Chemistry

LEARNING OBJECTIVES

After reading this unit, you will be able to:

•explain the difference between a decomposition reaction and combination reaction.

•explain single displacement, double displacement (precipitation and neutralization) and combustion reactions.

•write and balance ionic equations.

KNOWLEDGE GAIN

                       

A firework involves many different chemical reactions occurring at the same time.

ACTIVITY 5.1: Illustrating Chemical Reactions

Consider the following situations of daily life and think what happens when:

•milk is left at room temperature on a hot day.

•an iron pan/nail is left exposed to humid atmosphere.

•grapes get fermented.

•food is cooked.

•food gets digested in our body

.•we respire.

Explanation

In all these situations, the nature and the identity of the initial substance have somewhat changed. We know that whenever a chemical change occurs, a chemical reaction has taken place.

In a chemical reaction, the reactants are converted into products. The conversion of reactants into products in a chemical reaction is often accompanied by some features which can be observed easily. The easily observable features (or changes) which take place as a result of chemical reactions are known as characteristics of chemical reactions. The important characteristics of chemical reactions are:

(i) Evolution of a gas

(ii) Formation of a precipitate

(iii) Change in color

(iv) Change in temperature

(v) Change in state

Any one of these general characteristics can tell us whether a chemical reaction has taken place or not. For example, if on mixing two substances, any of the above characteristics occurs, then we can say that a chemical reaction has taken place.

5.1 TYPES OF REACTIONS

There are millions of known chemical reactions. Many of these chemical reactions have common aspects, so they can be grouped into specific classification. The majority of chemical reactions (not all) fall into the following major categories:

•Combination reactions

•Decomposition reactions

•Single replacement reactions

•Double displacement reactions

•Combustion reactions

5.1.1 Combination Reactions

ACTIVITY 5.2: Combination of Iron and Sulphide (Combination Reaction)

Materials Required

Iron powder (filings) (7 grams), sulphur (4 grams), test tube, Bunsen burner, and pair of test-tube tongs.

Procedure

•Prepare a mixture containing iron powder and sulphur powder in the ratio 7:4 by mass.

•Note the appearance of the pure elements and the mixture.

•Take about 0.5 g of the mixture into a hard glass test tube.

•Heat the mixture at the base of the test tube—gently at first and then more strongly (use a blue flame throughout)

•Allow the test tube to cool down.

•Observe the product.

          

Can you write the chemical equation for this reaction?

Safety

•It is advisable to wear protective gloves and heat the mixture using tongs.

•Eye protection required.

In this activity, you will observe that when we start heating a mixture of iron filing and sulphur, the sulphur melts and reacts with the iron filings to form the compound iron (II) sulphide. In this reaction, two different elements combine to form a single product. This is an example of combination reaction.

Combination reactions are those reactions in which a single product is produced from two (or more) reactants. The general equation for a combination reaction involving two reactants is

        

The reactants X and Y can be elements or compounds, or a compound and an element. The product XY is always a compound.

Combination reactions may involve

•the combination of two elements to form a compound.

•the combination of a compound and an element to form a new compound.

•the combination of two compounds to form a new compound.

We will now discuss some examples of combination reactions:

Element + Element→Compound

Example 1

When iron powder is heated with sulphur, iron sulphide is formed.

                   

In this reaction, two elements, iron and sulphur are reacting together to form a single compound—iron sulphide. So it is a combination reaction.

Example 2

When sodium metal reacts with chlorine, sodium chloride is formed.

                   

In this reaction, two elements, sodium and chlorine, combine together to form a single compound—sodium chloride. So, this is a combination reaction.

Example 3

Hydrogen combines with chlorine to form hydrogen chloride.

                       

In this reaction, hydrogen and chlorine react together to form a single compound, hydrogen chloride gas. So, this is an example of combination reaction. This combination reaction is used in industry for the manufacture of hydro chloric acid (Hydrogen chloride gas on dissolving in water forms hydro chloric acid).

Compound + Element→Compound

Example 4

Carbon monoxide reacts with oxygen to form carbon dioxide.

                         

In this reaction, carbon monoxide compound reacts with oxygen to form a new compound carbon dioxide. So, this is a combination reaction.

Example 5

Phosphorus trichloride reacts with chlorine to form phosphorus pentachloride.

          

In this reaction, phosphorus trichloride reacts with chlorine to form a new compound phosphorus pentachloride. So, this is a combination reaction

.Compound + Compound→Compound

Example 6

Calcium oxide reacts with carbon dioxide to form calcium carbonate.

                 

In this reaction, calcium oxide (quicklime) and carbon dioxide combine together to produce a new compound calcium carbonate. So, this is a combination reaction.

Example 7

Ammonia reacts with hydrogen chloride to form ammonium chloride.

                               

In this reaction, two compounds, ammonia and hydrogen chloride, combine together to produce a new compound, ammonium chloride. So, this is a combination reaction.

EXPERIMENT 1

Aim

Combination of ammonia and hydrogen chloride.

Materials Required Ammonia solution, hydrochloric acid, glass rods (two), cotton buds, long glass tube, stopper (two) and clamp.

