WBBSE Solutions For Class 8 School Science Chapter 2 Element Compound And Chemical Reaction Long Answer Type Questions

Chapter 2 Element Compound And Chemical Reaction Chemical Reaction Long Answer Type Questions

 

Question 1. What is meant by a positive catalyst and a negative catalyst? Give a suitable example.
Answer:

Positive catalyst And a Negative catalyst:

The substance which can increase the rate of a chemical reaction is called a positive catalyst.
When potassium chlorate is heated at 630°C, oxygen is produced very slowly. But when potassium chlorate is mixed with manganese dioxide (in 1: a 4 ratio i.e., 1 part MnO₂ with 4 part KClO₃), then oxygen is produced at a much lower temperature (240°C). So, in this reaction, MnO₂ acts as a positive catalyst and increases the rate of the reaction.

⇒ \(2 \mathrm{KClO}_3 \rightarrow 2 \mathrm{KCl}+3 \mathrm{O}_2\)

The substance which decreases the rate of a chemical reaction is called a negative catalyst Hydrogen peroxide dissociates slowly even at room temperature to produce oxygen and water.

But when a little phosphoric acid ( H₃PO₄) is added to H₂O₂, it slows down the rate of decomposition of H₂O₂. So H₃PO₄ acts as a negative catalyst in this reaction.

⇒ \(\mathrm{H}_2 \mathrm{O}_2 \rightarrow \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2\)

Question 2. Mention the characteristic properties of a catalyst
Answer:

Catalyst:

Many chemical reactions occur very slowly. But the addition of a minute quantity of another substance can speed up the reaction.

The substances which can increase the rate of a chemical reaction are called catalysts. There are certain catalysts which slow down the rate of a particular chemical reaction.

They are termed negative catalysts. The chemical reaction which involves the use of a catalyst is called a catalytic reaction.

Catalyst is a substance which is present in small quantities and increases the rate of a chemical reaction without itself undergoing any permanent change. [ A negative catalyst however decreases the rate of the chemical reaction.] Examples Of Some Catalytic Reactions

 

 

Characteristics of a catalyst

1. At a particular temperature, the addition of a catalyst increases the rate of a chemical reaction.
2. A catalyst cannot initiate a reaction. This means that a catalyst cannot “start” a reaction which otherwise does not occur. It only influences the rate of a reaction.
3. A catalyst participates in the reaction it catalyzes but is regenerated at the end of the reaction.
4. A catalyst present in a very small amount is able to influence the rate of a reaction significantly.
5. There is no universal catalyst which can enhance the rate of all the reactions. A suitable catalyst for a specific reaction must be found only by proper experimentation.

How Does A Catalyst Work

For various reactions, the mechanism by which a specific catalyst works is known. But this discussion is beyond the scope of this book.

At this point, we can only say that a catalyst helps a reaction to occur by a different route, which requires relatively less energy.

So more reactants can participate in a reaction, producing more products. Let us consider a situation where a catalyst is in the solid state within a reaction mixture which is in a liquid state (or gaseous state).

Generally in such cases, the solid catalyst provides a surface where at least one of the reactants can “seat” and only those reactants which are “seated” on the solid surface can react with other reactants which are also seated or which are not seated but remain very close to the solid surface.

So we realize that work more the surface area of a solid catalyst, the more molecules it can accommodate on its surface and the more wilt the rate of the reaction.

If a solid catalyst is crushed to powder, the surface area of the solid substances increases substantially and thus more reactants get the opportunity to react.

(The process by which the reactant molecules “sit” on the solid surface is called “adsorption”.) If the physical state of the catalyst is different from that of the reactants, then the catalyst is called a heterogeneous catalyst.

For example, in the production of ammonia from the reaction between nitrogen gas and hydrogen gas, a solid powder of iron is used as a heterogeneous catalyst. For a heterogeneously catalyzed reaction, adsorption always precedes the chemical reaction.

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Organic catalysts: Enzymes

Chemical reactions are constantly occurring within our bodies too. Formation or breaking of chemical bonds rarely happens on their own within biological systems.

Enzymes are biological molecules that act as catalysts and help complex reactions occur. The enzymes are basically proteins, but they, ey may be associated with non-protein substances (known as coenzymes or prosthetic groups) that are essential for the action of the enzyme.

