Introduction

Chemical reactions are the process in which new substances with new properties are formed.

Chemical reactions involve the breaking of bonds between the atoms of the reacting substances and making of new bonds between the atoms of products.

During chemical reactions a large variety of rearrangement of atoms can take place to produce new substances having entirely different properties.

The rusting of iron objects on exposure to moist air, the changing of milk into curd and digestion of food in our body are all examples of chemical reactions.

The following observations can help us to determine whether a chemical reaction has taken place or not:

• Changing in state
• Change in colour
• Evolution of a gas
• Change in temperature

A Chemical reaction is represented in short form by writing a chemical equation.

Chemical Equations

The method of representing a chemical reaction with the help of symbols and formulae of the substances involved in it is known as a chemical equation. Let us take an example to understand the meaning of a chemical equation clearly.

When a magnesium ribbon is burnt in oxygen, it gets converted to magnesium oxide.

The above chemical equation for this reaction can be represented in word form as:

$$ \underbrace{Magnesium + Oxygen}\rightarrow \underbrace{Magnesium oxide} $$

The substances that undergo chemical change are called the reactants. The new substances formed during the reaction are called the products.

In the above reactions, magnesium and oxygen are reactants and magnesium oxide is product.

The reactants are written on the left-hand side with a plus sign (+) between them. Similarly, products are written on the right-hand side with a plus sign (+) between them.

The arrow is put between reactants and products and the arrow head is pointed towards the products, which shows the direction of the reaction.

Writing a chemical Equation

Chemical equation can be made more concise and useful if we use chemical formula instead of words in word-equation.

Example:             Mg + O₂ → MgO

Mg, O₂ and MgO are the formulae of magnesium, oxygen and magnesium oxide respectively.

Illustration 1:

Write the chemical equation for the following word-equation:

Zinc + Sulphuric acid → Zinc sulphate + Hydrogen

Solution:              Zn + H₂SO₄ → ZnSO₄ + H₂

Exercise 1:

1. What change will you observe when lime is added to water?
2. When carbon is burnt in presence of oxygen, carbon dioxide is formed. Write the chemical equation for this change?
3. Write the chemical equation for the burning of methane.

Ans:

1. Heat will be generated and the temperature of the solution will increase.
2. C + O₂ → CO₂
3. CH₄ + O₂ → CO₂ + H₂O

Balanced Chemical Equation

A balanced chemical equation has an equal number of atoms of different elements in the reactants and products.

In other words, the number of atoms of each element remains the same, before and after the chemical reaction.

Consider the following chemical equation

Zn + H₂SO₄→ ZnSO₄ + H₂

Let us examine the number of atoms of different elements on both sides of the arrow:

As the number of atoms of each element is the same on both sides of the arrow, hence the chemical equation is called balanced chemical equation.

If we see a chemical equation Fe + H₂O → Fe₂O₃ + H₂

Then we can say that the equation is not balanced since number of Fe and O atom is not same in reactant and product both sides.

Following steps are involved in balancing the chemical equations:

Step 1:  First of all identify the atoms which are not balanced e.g. Fe and O in above mentioned equation.

Step 2:  Pick any one element and start balancing. It is convenient to start with the compound that contains the maximum number of atoms.

Step 3:  Then pick the 2^nd element to balance the partly balanced equation.

Step 4:  Pick the next element to get the equation balanced.

This procedure is to be repeated till the equation is finally balanced.

So the balanced equation is: 2Fe + 3H₂O → Fe₂O₃ + 3H₂

Note : To make a chemical equation more informative, the physical states of the reactants and products are mentioned along with their chemical formulae. The gaseous, liquid and solid states of reactants and products are represented by the notations (g), (l) and (s)respectively. The word aqueous (aq) is written if the reactant or product is present as a solution in water. Moreover an arrow ↓ for the information of precipitate and arrow ↑ for the evolution of a gas can be assigned. So the balanced chemical equation of the above example can be represented as: 2Fe (s) + 3H₂O (g) → Fe₂O3 (s) + 3H₂(g)

Note : Sometimes the reaction conditions such as temperature, pressure, catalyst etc; for the reaction are indicated above and or below the arrows in the chemical equations.