Procedure

• Fix a cotton bud on two glass rods.

• Mark the glass rods as A and B.

• Dip rod A into concentrated ammonia solution and rod B into concentrated hydro-chloric acid.

• Arrange all apparatuses as shown in diagram below.

                

Note: Insert the two stoppers at the end of glass tube (see diagram)

• Observe the glass tube after a few minutes.Write the balanced chemical equation for this reaction.

Safety

• Do not soak the cotton buds into chemicals with hand.

• The experiment must be done carefully in the presence of teacher.

• Make sure that the glass tube is clamped properly

Explanation

In the experiment, you will observe that when ammonia reacts with hydrochloric acid, a white solid appears inside the glass tube. This solid is ammonium chloride (NH4Cl). Here, two reactants react (or combine) to give one product. So this is a combination reaction.

EXERCISE 5.1

1. When iron powder is heated with sulphur, ______ is formed.This is an example of ______ .

2. Complete the reaction: H₂(g) + Cl₂ (g) →?

3. Calcium oxide reacts with sulphur dioxide to form calcium carbonate. (True or False)

4. Name the compound formed when ammonia reacts with hydrogen chloride.

5. Hydrogen chloride gas on dissolving in water forms ______ acid.

5.1.2 Decomposition Reactions

ACTIVITY 5.3: Decomposition of Silver Chloride

•Take about 2 g of silver chloride in a China dish.

•Note the color of silver chloride.

•Place the China dish in sunlight for 10–30 minutes

.•Observe the color of silver chloride after 30 minutes.

                        

In Activity 5.3, you will observe that white silver chloride turns grey in sunlight. This is due to the decomposition of silver chloride into silver and chlorine by light.

Reaction

                             

Decomposition reactions are those reactions in which a single reactant breaks down into two (or more) simpler substances (elements or compounds). The general equation for a decomposition reaction in which there are two products, is

The reactant XY is always a compound. The products X and Y may be elements or compounds.

In other words, decomposition reactions are opposite of combination reactions. These reactions often involve an energy source such as heat, light, or electricity which breaks apart the bonds of compounds.

The products of decomposition reactions may be

•two elements

•one (or more) elements and one (or more) compounds

•two (or more) compounds.

Some examples of decomposition reactions are:

Compound → Element+ Element

ACTIVITY 5.4: Electrolysis of Water

Take a plastic mug. Drill two holes at its base and fit rubber stoppers in these holes. Insert carbon electrodes in these rubber stoppers as shown in figure.

•Connect these electrodes to a 6-volt battery.

•Fill the mug with water such that the electrodes are immersed. Add a few drops of dilute sulphuric acid to the water.

•Take two test tubes filled with water and invert them over the two carbon electrodes.

•Switch on the current and leave the apparatus undisturbed for some time.

•You will observe the formation of bubbles at the electrodes. These bubbles displace water in the test tubes.

                 

•Is the volume of the gas collected the same in both test tubes?

•Once the test tubes are filled with the respective gases, remove them carefully.

•Test these gases one by one by bringing a burning candle close to the mouth of the test tubes.

Caution: This step must be performed carefully by the teacher.

•What happens in each case?

•Which gas is present in each test tube?

Example 1

When electricity is passed through acidified water, it decomposes into hydrogen and oxygen.

                

In this reaction, water splits up to form hydrogen and oxygen. This decomposition reaction takes place by the action of electricity. This reaction is called electrolysis of water.

Compound → Compound + Element

Example 2

In the presence of light, hydrogen peroxide decomposes into water and oxygen.

            

The rate of decomposition of hydrogen peroxide is increased by the catalyst MnO2.A 30% solution of hydrogen peroxide rapidly decomposes to oxygen and water when treated with manganese dioxide. This decomposition is exothermic.

               

Example 3

When Sodium nitrate is heated, it decomposes to produce sodium nitrite and oxygen. This reaction takes place at a temperature of 380–500°C.

                    

Example 4

When potassium chloride is heated in the presence of manganese dioxide catalyst, it decomposes to give potassium chloride and oxygen.

                 

In this reaction, potassium chlorate decomposes into potassium chloride and oxygen. So this is a decomposition reaction. In this reaction, catalyst manganese dioxide (MnO2) allows the decomposition to occur at a lower temperature. This reaction is used for preparing small amount of oxygen in the laboratory.

                    Compound → Compound + Compound

Example 5

When calcium carbonate is heated, it decomposes to give calcium oxide and carbon dioxide.

                   

In this reaction, calcium carbonate breaks up into two simpler compounds—calcium oxide and carbon dioxide. So, this is a decomposition reaction.

•Calcium carbonate is also called limestone.

•Calcium oxide is also called lime or quicklime.

•Calcium oxide obtained from the decomposition of calcium carbonate used in the manufacture of glass and cement.

ACTIVITY 5.5: Decomposition of Sodium Nitrate

•Take about 2 g sodium nitrate powder in a boiling tube.

•Hold the boiling tube with a pair of tongs and heat it over a flame, as shown in figure.