The term “enzyme” meaning “from yeast” was coined by German physiologist Wilhelm Kuhne in 1876. Eduard Buchner showed that fermentation, previously believed to depend on a mysterious “life force” contained only in living organisms, could be achieved by extracts from yeast (which is not “living”).

Enzymes are highly specific catalysts in the sense that a particular enzyme acts on only a particular reactant (generally called “substrate”) to complete a specific task.

Experiments have revealed the catalytic activity of an enzyme Some hands-on experiments can be done to observe the effect of an enzyme on a chemical reaction.

Urease	Catalyses hydrolysis of urea \(\begin{aligned} & \left(\mathrm{NH}_2 \mathrm{CONH}_2+\mathrm{H}_2 \mathrm{O} \rightarrow\right. \\ & \left.2 \mathrm{NH}_3+2 \mathrm{CO}_2\right) \end{aligned}\)
Invertase	Converts surcose to glucose and fructose
Pepsin (Found in our Stomach)	Converts proteins into smaller amino acids
Trypsin (Found in our intestine)	Smaller Amino acids
Catalase	decompose Hydrogen peroxide (H2O) to water and oxygen  \(2 \mathrm{H}_2 \mathrm{O}_2 \rightarrow 2 \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2\)

 

1. Let us mix equal volumes of hydrogen peroxide and water in a test tube and a small piece of fresh potato is added. The enzyme, catalase, is found in the fresh potato.

Almost immediately after the addition, bubbles of oxygen start evolving from the reaction mixture. We have already mentioned that catalase decomposes hydrogen peroxide into oxygen and water. The reaction can be repeated with fish or goat liver, which also contains the enzyme catalase.

Enzymes are indispensable for living things because all the biological reactions occurring within living things are catalyzed by enzymes.

Starting from digestion of food to synthesis is due to a relatively small region of the protein molecule present in the enzyme. This region is commonly referred to as an “active centre

2. Let us dissolve some urea in a small volume of water taken in a beaker and to it some amount of powdered ahar daal is added.

After 10 – 15 minutes, if we smell the reaction mixture, a faint pungent smell of ammonia is obtained. Since ammonia is a weak base, so if we add a few drops of phenolphthalein, the reaction mixture turns pink. Arhar daal and watermelon seeds contain the enzyme urease.

Urease catalyses the hydrolysis of urea. Microbes in urinals decompose urea present in urine and that is why we get the smell of ammonia in the urinal.

proteins, hormones, DNA etc., and enzymes are everywhere. Each enzyme performs only a particular action on a particular substrate. Enzymes are indispensable for living things because all the biological reactions occurring within living things are catalyzed by enzymes.

Starting from digestion of food to synthesis is due to a relatively small region of the protein molecule present in the enzyme.

This region is proteins, hormones, DNA etc., enzymes are everywhere. Each enzyme performs only a particular action on a particular substrate.

Question 3. What is an exothermic and endothermic reaction? Give examples.
Answer:

Exothermic And Endothermic Reaction:

Chemical reactions which proceed with the evolution of heat energy are called exothermic reactions. Reactants → Products + heat (Q) Burning of coal is an example of an exothermic reaction. Here, the carbon present in coal burns in oxygen to produce carbon dioxide with the evolution of large amounts of heat.

⇒ \(\mathrm{C}+\mathrm{O}_2 \rightarrow \mathrm{CO}_2+\text { heat (Q) }\)

Chemical reactions which proceed with the absorption of heat energy are called endothermic reactions.
Reactants + heat (Q) →Products Making quicklime from limestone is an example of an endothermic reaction, where heat is absorbed.

⇒ \(\mathrm{CaCO}_3 \rightarrow \mathrm{CaO}+\mathrm{CO}_2-\text { heat (Q) }\)

Question 4. Give one example of each of the exothermic physical processes and endothermic physical processes.
Answer:

When ammonium chloride (NH₄CI) is dissolved in water, then heat is absorbed and the solution becomes cold. This is an example of an endothermic physical process.

When sulphuric acid ( H₂SO₄) is slowly added dropwise to water, it becomes soluble and a large amount of heat is liberated and the solution becomes very hot. This is an example of an exothermic physical process.

Question 5. Describe a simple experiment to show that water and oxygen, both are necessary for the rusting of iron.
Answer:

Rusting:

Rusting of iron is a common form of corrosion in which the metal is eaten up gradually due to oxidation of the metal by the action of air, moisture or a chemical (such as acid) on its surface.

