ICSE Class 9 Chemistry Chapter 02 Chemical Changes and Reactions

Chemical Changes And Reactions

Syllabus

Types of chemical changes: Direct combination, decomposition, displacement, double decomposition. (The above to be taught with suitable chemical equations as examples).

Energy changes in a chemical change: Exothermic and endothermic reactions with examples - evolution/absorption of heat, light and electricity.

2.1 Chemical Reaction

A chemical reaction is the process of breaking the chemical bonds of the reacting substances (reactants) and making new bonds to form new substances (products).

A chemical bond is the force that holds the atoms of a molecule together, as in a compound.

A chemical change or chemical reaction occurs when particles collide. Collisions occur when reactants are in close contact or by supply of energy. Thus, one or more of the following conditions are necessary for a chemical change:

(i) Mixing (close contact)

In some cases, a chemical reaction occurs only when two substances are mixed in their solid states.

Example 1: Iodine and phosphorus react explosively when brought into close contact.

Example 2: Lead nitrate (white) and potassium iodide (white) react to make lead iodide (yellow).

Pb(NO₃)₂ (s) + 2KI (s) → 2KNO₃ (s) + PbI₂ (s)

(ii) Solution

In some cases, a chemical reaction occurs when substances are mixed in, i.e. molten or aqueous state.

Example 1: Oxalic acid crystals and sodium carbonate react in water solution only.

Example 2: Sodium chloride and silver nitrate also react in a solution state to form the precipitate of silver chloride and sodium nitrate.

NaCl (aq) + AgNO₃ (aq) → AgCl + NaNO₃ (aq) white ppt.

(iii) Heat

Some chemical reactions occur only on heating.

Example 1: Copper carbonate decomposes on heating (delta symbol for heating) into copper oxide and carbon dioxide.

CuCO₃(s) → CuO (s) + CO₂(g)

Example 2: Lead nitrate decomposes on heating leaving yellow residue lead monoxide, brown gas nitrogen dioxide and colourless gas oxygen.

2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂

(iv) Light

Some chemical reactions take place by the action of light. These are called photochemical reactions or photolysis. Molecules of the reactants absorb light energy to get activated, and then react rapidly.

Example 1: Plants form glucose from carbon dioxide and water in the presence of light.

6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂

Example 2: Hydrogen and chlorine react in the presence of sunlight.

H₂ + Cl₂ → 2HCl

Example 3: If chlorine water is exposed to sunlight, oxygen is evolved.

Cl₂ + H₂O → HCl + HClO

2HClO → Sunlight/UV radiation → 2HCl + O₂

Note: Solutions of silver nitrate, hydrogen peroxide are kept in brown bottles in the laboratory because they decompose in the presence of light.

Example 4: Silver nitrate decomposes in the presence of light.

2AgNO₃ → sun light → 2Ag + 2NO₂ + O₂

(v) Electricity

Certain chemical reactions, like the decomposition of certain compounds, occur only when electricity is passed through the substance.

Example 1: Electrolysis of acidulated water occurs only in the presence of electricity, to give hydrogen and oxygen.

2H₂O → electricity → 2H₂ + O₂

Example 2: On passing current through molten sodium chloride, sodium and chlorine are separately obtained.

2NaCl → electricity → 2Na + Cl₂

(vi) Pressure

Some chemical reactions take place only when the involved substances are subjected to high pressure.

Example 1: Mercuric chloride and potassium iodide when rubbed in a mortar, give a scarlet-coloured substance called mercuric iodide.

Example 2: Nitrogen and hydrogen, when subjected to high pressure, produce ammonia (in the presence of iron as catalyst)

N₂ + 3H₂ ⇌ 2NH₃ above 200 atm

(vii) Catalyst

Some chemical reactions need a catalyst to accelerate or decelerate the rate at which they occur. The catalysts themselves do not take part in the reaction.

Example 1: Potassium chlorate decomposes only at 700°C, and even then the rate of release of oxygen is very slow. But when potassium chlorate is heated in the presence of manganese dioxide, decomposition begins at much lower temperature, 300°C, and manganese dioxide remains unaffected. Thus, in this reaction, manganese dioxide acts as a catalyst.

2KClO₃ → MnO₂ → 2KCl + 3O₂

Example 2: Ammonia reacts with oxygen to produce nitric oxide and water vapour in the presence of (platinum) a catalyst.

4NH₃ + 5O₂ → Pt 800°C → 4NO + 6H₂O

(a) Positive catalyst: When a catalyst accelerates a reaction, it is known as a positive catalyst.

For example, the rate of decomposition of hydrogen peroxide gets increased in the presence of manganese dioxide.