•What do you observe? Note down the change, if any.

                    

Explanation

You will observe a white or slightly yellow residue of sodium nitrite in the test tube.

Reaction

                           

ACTIVITY 5.6: Heating Ferrous Sulphate Crystals

•Take about 2 g ferrous sulphate crystals in a dry boiling tube.

•Note the color of the ferrous sulphate crystals.

•Heat the boiling tube over the flame of a burner or spirit lamp as shown in figure.

•Observe the color of the crystals after heating.Notice that the green color of the ferrous sulphate crystals changes.

                      

In Activity 5.6, when ferrous sulphate is heated strongly, it decomposes into ferric oxide, sulphur dioxide and sulphur trioxide.

                          

In this reaction, you can observe that a single reactant breaks down to give simpler products. So this is a decomposition reaction. Ferrous sulphate crystals (FeSO4⋅7H2O) lose water when heated and the color of the crystals changes. It then decomposes to ferric oxide (Fe2O3), sulphur dioxide (SO2) and sulphur trioxide (SO3). Ferric oxide is a solid, whereas SO2 and SO3 are gases.

EXPERIMENT 2

Aim

Decomposition of Hydrated Copper (II) Sulphate.

Materials Required

Copper sulphate crystals (2 g), evaporating dish, burner, tripod, balance, wire gauze.

Procedure

• Measure and note the weight of evaporating dish.

• Take 2 g copper sulphate in a evaporating dish.

• Measure and note the weight again.

• Fix all apparatus as shown in figure on page 103.

• Heat hydrated copper sulphate gently until the color disappears.

Explanation

The heat causes the hydrated copper sulphate (blue) to split into anhydrous copper sulphate and water. Anhydrous copper sulphate is white in color.If we add water to the anhydrous copper sulphate, the white powder becomes blue again.

        

Safety

• Wear goggles, do not look directly into the evaporating dish.

• Keep away from the burner.

• Do not touch the warm dish.

• This experiment must be performed carefully in the presence of teacher

When hydrated copper sulphate is heated, it decomposes into anhydrous copper sulphate and water.

                      

In this reaction, when the blue crystals of hydrated copper (II) sulphate are heated, water molecules evaporate leaving anhydrous copper sulphate.

We have seen that the decomposition reactions require energy either in the form of heat, light or electricity for breaking down the reactants.

Note: Reactions in which energy is absorbed are known as endothermic reactions.

The digestion of food in the body is an example of decomposition reaction. The major constituents of our food such as carbohydrates, fats, proteins, etc. decompose to give simple glucose and amino acids. The glucose further combines with oxygen to release large amount of energy which keeps our body working.

Note: Reactions in which energy is released are known as exothermic reactions.

EXERCISE 5.2

1. What will you observe when 0.5 g lead nitrate is heated?

2. The color of anhydrous copper sulphate is ______ .

3. The digestion of food in the body is an example of combination reaction.

(True or False)

4. Write a balanced chemical equation of the following:

(a) Silver bromide exposed to light.

(b) Calcium carbonate is heated.

(c) When electricity is passed through acidified water.

5. Name the catalyst which increases the rate of decomposition of H₂O₂

5.1.3 Single Replacement Reactions

ACTIVITY 5.7: Reaction of Iron Nails with Copper Sulphate

Take three iron nails and clean them by rubbing with sand paper.

•Take two test tubes marked as (A) and (B). In each test tube, take about 10 ml of copper sulphate solution.

•Tie two iron nails with a thread and immerse them carefully in the copper sulphate solution in test tube B for about 20 minutes [figure (a)]. (If the solution is concentrated, leave it for 2 hours) keep one iron nail aside for comparison.

•After 20 minutes, take out the iron nails from the copper sulphate solution.

•Compare the intensity of the blue color of copper sulphate solutions in test tubes (A) and (B), [figure (b)].

•Also, compare the color of the iron nails dipped in the copper sulphate solution with the one kept aside[figure (b)].

                    

Single replacement reactions are those reactions in which one element displaces another element from a compound. The general equation for this type of reaction is                               

In reactants, X is an element and YZ is a compound. In products, XZ is a compound and Y is an element.

In single replacement reactions, more active metals displace less active metals (or hydrogen) from their compounds. Some common elements are arranged in decreasing order of their ability to replace element (metal ion) in aqueous solution. This series is known as activity series.

  

Note: Potassium (K) is the most reactive metal and gold (Au) is the least reactive metal.

Example 1

When a piece of iron metal (say, an iron nail) is placed in copper sulphate solution, then iron sulphate solution and copper metal are formed.

      

In this reaction, iron displaces copper from copper sulphate solution. The deep blue color of copper sulphate solution fades due to the formation of light green solution of iron sulphate. A red-brown coating (or layer) of copper metal is formed on the surface of iron metal (or iron nail).

Note that this displacement reaction occurs because iron is more reactive than copper.

Example 2

When a copper strip is placed in a solution of silver nitrate, then copper nitrate solution and silver metal are formed.