For example:

N₂ (g) + 3H₂ (g) high pressure & low temperature → Fe catalyst 2NH₃ (g)

6CO₂ (g) + 6H₂O (l) sunlight → chlorophyll C₆H₁₂O₆ (aq)+6O₂ (g)

   (Glucose)

Illustration 2:

Balance the following equation: H2SO4 + NaOH → Na2SO4 + H2O

Solution:              The unbalanced atoms are Na and H. We will start with balanced of Na.

                                H₂SO₄ + 2NaOH → Na₂SO₄ + H₂O

                                Now balance H atom

                                H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O

Exercise 2:

A. Balance the following equations:

1. C + CO₂ → CO
2. C + O₂ → CO
3. Fe + CI₂ → FeCI₃

B. Balance the following chemical equations:

1. Fe + O₂ → Fe₂O₃
2. AI(OH)₃ → AI₂O₃ + H₂O
3. BaCI₂ + H₂SO₄ → BaSO₄ + HCI

C. Balance and write down the physical state of the reactants and products in the following reactions:

1. CO₂ + H₂O → C₆H₁₂O₆ + O₂
2. BaCI₂ + Na₂SO₄ → BaSO₄ + NaCI

ANS:

A.

1. C + CO₂ → 2CO
2. 2C + O₂ → 2CO
3. 2Fe + 3CI₂ → 2FeCI₃ 
   

B.

1.  4Fe + 3O₂ → 2Fe₂O₃
2. 2A(OH)₃ →AI₂O₃ + 3H₂O
3. BaCI₂ + H₂SO₄ → BaSO₄ + 2HCI

C.

1. 6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(aq) + 6O₂(g)
2. BaCI₂(aq) + Na₂SO₄(aq) → BaSO₄↓ + 2NaCI(s)

Types of chemical reactions

1. Combination reaction

In a combination reaction, two or more substances combine to form a new substance. For example

C(s) + O₂ (g) → CO₂ (g)

2H₂ (g) + O₂ (g) $$ \overset{combination}{\rightarrow} $$ 2H₂O (g)

Combination reactions are also termed as synthesis reaction.

2. Decomposition Reaction

In a decomposition reaction, a single compound breaks down to produce two or more simpler substances. The decomposition reactions take place when energy is supplied in the form of heat electricity or light. When a substance is decomposed by passing electric current the process is called electrolysis.

2H₂O(l) $$ \overset{electric current}{\rightarrow} $$ 2H₂ (g) + O₂ (g)

When a substance decomposes on heating it is called thermal decomposition.

CaCO₃ (S) \overset{Heat}{\rightarrow} CaO (s) + CO₂ (g)

Lime stone                                                          Quick lime

When a substance is decomposed in presence of sunlight, it is called a photochemical decomposition.

AgBr $$ \overset{sunlight}{\rightarrow} Ag + Br

Before we discuss the displacement reactions, we should know about the reactivity of metals.

Reactivity series of metals

Reactivity series of metals is a series in which the metals are arranged in the decreasing order of their reactivity. A more reactive metal can displace the less reactive metal from its solution.

3. Displacement reaction

In a displacement reaction, a more active element displaces or removes another element from a compound, for examples Zn displaces Cu CuSO₄ solution because Zn is more active than Cu.

Zn (s) + CuSO₄ (aq) → ZnSO₄ (aq) +Cu(s)

Mg (s) + CuSO₄ (aq) → MgSO₄ (aq) + Cu(S)

4. Double Displacement reaction

The reactions in which two different atoms or groups of atoms are displaced by other atoms or groups of atoms are double displacement reactions, for eg.

