Chemical Reactions and Equations

Consider daily life situations where changes occur:

Milk left at room temperature during summers.
Iron tawa/pan/nail exposed to humid atmosphere.
Grapes get fermented.
Food is cooked.
Food is digested in our body.
Respiration.

In these situations, the initial substance's nature and identity change, indicating a chemical reaction has taken place. Chemical reactions involve physical and chemical changes. To understand chemical reactions, consider these activities:

Activity 1.1: Burning Magnesium Ribbon

Clean a 3-4 cm long magnesium ribbon with sandpaper.
Hold it with tongs and burn it using a spirit lamp or burner.
Collect the ash in a watch-glass.
Keep the burning ribbon away from your eyes.
Observation: Magnesium ribbon burns with a dazzling white flame and changes into a white powder, which is magnesium oxide. This is due to the reaction between magnesium and oxygen in the air.

Activity 1.2: Reaction of Zinc with Acid

Take a few zinc granules in a conical flask or test tube.
Add dilute hydrochloric acid or sulphuric acid.
Handle the acid with care.
Observe any changes around the zinc granules.
Touch the conical flask or test tube to check for temperature change.
Observation: Hydrogen gas is formed.

Activity 1.3: Reaction of Lead Nitrate with Potassium Iodide

Take lead nitrate solution in a test tube.
Add potassium iodide solution.
Observation: A color change and precipitate formation occur.

Observations Indicating Chemical Reactions

From the above activities, a chemical reaction is determined by:

Change in state
Change in color
Evolution of a gas
Change in temperature

Chemical Equations

Word Equations

Activity 1.1 can be described as: Magnesium ribbon burns in oxygen to form magnesium oxide. This can be written as a word equation:

Magnesium + Oxygen \rightarrow Magnesium \,oxide

Reactants: Substances undergoing change (magnesium and oxygen).
Product: New substance formed (magnesium oxide).
Reactants are written on the left-hand side (LHS) with a plus sign (+) between them.
Products are written on the right-hand side (RHS) with a plus sign (+) between them.
An arrow points towards the products, indicating the reaction's direction.

Writing Chemical Equations

Chemical equations use chemical formulae instead of words to be more concise.

Mg + O_2 \rightarrow MgO

Count and compare the number of atoms of each element on the LHS and RHS.
If the number of atoms is the same on both sides, the equation is balanced.
If not, the equation is unbalanced (a skeletal chemical equation).

Balanced Chemical Equations

Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.
The total mass of elements in the products equals the total mass of elements in the reactants.
The number of atoms of each element remains the same before and after the reaction.
Therefore, skeletal chemical equations must be balanced.

Consider:

Zinc + Sulphuric \,acid \rightarrow Zinc \,sulphate + Hydrogen

Zn + H2SO4 \rightarrow ZnSO4 + H2

The number of atoms of each element is the same on both sides; hence, it is a balanced chemical equation.

Consider:

Fe + H2O \rightarrow Fe3O4 + H2

Balancing Chemical Equations Step by Step

Step I: Draw boxes around each formula (do not change anything inside the boxes).
Fe + H2O \rightarrow Fe3O4 + H2
Step II: List the number of atoms of different elements present in the unbalanced equation.

Element	Reactants (LHS)	Products (RHS)
Fe	1	3
H	2	2
O	1	4

Step III: Start balancing with the compound that contains the maximum number of atoms (Fe3O4) and select the element with the maximum number of atoms (oxygen).
Step IV: Balance oxygen atoms.
Fe + 4H2O \rightarrow Fe3O4 + H2
Atoms of Reactants Products
Oxygen Initial 1 4
To balance 1 \times 4 4
Step V: Balance hydrogen atoms.
Fe + 4H2O \rightarrow Fe3O4 + 4H2
Atoms of Reactants Products
Hydrogen Initial 8 2
To balance 8 2 \times 4
Step VI: Balance iron atoms.
3Fe + 4H2O \rightarrow Fe3O4 + 4H2
Atoms of Reactants Products
Iron Initial 1 3
To balance 1 \times 3 3
Step VII: Check the correctness of the balanced equation.
3Fe + 4H2O \rightarrow Fe3O4 + 4H2
The numbers of atoms of elements on both sides are equal.
This method is called the hit-and-trial method.