       

In this case, copper displaces silver from silver nitrate compound. This displacement reaction occurs because copper is more reactive than silver. A shining grayish white deposit of silver is formed on the copper strip and the solution becomes blue due to the formation of copper nitrate.

Example 3

When a strip of zinc metal is placed in copper sulphate solution, then zinc sulphate solution and copper are obtained.

       

In this reaction, zinc displaces copper from copper sulphate compound so that copper is set free (or liberated). The blue color of copper sulphate solution fades due to the formation of zinc sulphate (which is colorless). A red-brown deposit of copper metal is formed on the zinc strip.Note that this displacement reaction takes place because zinc is more reactive than copper.

Example 4

When a piece of magnesium metal is placed in copper sulphate solution, then magnesium sulphate solution and copper metal are formed.

             

In this reaction, magnesium displaces copper from copper sulphate solution. The blue color of copper sulphate solution fades due to the formation of colorless solution of magnesium sulphate. A red-brown deposit of copper metal is formed on the magnesium piece. Here, magnesium is able to displace copper from copper sulphate solution because magnesium is more reactive than copper

Example 5

When a strip of lead metal is placed in a solution of copper nitrate, then lead nitrate solution and copper metal are formed.

                   

In this case, lead displaces copper from copper nitrate solution. The green color of copper nitrate solution fades due to the formation of colorless solution of lead nitrate. A red-brown layer of copper metal is deposited on the lead strip. Please note that lead is able to displace copper from copper nitrate solution because lead is more reactive than copper.Another point to be noted is that copper used in this reaction is actually copper (II) nitrate

Example 6

When copper oxide is heated with magnesium powder, then magnesium oxide and copper are formed.

                

This is a displacement reaction. In this displacement reaction, a more reactive metal, magnesium, is displacing a less reactive metal, copper, from its oxide, copper oxide.

Example 7

When iron (III) oxide is heated with aluminium powder, then aluminium oxide and iron metal are formed.

    

In this displacement reaction, a more reactive metal, (aluminium), is displacing a less reactive metal (iron), from its oxide, iron (III) oxide.

All the above examples of displacement reactions are ‘single displacement reactions’. This is because in all these reactions only ‘one element’ displaces ‘another element’ from its compound.

EXERCISE 5.3

1. Copper is more reactive than silver.

                                                                                                (True or False)

2. Why does blue color of copper sulphate solution fade when a piece of iron object is placed into it?

3. When a copper strip is placed in a solution of zinc sulphate solution, copper sulphate is formed.

                                                                                                        (True or False)

4. Write a balanced chemical reaction of the following:

(a) When copper oxide is heated with magnesium powder.

(b) When iron oxide is heated with aluminium powder.

5. When a strip of lead metal is placed in a solution of ______ , lead nitrate and ______ are formed.

5.1.4 Double Displacement Reactions

ACTIVITY 5.8: Formation of Precipitate

•Take about 3 ml of sodium sulphate solution in a test tube.

•In another test tube, take about 3 ml of barium chloride solution.

•Mix the two solutions (Figure).

        

•What do you observe?

In Activity 5.8, you will observe that a white substance, which is insoluble in water, is formed. This insoluble substance formed is known as a precipitate. Any reaction that produces a precipitate can be called a precipitation reaction.

Reaction

Double displacement reactions are those reactions in which two compounds react by exchange of ions to form two new compounds. These reactions are also known as metathesis reactions or exchange reactions.

The general equation for a double displacement reaction is

where A and B are positive ions (cations).X and Y are negative ions (anions).

Another example of double displacement reaction is when we mix a solution of silver nitrate (AgNO₃) and sodium chloride, silver chloride and sodium nitrate is formed.

    

Double displacement reactions result in the removal of ions from the solutions. The removal of ions can occur in the following two ways:

•By precipitation reaction.

Definition: When we mix two solutions both containing a soluble compound, the mixture immediately becomes cloudy by the formation of insoluble precipitate. The reaction which produces a precipitate is called precipitation reaction. Precipitation reactions can be used to test the presence of metals.

•By neutralization reaction.

Definition: When solutions of strong acids and bases are mixed, the H⁺ (aq) from the acid combines with the OH^– (aq) from the base to form water. The reaction of an acid and a base is known as a neutralization reaction.

Precipitation Reaction

ACTIVITY 5.9: Formation of Silver Iodide

Take silver nitrate solution in a test tube and add potassium iodide solution.What do you observe?

Explanation

You will observe that a yellow precipitate is formed at once. The yellow precipitate is of silver iodide which is formed as a result of double displacement reaction.

      

Some examples of precipitation reaction:

Example 1

If barium chloride solution is added to copper sulphate solution, then a white precipitate of barium sulphate is produced along with copper chloride solution:

        

In this double displacement reaction, two compounds, barium chloride and copper sulphate, react by an exchange of their ions to form two new compounds, barium sulphate and copper chloride.

Example 2

When hydrogen sulphide gas is passed through copper sulphate solution, then a black precipitate of copper sulphide is formed along with sulphuric acid solution:

             

In this double displacement reaction, two compounds, copper sulphate and hydrogen sulphide react by an exchange of ions to form two new compounds, copper sulphide and sulphuric acid.