BaCl₂ (aq) + Na₂SO₄ (aq) → BaSO₄ (s) + 2NaCl (aq)

AgNO₃ (aq) + NaCl (aq) → AgCl (s) + NaNO₃ (aq)

They can be further categorized into precipitation reaction and neutralization reaction.

Illustration 3:

What happens when a piece of iron metal is placed in copper sulphate solution? Name the type of reaction involved?

Solution:              Iron being more reactive will displace copper from copper sulphate solution

                                Fe(s) + CuSO₄ (aq) → FeSO₄ (aq) + Cu(s)

                                It is a displacement reaction.

Exercise 3:

A. Write balanced equation for the following reaction and identify the type of reaction:

1. Zinc carbonate (s) → Zinc oxide (s) + Carbon dioxide (g)
2. Zinc (s) + hydrochloric acid (dil.) → Zinc Chloride + Hydrogen

B. Write balanced equation for the following reaction and indentify the type of reaction:

1. Silver nitrate → Silver + Oxygen + Nitrogen dioxide
2. Sulphur dioxide + Oxygen → Sulphur trioxide

C. Identify the type of reactions:

1. AgCI $$ \overset{hv}{\rightarrow} $$ Ag + CI
2. Mg + CuSO₄ → MgSO₄ + Cu
3. 2FeCI₂ + CI₂ → 2FeCI₃

ANS:

A. 

1. ZnCO₃ → ZnO + CO₂,                        Decomposition reaction
2. Zn + 2HCI → ZnCI₂ + H₂,                  Displacement reaction

B.

1. 2AgNO₃ → 2Ag + 2O₂ + NO₂,            Decomposition reaction
2. 2SO₂ + O₂ → 2SO₃,                          Combination reaction

C.

1. photochemical decomposition
2. displacement reaction
3. combination reaction

5. Oxidation and Reduction Reaction

This type of reaction involves transfer of electrons.

Oxidation is

1. Addition of oxygen
2. Removal of hydrogen
3. Loss of electron

Eg.          Mg + O₂ → MgO

Any chemical substance following any of these rules is said to be oxidized.

Reduction is

1. Removal of oxygen
2. Addition of hydrogen
3. Gain of electron

Eg.          Cl + e^-→Cl^-

Reactions involving both oxidation and reduction process, occurring simultaneously are known as Redox reactions.

For eg.  H₂S + Cl₂→ 2HCl + S

This is a molecular equation. The representation of molecular equation in terms of its ions is known So, its ionic equation is:

………………… Pending ……………….

Here sulphide ion is losing its electrons and chlorine is accepting its electrons therefore the reaction involves oxidation and reduction both, therefore it’s a redox reaction. Since sulphur is losing its electrons, therefore sulphur is oxidized and Cl₂ is gaining electrons therefore chlorine is reduced. Reduction of Cl₂ is brought by sulphur therefore sulphur is reducing agent and oxidation of sulphur is brought by Cl₂ therefore Cl₂ is oxidizing agent.

Oxidation Number or Oxidation State

It is defined as the charge (real or imaginary) which an atom appears to have when it is compound state. In the case of electrovalent compounds, the oxidation number of an element or radical is the same as the charge on the ion. The following rules are followed in ascertaining the oxidation number in any type of compounds:

1. The oxidation number of an atom in elementary state is always zero.
2. The oxidation number of fluorine is always – 1.
3. The oxidation number of oxygen is – 2 in all compounds except in peroxides, super oxides and oxygen fluorides.
4. The oxidation number of hydrogen is +1 in all of its compounds except I metallic hydrides. In metallic hydrides, oxidation number of hydrogen is -1.
5. The oxidation number of aion is equal to the electrical charge present on it.
6. The oxidation number of alkali metal is +1 and that of alkaline earth metals is +2.
7. For complex ions, the algebraic sum of oxidation number of all the atoms is equal to the net charge on the ion.
8. In the case of neutral molecules, the algebraic sum of oxidation numbers of all the atoms present in the molecules in zero.

