WBCHSE Class 11 Chemistry Classification Of Elements And Periodicity In Properties Notes

Classification Of Elements And Periodicity In Properties Introduction

Element Classification Principles: At present, 118 elements are known to us. It is Almost an impossible task to remember the individual properties of these elements And A larger number of compounds derived from them.

Several attempts were made by former scientists to arrange the elements in a coherent and orderly manner.

After Dalton’s Atomic theory, attempts were made to establish a correlation between the atomic masses of various elements and their properties.

But until a method for the estimation of correct atomic masses of elements was innovated, the work on the proper classification of elements could not make any significant progress.

However, after the atomic masses of elements were correctly determined, the attempts for the classification of elements received particular attention.

The way of arranging similar elements together and separating them from dissimilar elements is called the classification of elements.

Historical Background Of The Classification Of Elements Based On Atomic Weight

Dobereiner’s Law Of Triads

In 11117, German scientist Doberelnor stated that in a group of three chemically similar elements, called a triad, the atomic weight of (the middle element of each triad Is very close to the arithmetic mean of those of the other two elements.

Element Classification Principles

This was called Oohereiner’s law of triads. Some familiar triads, based on lids law, are shown below:

From the table, it is observed that the atomic weight of sodium (Na) is the average of the atomic weights of lithium (Li) and potassium (K) \(\left[\frac{7+39}{2}=23\right]\)

This relationship is only applicable to a limited number of elements and hence fails to classify all the known elements.

However, it cannot be denied that it indicated the existence of an inter-relationship between the properties and atomic weights of elements.

Law of Telluric Screw

In 1862, Chancourtois attempted to classify the elements based on atomic mass. He took a vertical cylinder with 16 equidistant lines drawn on its surface (lines are parallel to the axis of the cylinder).

He drew a spiral line or helix on the surface making an angle of 45° to the axis of the cylinder.

The atomic weights were plotted vertically along the spiral line. He arranged the elements on the helix in order of their increasing atomic weights.

It was observed that in the telluric screw, the elements that differed from each other in atomic weight by 16 or multiples of 16 fell on the same vertical line.

The elements lying on the same vertical line showed nearly the same chemical properties. However, this concept did not attract much attention.

Element Classification Principles

Newlands’ Law Of Octaves

Arranging the known elements in the ascending order of their atomic weights, Newlands, observed (1865) that properties of the eighth element, starting from a given one, is a kind of repetition of the first, like the eighth note in an octave of music. He called this regularity the law of octaves.

Starting from Li, the eighth element is Na and the eighth element following Na is K.

There exists a striking resemblance in properties among these elements. Similarly, F shows similarity with the eighth element Cl following it in properties.

The law of octaves was found to be satisfactory in the case of lighter elements from hydrogen (H) to calcium (Ca). However in the case of heavier elements beyond calcium, it lost its validity and hence, the law was discarded.

Lothar Meyer Arragngement

In 1869, Lothar Meyer, a German scientist, studied the different physical properties of the known elements and plotted a graph of atomic volume (atomic weight divided by density) against the atomic weight of various elements.

He noticed that the elements with similar properties occupied similar positions on the curve.

Based on this observation, Lothar Meyer concluded that the physical properties of the elements are a periodic function of their atomic weights.

Periodic Law

In 1869, Russian chemist, Dmitri Mendeleev, examined the relationship between the atomic weights of the elements and their physical and chemical properties.

From his studies, Mendeleev pointed out that the physical and chemical properties of elements are periodic functions of their atomic weights. This generalisation is called Mendeleev-Lothar Meyer Periodic Law or simply Mendeleev’s Periodic Law.

Element Classification Principles

Mendeleev’s Periodic law: Physical and chemical properties of elements are a periodic function of their atomic weights. This law implied that if the elements are arranged in the order of increasing atomic weights, the physical & chemical properties of the elements change regularly from one member to another and get repeated after a definite interval. This recurrence of properties ofthe elements at definite intervals is called the periodicity of elements.

Periodic classification and periodic properties: Based on the periodic law, the classification of elements according to the increasing atomic weight is called periodic classification. The properties of the elements which are directly or indirectly related to their electronic configurations and show a regular gradation when we descend in a group or move across a period in the periodic table are called periodic properties.

For example—The size of atoms or atomic radii, ionic radii, atomic volume, metallic character, ionisation enthalpy, electron affinity, electronegativity, melting point, boiling point, valency etc.

Radioactivity is not a periodic property of elements: Radioactivity is neither directly nor indirectly related to the electronic configuration of atoms. It depends on the ratio between the number of neutrons and protons present in the atom.

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Mendeleev’s Periodic Table

Based on his periodic law, Mendeleev arranged the then-known elements in the form of a table (consisting of several rows and columns) which is known as Mendeleev’s periodic table.

Mendeleev’s original periodic table (1871) contained only 63 elements known at that time. There were no places for inert gases because these were not discovered at the time of publication of the table.

Mendeleev, however, left several blank places in the table and predicted that there must be some unknown elements which would be discovered in due course of time.

He even predicted their properties based on the properties of the adjacent elements.

Element Classification Principles

Mendeleev’s predictions were proved to be astonishingly correct when these elements were discovered later. Mendeleev’s table, now in use, is a modified version ofthe table originally designed by him. Important features of the modified form of Mendeleev’s periodic table are discussed below.

Periods and Groups: In Mendeleev’s periodic table, the elements were arranged in the increasing order of their atomic weights (but in the modified form these were arranged in increasing order of their atomic numbers) into several horizontal rows.

These horizontal rows were placed one below the other in such a way that chemically similar elements fell in the same vertical column. The horizontal rows are called periods and the vertical columns are called groups or families.

There is a gradual change in the properties of the elements with an increase in atomic mass across a period. However, elements belonging to the same group exhibit close chemical similarities.

In the modern version of Mendeleev’s table, there are seven periods (1 to 7) and nine groups (I to VIII and 0). Gr-0 consists of the inert gases (Mendeleev’s original table did not contain this group)

Main features of Mencleleov’s Periodic Table

The first period contains only 2 elements (II and He). This is called the shortest period.

The second period contains only 8 dements (Li-Bc-B-C-NO-F-Nc), beginning with alkali metal Li and ending with Inert gas No. This Is called the first short period.

The period also contains elements (Na-Mg-Al-Si-P-SCl-Ar), beginning with the alkali metal Na and ending with the Inert gas Ar. This is called the second period.

The elements of these two short periods occur in nature in very large amounts and they typify the properties of all the other members of the group to which they belong. So they are called typical elements.

Element Classification Principles

The fourth period contains 18 elements. It begins with the alkali metal K and ends with the inert gas Kr. This period is called the first long period.

The fourth period contains 10 additional elements than the second and third periods. These 10 elements (Sc to Zn) are called the transition elements. This period consists of two series (the even and the odd series).

The fifth period also contains 18 elements. It begins with the alkali metal Rb and ends with the inert gas Xe. This period is called the second long period. 10 elements from Y to Cd are called transition elements. The fifth period also consists of two series (the even and the odd series).

The sixth period contains 32 elements, and so it is called the longest period. It begins with the alkali metal Cs and ends with the inert gas Rn.

This period contains 10 transition elements (La and Hf to Hg) and 14 lanthanide elements (Ce to Lu ).

These 14 elements are also called rare earth elements because these elements were believed to be present in nature in negligible amounts.

The sixth period also consists of two series (the even and the odd series).

Element Classification Principles

The seventh period may contain a maximum of 32 elements (beginning with Fr), but all the elements have not yet been discovered. Till now 28 elements have been discovered. So it is an incomplete period.

All elements of this period are radioactive. The elements from Francium (Fr) to Uranium (U) are naturally occurring, while the elements beyond uranium are man-made.

In this period, the 14 elements beyond actinium, Ac [i.e., the elements from thorium (Th) to lawrencium (Lr) ] are called actinides, while the elements beyond uranium (U) are called transuranic elements.

WBCHSE Class 11 Chemistry Element Classification Principles

Even and odd series: Elements belonging to each of the 4th, 5th and 6th periods are divided into two series: the even and the odd series. The three even series begin with the alkali metals K, Rb and Cs, while the three odd series begin with the coinage metals Cu, Ag and Au respectively.

Subgroups: Except for the Gr-VHI and Gr-0, each of the other groups (Gr-I to VII) is divided into two subgroups designated as ‘A’ and ‘B! In long periods (4th, 5th and 6th), the elements of the even series areplacedinsubgroup-Aand those ofthe odd series are placed in subgroup-B.

In short periods (2nd and 3rd), elements of Gr-I and Gr-II are placed in subgroup-A, while those of the other groups are placed in subgroup-B. Within the same group, the properties of the elements of subgroups and B are altogether different, except for their valencies.

Element Classification Principles

However, elements of the same subgroup exhibit more or less similar properties.

For example, alkali metals of Gr-IA are closely alike. However, Gr-IA metals differ remarkably from the coinage metals of Gr-IB (Cu, Ag and Au), although they have a common valency of ‘1’.

Additional pieces of information about groups and sub-groups:

Elements of subgroup A are more electropositive than those of subgroup B. For example, Gr-VIIA elements (Mn, Tc, Re) are electropositive, while Gr-VIIB elements (F, Cl, Br and I) are electronegative characters.

Gr-VIII has no subgroups. It contains a total of 9 elements, belonging to periods 4, 5 and 6. These nine elements—Fe, Co, Ni (period-4); Ru, Rh, Pd (period-5) and Os, Ir, Pt (period-6) are arranged in this manner due to similarity in their properties and their, atomic weights are also close to v each other.

Each grip, of three elements, is called Mendeleev’s triad elements. Mendeleev coined the term ‘transitional element’ for these elements.

Element Classification Principles

Gr-0 has no subgroups. It contains inert gases: He, Ne, Ar, Kr, Xe and Rn. These elements are chemically inert and do not exhibit any tendency to combine with other elements. So they are zero-valent elements and placed in Gr-‘0′. This group acts as a bridge between highly electronegative halogens (VIIB) and highly electropositive alkali metals (IA).

Due to their similarity in chemical properties, La and 14 elements from Ce -Lu are placed together in Gr-IIIA of the 6th period. The 14 elements from Ce to Lu are called lanthanoids. For similar reasons, Ac and 14 elements from Th-Lr are placed together inGr-IIIA ofthe 7th period. The 14 elements from Th to Lr are ( called actinoids.

Importance & usefulness of Mendeleev’s periodic table Systematic Study of the elements: Mendeleev, for the first time, arranged a vast number of elements in such a way that the elements with similar chemical properties are placed in the same group.

This made the study of elements quite systematic because if the properties of one element (and its compounds) in a particular group are known, then the properties of the rest of the elements (and their compounds) can be predicted

WBCHSE Class 11 Chemistry Element Classification Principles

Prediction of new elements: Mendeleev left some gaps in the periodic table to accommodate new elements to be discovered in future. he even predicted the properties of those unknown elements based on their positions in the table.

when these elements were discovered, their properties were found to be similar as predicted By Mendeleev. For example, Mendeleev left two vacant places below b and al in gr-3 and one vacant place below.

Element Classification Principles

He named those elements eka-boron, ca-aluminium and ca-silicon respectively as he predicted that the properties of these elements would be similar to that of boron, aluminium & silicon.

In 1075, de Baisbaudron discovered eka-aluminium and named it gallium. in 1079, n. l. Nilson discovered eka-boron and named it scandium. In 1006, Winkler discovered eka-silicon and named it as germanium.