Example 3

When ammonium hydroxide solution is added to aluminium chloride solution, then a white precipitate of aluminium hydroxide is formed along with ammonium chloride solution:

               

In this double displacement reaction, two compounds, aluminium chloride and ammonium hydroxide, react by an exchange of their ions to form two new compounds, aluminium hydroxide and ammonium chloride.

Example 4

When potassium iodide solution is added to lead nitrate solution, then a yellow precipitate of lead iodide is produced along with potassium nitrate solution:

             

In this double displacement reaction, two compounds, lead nitrate and potassium iodide, react by an exchange of ions to form two new compounds, lead iodide and potassium nitrate. Note that lead nitrate, Pb(NO₃)₂, is also written as lead (II) nitrate.

Neutralization Reaction

ACTIVITY 5.10: Illustrating Neutralization Reaction

•Take about 2 ml of dilute NaOH solution in a test tube.

•Add 2–3 drops of phenolphthalein solution.What is the color of the solution?

•Add dilute HCl solution to the pink solution drop by drop.What do you observe?

•Now add a few drops of NaOH solution and again observe if there is any change in color.

Observations

The color of the solution becomes dark pink.

It is observed that when sufficient quantity of HCl solution has been added, the pink color of the solution disappears.

In Activity 5.10, you will observe that the pink color reappears.

Explanation

Phenolphthalein has pink color in basic medium. When sufficiently large quantity of HCl solution is added to the NaOH solution, the medium becomes acidic due to the presence of excess acid. In acidic medium, the phenolphthalein is colorless and hence the solution becomes colorless. When we again add NaOH solution, it neutralizes the excess acid and the solution again becomes alkaline. Therefore, the pink color reappears.In neutralization reaction, a salt is formed from the cation of the base and the anion of the acid. For example, sodium hydroxide and hydrochloric acid react to form sodium chloride and water.

                

In this double displacement reaction, two compounds, sodium hydroxide and hydrochloric acid, react by an exchange of ions to form two new compounds, sodium chloride and water. Note that no precipitate is formed in this double displacement reaction (This is because sodium chloride is soluble in water).

Some examples of neutralization reactions are:

           

Some applications of neutralization reaction

1. A person suffering from hyper acidity is advised to take antacid tablets or antacid suspension. Antacid preparations contain magnesium hydroxide as the active component which neutralizes the excess acid present in the stomach.

                

2. In acidic soils, slaked lime is added to reduce acidity.

                   

3. The sting of ants and bees contains formic acid. It is neutralized by rubbing soap or dilute ammonia solution.

4. The sting of yellow wasps contains an alkali. It is neutralized by rubbing dilute acetic acid.

EXERCISE 5.4

1. The insoluble substance formed during a chemical reaction is called ______.

2. The reaction of an acid and a base is known as ______ .

3. Precipitation reaction can be used to test the presence of metals. (True or False)

4. What happens when barium chloride solution is added to copper sulphate solution?

5. Complete the following equation:

                         

6. Write the general equation for a double displacement reaction.

5.1.5 Combustion Reactions

ACTIVITY 5.11: Burning of Magnesium Ribbon

•Take a piece of magnesium ribbon and hold it with a pair of tongs. Light magnesium ribbon.

•The magnesium ribbon starts burning with a dazzling flame.

Explanation

You would observe that magnesium ribbon soon changes into white powder. This white powdery substance is magnesium oxide which is formed as a result of combination reaction.

Reaction

                       

This type of reaction refers to the reaction of an element or compound with oxygen. Combustion usually releases a lot of heat energy. It is also referred to as burning.

Note:All combustion reactions are exothermic reactions.

Elements on Combustion

When elements undergo combustion, generally only one product is formed.

For example,

•Burning of coal

•Formation of water

•Burning of magnesium

                             

Compounds on Combustion

When compounds undergo combustion, two or more products are formed. When carbon-hydrogen (hydrocarbon) or carbon-hydrogen-oxygen compounds undergo combustion in an excess of oxygen, the products are carbon dioxide and water. For example,

•Burning of natural gas

•Glucose reacts with oxygen

When any hydrocarbon burns in insufficient oxygen, the products are carbon monoxide and water.

For example,

EXERCISE 5.5

1. All combustion reactions are exothermic.

(True or False)

2. What happens when glucose reacts with oxygen?

3. When any hydrocarbon burns in insufficient oxygen, the products are ______ and _______ .

4. Combustion reaction is also referred to as ______ .

5. Complete the reaction

                        

5.2CLASSIFICATION OF CHEMICAL REACTIONS AS ENDOTHERMIC AND EXOTHERMIC REACTIONS

During a chemical reaction, energy changes from one type to another. Energy is usually transferred either to or from the surroundings.

Depending upon the evolution or absorption of energy, the chemical reactions can be classified into two types: exothermic and endothermic.

5.2.1 Exothermic Reactions

The chemical reactions which proceed with the evolution of heat energy are called exothermic reactions.