Ion Electron Method of Balancing Redox Reaction

The following steps are followed:

1. Ionic equation of redox reaction is first written.
2. The ionic reaction is split into two half reaction, one for oxidation and the other for reduction.
3. Each half reaction is balanced for the number of atoms of each element. For this purpose:

1. First of all atoms other than H and O for each half reaction are balanced using simple multiples.
2. In acidic and neural mediums, H ions are added to the side deficient in hydrogen and water molecules to the side deficient in oxygen.
3. In alkaline medium, for each excess of oxygen atom, one water molecule is added to the same side and two OH^- ions to the other side. Ifhydrogen is still unbalanced, one OH^- ion is added for each excess of hydrogen on the same side and one water molecule to the other side.

1. Electron are added to the side deficient in electrons as to equalize the charges on both sides of the half reactions.
2. Electrons are made equal in both the half reactions by multiplying one or both the half reactions by a suitable number.
3. Both the balanced half reactions are added and any term common to both sides is cancelled.

Illustration 4:

Identify reducing and oxidizing agent in the given equation:

…………… Pending ………………

Solution:              Fe is gaining electrons so it is reduced therefore AI is reducing agent

                                AI is losing electrons so it is oxidized therefore Fe₂O₃ is oxidizing agent.

Exercise 4:

A. Find the substance oxidized, reduced, oxidizing agent and reducing agent in the following reactions:

1. 2AI + 6HCI → 2AICI₃ + 3H₂
2. H₂ + CI₂ → 2HCI
3. Zn + H₂SO₄ →ZnSO₄ + H₂

B. Find out the oxidation state of the following reactants and products:

1. C + O₂ → CO₂
2. CaCO₃ → CaO + CO₂

C. Balance the following equations by ion-electron method.

1. Fe³⁺ + H₂S → Fe²⁺ + H⁺ + S.

ANS:

A.

1.

1. $$ \overset{0}{C} + \overset{0}{O}_{2}\rightarrow \overset{+4}{C}\overset{-2}{O}_{2} $$
2. $$ \overset{+2}{Ca}\overset{+4}{C}\overset{-2}{O}_{3} \rightarrow \overset{+2}{Ca}\overset{-2}{O}+\overset{+4}{C}\overset{-2}{O}_{2} $$
3. Fe³⁺ + H₂S → Fe²⁺ + H⁺ + SH₂S → 2H⁺ + S                    ;               Fe³⁺ → Fe²⁺

2^nd step : Adding electrons to the side deficient in electrons,
3^rd step : Balancing electrons in both the half reactions,

1. H₂S → 2H⁺ + S + 2e           ;               Fe³⁺ + e → Fe²⁺
2. (oxidation half reaction)                (Reduction half reaction)
3. 1^st step : Splitting the redox reaction into two half reactions,

                H₂S → 2H⁺ + S + 2e           ;               2Fe³⁺ + 2e → 2Fe²⁺

4^th step : Adding both the half reactions.

                H₂S + 2Fe³⁺ → 2H⁺ + S + 2Fe²⁺

The effects of oxidation reactions in everyday life

Corrosion:

When a metal is attacked by substance around it such as moisture, acids etc, it is said to be corroded and this process is called corrosion. Due to corrosion the new iron articles get coated with a reddish brown powder when left for some time.

The black coating on silver and the green coating on copper are other examples of corrosion.

Corrosion causes damages to car bodies, bridges, iron railings, ships and to all objects of metals.

Corrosion of iron is called rusting and it’s a serious problem. Every year an enormous amount of money is spent to replace damaged iron.

The rusting of iron can be prevented by painting, oiling, greasing, galvanizing, anodizing or making alloys.

Galvanization is a method of protecting steel and iron from rusting by coating them with a thin layer of zinc.

