the air and forms Na₂CO₃. Thus, sodium hydroxide cannot be accurately weighed and so a standard solution of NaOH cannotbe prepared.

Question 29. How the group-2 elements are commonly known? Why are they so called?
Answer:

Group-2 elements are commonly known as ‘alkaline earth metals’. These are so-called because their oxides are basic and are found in the earth’s crust.

Class 11 S-Block Notes

Question 30. Write the general electronic configuration of group-2 elements.
Answer:

Electronic configuration: [Inert gas] ns²

Question 31. Why group-2 elements are called s -s-block elements?
Answer:

Group-2 elements are called s -s-block elements because the last electron enters the s -s-orbital

Question 32. Which element of group 2 shows abnormal behavior?
Answer:

Beryllium (Be) shows abnormal behavior

Question 33. Which group-2 element has a slightly different electronic configuration than the rest of the elements?
Answer:

Beryllium has 2 electrons in its penultimate shell while the rest have 8 electrons in their penultimate shells.

Question 34. Explain why the atomic and ionic radii of Mg is less than those of Na and Ca.
Answer:

The electrons of Mg having a higher nuclear charge are more strongly attracted towards the nucleus. Thus, the atomic and ionic radii of Mg are less than Na. Again on moving down the group (from Mg to Ca), the atomic and ionic radii increase due to the addition of new shells and the increasing screening effect joindy overcomes the effect of increasing nuclear charge. Thus, the atomic and ionic radii of Mg is less than Ca

Question 35. Arrange Mg²⁺, Ba²⁺, Sr²⁺, Be²^+, and Cav according to decreasing order of their hydration enthalpies. Explain your answer.
Answer:

The correct order is  → Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺

Class 11 S-Block Notes

Question 36. Alkaline earth metals predominantly form ionic compounds. However, the first member of the group forms covalent compounds—Explain why.
Answer:

Due to relatively high electronegativity, the first member of each of these groups tends to form covalent compounds

Question 37. A white residue is obtained when metallic Mg is burnt in air. This residue when heated with water emits an ammoniacal smell. Explain these observations.
Answer:

On burning metallic magnesium in the air, magnesium oxide (MgO) and magnesium nitride (Mg₃N₂) are formed. So, the white residue obtained is a mixture of MgO & Mg₃N₂.

This residue when heated with water, results in the hydrolysis of Mg₃N₂ which emits an ammoniacal smell.

Question 38. Why calcium is better than sodium in eliminating a small amount of water from alcohol?
Answer:

Both Na and Ca react with water to form the corresponding hydroxide. However, sodium rapidly reacts with alcohol to form sodium ethoxide (NaOC₂H₅) but calcium reacts quite slowly with alcohol

2C₂H₅OH + 2Na → 2C₂H₅ONa + H₂↑

Thus, for eliminating a small amount of water from alcohol, calcium is better than sodium

Question 39. Why Mg-ribbon continues to burn in the presence of SO₂ gas
Answer:

During the combustion of Mg-ribbon, the amount of heat generated leads to the decomposition of sulfur dioxide into sulfur and oxygen. This oxygen is responsible for the continued burning of Mg-ribbon.

2Mg + SO₂→ 2MgO + S

Question 40. Except Be(OH)₂, all other alkaline earth metal hydroxides are basic and their basic strength increases down the group.
Answer:

Be has high ionization enthalpy, for which Be(OH)₂ is amphoteric.

Question 41. Why is Mg(OH)₂ less basic than NaOH?
Answer:

Due to the larger ionic size and lower ionization enthalpy of Na, the Na — OH bond in NaOH is weaker than the Mg — OH bondin Mg(OH)₂. Thus, NaOH is more basic than Mg(OH)₂

Question 42. Sparingly soluble carbonate salts of alkaline earth metals become easily soluble in water in the presence of CO₂. Why?
Answer:

Sparingly soluble carbonate salts of alkaline earth metals are converted into soluble bicarbonate salts in the presence of CO₂ So, these salts become easily soluble in water.

CaCO₃(s) + CO₂(g) + H₂O(l)→ Ca(HCO₃)₂ (aq)

Question 43. The anhydrous chloride salt of which alkaline earth metal is used as a drying agent in the laboratory?
Answer:

Anhydrous chloride salt of calcium metal, i.e., calcium chloride (CaCl₂) is used as a drying agent in laboratory

Class 11 S-Block Notes

Question 44. Why BeCO₃ is kept in an atmosphere of CO₂?
Answer:

BeCO₃ being highly unstable easily decomposes to give off CO₂ when kept in an open atmosphere.

Question 45. Name a bivalent element whose oxide is soluble in excess NaOH solution.
Answer:

Beryllium; BeO + 2NaOH→ Na₂BeO₂ + H₂O

Question 46. Why CaF₂ is considered the most important fluoride salt among all the fluoride salts of alkaline earth metals?
Answer:

CaF₂ is the most important source for the preparation of fluorine

Question 47. Write down some important points of difference between beryllium and magnesium.
Answer:

• Beryllium is harder than magnesium,
• Compounds of Be are largely covalent, whereas most of the compounds of Mg are ionic,
• Be does not exhibit a coordination number of more than 4, while Mg exhibits a coordination number of 6

Question 48. Be usually forms covalent compounds but other elements of group ionic compounds. Why?
Answer:

Be usually forms covalent compounds due to its high ionisation enthalpy and small size. However, due to comparatively low ionization enthalpy and large size, other elements of group-2 form ionic compounds

Question 49. Which compounds of the alkaline earth metals are used as refractory substances?
Answer:

The oxides of alkaline earth metals (MO) have high melting points and so are used as refractory substances

Question 50. Carbonaceous impurities in gypsum and any fuel are avoided during the preparation of Plaster of Paris. Explain.
Answer:

Carbonaceous impurities in gypsum and any fuel are avoided during the preparation of Plaster of Paris because carbon will reduce CaSO₄ to CaS.

CaSO₄ + 4C →  CaS + 4CO ↑