2H₂O₂ → MnO₂ → 2H₂O + O₂

Promoters: Substances that influence the rate of a chemical reaction by improving the efficiency of the catalyst are called promoters. For example, in the manufacture of ammonia, iron acts as a catalyst and molybdenum as a promoter. Molybdenum increases the efficiency of the catalyst iron.

iron + N₂ + 3H₂ ⇌ 2NH₃ molybdenum

(b) Negative catalyst: A catalyst employed to retard a reaction is known as a negative catalyst.

Example 1: Phosphoric acid retards the rate of decomposition of hydrogen peroxide.

Example 2: The rate of oxidation of chloroform decreases in the presence of alcohol.

Note: Certain chemical reactions proceed by absorption of sound energy. For example, Acetylene breaks up into carbon and hydrogen by absorbing sound energy

C₂H₂ → Sound energy → 2C + H₂

2.2 Characteristics Of Chemical Reactions

Certain chemical reactions are characterized by changes that are quite easily observed. Some of these typical changes are:

(i) Evolution of gas

In many chemical reactions, one of the products is a gas.

Examples:

(a) When zinc reacts with dilute sulphuric acid, hydrogen gas is evolved, with an effervescence.

Zn + H₂SO₄ → ZnSO₄ + H₂ [zinc] [dil. sulphuric acid] [zinc sulphate] [hydrogen]

Effervescence: The formation of gas bubbles in a liquid during a reaction is called effervescence. In the reaction given above, one of the reactants is a liquid. In such cases, i.e. when one of the reactants is a liquid, the gas produced forms bubbles in the liquid, i.e. effervescence takes place.

(ii) Change of colour

Certain chemical reactions are characterized by a change in the colour of the reactants.

Example: When a few pieces of iron are dropped into a blue coloured copper sulphate solution, the blue colour of the solution fades and eventually turns into light green due to the formation of ferrous sulphate.

Fe + CuSO₄(aq) → FeSO₄ + Cu [iron] [blue solution] [green] [copper] (red deposit)

(iii) Formation of precipitates

Certain chemical reactions are characterized by the formation of insoluble solid substances called precipitates.

Examples:

(a) When a solution of silver nitrate is added to a solution of sodium chloride, a white insoluble substance (precipitate), silver chloride, is formed.

AgNO₃(aq) + NaCl(aq) → AgCl(ppt) + NaNO₃(aq) [silver nitrate solution] [sodium chloride solution] [silver chloride (white ppt)] [sodium nitrate solution]

(b) When ferrous sulphate solution is added to sodium hydroxide solution, a dirty green precipitate of ferrous hydroxide is formed.

FeSO₄(aq) + 2NaOH(aq) → Fe(OH)₂ + Na₂SO₄(aq) [ferrous sulphate] [sodium hydroxide] [ferrous hydroxide] (dirty green ppt) [sodium sulphate]

(iv) Change of state

In many chemical reactions, a change of state is observed. The reaction might start with gaseous or liquid reactants and end up with solid products, and vice versa.

Example: Ammonia gas reacts with hydrogen chloride gas to produce solid ammonium chloride.

NH₃(g) + HCl (g) ⇌ NH₄Cl(s)

Exercise 2(A)

Questions

1. (a) What is a chemical reaction?

(b) State the conditions necessary for a chemical change or reaction.

2. Define the following terms

(a) Chemical bond

(b) Effervescence

(c) Precipitate

3. Give an example of a reaction where the following are involved

(a) Heat (b) Light (c) Close contact (d) Electricity (e) Solution (f) Pressure (g) Catalyst

4. Define:

(a) Photochemical reaction

(b) Electrochemical reaction. Give one example in each case.

5. Give an example of each of the following chemical changes.

(a) A photochemical reaction involving (i) silver salt (ii) water

(b) A reaction involving (i) blue solution (ii) formation of dirty green precipitate

(c) Two gases combine to form white solid.

(d) A reaction where colour change is noticed.

6. Write the chemical reaction where the following changes are observed.

(a) Gas is evolved

(b) Colour change is noticed

(c) Precipitate is formed

(d) Physical state is changed

7. Give reason for the following:

(a) Silver nitrate solution is kept in coloured bottles.

(b) Molybdenum is used in the manufacture of ammonia.

(c) Blue solution of copper sulphate changes to green when a piece of iron is added to this solution.

Teacher's Note

Chemical reactions occur all around us daily - from cooking food to our bodies digesting meals, understanding what causes reactions helps us appreciate everyday processes.

2.3 Types Of Chemical Change Or Chemical Reaction

Types

1. Direct combination (or synthesis)

2. Decomposition

3. Displacement

4. Double decomposition

1. Direct combination or synthesis

A reaction in which two or more substances combine together to form a single substance is called a combination reaction.

A + B → AB

In the above reaction, a combination of substances A and B takes place to give a molecule of a new substance, AB.