Atoms of		Reactants	Products
Oxygen	Initial	1	4
	To balance	1 \times 4	4

Atoms of		Reactants	Products
Hydrogen	Initial	8	2
	To balance	8	2 \times 4

Atoms of		Reactants	Products
Iron	Initial	1	3
	To balance	1 \times 3	3

Writing Symbols of Physical States

To make a chemical equation more informative, include physical states:

(g) for gaseous state
(l) for liquid state
(aq) for aqueous solution
(s) for solid state

3Fe(s) + 4H2O(g) \rightarrow Fe3O4(s) + 4H2(g)

The symbol (g) with H2O indicates steam.
Physical states are usually not included unless necessary.
Reaction conditions (temperature, pressure, catalyst) are indicated above and/or below the arrow.

CO(g) + 2H2(g) \xrightarrow[340 atm]{} CH3OH(l)

6CO2(aq) + 12H2O(l) \xrightarrow[Chlorophyll]{Sunlight} C6H{12}O6(aq) + 6O2(aq) + 6H_2O(l)

Types of Chemical Reactions

During a chemical reaction, atoms of one element do not change into those of another element. Chemical reactions involve the breaking and making of bonds between atoms to produce new substances.

Combination Reaction

Activity 1.4: Formation of Slaked Lime

Take a small amount of calcium oxide (quick lime) in a beaker.
Slowly add water.
Touch the beaker to feel the temperature change.
Calcium oxide reacts vigorously with water to produce slaked lime (calcium hydroxide), releasing heat.

CaO(s) + H2O(l) \rightarrow Ca(OH)2(aq) + Heat

In this reaction, calcium oxide and water combine to form a single product, calcium hydroxide.
A reaction in which a single product is formed from two or more reactants is a combination reaction.

More Examples of Combination Reactions

Burning of coal:
C(s) + O2(g) \rightarrow CO2(g)
Formation of water:
2H2(g) + O2(g) \rightarrow 2H_2O(l)
When two or more substances combine to form a single product, it is a combination reaction.
Reactions in which heat is released are called exothermic chemical reactions.

Exothermic Reactions

Burning of natural gas:
CH4(g) + 2O2(g) \rightarrow CO2(g) + 2H2O(g)
Respiration:
C6H{12}O6(aq) + 6O2(aq) \rightarrow 6CO2(aq) + 6H2O(l) + energy
Decomposition of vegetable matter into compost.
A solution of slaked lime is used for whitewashing walls; calcium hydroxide reacts with carbon dioxide to form a thin layer of calcium carbonate.
Ca(OH)2(aq) + CO2(g) \rightarrow CaCO3(s) + H2O(l)

Decomposition Reaction

Activity 1.5: Heating Ferrous Sulphate Crystals

Take about 2 g of ferrous sulphate crystals in a dry boiling tube.
Note the color of the crystals.
Heat the boiling tube over a flame.
Observe the color change and smell the odor.

2FeSO4(s) \xrightarrow{Heat} Fe2O3(s) + SO2(g) + SO_3(g)

A single reactant breaks down to give simpler products.
Ferrous sulphate crystals lose water when heated and decompose to ferric oxide, sulphur dioxide, and sulphur trioxide.
Decomposition of calcium carbonate is used in various industries.
CaCO3(s) \xrightarrow{Heat} CaO(s) + CO2(g)
When a decomposition reaction is carried out by heating, it is called thermal decomposition.

Activity 1.6: Heating Lead Nitrate

Take about 2 g of lead nitrate powder in a boiling tube.
Heat it over a flame.
Observe the emission of brown fumes (nitrogen dioxide).

2Pb(NO3)2(s) \xrightarrow{Heat} 2PbO(s) + 4NO2(g) + O2(g)

Activity 1.7: Decomposition of Silver Chloride

Take about 2 g of silver chloride in a china dish.
Place it in sunlight.
Observe the color change.