The heat energy produced during the reaction is indicated by writing +q or more precisely by giving the actual numerical value along with products. In general, exothermic reactions may be represented as:

                                             

The heat evolved is expressed in the units of Joules (J) or kilo Joules (kJ) .For example, when one mole of carbon (coal) is burnt in oxygen, there is an evolution of 393.5 kJ of energy. It can be expressed as:

Some more examples of exothermic reactions are:

                                    

It may be mentioned here that:

•All combustion reactions are exothermic.

•Most of the reactions are exothermic.

•During exothermic reactions, a part of the potential energy possessed by the reactants is released. Hence, products of an exothermic reaction have less potential energy than the reactants as illustrated.

5.2.2 Endothermic Reactions

The chemical reactions which proceed with the absorption of heat energy are called endothermic reactions.

The heat energy absorbed during the reaction can be indicated by writing +q (or the actual numerical value) with the reactants. It can be indicated by writing –q (or the actual numerical value) with the products. In general, an endothermic reaction can be represented as:

A + B + q (heat) → C + D

A + B → C + D – q (heat)

where q is the heat absorbed.For example, formation of nitric oxide from nitrogen and oxygen proceeds with the absorption of 180.5 kJ of heat. It can be represented as:

                             

Some more examples of endothermic reactions are:

                                

           

It may be mentioned here that:

•Decomposition reactions are generally endothermic.

•The number of endothermic reactions is much less than the exothermic reactions.

•During endothermic reactions, reactants gain energy. Hence, products of an endothermic reaction have more potential energy than the reactants.

5.2.3 Explanation for the Energy Changes during Chemical Reactions

A chemical reaction involves the rearrangement of atoms. During the reactions certain bonds are broken while certain new bonds are formed between the atoms. Energy is absorbed for breaking the bonds and released during formation of bonds. If the energy required to break the bonds is more than the energy released during formation of bonds then there is net absorption of energy and the reaction is endothermic. On the other hand, if the energy released during formation of bonds is more than the energy absorbed for breaking the bonds, then there is net release of energy and the reaction is exothermic.

For example, let us consider the following reaction:

            

In this reaction, energy is required to break H—H and Cl—Cl bonds and is released during formation of two H—Cl bonds

Bond energy of H—H bond = 433 kJ/mole

Bond energy of Cl—Cl bond = 242 kJ/mole

Bond energy of H—Cl bond = 430 kJ/mole

Energy required to break two bonds

       = Bond energy of H—H bond + Bond energy of Cl—Cl bond

       = 433 + 242 = 675 kJ

Energy released

= 2 × Bond energy of H—Cl bond = 2 × 430 = 860 k

Since energy released during the reaction is more than the energy absorbed, there would be net release of energy and thus the reaction would be exothermic.

Net energy released = 860 – 675 = 185 kJ.

EXPERIMENT 3

Aim

To differentiate between exothermic and endothermic reactions.

Apparatus and equipment (per group)

•Beaker

•Thermometer

Chemicals (per group)

•Sodium hydroxide solution

•Dilute hydrochloric acid

•Sodium hydrogen carbonate solution

•Four spatula measures of citric acid

•Copper (II) sulphate solution

•Four spatula measures of magnesium powder (Highly flammable)

•3 cm Magnesium ribbon (Highly flammable)

•Dilute sulphuric acid

Introduction

Some reactions give out heat and others take in heat. In exothermic reactions, the temperature goes up, in endothermic reactions the temperature goes down. In this experiment, various reactions are examined. Temperatures are measured to decide whether a particular reaction is exothermic or endothermic.

Procedure

1. Use the apparatus as shown in figure.

2. Put 10 cm³ of sodium hydroxide solution in the beaker, record the temperature then add 10 cm³ of dilute hydrochloric acid, stirring with the thermometer. Record the maximum or minimum temperature.

3. Repeat the procedure for the following reactions:

(a) sodium hydrogen carbonate solution and citric acid;

(b) copper(II) sulphate solution and magnesium powder; and

(c) dilute sulphuric acid and magnesium ribbon.

                          

What to record

   

Safety

1. Wear eye protection. Some of the solutions are irritant.

2. Use the thermometer properly to measure changes in temperature.

EXERCISE 5.6

1. Distinguish between exothermic and endothermic reactions.

2. Energy is absorbed for ______ the bonds and released during the ______ of bonds.

3. Decomposition reactions are generally exothermic. (True or False)

4. Give an example of exothermic reaction.

5. Give an example of endothermic reaction.

5.3 IONIC EQUATIONS

Most of the equations that we have studied thus far are condensed equations. In condensed equation, all reactants and products are written as electrically neutral compounds and molecules. The reactions which occur in aqueous solution can also be represented by ionic equations. Generally, all single-replacement reactions and double-displacement reactions occur in aqueous solution.

           •In single replacement reaction, only cations are involved.

           •In double displacement reaction, both cations and anions are involved.

Let us consider an example.

The reaction of aqueous barium chloride with aqueous sodium sulphate

                        

In this reaction, three out of four components are strong electrolytes. Therefore a better description of this reaction on the atomic or molecular level might be given by an ionic equation.