In combination reactions:

(i) two elements combine to form a compound [this reaction is also called synthesis].

Carbon burns in oxygen to form a gaseous compound, carbon dioxide.

C + O₂ → heat → CO₂ + Heat [carbon] [oxygen] [carbon dioxide]

When magnesium is burnt, it combines directly with the oxygen of air to form magnesium oxide.

2Mg + O₂ → 2MgO

(ii) an element and a compound combine to give a new compound.

Example: Carbon monoxide, a compound, burns in the presence of oxygen, an element, to form a single product, carbon dioxide.

2CO (g) + O₂ (g) → heat → 2CO₂ (g) [carbon monoxide] [oxygen] [carbon dioxide]

Sulphur dioxide and oxygen combines under certain conditions to form sulphur trioxide.

2SO₂ + O₂ → 2SO₃

(iii) two or more compounds combine to form a single product.

Example: Ammonia and hydrogen chloride, both compounds, combine to form a new compound, ammonium chloride.

NH₃ (g) + HCl (g) → NH₄Cl (s) [ammonia] [hydrogen chloride] [ammonium chloride]

Experiments: How to perform direct combination reactions.

(i) Take some black lead sulphide in a test tube and heat it. The black sulphide reacts with oxygen to form white lead sulphate.

PbS + 2O₂ → PbSO₄ (combination)

(ii) Hold a piece of magnesium ribbon over a flame. It burns with a dazzling light, forming magnesium oxide.

2Mg + O₂ → 2MgO (synthesis)

2. Decomposition

It is the breaking up of a compound either into elements or simpler compounds, such that these products do not recombine to form the original compound.

Decomposition may occur in the presence of heat or light, or by the passage of an electric current.

A decomposition reaction that is brought about by heat is known as thermal decomposition.

In a decomposition reaction:

(i) a compound breaks up into two or more elements.

Examples:

(a) The compound mercuric oxide, when heated, decomposes to form two elements, mercury and oxygen.

2HgO (s) → delta → 2Hg (l) + O₂ (g) [mercuric oxide] [mercury] [oxygen]

Note: The symbol delta (Δ) is used to signify that heat has caused the reaction.

(b) When electric current is passed through acidulated water, the latter decomposes into hydrogen and oxygen.

2H₂O (l) → electric current → 2H₂ (g) + O₂ (g)

(ii) a compound can break up to form both elements and compounds.

Example: The compound potassium nitrate on heating decomposes to produce a compound, potassium nitrite, and an element, oxygen.

2KNO₃(s) → heat → 2KNO₂(s) + O₂(g) [potassium nitrate] [potassium nitrite] [oxygen]

(iii) a compound can break up to form two or more new compounds.

Example: The compound calcium carbonate on strong heating decomposes to form two compounds, calcium oxide and carbon dioxide.

CaCO₃(s) → heat 1000°C → CaO(s) + CO₂(g) [calcium carbonate] [calcium oxide] [carbon dioxide]

Thermal decomposition of metal compound

(i) Metal hydroxide

Hydroxides

(ii) Metal carbonates

Carbonates

(iii) Metal bicarbonates (metal hydrogen carbonate)

Metal bicarbonates or metal hydrogen carbonates decomposes to give metal carbonate, water vapour and carbon dioxide.

2NaHCO₃ → delta → Na₂CO₃ + H₂O + CO₂ [sodium hydrogen carbonate] [sodium carbonate]

Ca(HCO₃)₂ → delta → CaCO₃ + H₂O + CO₂ [calcium hydrogen carbonate] [calcium carbonate]

Mg(HCO₃)₂ → delta → MgCO₃ + H₂O + CO₂ [magnesium bicarbonate] [magnesium carbonate]

(iv) Metal nitrates

Nitrates

Decomposition reactions in our body: Digestion of food by our body is an example of a decomposition reaction. The starch present in the food we eat decomposes into glucose and sugar. Proteins undergo decomposition to form amino acids. Fats and oils are decomposed to fatty acids and finally oxidized by respiration into carbon dioxide and water.

Starch → Enzymes → Glucose → [O] → CO₂ + H₂O

Experiments: How to perform decomposition reactions.

(i) Take some lead nitrate crystals in a test tube and heat them. The crystals first melt and, on further heating, give out both nitrogen dioxide, a reddish brown gas, and oxygen. A yellow solid (lead monoxide) is left behind in the test tube.

2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂

(ii) Put some zinc carbonate in a test tube fitted with a cork and a bent glass tube. On heating, carbon dioxide is given out, which will turn lime water milky. The residue, i.e. zinc oxide, is yellow when hot, but it turns white on cooling.