It was observed that these newly discovered elements had properties similar to those already predicted by Mendeleev before their discovery.

Correction of doubtful atomic weights: with the help of Mendeleev’s periodic table, doubtful atomic weights of some elements are rectified.

For example, Be was assigned an atomic weight of 13.5 based on its equivalent weight (4.5) and valency (wrongly taken as ‘3’ because Be had certain similarities with trivalent metal Al).

With an atomic weight of 13.5, Be should be placed between carbon (At. weight 12) and nitrogen (At. weight 14).

Element Classification Principles

But no vacant place was available In between C and N. Mendeleev asserted that Be must be bivalent because of its similarity with Mg, Ca etc. Thus he corrected its atomic weight as 4.5 X 2 = 9.0.

Defects Of Mendeleev’s Periodic Table

Discrepancy or anomaly in periodicity: Mendeleev arranged the elements in increasing order of their atomic weights. But he violated this principle in certain cases to give appropriate positions to some elements based on their properties i.e., he laid more emphasis on the properties of those elements rather than their atomic weights.

In the following four pairs of elements, elements with higher atomic weight have been placed before elements with lower atomic weight.

Position of hydrogen: Controversial position of hydrogen in the periodic table also hints at discrepancies within the table. Like the alkali metals ofGr-IA, it exhibits univalency, high reactivity, electropositive character, strong affinity for non-metals and reducing character.

On the other hand, like the halogens of Gr- VIIB, it has atomicity, high ionisation energy, non-metallic character, existence in the gaseous state at normal temperature and pressure ability to combine with milk-forming hydrides (e.g., Nall).

Element Classification Principles

Since hydrogen exhibits similarities as well as dissimilarities with both the alkali metals and the halogens, the placement of hydrogen in any one of these two groups will naturally create difficulties. So it is desirable to fix a separate position for hydrogen in Mendeleev’s periodic table.

Placement of similar elements in different groups and dissimilar elements in the same group: in some cases, elements with almost similar properties have been placed in different groups.

Example: Cu and Hg resemble in properties but Cu is Gr-IB while Hg has been placed in Gr-IIB. Likewise, elements like Ba (Gr-IIA) and Pb (Gr-IVB) have been placed in different groups.

Again, some elements with dissimilar properties have been placed in the same group.

WBCHSE Class 11 Chemistry Element Classification Principles

Example: Highly reactive alkali metals such as Li, Na, K etc., have been placed together with almost inactive coinage metals such as Cu, Ag and Au in Gr-I. Likewise, Mn, Te
and Re having no similarity with F, Cl, Br etc. have been placed together

Lack of separate positions for Gr-VIII elements: No proper place has been allotted to nine elements belonging to Gr-VIII although they have many similarities in properties.

These are arranged in three triads, one in each of the 4th (Fe, Co, Ni ), 5th(Ru, Rh, Pb) and (Os, Ir, Pt)periods

Lack of suitable positions for Lanthanoids and Actinoids: The 14 elements following La from Ce to Lu (lanthanoids) and the 14 elements following Ac from Th to Lr (actinoids) have not been allotted separate positions in the main skeleton of the periodic table.

They have been placed in two separate rows at the bottom of the table. Besides, the number of elements in the lanthanoid and actinoid series cannot be determined from Mandeleev’s periodic table.

Position of isotopes: Isotopes of an element have different atomic weights. So they should be placed at different positions in the periodic table. However, all the isotopes of any specific element are placed in a single position (i.e., same period and same group)in Mendeleev’s periodic table.

Moseley’s experiment: Atomic number determines the fundamental property of an element.

In 1913, Moseley measured the frequencies of X-rays emitted by different metals when bombarded with high-speed electrons.

He observed that the frequencies ofthe prominent X-rays emitted by different metals were different but for each metal, there was a fixed value.

Element Classification Principles

He observed further that the square root of the frequency (v) of the X-rays emitted by a metal was proportional to the atomic number but not to the atomic mass of the metal, Le., Jv = a(Z- b) where ‘a’ is the proportionality constant and is a constant for all the lines in a given series of X-rays.

Thus a plot of Tv vs Z gave a straight line but a plot of Jv vs atomic mass does not bear such alinear relationship. This led Moseley to conclude that atomic number was a better fundamental property of an element than atomic mass.

Modification Of Mendeleev’s Periodic Law

Mendeleev regarded atomic weight as the fundamental property of an element and so he considered atomic weight as the basis of periodic classification of elements.

But Moseley, from his experimental results, showed clearly that atomic number is a better fundamental property of an element than its atomic weight.

This led Moseley to suggest that atomic number (Z) should be the basis of the classification of elements. This gave birth to a new periodic law known as the modem periodic law.

Element Classification Principles

Modern Periodic Law: The physical and chemical properties ofthe elements are aperiodic functions of their atomic numbers.

This implies that, if elements are arranged in order of increasing atomic numbers, the elements with similar chemical properties are repeated after certain regular intervals.

Rectification of the discrepancy in periodicity with the help of modern periodic law:

The original periodic law, based on atomic weight, was violated in the case of four pairs of elements [(Ar, K), (Co, Ni), (Te, I), (Th, Pa)].

In each pair, an element with a higher atomic weight is placed before the element having a lower atomic weight.

Element Classification Principles

In the modern periodic table (based on the atomic number), this discrepancy disappears because the atomic numbers of K, Ni, I and Pa are greater than those of Ar, Co, Te and Th respectively.

Placement of all the isotopes of any specific element in the same position of the periodic table is quite justified as the isotopes of elements have the same atomicnumber although they have differentatomic weights.

Theoretical justification of modern periodic law: Only nuclear electrons (or more specifically valence shell electrons) take part in chemical reactions, while the atomic nucleus remains unaffected.

So it is understandable that the properties of the elements will depend upon their atomic numbers (equal to the number of electrons) rather than their atomic weight or mass numbers (equal to the total number of protons and neutrons).

Periodicity of elements

The periodic repetition of elements having similar properties after certain regular intervals when the elements are arranged in the increasing order of their atomic numbers is called periodicity.

Cause of periodicity: According to modern periodicals, there is a repetition of properties of the elements after certain regular intervals when they are arranged in order of their increasing atomic numbers.

Element Classification Principles

Again from a close study of electronic configurations of various elements, it is observed that with successive increases in atomic number, there occurs a repetition of similar outermost shell electronic configuration (valence shell electronic configuration) after certain regular intervals.

By correlating these two observations, it can be concluded that periodicity in properties is due to the recurrence of similar valence shell electronic configuration after certain regular intervals when the elements are arranged in order of increasing atomic numbers.

This can be illustrated by the following examples—

Elements of Gr-IA have outermost electronic configuration ns1 (where n = outermost principal quantum number).

These elements exhibit similar chemical properties due to their similarity in the valence shell electronic configuration.

Elements of Gr-VIIB have outermost electronic configuration ns2np5.

All the halogens exhibit similar chemical properties due to their similarity in valence shell electronic configuration.

WBCHSE Class 11 Chemistry Element Classification Principles

Inert gases belonging to group possess similar chemical properties because they have similar valence shell electronic configurations (ns2np6).

It should be noted that properties of elements get repeated only after intervals of 2, 8,18 or 32 in the atomic numbers of the elements because similar electronic configurations recur only after such intervals.

Element Classification Principles

The numbers 2, 8, 18 and 32 are called magic numbers. These numbers are very useful in locating elements with similar properties

Moderntableor Long Form Of Periodic Table Bohr’s Table:

This is an improved form of the periodic table based on modern periodic law. It is also called Bohr’s table since it follows Bohr’s scheme for the classification of the element based on the outermost electronic configuration governed by the Aufbau principle. It consists of periods and 18 groups.

Structural Features Of Long Form Of Periodic Table

Description of periods: Like Mendeleev’s modified table, it also consists ofsevenperiodswhich are numbered from as1 to 7 from top to bottom.

The period number is equal to the value of i.e., the principal quantum number corresponding to the outermost shell of the atoms of the elements belonging to that period.

Each period begins with the filling of electrons in a new energy level. Several elements in each period are twice the total number of atomic orbitals available in the energy level that are being filled.

First period: This period begins with the filling of the first energy level (n = 1). Since the first shell has only one orbital [i.e., Is), which can accommodate a maximum of two electrons, there can be only two elements in the first period. These are hydrogen (Is1) and helium (Is2).

Second period: It starts with the filling of the second energy level (n = 2). Since the second shell contains four orbitals (one 2s and three 2p), it can accommodate a maximum of (2×4) = 8 electrons. So, there are eight elements in the second period.

Element Classification Principles

It begins with lithium (Li) in which 1 electron enters the 2s -orbital (3Li: 2s1) and ends up with neon (Ne) in which the second shell gets filled (10Ne: 2sz2p6).

Third period: The third period begins with the filling of the third energy level (n = 3). This energy level contains nine orbitals (one 3s, three 3p and five 3d).

According to the Aufbau principle, 3d -orbitals will be filled up only after filling the 4s -orbital.

Consequently, the third period involves filling only four orbitals (one 3s and 3p ) which can accommodate a maximum of (2 x 4) = 8 electrons. So, there are 8 elements in the third period from (Nas1) to 18Ar(3s23p6).

Fourth period: This period corresponds to the filling of the fourth energy level (n = 4). Out of 4s,4p,4d and 4f-orbitals belonging to this shell, filling of 4d -and 4f-orbitals does not occur in this period since their energies are higher than that of even 5s -orbital.

It must however be remembered that, after filling 4s -orbital, filling of five 3d -orbitals begins since energy yf -orbital is greater than that of 4s orbital but less than it of 4p -orbital.

So the fourth period involves filling of or 9 Orbitals (one 4s, five 3d and three 4p), which can accommodate (2×9) = 18 electrons.

Therefore, the fourth period contains 18 elements from potassium (19K: 4s1) to krypton (36Kr: 4sz3d104p6).

Element Classification Principles

This period contains 10 elements more than the third period corresponding to filling off 3d -orbitals. These 10 elements [2iSc(3d14s2) to 30Zn(3d104s2)] are called the first series of transition elements.

Fifth period: The fifth period corresponds to the filling of electrons in the fifth energy level (n = 5). Like the fourth period, it also accommodates 18 elements since only nine orbitals (one 5s, five 4d and three 5p) are available for filling with electrons.

It starts with rubidium in which one electron enters 5s -orbital (37Rb: 5s1) and ends up with xenon in which the filling of 5p -orbital is complete (54Xe: 5s24d105p6). 10 elements from 39Y(5s24d1) to 48Cd(5s24d10) corresponding to filling of five 4dorbitals are called second series oftransition elements.

Sixth period: The sixth period corresponds to the filling of electrons in the sixth energy level (n = 6).

This period involves the filling of sixteen orbitals (one 6s, seven4f, five 5d and three 6p) which can accommodate a maximum of (2 X 16) = 32 electrons. So there are 32 elements in the sixth period.

It begins with caesium (Cs) in which one electron enters 6s -orbital (55Cs: 6s1) and ends up with radon in which filling of 6p -orbital is complete (86Rn: 4f145d106s26p6).

Filling up of 4f-orbitals begins with cerium(58Ce) and ends with lutetium (71Lu). These 14 elements constitute the first inner-transition series, also called lanthanoids or rare earth elements.

Element Classification Principles

These are separated from the main frame periodic table and are placed in a horizontal row at the bottom of the table Again, 10 elements lanthanum (57La), hafnium (72Hf) to mercury (80Hg), corresponding to successive filling of10 5d -orbitals, constitute the third transition series.