Ionic equations are written by assuming that strong electrolytes dissociate in aqueous solution into corresponding ions. When the cations and anions of a compound in solution are shown separately, the equation is known as ionic equation. The total ionic equation for the reaction between BaCl₂ and Na₂SO₄ would be written as

               

Notice that in this equation, Na+ ions and Cl– ions appear on both sides of the ionic equation. These ions remain unchanged during the chemical reaction. Ions that are in an identical state on both sides of ionic equations are called spectator ions. If spectator ions are subtracted from both sides of the equation, the remaining equation is known as the net ionic equation. The net ionic equation focuses only on the species that have undergone a change in the reaction. The net ionic equation for the reaction between BaCl₂ and Na₂SO₄ is

                                

Net ionic equations contain all information necessary to understand the complete reaction.

5.4 RULES FOR WRITING IONIC EQUATIONS

The rules for writing ionic equations are as follows:

Rule 1:Write the full equation and balance it. Then write another equation using rules 2, 3 and 4.

Rule 2: For dissolved ionic substances, write the ions separately

Rule 3: For all solids (whether ionic or not), all liquids and all gases, write the full formula.

Rule 4: Cross out all the ‘spectator’ ions, i.e. those that appear separately on both sides of the equation.

Example 1

Write ionic equation for the reaction of magnesium with dilute hydrochloric acid.Solution

Rule 1:

Balance the equation:

  

Rule 2 and 3: Split dissolved ionic substances into separate ions:

          

Rule 4: Cancel the spectator ions (in bold), which are the 2Cl^– ions on each side, to give:

                      

Example 2

Write ionic equation for the reaction between copper(II) sulfate and sodium hydroxide.

Solution

Rule 1 gives:

               

From this full equation, the application of Rules 2 and 3 results in:

                     

he spectator ions are in bold type. Apply Rule 4 by crossing out the SO₄ ^2– ion on the left with the SO₄ ^2– ion on the right and the 2Na⁺ ions on the left with the 2Na⁺ ions on the right. The net ionic equation is:

              

The most difficult task in writing net ionic equations is determining which components of the reaction are strong electrolytes and should be written as ions. Strong electrolytes include soluble ionic compounds. Solubility rules for ionic compounds are given in Table 5.1. Some examples of strong acids are

   

Example 3

Write the condensed, ionic and net ionic equations when metallic copper is added to a solution of silver nitrate.

Solution

Example 4

Write the condensed, ionic and net ionic equations when aqueous hydrochloric acid reacts with aqueous sodium hydroxide.

Solution

Example 5

Write the molecular equation (condensed), ionic equation and net ionic equation when a strip of zinc metal is dipped into copper sulphate solution.

Solution

Example 6

Write the condensed ionic, and net ionic equations that describe the following chemical reactions.Aqueous solutions of sodium chromate and lead(II) nitrate react to form a yellow precipitate of lead(II)chromate and an aqueous solution of sodium nitrate.

Solution

  

Example 7

Write molecular, ionic and net ionic equations when aqueous silver nitrate reacts with copper(II) chloride.

Solution

 

EXERCISE 5.7

1. The reactions which occur in aqueous solution can also be represented by ______.

2. What do you mean by net ionic equation?

3. In double displacement reaction, both cations and anions are involved. (True or False)

4. Write net ionic equation when metallic copper is added to a solution of silver nitrate.

5. Write the net ionic equation for all reactions of strong acids with strong bases that form salt and water.

        5.5 SUMMARY

•A complete chemical equation represents the reactants, products and their physical states symbolically.

•In a combination reaction two or more substances combine to form a new single substance.

•Decomposition reactions are opposite to combination reactions. In a decomposition reaction, a single substance decomposes to give two or more substances.

•Decomposition reactions are usually endothermic because energy is required to break the bonds present in reactants.

•When an element displaces another element from its compound, a displacement reaction occurs.

•Two different atoms or groups of atoms (ions) are exchanged in double displacement reactions.

•Precipitation reactions produce insoluble salts.

•The reaction between an acid and a base to give a salt and water is known as neutralization reaction.

•Reactions in which heat is given out along with the products are called exothermic reactions.

•Reactions in which energy is absorbed are known as endothermic reactions.

•In single-replacement reaction only cations are involved.

•In double-displacement reaction both cations and anions are involved.

•When cations and anions of a compound are shown separately, in a chemical equation, the equation is called ionic equation.

•The ions remained unchanged during the chemical reaction and present on both sides of an ionic equation is called spectator ions.

•If spectator ions are subtracted from both sides of ionic equation, the remaining equation is called net ionic equation.

      5.6 GLOSSARY

•Absorption: a chemical or physical process by which one thing takes in or soaks up energy or a liquid or other substance from another

•Alkaline: having a pH greater than 7

•Anhydrous: containing no water especially a crystalline compound•Antacid: a medicine that prevent or correct acidity; especially in the stomach

•Burning: on fire

•Catalyst: a substance that increases the rate of a chemical reaction without taking part in the reaction •Combustion: the process of burning something

•Electrolysis: a chemical decomposition produced by passing an electric current through a liquid or solution containing ions.