When metallic iron is exposed to moist air (i.e., in presence of oxygen and water), it is converted to hydrated iron oxide (Fe₂O₃.n H₂O) [where n is the number of water molecules].

This is brittle and the mechanical strength of metallic iron is absent in this hydrated iron oxide. In other words, metal is “degraded”.

This is an example of an oxidation reaction, where metallic iron is oxidized to hydrated iron oxide.

⇒ \(4 \mathrm{Fe}+3 \mathrm{O}_2+2 \mathrm{nH}_2 \mathrm{O} \rightarrow 2 \mathrm{Fe}_2 \mathrm{O}_3 \cdot \mathrm{nH}_2 \mathrm{O}\)

Iron oxygen water hydrated iron oxide or rust.

The number of water molecules (n) in the rust varies, it is not fixed. Rusting involves unwanted oxidation of iron which occurs in nature on its own. It is a continuous process.

This single class of reaction is responsible for the destruction of various materials, instruments and infrastructures made of iron.

Everything made of iron and which is exposed to moist air is vulnerable to corrosion. Every year, crores of rupees are required for round-the-clock maintenance of costly instruments and infrastructures.

Corrosion control can be achieved by recognizing and understanding the corrosion mechanism, by using corrosion-resistant materials and designs and by using protective systems and devices and treatments.

For example, it is found that the more the exposed area of metallic iron is too moist air, the more the extent of rusting. Hence, if the exposed surface of iron is coated with some paints (such as coal tar or some synthetic paints) which prevent direct contact between iron and moist air, the rate of corrosion can be significantly reduced.

Probably this is the cheapest way to prevent corrosion Another more sophisticated way to prevent rusting is to coat an iron surface with another metal.

For example, metallic zinc can be electroplated on the iron surface to prevent rusting. The process of coating the surface of any substance with metallic zinc is known as galvanization.

We have mentioned that during rusting metallic iron is converted to hydrated iron oxide. Alternatively, we can say that Fe is converted to Fe²⁺ ion.

But we also know that oxidation and reduction occur simultaneously. So, if Fe-2e → Fe²⁺ is the oxidation reaction, which one is the reduction reaction? In fact, in an acidic solution, there can occur two types of reduction reactions, as follows:

⇒ \(\begin{aligned}
& 2 \mathrm{H}^{+} \text {(aqueous) }+2 \mathrm{e} \rightarrow \mathrm{H}_2 \text { (gas) } \\
& 4 \mathrm{H}^{+} \text {(aqueous) }+\mathrm{O}_2 \text { (gas) }+4 \mathrm{e} \rightarrow 2 \mathrm{H}_2 \mathrm{O} \\
& \text { (liquid) }
\end{aligned}\)

A simple experiment can be performed to show that water and oxygen, both are necessary for the rusting of iron.

Experiment: Let us take three beakers. In the first beaker, some iron nails are placed and it is left open in the air for a few days.

In the second beaker, some normal water is taken. A few iron nails are immersed in it, and then some molten wax is poured into the beaker in such a way, so as to create a layer of wax on the surface of the water, which prevents the passage of air through it into the water.

In the third beaker, instead of normal water some amount of boiled water is taken and the iron nails are immersed in it. In this case, also, the surface of the water is covered with a layer of wax,

so that water does not come in direct contact with the air outside. The second and third beakers are also left undisturbed for a few days.

Observation: After some days, it will be found that the nails in the first beaker are rusted. Some rusting takes place in the nails kept in the second beaker. But no rusting takes place in the nails kept in the third beaker.

Inference: This is because the nails are in direct contact with moisture and oxygen present in the air in the first beaker.

So rusting occurs. In the second beaker, nails are in direct contact with water and oxygen (which remains dissolved in water under ordinary temperature and pressure).

So here also rusting of the nails occurs. But in the third beaker, boiled water is taken. When water is properly boiled, the dissolved oxygen is driven out.

So, in absence of any oxygen rusting does not occur, although the nails are in direct contact with water.

This conclusively proves that for rusting to occur, the presence of both water and oxygen is necessary.

Question 6. What is oxidation? Give a suitable example.
Answer:

Oxidation:

Oxidation is a chemical reaction, which involves the addition of oxygen or addition of chlorine or the elimination of hydrogen.