2AgCl(s) \xrightarrow{Sunlight} 2Ag(s) + Cl_2(g)

White silver chloride turns grey in sunlight due to decomposition into silver and chlorine.

Activity 1.8: Electrolysis of Water

Set up the electrolysis of water apparatus.
Add a few drops of dilute sulphuric acid to the water.
Invert test tubes over the carbon electrodes.
Switch on the current and observe the formation of bubbles.
Test the gases with a burning candle.
Silver bromide also behaves in the same way:
2AgBr(s) \xrightarrow{Sunlight} 2Ag(s) + Br_2(g)
Decomposition reactions require energy in the form of heat, light, or electricity.
Reactions in which energy is absorbed are endothermic reactions.

Displacement Reaction

Activity 1.9: Iron Nails in Copper Sulphate Solution

Take three iron nails and clean them with sandpaper.
Take two test tubes (A and B) with copper sulphate solution in each.
Immerse two iron nails in the copper sulphate solution in test tube B for 20 minutes.
Keep one iron nail aside for comparison.
Compare the intensity of the blue color and the color of the iron nails.

Fe(s) + CuSO4(aq) \rightarrow FeSO4(aq) + Cu(s)

Iron displaces copper from copper sulphate solution.
This reaction is known as a displacement reaction.

Other Examples of Displacement Reactions

Zn(s) + CuSO4(aq) \rightarrow ZnSO4(aq) + Cu(s)

Pb(s) + CuCl2(aq) \rightarrow PbCl2(aq) + Cu(s)

Zinc and lead are more reactive than copper and displace copper from its compounds.

Double Displacement Reaction

Activity 1.10: Sodium Sulphate and Barium Chloride

Take 3 mL of sodium sulphate solution in a test tube.
Take 3 mL of barium chloride solution in another test tube.
Mix the two solutions.
Observe the formation of a white substance (precipitate).

Na2SO4(aq) + BaCl2(aq) \rightarrow BaSO4(s) + 2NaCl(aq)

The insoluble substance formed is a precipitate.
Any reaction that produces a precipitate is a precipitation reaction.
Reactions involving the exchange of ions between reactants are double displacement reactions.

Oxidation and Reduction

Activity 1.11: Heating Copper Powder

Heat a china dish containing about 1 g of copper powder.
Observe the surface turning black (copper(II) oxide).

2Cu + O_2 \xrightarrow{Heat} 2CuO

Oxygen is added to copper, and copper oxide is formed.
If hydrogen gas is passed over the heated material, the black coating turns brown as copper is obtained.

CuO + H2 \xrightarrow{Heat} Cu + H2O

If a substance gains oxygen during a reaction, it is oxidized.
If a substance loses oxygen during a reaction, it is reduced.
Reactions where one reactant gets oxidized and the other gets reduced are oxidation-reduction reactions or redox reactions.

Other Examples of Redox Reactions

ZnO + C \rightarrow Zn + CO

MnO2 + 4HCl \rightarrow MnCl2 + 2H2O + Cl2

If a substance gains oxygen or loses hydrogen, it is oxidized.
If a substance loses oxygen or gains hydrogen, it is reduced.

Effects of Oxidation Reactions in Everyday Life

Corrosion

Iron articles get coated with a reddish-brown powder (rusting).
Other metals also get tarnished.
Corrosion is the process where a metal is attacked by substances like moisture and acids.
Examples include the black coating on silver and the green coating on copper.
Corrosion causes damage to car bodies, bridges, iron railings, and ships.

Rancidity

Fats and oils in food materials become rancid when oxidized, changing their smell and taste.
Antioxidants are added to foods containing fats and oils to prevent oxidation.
Keeping food in airtight containers slows down oxidation.
Chips manufacturers flush bags with nitrogen gas to prevent oxidation.

Is magnesium oxidized or reduced when a magnesium ribbon burns in air? Magnesium is oxidized.