•Electrolytes: a liquid which contains ions and can be decomposed by electrolysis

•Endothermic: a chemical reaction in which heat energy is absorbed

•Exothermic: a chemical reaction in which heat energy is released

•Fermentation: the chemical breakdown of a substance by bacteria, yeasts, or other micro-organisms

•Flammable: easily set on fire

•Hydrated: combine chemically with water molecules, compounds containing water molecules

•Hydrocarbon: a compound of hydrogen and carbon

•Immersed: dip or submerge in a liquid•Insufficient: not enough; inadequate

•Liberated: to release gas during a chemical reaction

•Neutralize: to make an acidic or alkaline substance chemically neutral

•Precipitate: a solid substance deposited from a solution after a chemical reaction

5.7 UNIT ASSESSMENT

I. Multiple Choice Questions

1. Which of the following is/are characteristics of chemical reactions?

(a) Change in color (b) Evolution of gas (c) Formation of precipitate (d) All of these

2. When a single product is produced from two or more reactants, the reaction is

(a) Metathesis reaction (b) Decomposition reaction (c) Combination reaction (d) Displacement reaction

3. Combination reactions may involve

(a) Combination of two elements (b) Combination of two compounds (c) Combination of one element and one compound (d) All of the above

4. Electrolysis of sodium chloride is an example of

(a) Combination reactions (b) Decomposition reactions (c) Exchange reactions (d) None of these

5. Name the reaction in which energy in the form of heat, light and electricity is required to complete the reaction.

(a) Combination (b) Decomposition (c) Single replacement (d) Double replacement

6. In which reaction, more active metals displace less active metals?

(a) Combustion reaction (b) Exchange reaction (c) Single replacement reaction (d) Decomposition reaction

7. Which of the following statement(s) is/are not true?

(a) Calcium is less reactive than copper(b) Aluminium is more reactive than sodium (c) Iron is more reactive than zinc (d) All of these

8. Choose the correct statement(s).

(a) Magnesium is able to displace copper from copper sulphate(b) Zinc cannot displace copper form copper sulphate (c) Iron cannot displace copper from copper sulphate (d) Silver is able to displace copper from copper sulphate

9. When hydrocarbon compounds undergo combustion in excess of oxygen, the products are

(a) Carbon dioxide and carbon monoxide (b) Carbon monoxide and water (c) Carbon dioxide and water (d) Either (b) or (c)

10. Which of the following is a net ionic equation?

         

II. Open Ended Questions

1. Translate the following statements into chemical equations and balance them.

(a) Hydrogen gas combines with chlorine gas to form hydrogen chloride.

(b) Calcium oxide reacts with carbon dioxide to form calcium carbonate.

(c) Ammonia reacts with hydrogen chloride to form ammonium chloride.

2. Identify the type of reaction in each case.

(a) Hydrogen gas combines with nitrogen to give ammonia.

(b) On heating, calcium carbonate breaks up into calcium oxide and carbon dioxide.

(c) Hydrochloric acid reacts with sodium hydroxide to give salt and water.

3. Complete the following reactions:

      

4. What do you mean by exothermic and endothermic reaction?

5. A shining brown colored element X on heating in air becomes black. Name the element X and the black colored compound.

6. State the important use of decomposition reaction.

7. What happens when

(a) a piece of iron metal is placed in copper sulphate solution?

(b) a strip of copper is dipped in a solution of silver nitrate?

(c) copper oxide is heated with magnesium?

(d) iron powder is heated with sulphur?

(e) silver metal is placed in copper sulphate solution?

8. A colorless lead salt produces a yellow residue and brown fumes when heated

.(a) Name the brown fumes

(b) Name the yellow compound

(c) Name the lead salt

9. How can you say that “A decomposition reaction is opposite of a combination reaction”?

10. Distinguish between exothermic and endothermic reactions. Write a general equation for each of the reactions.

11. Write net ionic equations for the reaction of

(a) magnesium with dilute hydrochloric acid

(b) zinc with copper sulphate

(c) aqueous hydrochloric acid and aqueous sodium hydroxide

  12.

III. Practical-based Questions

1. Which statement is true for the following reaction?

A + BC =B + AC

(a) A is less reactive than B

(b) A is more reactive than C

(c) A is more reactive than B

(d) B is more reactive than A

2. What happens when an iron knife is dipped into aqueous solution of copper sulphate?

               

(a) Copper gets collected on the surface of iron knife

(b) Iron gets dissolved in copper sulphate

(c) No reaction takes place

(d) All options are wrong

3. The following reaction is an example of:

AB + CD =CB + AD

(a) Combination reaction

(b) Replacement reaction

(c) Double displacement reaction

(d) Decomposition reaction

4. The following figure illustrates

   

(a) Electrolysis of Sodium chloride

(b) Electrolysis of Water

(c) Electrolysis of Aluminium

(d) Electrolysis of Copper

5. When magnesium ribbon burns, the product formed is

        

              (a) Acidic              (b) Basic             (c) Amphoteric               (d) Neutral