Seventh period: This period corresponds to the filling of electrons in the seventh energy level (n = 7). Like the sixth period, it is expected to accommodate 32 elements corresponding to the filling of 16 orbitals (one 7s, seven5f, five 6d and three 7p ).

However, at present this period is incomplete consisting of 28 elements. The last element of this period will have an atomic number of of118 and will position theinert gas family

In this period, after filling of 7s -orbital [87Fr: 7s1 and 8S1Ra: 7s2 ], the next two electrons enter the 6rf-orbital (this is against the Aufbau principle) corresponding to the elements 8gAc and goTh.

Thereafter, the filling up of 57- orbital begins with giPa and ends from Th to Lr are commonly called actinoids, which constitute the second Inner-transition series, Although Th does not contain any electron.

‘if orbital, it is considered to be a member of the actinoid series. Like lanthanoids, 14 members of the actinoid series are placed separately in a horizontal row at the hotter of the periodic table.

Description of groups: Each of the 18 groups in the long form of the periodic table consists of many elements whose atoms have similar electronic configurations ofthe outermost shell (valence shell).

The members of each group exhibit similar properties. Successive members in a group are separated by magic numbers of either 8 18 or 32.

Element Classification Principles

According to the recommendation of IUPAC (1988), the groups are numbered from 1 to 18. Designations of these groups in different systems are presented in the following table-

WBCHSE Class 11 Chemistry Element Classification Principles

Element Classification Principles

The superiority of the long form of the periodic table over Mendeleev’s periodic table

1. In the long form of the periodic table, it is easy to remember and reproduce all the elements more easily in a sequence of atomic numbers.
2. it relates the positions of the elements in the table to their electronic configurations more dearly.
3. Gradual change In properties along the periods or similarity in properties along the groups can be interpreted by considering electronic configurations of the elements.
4. For example, elements of the same group exhibit marked similarities due to similar outer electronic configurations.
5. Splitting the periodic table into s-,p-, cl- and f-blocks has made the study of the elements easier.
6. The maximum capacity of each period to accommodate a specific number of elements is related to the capacity of different electronic shells to accommodate the maximum number of electrons.
7. Due to the elimination of sub-groups, dissimilar elements do not fall In the same group. Each vertical column (group) accommodates only those elements which have similar outer electronic configurations, thereby, showing similar properties.
8. Group-VIIl elements (involving triads) of Mendeleev’s table, have been provided separate positions in groups-8, 9 and 10.
9. Elements belonging to 1, 2, and 13-17 groups are classified as representative elements, while those belonging to 3-12 groups are classified as transition elements.
10. Elements are further classified as active metals (belonging to groups 1 and 2), heavy metals (belonging to groups 3-12) and non-metals (belonging to groups 13-18).
11. Transition elements of the 4th, 5th, 6th and 7th periods are assigned appropriate positions in this periodic table.
12. The completion of each period with an inert gas element is more logical. In a period as the atomic number increases, the quantum shells are gradually filled up until an inert gas configuration is achieved at group 18.
13. It thus eliminates the even and odd series belonging to the periods 4, 5 and 6 of Defects of the long form of the periodic table

If the Position of hydrogen: The position ofhydrogen is not settled. It can be placed along with alkali metals in group 1 or with halogens in group 17, as it resembles the alkali metals as well as the halogens.

Position of helium: Based on electronic configuration, He (Is2) should be placed in group 2. But, it is placed in group 18 along with the p -block elements. No other p -block element has the electronic configuration ofthe type ns2.

Element Classification Principles

Position of lanthanoids and actinoids: Lanthanoids and actinoids have not been accommodated in the main frame of the periodic table.

Position of isotopes: Isotopeshavenotgot separate places.

Properties of isotopes of heavier elements are more or less the same, but isotopes of lighter elements differ drastically in their physical, kinetic and thermodynamic properties.

So it is not desirable to place the isotopes in the same position. Despite these limitations, the long form of the periodic table, based on electronic configurations, is much more scientific and thus finds extensive use.

Classification Of Elements Into Different Blocks

Amount In tin long form of the periodic table it has been divided Into four blocks viz, s -block, p -block, d -block and f- block.

It Is done based on the nature of atomic orbitals into which the Inst electron (the differentiating electron) gets accommodated.

Elements of s and p -blocks except Inert gases, are called representative elements, and d -block elements, on the other hand, are called transition elements.

S -block elements

Elements In which the last electron enters the -subshell of their outermost energy level (n) are called s -block elements.

Since s -subshell can accommodate a maximum of 2 electrons, only two groups are included in this block.

Elements of group-1 (alkali metals) and group-2 (alkaline earth metals) which have outermost electronic configurations rui1 and ns2 respectively constitute the s -block. This block is situated at the extreme left portion of the periodic table.

Outermost electronic configuration of t-block elements: ns1.2 Inert gas element, helium (He, Is2) Is also considered as an s -block element.

Characteristics of s-block elements:

In the case of these elements, all shells except the outermost one, are filled with electrons.

Except for H and, all other elements of this block are metals. Because of their low ionisation potential, these metals are very reactive and do not occur freely in nature.

Element Classification Principles

All the metals of this block are good reducing agents because of the value of ionisation potential.

They are good conductors of heat and electricity.

WBCHSE Class 11 Chemistry Element Classification Principles

They are soft metals. They have low melting points, boiling points and low densities as compared to the adjacent transition elements.

Cations of group-IA and group-DA elements are diamagnetic and colourless since their orbitals do not contain odd electrons.

Except for Be and Mg, -block elements impart specific colour to the flame (flame test).

Salts of these elements except dichromate, permanganate arid chromate, are colourless.

Compounds of these elements are mainly ionic (only Li and Be can form covalent compounds in many cases).

They form stable oxides with oxygen (Na1, CaO), produce chlorides with chlorine (NaCl, CaCl2) and also form salt-like hydrides (NaH, KH, CaH2) with hydrogen.

Hydroxides of these elements [except Ca(OH)2, Mg(OH)2 and Be(OH)2] are soluble in water at ordinary temperature.

Element Classification Principles

The non-luminous flame of the Bunsen burner is rich in electrons. During the flame test, metal ions are converted into short-lived neutral atoms by accepting electrons from the flame.

Valence electrons ofthese neutral atoms absorb energy from the flame and get promoted to higher energy levels.

When the electrons return to lower energy levels, the absorbed energy is emitted in the form of radiation of different wavelengths in the visible range and as a consequence, different colours, depending.

The wavelengths of emitted light radiations are Imparted to the flame. For instance, the generation of golden-yellow flame during the flame test with sodium salt is due to the transition of one electron of Na -atom from 3s -orbital to 3p -orbital and its return to 3sorbltal after a very short interval.

The ionisation potentials of Be and Mg are sufficiently high because of their smaller size. So, their electrons cannot be excited to higher energy levels by absorbing energy from the flame. As a result, they fail to respond to the flame test.

P -block elements

Elements in which the last electron enters p -subshell of their outermost energy level (n) are called p-block elements.

Since p -subshell can accommodate a maximum of six electrons, 6 groups are included in this block.

Element Classification Principles

Elements of group-13, 14, 15, 16, 17 and 18 (excluding helium) having the outermost electronic configurations: ns2np1, ns2np2, ns2np3, ns2np4, ns2np5 and ns2np6 respectively, constitute the p -block. This block is situated at the extreme right portion ofthe periodic table

The elements of group 18 are balled noble gases or inert gases. They have the shell electronic configuration ns2np6 in the outermost shell. Group-17 elements are called halogens (salt producers)’ and group-16 elements are called chalcogens (ore-forming).

These two groups of elements have high electron-gain enthalpies (high negative values of A’H) and hence readily accept one or two electrons respectively to attain the stable noble gas configuration thereby forming negative and negative anions respectively.

 

WBCHSE Class 11 Chemistry Element Classification Principles

The elements of s- and p -blocks taken together are called representative normal or main group elements. Outermost electronic config. of-flock elements: ns2np1’6

Based on electronic configuration, helium (Is2) should not be considered as a p -block element, but from the standpoint of its chemical inertness (owing to the presence of a filled valence shell) it is justified to place group-18 along with other noble gas elements.

Characteristics of block elements:

Ionisation enthalpies of p -block elements are higher as compared to those of -block elements.

Most of the p -block elements are non-metals, some are metals and a few others are metalloids and inert gases.

Element Classification Principles

The metallic character increases from top to bottom within a group non-metallic character increases from left to right along a period. Hence, metals exist at the bottom ofthe left side ofthe p -block whereas non-metals lie at the top of the right ofthe p -block. Metalloids (B, Si, Ge, As, Sb ) stand midway between them.

The oxidising character of p -block elements increases from left to right in a period and reducing character increases from top to bottom in a group.

Most of them form covalent compounds, although ionic character increases continually down the group.

Elements of this block are non-conductors of heat and electricity, except metals and graphite.

Elements of this block are mostly electronegative.

Some of them exhibit variable oxidation states or valence states. Oxidation states may be both positive and negative.

Non-metallic elements of this block form acidic oxides.

They can form both coloured and colourless compounds.

Element Classification Principles

4th, 5th and 6th-period elements can form complex compounds by coordinate covalency due to the presence of vacant d-orbitals. oa aii.

Some of the p -block elements Fe.g.-Q Si, P, S, B, Ge, Sn, As etc.) show the phenomenon of allotropy.

Carnation property is shown by some- block elements (e.g., C, Si, Ge, N, S etc.)

d-block elements (Transition elements) Elements in which the last electron enters d -the subshell of their penultimate shell (i.e., the second from the outermost) are called d -block elements, d -subshell can accommodate a maximum of 10 electrons.

Therefore, ten groups are included in d -the block. Elements of group-3 [(n-1)d1ns2], 4, 5, 6, 7, 8, 9, 10, 11 and 12 [(n- l)d10ns2] constitute the d -block.

Atoms of the elements belonging to these groups usually contain or 2 (sometimes zero) electrons in the s -s-orbital of their outermost shell (i.e., n -th shell), while the differentiating electrons are being progressively filled in, one at a time, in the d -subshell of their penultimate shell [i.e., (n- 1) -th shell].

Electronic configuration of outer shell: (n-1) d1-10ns1-2

Element Classification Principles

This block is situated in between s -and p -blocks. In fact, d -block elements form a bridge between the chemically active metals of groups 2 on one side and the less reactive elements of groups 13 and 14 on the other side.

Hence, d -block elements are called transition elements These elements have been divided into four series called the first, second, third and fourth transition series.

First transition series or 3d-series: First transition series consists of 10 elements, belonging to the 4th period, from scandium (21Sc) to zinc (3QZn) in which 3d -orbitals are being progressively filled in. Zn is not a transition element.

Second transition series or 4d-series: Second transition series also consists of 10 elements, belonging to the 5th period, from yttrium (3gY) to cadmium (48Cd) in which 4d -orbitals are being progressively filled in. Cd is not a transition element.

Third transition series or 5d-series: Third transition series also consists of 10 elements, belonging to the 6th period. These are lanthanum (57La) and elements from hafnium (72Hf) to mercury (80Hg). In all these elements, 5d orbitals are being successively filled in. Hg is not a transition element.

Element Classification Principles

Fourth transition series (6d-series: The fourth transition series is formed from a part of the seventh period and it contains 10 elements. These, are actinium (89Ac) and elements from rutherfordium (104Rf) to ununbium (112Uub), in which 6d -orbitals are being progressively filled in.