Rancidity

The most important cause of deterioration in fast and fatty foods is oxidation. Oxidation of fats results in the replacement of an oxygen ion for a hydrogen ion in the fatty acid molecule. This substitution destabilizes the molecule and makes it possible for other odd chemical fragments to find a place along the chain. What we perceive is an unpleasant change in the flavor and odour of a food is called rancidity. Unsaturated fats are more susceptible to oxidation than are saturated fats. Factors which accelerate fat oxidation include trace metals (iron, zinc, etc.), salt, light, water, bacteria, and molds. Fat oxidation can be retarded by use of antioxidants, by use of spices such as sage and rosemary, and by use of light and/ or air tight wrapping. Some high fat foods such as potato chips are packaged in materials that protect them from oxygen and the containers are flooded with nitrogen to further exclude oxygen. The nitrogen also serves as a cushion to minimize breakage of the chips during transport.

As soon as a food, feed, or ingredient is manufactured, it begins to undergo a variety of chemical and physical changes. Oxidation of fats is one common and frequently undesirable chemical change that may impact flavor, aroma, nutritional quality and in some cases, even the texture of product. The chemicals produced from oxidation of fats are responsible for rancid flavors and aromas. Vitamins and other nutrients may be partially or entirely destroyed by highly reactive intermediates in the fats oxidation process. Oxidized fats can interact with proteins ad carbohydrates causing changes in texture. Of course, not all fats oxidation is undesirable. Enzymes, for example, promote oxidation of fats membranes during ripening of fruit. For most products, though, predicting and understanding oxidation of fats is necessary to minimize objectionable flavors and aromas arising from fat rancidity.

Illustration 5:

Aluminium burns in chlorine to form aluminium chloride, AICI3. Write a balanced equation for this reaction.

Solution:              2AI + 3CI₂ → 2AICI₃

Illustration 6:

What type of reactions are represented by the following equations?

1. 2FeSO4 → Fe2O3 + SO2 + SO3
2. 2AgNO3 + Cu → Cu(NO3)2 + 2Ag

Solution:              (a) decomposition;          (b) displacement

Illustration 7:

Balance the following reactions and indicate which types which types of chemical reaction are being represented:

1. NaBr + Ca(OH)₂ → CaBr₂ + NaOH
2. NH₃ + H₂SO₄ → (NH₄)₂ SO₄
3. C₅H₉O + O₂ → CO₂ + H₂O

Solution:

1. 2 1 1 2           Double displacement reaction
2. 2 1 1              Synthesis reaction
3. 4 27 20 18      Combustion reaction

Illustration 8:

Write the formula for each material correctly and then balance the equation. For each reaction tell what type of reaction it is

1. Calcium carbonate heated to leave calcium oxide and carbon dioxide.
2. Ammonia gas when it is pressed into water will make ammonium hydroxide.

Solution:

(a)          CaCO₃ → CaO + CO₂             Decomposition

(b)          NH₃ + H₂O → NH₄OH             Combination

Illustration 9:

Write the formula for reaction material correctly and then balance the equation. For each reaction tell what type of reaction it is

1. Zinc sulphide and oxygen become zinc oxide and Sulphur.
2. Lithium oxide and water make lithium hydroxide.
3. Aluminium hydroxide and sulphuric acid neutralize to make water and aluminium sulphate

Solution:

1. 2ZnS + O₂ → 2ZnO + 2S
2. Single displacement
3. Li₂O + H₂O → 2LiOH
4. Synthesis
5. 2A(OH)₃ + 3H₂SO₄ → 6H₂O + AI₂(SO₄)₃
6. Double displacement or Acid-Base Neutralization

Illustration 10:

Write the formula for each material correctly and then balance the equation. Fir each reaction tell what type of reaction it is

1. Sulphur burns in oxygen to make Sulphur dioxide.
2. Barium hydroxide and sulphuric acid make water and barium sulphate.
3. Aluminium sulphate and calcium hydroxide become aluminum hydroxide and calcium sulphate.