ZnCO₃ → ZnO + CO₂

(iii) Heat orange-coloured ammonium dichromate in a test tube. Upon heating, it swells and decomposes, evolving nitrogen and water vapours and a green solid, chromium oxide is left behind.

(NH₄)₂Cr₂O₇ → Cr₂O₃ + 4H₂O + N₂ [ammonium dichromate] [chromium oxide]

(iv) When hydrated copper(II) sulphate is heated in a test tube, the blue-coloured crystals change into white anhydrous salt. The change may be represented by the following equation:

CuSO₄ - 5H₂O(s) → CuSO₄(s) + 5H₂O(g) [hydrated salt (blue)] [anhydrous salt (colourless)]

However, anhydrous copper(II) sulphate may be changed to the blue hydrated form by taking a sample of the anhydrous salt and adding water to it (this is the test to detect the presence of water).

CuSO₄(s) + 5H₂O(l) → CuSO₄.5H₂O(g) [anhydrous salt (colourless)] [hydrated salt (blue)]

So far, most of the reactions that we have considered proceed quite definitely in a certain direction, and it is possible to identify the reactants and the products. However, there also exists a group of reactions in which the direction of chemical change can be reversed by changing the conditions under which the reaction is taking place. Such reactions are called reversible reactions.

Thus, this is a reversible reaction, and the equation for the reaction is:

CuSO₄.5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(g)

To show that the reaction is reversible, we put the sign between the reactants and the products.

If steam is passed over red hot iron, the latter is partially converted into magnetic oxide, and hydrogen is released. If, on the other hand, hydrogen is passed over the heated oxide, it partially changes back to steam.

3Fe + 4H₂O ⇌ Fe₃O₄ + 4H₂

In either case, the reaction ends with a mixture of reactants and products.

Experiments: To show thermal dissociation

(v) Heat some solid ammonium chloride in a test tube. Two colourless gases ammonia and hydrogen chloride are produced. As these gases move up at the upper part of the test tube which is cooler, they combine to form ammonium chloride, which appears as a white sublimate on the cooler upper side of the test tube.

NH₄Cl ⇌ NH₃ + HCl

A simultaneous reversible decomposition reaction brought about only by heat is thermal dissociation.

Example of thermal dissociation. On heating, nitrogen tetraoxide changes to nitrogen dioxide, a reddish brown gas. On cooling, nitrogen dioxide changes into the original compound, nitrogen tetraoxide.

N₂O₄ ⇌ heat/cool ⇌ 2NO₂

3. Displacement

It is a chemical change in which a more active element displaces a less active element from its salt solution.

Experiments: To show displacement reactions.

(i) Take a solution of copper sulphate in a beaker, add a few pieces of zinc, and stir with a glass rod. The blue colour of the solution gradually fades, and soon the solution becomes colourless. At the same time, reddish brown particles of copper aciton appear in the beaker.

CuSO₄ + Zn → ZnSO₄ + Cu

(ii) In a test tube, take some dilute sulphuric acid and drop a small piece of magnesium ribbon. Brisk effervescence takes place, and hydrogen is evolved, which burns with a pop sound on bringing a burning match stick near the mouth of the test tube.

Mg + H₂SO₄ → MgSO₄ + H₂

(iii) Pass chlorine gas through a solution of potassium iodide. The colourless solution turns yellow brown as iodine appears.

2KI + Cl₂ → 2KCl + I₂

From the above examples, it can be noticed that in metal, zinc is more active (reactive) than copper, and magnesium is more active compared to hydrogen. In non-metals chlorine is more active as compared to iodine.

By taking similar examples, the following activity series can be prepared.

Activity (reactivity) of elements

Note: The more reactive element displaces the less reactive element from its salt solution.

4. Double decomposition

This is a type of chemical change in which two compounds in a solution react to form two new compounds by mutual exchange of radicals. Double decomposition reaction is also called double displacement reaction.

AB + CD → AD + CB

These reactions are of two types: (a) precipitation reactions (b) neutralization reactions.

(a) Precipitation reaction

A chemical reaction in which two compounds in their aqueous state react to form an insoluble salt (a precipitate) as one of the products is known as a precipitation reaction.

For example:

BaCl₂ (aq) + Na₂SO₄ (aq) → BaSO₄(s) + NaCl (aq) white ppt.

CuSO₄ (aq) + H₂S (g) → CuS(s) + H₂SO₄ (aq) black ppt

Experiments: To show double decomposition reactions.

(i) Take a solution of silver nitrate in a test tube and add dilute hydrochloric acid or a solution of sodium chloride. A white, curdy precipitate of sodium chloride. A white, curdy precipitate of sodium chloride is formed.

AgNO₃ + HCl → AgCl + HNO₃

AgNO₃ + NaCl → AgCl + NaNO₃

(ii) Fill