All d -block elements are not transition elements. Only those d -block elements in which atoms in their ground state or any stable oxidation state contain incompletely filled subshells are considered transition elements.

WBCHSE Class 11 Chemistry Element Classification Principles

Characteristics of d-block elements:

1. All d -block elements are metal. Their ionisation potential lies mid-way between those of s and p -block elements.
2. Elements of the 5d series (especially Pt. Au and Hg) are inert under ordinary conditions. Thus, they are known as noble metals.
3. Elements of this block exhibit variable oxidation states and valencies because ofthe presence of partially filled d orbitals in their atoms, ns -electrons and different numbers of(n-l)d electrons participate in bonding at the time of reaction with atoms of other elements.
4. They are solids (except Hg), hard and have high melting and boiling points.
5. They can form both ionic and covalent compounds.
6. They exhibit paramagnetic character due to the presence of one or more unpaired electrons in their atoms or ions (exception-Sc3+, Ti4+, Zn2+, and Cu+ which do not contain odd electrons and are diamagnetic). Fe and Co can be converted into magnets and hence, they are ferromagnetic.
7. They frequently form coloured ions in solids or solutions. With the change in their oxidation numbers, there also occurs a change in the colour ofthe formed ions.
8. d -block elements exhibit a very distinctive property of forming coloured coordination complexes.
9. This tendency may be ascribed to the small size of the atom or ion, a high nuclear charge of the ion and the presence of an incomplete d -d-orbital, capable of accepting electrons from the ligands.
10. They are less electropositive than s -block elements but more electropositive than p -block elements.
11. Several transition metals such as Cr, Mn, Fe, Co, Ni, Cu etc., and their compounds are used as catalysts.
12. Many transition metals form alloys.

F-block elements (Inner-transition elements)

Elements in which the differentiating electron (i.e., the last electron) enters the f-subshell of their antepenultimate shell (i.e., the 3rd from the outermost) are called f-block elements.

All the F-block elements belong to group 3 (3B) of the periodic table. In these elements, s -orbital last shell (n) is filled, d -subshell of the penultimate shell [i.e., (n- 1) th shell] contains 0 or 1 electron, while f-subshell of the antepenultimate shell [i.e., (n-2)th shell] gets progressively filled in.

General electronic config.: (n-2)f1-14(n-l)d0-1ns2

Lanthanide series or 4f-Series: The first series follows lanthanum (La) in the 6th period and consists of 14 elements from cerium (58Ce) to lutetium (71Lu).

These 14 elements are collectively called lanthanoids because they closely resemble lanthanum in their properties.

Element Classification Principles

These are also called rare-earth elements since most of these elements occur in very small amounts in the earth’s crust.

Actinoid series or 5f-series: The second series follows actinium (sgAc) in the 7th period and consists of 14 elements from thorium (goTh) to lawrencium (103Lr).

These 14 elements are collectively called actinoids because they closely resemble actinium in their properties.

All the actinoids are radioactive elements. 4fand 5f-series of elements are also called inner-transition elements because they form transition series within the transition elements of d -block.

Characteristics of f-block elements:

1. They are all heavy metals.
2. They exhibit variable valency. +3 oxidation state is most common. Few elements are found to occur in +2 and +4 oxidation states.
3. Some members exhibit paramagnetism due to the presence of odd electrons.
4. They form complex compounds, most of which are coloured.
5. They have high densities.
6. They generally have high melting and boiling points.
7. Within each series, the properties of the elements are quite similar. It is very difficult to separate them from a mixture.
8. Actinoids are radioactive. The first three members (Th, Pa, U ) occur in nature, while the others are man-made. The elements after uranium are called transuranic elements.

Stair-step diagonal

The right side of the long form of the periodic table is composed of p -block elements belonging to groups 13 (3A), 14(4A), 15(5A), 16(6A), 17(7A) and 18 (WlA or 0).

This segment includes four types of elements viz., metals, nonmetals, metalloids and inert gases.

Element Classification Principles

There is no sharp line of demarcation to classify the metals and non-metals, but the zig¬ zag diagonal line (looking like stair-steps) running across the periodic table from boron (B) to astatine (At) is considered as a separation between the metals and non-metals.

WBCHSE Class 11 Chemistry Element Classification Principles

This line is called the stair-step diagonal. The elements B, Si, Ge, As, Sb and Te bordering this line’ -aii d- running diagonally across the periodic table are 8 known as metalloids (which exhibit properties that are characteristics of both metals and non-metals).

The elements (except A1 ) lying between the stair-step diagonal line and the d -block elements are referred to as post-transition elements.

Classification Of Elements based on Outer Electronic Configurations

Based on electronic configurations of the ultimate and penultimate shell of the atoms, Bohr divided the elements into four classes viz., gas elements,

1. Representative elements,
2. Transition elements and
3. Inner-transition elements.

Inert gas elements

S and p -subshells of the outermost shell of the elements of this class are filled.

Except He (electronic configuration: Is2), all other inert gas elements have the valence shell electronic configuration: ns2np6.

All these elements are stable and chemically inert as their outermost shells contain octets of electrons.

Element Classification Principles

They do not normally participate in chemical reactions because the gain or loss of electrons by their atoms would disturb their stability. So, they are called inert gas elements.

Their valency being zero, they find a place in group ‘0’ or ’18 These elements act as a bridge between highly electropositive alkali metals and strongly electronegative halogens.

Representative Elements

Elements present in s – and p -blocks (except group-1) of the periodic table are known as representative elements. The electronic configuration of the outermost shell of these elements varies from ns1 to ns2np5. These consist of some metals, all non-metals and metalloids.

The name ‘representative’ has been assigned to the elements because of their frequent occurrence nature and because they typify the properties of all other members of the group to which they belong.

All the elements of groups, IIA and from 3A to VILA are included in this class.

These elements are very reactive Chemical reactivity of these elements can be ascribed to the ability of their atoms to attain inert gas electronic configuration (ns2np6 or Is2) either by gaining or losing electron(s) or by sharing one or more electron pairs with other atoms. These elements are also known as typical elements.

Transition Elements

Elements of this class are characterised by the presence of atoms in which the inner d -subshell is not filled. According to the modified definition, the elements in which atoms in their ground state or any stable oxidation state contain incompletely filled d -subshell are known as transition elements. Atoms of the elements in this class have the general electronic configuration: (n-I)dl-10 ns1-2.

Element Classification Principles

Cu, Ag and Au, despite having filled d orbitals, are regarded as transition elements. This is because at least in one stable oxidation state of these elements, d subshell remains incompletely filled.

There are four transition series corresponding to the filling of 3d,4d,5d and 6d orbitals These four series belong to the 4th. 5th. 6th and the period of the periodic table.

Each series begins with a member ofthe group-3 and ends with a member of the group-12.

Characteristics:

1. All transition elements are metallic.
2. They have more than one oxidation state or valency.
3. Their ions are coloured.
4. They form complex compounds.
5. Elements of group-12 (11B) (Zn, Cd, Hg) are not considered astransition elementsbecausetheyhaveno partially filled d -orbitals in any of their oxidation states.
6. Moreover, they do not form stable complexes and do not show characteristic colour and paramagnetism.
7. However, their tendency to form complex is much greater than that of the representative elements.
8. They exhibit properties of both transition representative elements.

Differences between typical and transition elements:

1. During the building up of an atom of a typical element by the filling of electrons in its various orbitals, the last electron goes to s -or p -orbital of the outermost shell (n).
2. However, in the case of transition elements, the last electron enters the inner d -d-orbital of(n- 1) th shell.
3. For the representative elements, atomic volume or radius decreases but ionisation enthalpy and electro negativity go on increasing with the increase in atomic number across a period.
4. In the case of the transition elements, as the last electron enters the inner (n- l)d -orbital, the extent of change is relatively small.
5. Most of the representative elements exhibit only one valency. Some elements, of course, show more than one valency.
6. But transition elements show 2 or more valencies through the participation of inner d -d-orbital electrons
7. In the case of representative elements, the tendency to form complex compounds is almost negligible while transition elements are found to show a strong tendency to produce complex compounds due to the presence of incompletely filled d -d-orbital.
8. Compounds formed by representative elements are, in general, colourless but the compounds of transition elements are mostly coloured.
9. Due to the absence of odd electron(s), compounds formed by representative elements are diamagnetic while transition metal compounds, because of the presence of odd electrons, are paramagnetic.
10. Many of the transition metals and their compounds act as catalysts in chemical reactions. Such a tendency is seldom observed in the case of representative elements.

Inner-transition elements

Elements of this class are also transition elements, although they may be distinguished from the regular transition series by their electronic configurations.

Atoms of these elements not only contain incompletely filled d -subshell [(n-l)d] but also contain incompletely filled /-subshell [(n- 2)/].

These elements comprise a transition series within a transition series and hence, they are called Inner-transition elements.

Element Classification Principles

The two series of inner-transition elements are O lanthanoids (rare earth elements) and actinoids.

In the case of 14 elements i.e., cerium (Cel to lutetium (71Lu) following lanthanum (57La), 4/- and 5d subshells remain incompletely filled. These are called lanthanoids. Their general electronic configuration is:

4f1- 14 5(io- 1 6sz Wlth increase in atomic number (58-71). the differentiating electrons of these elements enter the 4f- subshell, despite the presence of a partially filled 5d -subshell. The total electron-accommodating capacity of /-subshell Is 14.

So the number of lanthanoids is also 14. Likewise, 14 elements after actinium (89Ac), from thorium (90Th) to lawrenclum (103Lr) are called actinoids. Their general electronic configuration is 5f1’14 6d01 7s2. With the increase in atomic number (90-103), the differentiating electrons enter the 5f-subshell, despite the presence of an incompletely filled 6d -subshell. Hence, like the lanthanoids, the number of actinoids Is also 14.

Lanthanoid contraction

In the case of lanthanoids (58Ce – 71Lu), it is observed that with an increase in atomic numbers, atomic and ionic size (M3+) go on decreasing, although the decrease in Ionic radii is much more regular than that of atomic radii.

This decrease in atomic and ionic radii with an increase in atomic number in the case of lanthanoids, is known as lanthanoid contraction.

Cause Of lanthanoid contraction: The general electronic configuration of lanthanoids is: 4/’“ 5rf01 (is2. The differentiating electrons of these elements enter the 4f-subshell.

Now due to their diffused shape, f-orbitals have a very poor shielding effect. Thus with the gradual addition of the f- electrons, the atomic number increases by one unit while the shielding effect does not increase appreciably; i.c., there is a gradual increase in the effective nuclear charge acting on the outermost electrons.

Consequently, the attraction of the nucleus for the electrons in the outermost shell increases, causing the electron cloud to shrink although it’s magnitude is small. Thus, there is a gradual shrinkage in the atomic and ionic radii with an increase in atomic number.

Element Classification Principles

Precisely speaking, f-orbitals are too diffused to screen the outermost electrons effectively against the attractive force of the nucleus. Thus, there is a slow contraction in atomic and ionic radii (lanthanoid contraction).

In the same way, the d -contraction due to the accommodation of die electrons in (n- 1) d -subshell in the transition series can be interpreted. But d -orbitals are more effective in screening compared to tire f-orbitals. So this effect is less pronounced in the case of transition elements.

Element Classification Principles

Classification of elements as metals, non-metals and metalloids

All the known elements can be divided into three classes— metals, non-metals and metalloids based on their properties.

Metals: About 78% of the known elements are metals. They appear mainly on the left side and central portion of the long form ofthe periodic table.