Solution:

1. S + O₂ → SO₂
2. Synthesis
3. Ba(OH)₂ + H₂SO₄ → 2H₂O + BaSO4
4. Double displacement or Acid-Base Neutralization
5. AI₂(SO₄)₃ + 3Ca(OH)₂ → 2AI(OH)₃↓ + 3CaSO₄↓(Both calcium sulphate and aluminum hydroxide are precipitates.)
6. Double displacement

Illustration 11:

Balance the following chemical equations:

1. H₃PO₄ + KOH → K₃PO₄ + H₂O
2. K + B₂O₃ → K₂O + B

Solution:             

1. H₃PO₄ + 3KOH → K₃PO₄ + 3H₂O
2. 6K + B₂O₃ → 3K₂O + 2B

Exercise 5:

A. When magnesium is burnt in chlorine, it forms MgCI2, state which element is oxidized and which is reduced.

B. In the redox reaction:Which one is the oxidizing agent?

         A²⁺ + B → B²⁺ + A

c. Balance the following reactions and indicate which types of chemical reaction are being represented:

1. Pb + H₃PO₄ → H₂ + Pb₃(PO₄)₂
2. Li₃N + NH₄NO₃ → LiNO₃ + (NH₄)₃N
3. HBr + AI(OH)₃ → H₂O + AIBr₃

D. Write the formula for each material correctly and the balance the equation. For each reaction tell what type of reaction it is

1. Phosphoric acid plus sodium hydroxide.
2. Propane burns (with oxygen)
3. Zinc and copper II sulphate yield zinc sulphate and copper metal

E. Balance the following chemical equations:

1. Na + NaNO₃ → Na₂O + N₂
2. C + S → CS₂

F. Balance the following chemical equations:

1. FeS₂ + O₂ → Fe₂O₃ + SO₂
2. C + SO₂ → CS₂ + CO

G. Write the formula for each material correctly and then balance the equation. For each reaction tell what type of reaction it is

1. sulphuric acid reacts with zinc
2. acetic acid ionizes.
3. steam with methane to get hydrogen and carbon dioxide

H. Write the formula for each material correctly and then balance the equation. For each reaction tell what type of reaction it is

1. copper metal and silver nitrate react to form silver metal and copper II nitrate.
2. sodium metal and chlorine react to make sodium chloride.
3. calcium phosphate and sulfuric acid make calcium sulfate and phosphoric acid.

ANS:

A.  Mg is oxidized and CI2 is reduced.

B.  A2+ is the oxidizing agent

C.  

1. 3 2 3 1                        Single displacement reaction
2. 1 3 3 1                        Double displacement reaction
3. 3 1 3 1                        Acid-base reaction

D.

1. H₃PO₄ + 3NaOH → Na₃PO₄ + 3H₂O                  Double displacement (Neutralization)
2. C₃H₈ + 5O₂ → 4H₂O + 3CO₂                             Burning of A Hydrocarbon
3. Zn + CuSO₄ → ZnSO₄ → ZnSO₄ + Cu              Single displacement

E.

1.  10Na + 2NaNO₃ → 6Na₂O + N₂
2.  C + 2S → CS₂

F.

1. 4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂
2. 5C + 2SO₂ → CS₂ + 4CO

G.

1. $$ H_{2}SO_{4}+Zn\rightarrow ZnSO_{4}+H_{2} $$                   Single displacement
2. $$ CH_{3}COOH\rightleftharpoons CH_{3}COO^{-}+H^{+} $$  lionization (Notice that it is reversible)
3. $$ 2H_{2}O+CH_{4}\rightarrow 4H_{2}+CO_{2} $$

H.

1. $$ Cu +2AgNO_{3}\rightarrow 2Ag+Cu(NO_{3})_{2} $$                   Single displacement
2. $$ 2Na+CI_{2}\rightarrow 2NaCI $$                                                 Synthesis
3. $$Ca_{3}(PO_{4})_{2}+3H_{2}SO_{4}\rightarrow3CaSO_{4}+2H_{3}PO_{4}$$Double displacement