1. Addition of oxygen: Burning of sulphur in air or oxygen produces sulphur dioxide. Here, oxygen combines with sulphur.

So, it is an example of an oxidation reaction. \(\mathrm{S}+\mathrm{O}_2 \rightarrow \mathrm{SO}_2\)

2. Addition of chlorine: Chlorine combines with iron to produce ferric chloride. So this is an oxidation reaction.

3. Elimination of hydrogen: When hydrogen sulphide ( H₂S) gas is passed through chlorine water, some amount of sulphur is precipitated and hydrogen chloride (HCI) is produced. Here hydrogen is eliminated from hydrogen sulphide. So this is an example of oxidation.

Question 7. What is reduction? Give a suitable example. Reduction is a chemical reaction which involves, the elimination of oxygen or elimination of chlorine or the addition of hydrogen.
Answer:

1. Elimination of oxygen:

When hydrogen gas is passed over hot, black, cupric oxide, reddish-brown metallic copper is formed as residue. Here, oxygen is eliminated from CuO. This process is a reduction.

 

 

2. Elimination of chlorine: When HCI gas is irradiated with light, it is dissociated into hydrogen and chlorine gas. Here, chlorine is eliminated from HCI. So it is an example of reduction.

 

3. Addition of hydrogen: When hydrogen sulphide ( H₂S) gas is passed through chlorine waterborne amount of sulphur is precipitated and hydrogen chloride (HCI) is produced. So hydrogen is eliminated from H₂S, so reduction has occurred in this case.

 

Question 8. What is the electronic theory of oxidation and reduction? Explain with suitable examples.
Answer:

Electronic theory of oxidation and reduction:

Oxidation is defined, as, a chemical reaction involving the loss of electrons, from an atom or ion and reduction, is a chemical reaction involving the gain of electrons by an atom or ion.

The substance which accepts electron(s) is reduced. It is called the oxidising agent or oxidant. The substance that loses an electron(s) is itself oxidised. That is known as a reducing agent or reductant.

Let us consider the following chemical reaction:

⇒ \(\mathrm{CuSO}_4+\mathrm{Fe} \rightarrow \mathrm{FeSO}_4+\mathrm{Cu}\)

In copper sulphate, copper exists as a cupric ion (Cu²⁺) whereas, in ferrous sulphate, iron is present as a ferrous ion (Fe²⁺). Hence, the above reaction can be alternatively represented as follows:

⇒ \(\mathrm{Cu}^{2+}+\mathrm{Fe} \rightarrow \mathrm{Cu}+\mathrm{Fe}^{2+}\)

Here, uncharged metallic iron (Fe) loses two electrons and becomes positively charged Fe²⁺ and these two electrons are gained by positively charged Cu²⁺ and are converted into uncharged metallic copper (Cu).

⇒ \(\mathrm{Fe}-2 \mathrm{e} \rightarrow \mathrm{Fe}^{2+}\)

⇒ \(\mathrm{Cu}^{2+}+2 \mathrm{e} \rightarrow \mathrm{Cu}\)

So, as per the electronic theory, Fe is oxidized and Cu²⁺ is reduced.

⇒ \(\mathrm{Fe}-2 \mathrm{e} \rightarrow \mathrm{Fe}^{2+}\) Loss of two electrons by Fe atom and formation of Fe²⁺ This is an oxidation reaction and Fe is the reducing agent.

⇒ \(\mathrm{Cu}^{2+}+2 \mathrm{e} \rightarrow \mathrm{Cu}\) Gain of two electrons by Cu²⁺ atom and formation of Cu.

This is a reduction reaction and From the above discussion we find that an oxidising agent (or oxidant) accepts electrons) and is reduced (i.e. converted to reduced species) while a reducing agent (or reductant) loses an electron(s) and is oxidised (i.e., converted to oxidised species)

 

Question 9. Show that oxidation and reduction take place simultaneously.
Answer:

Let us consider the following reaction.

We know that in the case of reduction removal of oxygen takes place and in the case of oxidation addition of oxygen takes place.

Here, oxygen is removed from cupric oxide, the reduction has occurred and oxygen is added to hydrogen to form water, so oxidation has occurred. It is also clear that both oxidation and reduction have occurred simultaneously.

Question 10. Using the electronic theory of oxidation and reduction, show with a suitable example that oxidation and reduction take place simultaneously.
Answer:

When some pieces of metallic zinc are dropped in an aqueous solution of copper sulphate, then after some time it is found that reddish-brown patches of metallic copper are deposited on the silver or grey-coloured pieces of zinc.