Examples are—

1. Alkali metals,
2. Alkaline earth metals,
3. D -block elements,
4. F-block elements,
5. Higher members of p -block elements.

Metals have the following characteristics—

1. They are solids at room temperature. Mercury is an exception, which is a liquid at ordinary temperature.
2. Gallium (melting point 30°C) and caesium (melting point 29°C) are also liquids above 30°C.
3. They usually have high melting and boiling points.
4. They are good conductors of heat and electricity.
5. They are malleable (can be flattened into thin sheets) and ductile (can be drawn out into wires).

Non-metals: There are only about 20 non-metals discovered so far. They are located towards the top right-hand side of the periodic table. Hydrogen and some p-block elements are non-metals.

• Six of the non-metals (C, B, P, S, Se and I) are solid.
• Bromine is the only liquid non-metal.
• The remaining non-metals (N, O, F, Cl, H and inert gases) are gases.
• Non-metalshavelowmelting and boiling points (boron and carbon are exceptions).
• They are poor conductors of heat and electricity (graphite is a good conductor of electricity).
• Nonmetallic solids are usually brittle and are neither malleable nor ductile.

Element Classification Principles

Metalloids: There are some elements which have certain characteristics common to both metals and non-metals.

These are called semimetals or metalloids. Examples are—silicon (Si), germanium (Ge), arsenic (As), antimony (Sb) and tellurium (Te).

In most of their properties (both physical and chemical), metalloids behave as non-metals. However, they somewhat resemble the metals in their electrical conductivity. They tend to behave as semiconductors.

This property is found particularly in the case of silicon and germanium. These two metals are mainly responsible for the remarkable progress in the past five decades in the field of solid-state electronics.

Determination Of The Position Of An Element In Long Form Of Periodic Table

Since there is a close relationship between the long form of the periodic table and the electronic configuration of elements, the serial numbers of periods and groups and the type of block to which an element belongs can be predicted by following the guidelines given below:

Period: Serial number of the period = principal quantum number (n) ofthe valence shell.

Example: Mg (ls22s22pfi3s2) belongs to the third period because the principal quantum number of its valence shell is 3.

Block: The publicly into which the differentiating electron [i.e., the last electron) enters, represents the block to which the given element belongs (except He ).

Example: Sc (ls22s22p63s23pa4s23dl) belongs to d -block because the last electron [i.e., the 21st electron) enters the 3d -subshell.

Group: The group to which an element belongs can be predicted based on the number of electrons present in the outermost [i.e., fth) and the penultimate [i.e., [n —1) th] shell.

Element Classification Principles

For .s -block elements: Group-number = Number of valence electrons i.e., no. of electrons in the ns -orbital.

For p -block elements: Group-number = 10 + no. of valence electrons = 10 + no. of ns -electrons + no. of np electrons.

For d -block elements: Group-number = no. of ns- electrons + no. of (n- l)d -electrons.

For f- block elements: Group number= 3 (fixed).

Examples: Determination of the position of the elements with the following electronic configurations in the long form of the periodic table—

1. ls22s22p63s1
2. ls22s22p4
3. ls22s22p63s23p63d24s2
4. ls22s22p63s23p64s2
5. ls22s22p63s23p63d104s1

The given element belongs to s -the block because the differentiating electron (i.e., the the 11th electron) entering 3s orbital,

For this -block element, group-number = no. of electrons in the 3s -orbital =1.

Serial no. of the period = principal quantum number ofthe valence shell =

The differentiating electron [i.e., the 8th election) enters the p-subshell. So, tile given element belongs to p block, [b] Serial no. of the period = principal quantum number of the valence shell =

For this p -block element, group number= 10+ no. ofvalence electrons = 10 + number of ns electrons + no. of np -electrons =10 + 2 + 4 =16.

The differentiating electron [i.e., the 22nd electron) enters the 3d -subshell. So, the given element belongs to d -the block,

Element Classification Principles

Serial no. of the period = principal quantum number of the valence shell = 4.

For die d block element, group-number = no. of ns -electrons + no. of (n- 1)d -electrons = 2 + 2 = 4.

The differentiating electron {l.e., the 20th electron) enters the 4s -subshell. So, the given element belongs to s -block,

Serial no. of the period = principal quantum number of the valence shell =4.

For this -block element, group-number = no. of electrons in outermost shell = 2.

The differentiating electron [i.e., the 29th electron) enters the 3d -subshell. So, the given element belongs to d -block, [b] Serial no. of the period = principal quantum number of the valence shell = 4.

For this d -block element, group no. = no. of ns electrons + no. of[n- 1)d -electrons = 1 + 10 = 11.

IUPAC Nomenclature Of Transuranic Elements (Atomic Number More Than 100)

The elements beyond fermium (100) are called transfermium elements. They have atomic numbers above 101.

Fermium (100), mendelevium (101), nobelium (102), and lawrencium (103) are named after eminent scientists. Some of the elements with atomic numbers higher than 103 were synthesized and reported simultaneously by scientists from the USA and the Soviet Union.

Each group proposed different names for die same element, e.g., an element with atomic number 104 was named Rutherfordium by USA scientists while Soviet scientists named it Kurchatovium. To overcome such controversies, the

IUPAC (1977) has recommended a new method of naming these elements. This is discussed here.

Element Classification Principles

The digits expressing the atomic number of an element are represented serially (from left to right) by using the numerical roots given below.

The successive roots are written together and the name is ended by ‘ium! To avoid repetition of some letters, the following procedure is adopted.

If ‘enn’ occurs before ‘nil; the second ‘n’ of ‘enn’ is dropped.

Similarly the letter ‘i’ of ‘bi’ and ‘tri’ are dropped when they occur before ium bi+ium= bium, tri+ium= trium, enn+nil= ennil etc.

The symbol of an element is derived by writing successively the initial letters (z.e., abbreviations) of the numerical roots which constitute the name.

Element Classification Principles

 

Valency

The valency of an element is defined as the combining capacity of that element. The valency of an element is usually expressed in terms of the number of H-atoms that combine with an atom of the element.

The chemical properties of an element depend upon the number of electrons present in the outermost shell ofthe atom.

Electrons present in the outermost shell are called valence electrons and these electrons determine the valency ofthe atom.

In the case of the representative elements the valency of an atom is generally equal to either the number of valence electrons or equal to eight minus the number of valence electrons,.

However, transition and inner-transition elements exhibit variable valency involving electrons of the outermost shell as well as d- or f-electrons present in penultimate or antepenultimate shells.

Variation of valency in a period: In the case of the representative elements, the number of valence electrons increases from 1 to 7 from left to right in a period.

Oxygen-based valency increases from 1 to 7 and it becomes a zero noble gas series (because of its inertness). The maximum valency of ‘8’ is shown only by Os and Ru in 0s04 and Ru04 respectively. These two elements (transition elements) belong to group- 8 (VmB)in the periodic table.

Element Classification Principles

However, hydrogen-based and chlorine-based valency of representative elements along a period first increases from group-1 to 4 (valency= group no.) and then decreases from group-4 to 0 (valency= 8- group no).

Variation of valency in a group: On moving down a group, the number of valence electrons remains the same. Therefore, all the element groups exhibit the same valency.

Example: All the elements of group-IA (Li, Na, K, Rb etc.) have alencyT’ and that of group-2A (Mg, Ca, Sr etc.) exhibit avalencyof’21Noble gases present in group-VIIIA are zerovalent since these elements are chemically inert.

Ionisation Enthalpy or inonisation potential.

If energy is supplied to an atom, electrons may be promoted to higher energy states. If sufficient energy is supplied, one or more electrons may be removed completely from the atom leading to the formation of a cation. This energy is referred to as ionisation energy orionisation enthalpy (A/T).

Element Classification Principles

Ionisation Enthalpy or ionization potential Dentition: Ionisation enthalpy or more accurately first ionisation enthalpy of an element is defined as the amount of energy required to remove the most loosely bound electron from the valence shell of an isolated gaseous atom existing in its ground state to form a cation in the gaseous state.

Explanation: If AH1 (or I1) is the minimum amount of energy required to convert any gaseous atom in its ground state into gaseous ion M+, then the ionisation enthalpy or more accuratelyfirstionisation enthalpyofM is AH1 (orI1).

\(\begin{aligned}
& \mathrm{M}(\mathrm{g})+\Delta H\left(I_1\right) \longrightarrow \mathrm{M}^{+}(\mathrm{g})+e \\
& \text { (Isolated gaseous (Energy) (Gaseous (Electron at } \\
& \text { atom) cation) infinite distance) } \\
&
\end{aligned}\)

Element Classification Principles

Importance: The ionisation enthalpy of an element gives an idea about the tendency of its atoms to form gaseous cations.

Energy is always required to remove electrons from an atom and hence, ionisation enthalpies are always positive.

Units: It is expressed in kj per mole of atoms (kj. mol-1).

Formerly, it was expressed in electron-volt per atom (eV- atom-1) or kcal per mole of atoms (kcal. mol-1)

1ev per atom =23.06 kcalmol-1 =96.5 kl-mol¯¹.

‘Ionisation enthalpy is also called ‘ionisation potential’ because it is the minimum potential difference required in a discharge tube to remove the most loosely bound electron from an isolated gaseous atom to form a gaseous cation.

Successive ionisation enthalpies: Like the removal of the first electron from an isolated gaseous atom, it is possible to second, third etc., electrons successively from cations one after another.

The minimum amount of energy required to remove the second, third etc., electrons from unipositive, dipositive etc., ions to form M2+, M3+ etc., ions of the element are called second AH2 (or /2), tlirid AH3 (or I3 ) etc., ionisation enthalpies respectively.

Element Classification Principles

\(\begin{aligned}
& \mathrm{M}^{+}(g)+\Delta H_2\left(\text { or } I_2\right) \rightarrow \mathrm{M}^{2+}(g)+e \\
& \mathrm{M}^{2+}(g)+\Delta H_3\left(\text { or } I_3\right) \rightarrow \mathrm{M}^{3+}(g)+e
\end{aligned}\)

The second ionisation enthalpy is higher than the first ionisation enthalpy as it is more difficult to remove an electron from a cation than from a neutral atom. ] Similarly, the third ionisation enthalpy is higher than the second and so on i.e.,

AHj(or l1) < AH2(or l2) < AH3(or l3) <

If not mentioned, the term ‘ionisation enthalpy is always used to mean the first ionisation enthalpy of an element.

Formerly, first, second, third etc, ionisation enthalpies were denoted by the symbolsI, IZ, I3 etc. Such symbols will be used in many places in this book.

Factors governing ionisation enthalpy:

Atomicsize: Ionisation enthalpy decreases as the atomic size increases and vice-versa.

The attractive force between the electron (to be removed) and the nucleus is inversely proportional to the distance between them.

Thus, as the size of the atom increases, the hold of the nucleus over valence electrons decreases and consequently ionisation enthalpy decreases. For example, l1(Li)>l1(Na)>L1(K).

Element Classification Principles

The magnitude of nuclear charge: Ionisation enthalpy increases with an increase in nuclear charge and vice-versa.

This is due to the fact the force of attraction between the valence electron (to be removed) and the nucleus increases with an increase in the nuclear charge provided that the outermost electronic shell remains the same.

Screening effect of inner-shell electrons: As the screening effect or shielding effect of the inner electrons increases, the ionisation enthalpy decreases

In multi-electron atoms, the inner electronic shells act like a screen between the nucleus and the outermost electronic shell.