Further analysis confirms the presence of Zn²⁺ ions in the aqueous solution. The overall chemical reaction is,  Alternatively, we can write, \(\mathrm{CuSO}_4+\mathrm{Zn} \rightarrow \mathrm{ZnSO}_4+\mathrm{Cu}\) Here, Zn has lost two electrons, so it is oxidised and Cu²⁺ has gained two electrons, and it is reduced.

So, we find that in this chemical reaction, both oxidation and reduction have occurred simultaneously.

 

 

Question 11. For the following reactions, predict with an explanation, which substance has been oxidised and which substance has been reduced.

⇒ \(2 \mathrm{KI}+\mathrm{Cl}_2 \rightarrow 2 \mathrm{KCl}+\mathrm{I}_2\)

⇒ \(\mathrm{Fe}_2 \mathrm{O}_3+3 \mathrm{CO} \rightarrow 2 \mathrm{Fe}+3 \mathrm{CO}_2\)

⇒\(\mathrm{ZnO}+\mathrm{C} \rightarrow \mathrm{Zn}+\mathrm{CO}\)

⇒ \(2 \mathrm{SO}_2+\mathrm{O}_2 \rightarrow 2 \mathrm{SO}_3\)

⇒ \(\mathrm{FeO}+\mathrm{H}_2 \rightarrow \mathrm{Fe}+\mathrm{H}_2 \mathrm{O}\)

⇒ \(\mathrm{C}_2 \mathrm{H}_4+\mathrm{H}_2 \rightarrow \mathrm{C}_2 \mathrm{H}_6\)

⇒ \(2 \mathrm{FeCl}_3+\mathrm{H}_2 \mathrm{~S} \rightarrow 2 \mathrm{FeCl}_2+2 \mathrm{HCl}+\mathrm{S}\)

Answer:

 

Question 12. For the following reactions, predict which substance has been oxidised and which substance has been reduced. Also, identify the oxidant and reductant.

⇒ \(\mathrm{H}_2 \mathrm{~S}+2 \mathrm{FeCl}_3 \rightarrow 2 \mathrm{HCl}+\mathrm{S}+2 \mathrm{FeCl}_2\)

⇒ \(2 \mathrm{Na}+\mathrm{H}_2 \rightarrow 2 \mathrm{NaH}\)

⇒ \(\mathrm{Zn}+\mathrm{CuSO}_4 \rightarrow \mathrm{ZnSO}_4+\mathrm{Cu}\)

⇒ \(\mathrm{CH}_4+2 \mathrm{O}_2 \rightarrow \mathrm{CO}_2+2 \mathrm{H}_2 \mathrm{O}\)

⇒ \(\mathrm{FeO}+\mathrm{CO} \rightarrow \mathrm{Fe}+\mathrm{CO}_2\)

Answer:

 

Question 13. Prove by an example that oxidation and reduction are two complementary processes of electron loss and electron gain in a reaction.
Answer:

In an oxidation reaction, electrons are given up by an atom or ion and are thereby oxidised. These electrons can not escape out of the reaction medium or accumulate in it, but they should be consumed up or gained by some other atoms or ions that remain in the reaction medium. These latter atoms or ions will thereby be reduced.

An atom of magnesium is changed into an Mg2+ ion by the loss of two electrons: Mg- 2e →Mg²⁺. Again, an atom of oxygen is changed into an oxide ion (O_2-) by the gain of 2 electrons: O + 2e→O₂^–. So in this case, the magnesium atom is oxidised and the oxygen atom is reduced.

As a result, the first process is oxidation and the second one is reduction.

Question 14. Water is added to a test tube containing dry baking soda and dry crystals of tartaric acid. Again kerosene oil is added to another similar test tube containing an equal quantity of each baking soda and dry crystals of tartaric acid. What would be your observations in these two cases and why?
Answer:

We know that “like dissolves like”. Baking soda is not soluble in organic solvents like kerosene. The second test tube will show no reaction.

When water is added to the first test tube, an aqueous solution of baking soda and an aqueous solution of tartaric acid react together to evolve bubbles of carbon dioxide.

Since baking soda is soluble in water, hence the aqueous solution of the salt allows sufficient contact with the aqueous solution of another reactant for the chemical reaction to occur.

The presence of water as a solvent produces ions and these ions react between themselves to make the chemical reaction happen.