As a result, the nuclear attractive force acting on the electrons in the outermost shell is somewhat reduced i.e., the effective nuclear charge gets reduced to some extent.

Thus, the inner-electronic shells shield the electron (to be removed) from the nuclear attractive force, resulting in a reduction of ionisation enthalpy.

If other factors do not change, the ionisation enthalpy decreases with an increase in the number of inner electrons.

In multi-electron atoms, the ability of the electrons present in the inner shells to shield or screen the outer electrons from the attractive force of the nucleus is called the shielding effect or screening effect.

Naturally, the magnitude of the screening effect depends on the number of electrons present in the inner shells. In a particular energy level, the screening effect of the electrons presenting different subshells follows the sequencers >p> d> f.

Due to the screening effect, the valence shell electrons do not feel the full charge ofthe nucleus. The actual nuclear charge experienced by the valence shell electrons is called the effective nuclear charge.

This is given by the relation, Effective nuclear charge (Z)= total nuclear charge (Z) – screening constant (cr) where the screening constant (cr) takes into account the screening effect ofthe electrons present in the inner shells.

Element Classification Principles

Penetration effect of electronic subshells: Ionisation enthalpy increases as the penetration effect ofthe electron (to be removed) increases. It is known that in the case of multielectron atoms, the electrons in the s -s-orbital have the maximum probability of being found near the nucleus. In a given quantum shell this probability goes on decreasing in the sequence s->p-> d-> f.

This means that in a given shell, the penetration power of different subshells decreases in the order: of s->p-> d-> f-.

Now, if the penetration power of an electronic is greater, it is closer to the nucleus and held more firmly by it.

So it is more difficult to remove such an electron from the atom and consequently, ionisation enthalpy will be high.

Thus for the same shell, the energy needed to knockout an s -s-electron is greater than that required for a p-electron, which in turn will be more than that required to remove a d-electron and so on. In other words, ionisation enthalpy follows the sequence, s>p> d> f.

Effect of half-filled and filled subshells: It is known that half-filled and filled subshells have extra stability associated with them. Therefore, the removal of electrons from such subshells (having extra stability) requires more energy than expected.

Consequently, atoms having half-filled or filled subshells in their valence-shell have higher values of ionisation enthalpies.

Example: Be(ls22s2) has higher ionisation enthalpy than B(ls22s22p1) because ionisation of Be requires the removal of one electron from its filled 2s -orbital in the valence shell.

For similar reasons Mg(ls22s22p63s2) has higher ionisation enthalpy than Al(ls22s22p63s23p1)

N(ls22s22p3) has higher ionisation enthalpy than 0(ls22s22p4) because ionisation of nitrogen requires the removal of one electron from its half-filled 2p -the valence shell. Similarly, the ionisation enthalpy of P(3s23p3) is greater than that of S(3s23p4).

Effect of electronic configuration of the outermost shell:

Atoms, having the outermost electronic configuration ns2npG, are exceptionally stable because of their filled octet.

Removal of an electron from an atom having such a stable electronic configuration requires a large amount of energy.

Consequently, the noble gases He, Ne, Ar, Kr, Xe etc. (with outermost electronic configuration ns2np6) have very high ionisation enthalpy.

Variation of ionisation enthalpy in the periodic table: The periodic trends of the first ionisation energy of the elements are quite striking as seen from the graph.

The graph consists of several maxima and minima. In each period maxima occur at the noble gases which have filled the octet with the electronic configurations (ns2np6).

Element Classification Principles

Due to very high ionisation enthalpies, these elements are almost inert and show extremely low chemical reactivity.

In each period minima occur at the alkali metals which have only one electron in the outermost s -orbital. Due to very low ionisation enthalpies, these elements are highly reactive.

Variation of ionisation enthalpy across a period: For representative elements (s and p -block elements), ionisation enthalpy usually increases with increasing atomic number across a period. This is because as we move from left to right across a period—

The nuclear charge increases regularly, several shells remain the same and the addition of different electrons occurs in the same shell, and atomic size decreases.

As a result of a gradual increase in nuclear charge and a simultaneous decrease in atomic size, the valence electrons are more and more tightly held by the nucleus.

Therefore, more and more energy is needed to remove one valence electron and hence, ionisation enthalpy increases with an increase in atomic number across a period.

In any period, alkali metal has the lowest ionisation enthalpy and inert gas has the highest ionisation enthalpy.

Element Classification Principles

On careful examination of ionisation enthalpy values, some irregularities in the general trend are noticed. Can each period be explained based on different factors governing ionisation enthalpy?

Examples:

1. l1 of Be>I 1of B: Forionisation of boron (ls22s22p1), one electron is to be removed from the singly filled 2p orbital and this requires lesser energy, while for the ionisation of beryllium (ls22s2) one electron is to be removed from the more penetrating filled 2s orbital.

Furthermore, the Removal of an electron from Batom gives B+ a stable electronic configuration with a filled 2s -subshell (ls22s2) and so it requires a smaller amount of energy.

On the other hand removal of an electron from the filled 2s -orbital of Be -atom to give Be+ (1S22S1) requires a greater amount of energy.

Consequently, the first ionisation enthalpy of B is less than that of Be. of N > l1 of O: Electronic configuration of nitrogen (ls22s22p3) in which the outermost 2p -subshell is exactly half-filled is more stable than the electronic configuration of oxygen (ls22s22p4)in which the 2psubshell is neither half-filled nor filled.

Removal of 1 electron from the O -atom gives 0+ with a stable electronic configuration having a half-filled 2p -subshell (ls22s22p3), but this is not so in the case of the N -atom because N+ has the electronic configuration ls22s22p2.

In other words, the removal of an electron from the O -atom gives a cation with a more stable electronic configuration than that obtained by the removal of one electron from the N -atom. Thus, the first ionisation enthalpy of oxygen is less than that of nitrogen.

\(\mathrm{N}\left(1 s^2 2 s^2 2 p^3\right) \stackrel{-e}{\longrightarrow} \mathrm{N}^{+}\left(1 s^2 2 s^2 2 p^2\right)\) \(\mathrm{O}\left(1 s^2 2 s^2 2 p^4\right) \stackrel{-e}{\longrightarrow} \mathrm{O}^{+}\left(1 s^2 2 s^2 2 p^3\right)\)

The very high l1x value of the Exceptionally high value of the first ionisation enthalpy of neon (noble gas) amongst the elements of the 2nd period is due to its stable electronic configuration(ns2np6) ofthe outermost shell.

Variation of ionisation enthalpy down a group: For representative elements, ionisation enthalpy decreases regularly with an increase in atomic number on moving down a group from one element to the other.

Element Classification Principles

Explanation: The regular decrease in ionisation enthalpy (1.£.) may be attributed to the following factors:

On moving down a group, the atomic size increases successively due to the addition of one new electronic shell at each succeeding element.

Thus, the distance of valence shell electrons from the nucleus increases. Consequently, the nuclear attractive force on the valence electrons decreases and this, in turn, decreases the ionisation potential.

There is an increase in the shielding effect on the outermost electrons due to an increase in the number of inner electronic shells. This increased shielding effect tends to decrease the ionisation potential on moving down a group.

On moving down a group, the nuclear charge increases regularly and this increases the force of attraction of the nucleus on the valence electrons; this tends to increase the ionisation potential.

The combined effect ofthe increase in size and the shielding effect outweighs the effect of the increased nuclear charge.

Consequently, the ionisation enthalpies of the elements decrease regularly on going down a group. This is evident from the values of the first ionisation enthalpies of the elements of group-1 (alkali metals) as given in the adjacent table.

Element Classification Principles

Periodic variation of first ionisation enthalpies (eV) of the elements is evident from the following table.

Element Classification Principles

Some important facts about ionisation enthalpy:

1. The ionisation enthalpy of representative elements (s and block elements) increases from left to right across the period.
2. Exceptions are observed for some pairs of elements,
3. \(\text { e.g. } I_1(\mathrm{Be})>I_1(\mathrm{~B}) ; I_1(\mathrm{Mg})>I_1(\mathrm{Al}) ; I_1(\mathrm{~N})>I_1(\mathrm{O})\)
4. In any period, alkali metal has the least ionisation enthalpy.
   Cesium (Cs) has the lowest value of I. All the elements.
5. In each period, inert gas elements show the highest value of first ionisation enthalpy. Helium (He) has the maximum value ofI.E. of all the elements.
6. Among the representative elements, metals have low I.E., while non-metals have high values of I.E.
7. Generally, first ionisation enthalpies of transition elements (d -block elements) increase slowly from left to right in a period. This is partly due to the poor screening effect of d orbitals and partly due to electron-electron repulsive forces.
8. f-block elements also show a small change in their ionisation enthalpies on increasing atomic number.
9. From Pd to Ag, from Cd to In and also from Hg to Tl, there is a sudden decrease in ionisation enthalpy even though the atomic number increases.

Electron-gain enthalpy or electron affinity: energy is released when an electron is added to an isolated gaseous atom to convert it into a negative ion.

This energy is called electron-gain enthalpy. Electron-gain enthalpy of an atom is thus a measure of its tendency to form an anion. It is denoted by AHgg or EA.

Electron-gain enthalpy or electron affinity Definition: Electron-gain enthalpy is defined as the enthalpy change involved when an electron is added to an isolated gaseous atom in its lowest energy state (ground state) to form a gaseous ion carrying a unit negative charge.

\(\begin{aligned}
& \mathrm{X}(\mathrm{g})+e \longrightarrow \mathrm{X}^{-}(g)+\operatorname{Energy}(q) \\
& \text { Gaseous atom Electron Gaseous anion } \\
& \text { or, } \mathrm{X}(g)+e \longrightarrow \mathrm{X}^{-}(g), \quad \Delta H=-q \\
& \text { Gaseous atom Electron Gaseous anion } \\
&
\end{aligned}\)

Explanation: If q is the amount of energy released when an electron is added to the isolated gaseous atom ‘X’ in its ground state to convert to the gaseous ion X-, then the electron-gain enthalpy (electron affinity) of X is given by,AHeg = -q

Element Classification Principles

\(\begin{aligned}
\mathrm{F}(g)+e & \longrightarrow \mathrm{F}^{-}(g)+328 \mathrm{~kJ} \cdot \mathrm{mol}^{-1} \\
\text { or, } \mathrm{F}(g)+e & \longrightarrow \mathrm{F}^{-}(g), \Delta H=-328 \mathrm{~kJ} \cdot \mathrm{mol}^{-1} \\
\mathrm{Be}(g)+e & \longrightarrow \mathrm{Be}^{-}(g)-66 \mathrm{~kJ} \cdot \mathrm{mol}^{-1} \\
\text { or, } \mathrm{Be}(g)+e & \longrightarrow \mathrm{Be}^{-}(\mathrm{g}), \Delta H=66 \mathrm{~kJ} \cdot \mathrm{mol}^{-1}
\end{aligned}\)

When an F -atom combines with an electron to form an F- ion, energy is released. So enthalpy change has a negative value. Thus electron-gain enthalpy of fluorine is given by AHeg = be supplied to convert a Be -atom to a Be ion.

So enthalpy change has a positive value. Thus electron-gain enthalpy of beryllium is givenby, by AHeg = +66 kj. mol-1.

Points to remember: QElectron-gain enthalpy of an atom is a measure of its tendency to form an anion.

Electron-gain enthalpy has usually a negative value, but it may also have a positive value, especially for noble gases.

The numerical value of the ionisation enthalpy of an I negative ion (X-) is equal to the electron-gain enthalpy of the neutral atom (X).

However, energy is usually evolved during the process of electron acceptance but energy is usually absorbed during the expulsion of electrons from an atom. So electrongain enthalpy of X and ionisation enthalpy of X- have opposite signs.

Electron-gain enthalpy with a -ve sign indicates that energy is released when the neutral atom accepts an electron (only numerical values are taken for comparison when periodicity or otherproperties are considered).

The high value of electron-gain enthalpy indicates that an added electron is strongly bound, while a low value indicates that a new electron is weakly bound to the atom.

Units: Electron-gain enthalpies are expressed in kilojoule per mole (kj. mol-1) or in electronvolt (eV) per atom.

Successive electron-gain enthalpies: Like the second and higher ionisation enthalpies, second and higher electrongain enthalpies are also possible. However, the addition of a second electron to a negative ion (X-) is opposed by the electrostatic force of repulsion.

Element Classification Principles

So energy is to be supplied for the addition ofthe second electron. Thus, the second electron-gain enthalpy of an element is positive, and so is the third, and so on.

For example, when an electron is added to an oxygen atom to form an O-ion, energy is released. However, when another electron is added to the O- ion to form the O2- ion, energy is absorbed.

First electron-gain enthalpy:

\(\begin{array}{r}
\mathrm{O}(g)+e \longrightarrow \mathrm{O}^{-}(g), \quad \Delta H_{e g}=-141 \mathrm{~kJ} \cdot \mathrm{mol}^{-1} \\
\text { Energy released }
\end{array}\)

Second electron-gain enthalpy:

\(\begin{array}{r}
\mathrm{O}^{-}(\mathrm{g})+e \longrightarrow \mathrm{O}^{2-}(\mathrm{g}), \quad \Delta H_{e g}=\underset{ }{780 \mathrm{~kJ} \cdot \mathrm{mol}^{-1}} \\
\text { Energy absorbed }
\end{array}\)

Similarly, the first and second electron-gain enthalpies of sulphur are -200 kj mol-1 and +590 kj mol-1 respectively.

Factors governing electron-gain enthalpy: in general, the factors favouring the ionisation process disfavour the electron-gain process.

Effective nuclear charge: As effective nuclear charge (Z+) increases, the force of attraction between the nucleus and the incoming electron increases and hence, the numerical value of electron gain enthalpy increases.

Thus, the numerical magnitude ofelectrongain enthalpyof carbon (Z = 6, IE = -122 kj.mol-1 ) is greater than that ofboron(Z = 5, IE = -27 kj-mol-1).

Atomic size: As the size of the atom increases, the distance between the nucleus and the outermost shell (which receives the incoming electron) increases.

If the effective nuclear charge (Z+) per electron at the periphery is more or less the same for different species (e.g., in a group of representative elements), the force of attraction towards the nucleus of the electrons at the periphery is less for the larger species.

Consequently, the numerical magnitude of electron-gain enthalpy decreases as the atomic size increases. Thus for representative elements, the numerical value of electron-gain enthalpy decreases as the atomic number increases on moving down a group.

Nature of the orbital into which new electron gets accommodation: Orbitals which can penetrate more towards the nucleus are more suitable to accommodate the incoming electron.

Thus the ease of accommodation of Incoming eLectron follows the order ns> np> nd > nf, as the penetration effect of different orbitals follows this sequence. So the numerical magnitude of electron-gain enthalpy decreases in the sequence ns > np> nd> nf.

Nature of the outer electronic configuration: If the atoms of an element bear extra stability due to either the half-filled or full-filled subshell in their outermost level, then such atoms are very much reluctant to accept the incoming electron.

Element Classification Principles

On the other hand, if the newly added electron creates a half-filled or full-filled subshell, then the process is favoured.

Thus some ofthe elements of GrIIA(ns2), Gr-IIB[(n- l)d10ns2], Group-VA (ns2np3) and all the noble gas elements (ns2np6) have positive electron-gain enthalpies (AHeg).

On the other hand, elements of Gr-VIIA have very high electron-gain enthalpies with negative signs, because they can attain inert gas configuration accepting one electron.

Variation of electron-gain enthalpy across a period: On moving from left to right in a period, effective nuclear charge, Z nuclear charge (Z)- shielding effect of the inner shells) increases and size decreases with the increase in atomic number.

Both these factors tend to increase the nuclear attraction experienced by the incoming electron and hence, the numerical value of electron-gain enthalpy, in general, increasesin a period from left to right. It reaches a maximum value at Gr-VIIA (halogens).

Due to some characteristic electronic configuration, the general trend is violated in some cases {e.g., Be and N in the 2nd period; Mg and P in the third period).

Variation of electron-gain enthalpy down a group: For the representative elements, on moving down a group, the effective nuclear charge Z per electron at the periphery (outermost shell) remains more or less constant because the effect of increased nuclear charge is counterbalanced by the shielding effect of the inner electronic shells.

However, the atomic size gradually increases due to the addition of new quantum levels. Thus the nuclear attractive force experienced by any added electron (incoming electron) decreases as the atomic number increases, and consequently, the numerical value of electron-gain enthalpy decreases down a group.

Element Classification Principles

Some typical trends in electron-gain enthalpy & their explatition:

Experiences significant electron-electron repulsion from the other electrons present in the small-sized 2p -subshell.

On the other hand, in a chlorine atom, the added electron goes to the large-sized 3p -subshell. Hence, it experiences less electron-electron repulsion.

Another factor that favours the uptake of electrons by the Cl -atom, is that there is the possibility of the delocalisation of the increased electron density in the vacant 3d -orbital of Cl-atom.

This mechanism is not operative in F-atom because of the absence of d -orbital in the second shell. Consequently, the numerical value of electron-gain enthalpy of Cl is greater than that of F.

Electron-gain enthalpy is greater than that of O: The reason for this anomaly is similar to that of Cl versus F.

The added electron experiences considerable electron-electron repulsion from the other electrons present in the small-sized 2p -subshell of O.

This repulsion outweighs the increased attractive force of the nucleus acting on the added electron. In the S-atom, the added electron goes to the large-sized 3p -subshell.

Hence, it experiences less electron-electron repulsion. Another factor that favours the uptake of electrons by S -S-atoms is that there is a possibility ofthe delocalisation of the increased electron density in the vacant-3d -orbital of S-atom.

This mechanism is not operative in the O -atom because of the absence of any d orbital in the 2nd shell.

Element Classification Principles

Consequently, the numerical value of electron-gain enthalpy S is greater than that of O.

Gr-llA metals (e.g., Be, Mg etc.) have lower electron-gain enthalpies than Gr-IA metals (e.g., Li, Na, K efc: Gr- A metals have outer electronic configuration ns2. Hence, the addition of an extra electron brings the configuration ns2np1.

This process is disfavoured in two ways: The addition of a new electron destroys the full-filled subshell structure and 0 accommodation of the new electron occurs in the p -orbital which is less penetrating.

For alkali metals (ns1), however, accommodation of the new electron occurs in the ns -subshell giving rise to a filled ns2 configuration.

Thus, the electron-gaining process is more favourable for Gr-IA elements compared to Gr-IIA elements. Be and Mg ofGr-IIA have positive electron-gain enthalpy.

Halogens have the highest electron-gain enthalpies: This is because of the valence-shell electronic configuration of the halogensis ns2np5 and so theyrequire only one more electron to acquire the stable inert gas-like electronic configuration (ns2np&).

As a result, halogen atoms have a strong tendency to accept an additional electron. Consequently, the numerical values of their electron-gain enthalpies are very high.

Phosphorous (3s23pi3) has relatively low electron-gain enthalpy: This is because the P -atom has a relatively stable outer electronic configuration with exactly half-filled p -orbital.

Hence, it is reluctant to accept an extra electron. Consequently, it has low electron-gain enthalpy.

The electron-gain enthalpy of noble gas is high and positive:

The atoms of noble gases have a very stable outermost electronic configuration with filled subshells (ns2np6).

Any additional electron would have to be placed in an orbital ofthe next higher energy level.

The shielding effect of the inner electrons and the large distance from the nucleus makes the addition of an electron highly unfavourable. So, noble. high and adjaisitive values ofelectron-gain enthalpy.

Electronegativity

This topic will be discussed elaborately in the chapter ‘Chemical Bonding’ Here we will briefly discuss only the definition of electronegativity and its periodicity.

Electronegativity Definition: Electronegativity is defined as the tendency of an atom to attract the shared pair of electrons towards its nucleus when the atom is covalently bonded in a molecule.

Consequently, the more electronegative atom withdraws the bonding electron cloud more towards its nucleus giving rise to an accumulation of negative charge on it.

The electronegativity of an element is not its inherent property. It depends on its surrounding environment in the molecule in which the electronegativity of the element is being considered.

Element Classification Principles

Thus the electronegativity of S is different in different compounds such as H2S, S02, SFg etc.

Further, it is to be noted that unlike ionisation enthalpy and electron-gain enthalpy, electronegativity is not a measurable quantity.

Factors controlling electronegativity: Electronegativity of the elements depend on—

The atomic number of an element, i.e., the total quantity of positive charge in the nucleus of an atom,

Size of atom Or Atomicradius, number of electronicshes in an atom, oxidation state ofthe atom, state of hybridisation of the atom in the molecule under consideration. Note the electronegativity of elements.

Variation of electronegativity across a period: As Cs 0.7 At 2.2 Ford -block element, on moving down from 3d- to from left to right along a period, nuclear charge Increases while the atomic radius or size decreases.

Hence the attraction between the outer (or valence) electrons and the nucleus increases with increasing atomic number.

Consequently, the electronegativity of the atom increases from left to right across a period. Thus alkali metals of group-1 on the extreme left have the lowest electronegativity whereas, the halogensin group-17 on the right have the highest values of electronegativity in their respective periods.

This is evident from the electronegativity values ofthe elements of the second and third periods.

Variation of electronegativity down a group: As we move down a group, atomic size (radius) as well as the magnitude of nuclear charge increases but the effect of increased nuclear charge on the outer electrons is mostly counterbalanced by the screening effect of a larger number of inner electronic shells.

Hence the nuclear pull on the outer (valence) electrons decreases due to the increase of atomic size on moving down a group.

Consequently, the electronegativity of an atom decreases from top to bottom in a group. This is evident from the electronegativity values of the alkali metals of group-1 (IA) and halogen elements of group-17(VIIA)

Element Classification Principles

Ford -block element, on moving down from 3d- to 4d -series, electronegativity falls slightly but on reaching 5d series, electronegativity increases due to lanthanide contraction.

Relationship between electronegativity and non-metallic (or metallic) character of elements: Non-metallic elements have a strong tendency to gain electrons. So, electronegativity is directly related to the metallic character elements.

It can be further extended to say that electronegativity is inversely related to the metallic character of elements.

Thus the increase in electronegativity along a period is accompanied by an increase in non-metallic character (or decrease in metallic character) of elements.

Likewise, the decrease in electronegativity down a group is accompanied by a decrease in the non-metallic character (or increase in metallic character) of elements.

All these periodic trends are summarised in the given figure: (Direction arrows indicate increasing trend of the respective properties)

Element Classification Principles

Periodicity in density, melting point and boiling point

Different elements exhibit periodicity in various physical properties such as density, melting point, boiling point etc.

Periodic variation of density: On moving along a period from left to right, the density of representative elements first increases, reaches the maximum value at group-IIIA orIVA and then decreases with an increase in atomic number.

This trend is observed particularly in the case of representative elements. In a group, density generally increases from top to bottom with a rise in atomic number.

Periodic variation of melting and boiling points: On moving along a period from left to right, the melting and boiling points of representative elements first increase, reach maximum values at group IVA and thereafter go on decreasing. Minimum melting and boiling points are shown by the noble gas in the respective period.

Periodicity In Properties Of Oxides And Hydrides

Nature of oxides of the elements: On moving from left to right across a period, the basic properties and electrovalent character of oxides of elements decrease while their acidic property and covalent character gradually increase.

On the other hand, in a group, the basic property of oxides increases from top to bottom.

The nature of the oxides of transition metals depends on the oxidation state of the metals. With the increase in the oxidation state of transition metals, the acidic properties of their oxides increase.

Element Classification Principles

Example: CrO is a basic oxide, Cr203 is amphoteric and Cr03 I -a is an acidic oxide. In the case of oxides of the elements of the second period, it is observed that lithium oxide (Li20) is strongly basic. It reacts with water to produce a strong base namely lithium hydroxide (LiOH)

\(\mathrm{Li}_2 \mathrm{O}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons 2 \mathrm{Li}^{+}+2 \mathrm{OH}^{-} \rightleftharpoons 2 \mathrm{LiOH}\)

BeO is an amphoteric oxide. It reacts with both acids and bases to form salt and water.

\(\text { Basic property: } \mathrm{BeO}+2 \mathrm{HCl} \rightarrow \mathrm{BeCl}_2+\mathrm{H}_2 \mathrm{O}\) \(\text { Acidic property: } \mathrm{BeO}+2 \mathrm{NaOH} \rightarrow \mathrm{Na}_2 \mathrm{BeO}_2+\mathrm{H}_2 \mathrm{O}\)

B203 is an acidic oxide though it possesses a slight basic property. It reacts with water to form orthoboric acid and with alkali to yield borate salt.

CO2 is an acidic oxide and it reacts with alkali to produce carbonate salt. N2O is an acidic oxide. It reacts with alkali to produce salt and water.

⇒ \(\begin{gathered}
\mathrm{B}_2 \mathrm{O}_3+3 \mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{H}_3 \mathrm{BO}_3 ; \mathrm{CO}_2+2 \mathrm{NaOH} \rightarrow \mathrm{Na}_2 \mathrm{CO}_3+\mathrm{H}_2 \mathrm{O} \\
\mathrm{N}_2 \mathrm{O}_5+2 \mathrm{NaOH} \rightarrow 2 \mathrm{NaNO}_3+\mathrm{H}_2 \mathrm{O}
\end{gathered}\)

Nature of hydrides of elements: As we move from left to right across a particular period, the tendency of the elements to form hydrides and the thermal stability, covalent character, acidic property, and volatility of the hydrides increases while the reducing property progressively decreases.

The hydrides of the strongly electropositive metals towards the left of a period are ionic having high melting points. On ionisation, they produce hydride ions (H-). Again, hydrides of non-metals towards the right of the period are covalent and have low melting and boiling points.

On moving down a group, the tendency of the elements to form hydrides decreases. The stability of the hydrides also decreases in the same sequence. Variation of other properties along any group depends on the group to which the hydride-forming element belongs.

Hydrides of alkali metals in group IA and alkaline earth metals in group 2 are salt-like polar or ionic.

These compounds are composed of positive metallic ions and negative hydride ions (H¯). On electrolysis of these ions and negative hydride ions (H-).

On electrolysis of these discharged at the cathode and anions (H- ions) at the anode; e.g., electrolysis of molten sodium hydride leads to the formation of metallic sodium at the cathode and H2 gas at the anode.

Element Classification Principles

Hydrides of the elements of groups IVA to VIIA are covalent and nonpolar; e.g., CH4, SiH4, PH3 etc., are gaseous and insoluble in water. NH3 and H2S are gaseous but soluble in water. An aqueous solution of NH3 is feebly basic and the aqueous solution of H2S is weakly acidic.

On the other hand, HC1, HBr and HI, despite being covalent compounds more soluble in water and dilute aqueous solutions, dissociate almost completely.

\(\mathrm{HX}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}+\mathrm{X}^{-} \quad[\text { where } \mathrm{X}=\mathrm{Cl}, \mathrm{Br}, \mathrm{I}]\)

Aqueous solution of HX is strongly acidic. In electrolysis of their aqueous solutions hydrogen ions (H+) are liberated at the cathode and halide ions (X-) at the anode e.g., when an aqueous solution of hydrochloric acid is electrolysed, H2 gas is evolved at the cathode and Cl2 gas at the anode.

The trend of variation in properties of different elements in the periodic table from left to right across a period and from top to bottom in a group is shown in the given table-

Element Classification Principles

Diagonal Relationship Definition: Some elements of certain groups in the second period show similarity in properties with the diagonally opposite elements of the third period, and such similarity in properties is referred to as a diagonal relationship.

Reason for diagonal relationship: The Reason for diagonal relationship is due to opposing trends in periodic properties along a period from left to right and down the group.

For example, the atomic and ionic radius of elements decrease a periodic but increase down a group Ionisation enthalpy, electron gain enthalpy and electron negativity increase along a period but decrease down a group.

On moving diagonally, two opposite trends mutually cancel, so the elements of the period- 2 and 3 listed above are related to each other diagonally and they show similar chemical properties. Thus Li resembles Mg; Be resembles A1; and B resembles Si.

The diagonal relationship is also explained based on polarising power ofcation. On moving along the period from left to right, the charge on the cation increases, while ionic size decreases and hence polarising power increases.

Again on moving down a group, the charge on the cation remains the same, while ionic size increases. Hence polarising power decreases. So on moving diagonally, polarising power remains more or less the same and the elements exhibit similar properties.

Absence of diagonal relationship in case of long periods: Because of the intervening d – and /-series, the diagonal relationship does not hold well for long-period elements (4th, 5th… period elements).

Because the group trend of many properties in the transition series is opposite compared to that in the representative elements. However, the trend along the periods remains the same for both the representative and d -d-block elements.

Position Of Hydrogen And Inert Gases In The Periodic Table

The position of hydrogen in the periodic table is controversial. Given its chemical analogy with both the elements of group and that of group- VIIA, it can either be placed in group 2A or group VIIA. Resemblances ofhydrogen with the element.

Arguments in favour of placing hydrogen in group IA

Valency: The electronic configuration ofhydrogen is Is1 and the general electronic configuration of the elements of group-IA is ns1, i.e., like the elements of group-IA, hydrogen has only one valence electron and its valency is 1.

Element Classification Principles

Electropositive character: Like all group-IA elements, hydrogen tends to form cations by losing one electron.

\(\mathrm{Na}-e \longrightarrow \mathrm{Na}^{+} ; \mathrm{K}-e \longrightarrow \mathrm{K}^{+} ; \mathrm{H}-e \longrightarrow \mathrm{H}^{+}\)

Like elements of group-IA, hydrogen reacts with electronegative elements such as chlorine, oxygen and sulphur to produce similar type of compounds, e.g

HC1 , H20 , H2S ; NaCI , Na20 , Na2

Electrolysis of chloride compounds: Electrolysis of molten NaCl results in the deposition of metallic sodium at the cathode. Likewise, when an aqueous solution of HCl is electrolysed, H2 gas is liberated at the cathode.

\(\mathrm{NaCl} \rightleftharpoons \mathrm{Na}^{+}+\mathrm{Cl}^{-}\)

1. Cathode : Na+ + e→Na
2. Anode : Cl-e →Cl ; CI + C1→C12T
3. HCl H+ + Cl Cathode : H+ + e →H ; H + H→H2?
4. Anode : Cl–e →C1 ; C1 + C1→C12T

Reducing property: Like the elements of Gr-IA, hydrogen loses electrons easily and exhibits a reducing property.

Formation of alloy: Hydrogen dissolves in metals like Pd, Pt etc., by adsorption. This occlusion of hydrogen is comparable to the formation of alloys by elements of group IA.

Mutual displacement: Hydrogen atom(s) of hydrochloric, sulphuric or nitric acids can be displaced by the same number of atoms of group-IA elements. Again, atoms of the group-IA elements can be replaced by hydrogen atoms from the salts produced.

Formation of stable oxide: Oxides of group-IA elements are highly stable (e.g., Na20, K20 etc.). Similarly, oxide of hydrogen (H20) is also highly stable.

Formation of peroxide: Like the elements of group-IA, hydrogen also forms peroxide (H202). The analogous peroxides of group-IA elements are Na202, K202 etc.

Electron affinity: Hydrogen and the group-IA elements have comparable values of electron affinity.

In light of the above similarities between the elements of group IA and hydrogen, hydrogen can be placed along with the elements of group IA. However, the placement of hydrogen in group IA leaves six vacant places in between H and He in the first period.

Arguments in favour of placing hydrogen in group-VIIA

Electronic configuration: The electronic configuration of hydrogen is Is1 and the electronic configuration of the outermost orbit of the elements of group-VIIA is ns2np5, i.e., the outermost orbit of both hydrogen and elements of group-VILA has 1 electron less than the electronic configurations of the nearest inert gas. So, their valency 1.

Non-metallic character and atomicity: hike the dements of group-VIIA, hydrogen is also it non-metal and forms a diatomic molecule.

Formation of anion: Like the elements of group VIIA, the hydrogen atom also tends to attain the electronic configuration of Its nearest Inert gas (Me) by accepting I electron and forming anion (H” ); e.g.,

\(\mathrm{H}\left(1 s^1\right) \stackrel{+e}{\longrightarrow} \mathrm{H}^{-}\left(1 s^2\right) ; \mathrm{X}\left(n s^2 n p^5\right) \stackrel{+e}{\longrightarrow} \mathrm{X}-\left(n s^2 n p^6\right)\)

Both VIIA elements and hydrogen form electrovalent halide and hydride respectively. During the electrolysis of metallic hydrides, like halogens, hydrogen is also liberated at the anode.

Formation of covalent compounds: Just like elements of group VIIA, hydrogen reacts with different non-metals to produce covalent compounds with analogous formulas.

Compoundsinvolving11: CH4, NH3, H20, HF, SiH4. CompoundsinvolvingCl: CC14, NC13, C120, CIF, SiCl4

Substitution by halogens: H-atoms of the hydrocarbons can be substituted by Gr-VIIA elements, partially or completely.

Ionisation potential: Just like the elements of group VIIA, the ionisation, potential of hydrogen is very high but the ionisation potential of alkali metals is quite low. The following table of ionisation potentials shows the comparative picture of ionisation potential quite explicitly.

Maintenance of continuity in the periodic table: If H is placed in group YIIA, no vacant space remains between H and H. So, continuity in the periodic table is not disturbed.

From the above discussion, it is apparent that hydrogen is a unique element characterised by peculiar and distinctive unique element characterised by peculiar and distinctive position to it in the periodic table. It is reasonable to set aside a separate position for hydrogen in the periodic table. In the modern periodic table, hydrogen has been J given a completely separate place, at the top ofthe table.

Position of Inert Gases In The Periodic Table

Inert gas elements have very stable electronic configurations of their outermost or valence shell (ns2 for He and ns2np6 for others).

For this reason, these elements show little or no tendency to lose or gain electrons to form ions to give electrovalent bonds or do not share electrons with other elements to form covalent bonds. So the combining capacity or valency of these elements is zero.

Thus they are placed in group ‘zero’ ofthe periodic table. This group forms a bridge between the most electropositive alkali metal elements of group-IA and the most electronegative halogen elements of group-52A.