A chemical equation is the symbolic representation of a chemical change.
Reactants and Products
The substances, in which the chemical change is brought, are called reactants and the substances that are formed as a result of the chemical reaction are called products.
The transformation of a chemical substance into a new chemical substance by making and breaking of bonds between different atoms is known as Chemical Reaction.
We can observe or recognize a chemical reaction by observing
- Change in state
- Change in colour
- Evolution of a gas
- Change in temperature
A word-equation shows the change of reactants to products through an arrow placed between them. The reactants are written on the left-hand side (LHS) with a plus sign (+) between them. Similarly, products are written on the right-hand side (RHS) with a plus sign (+) between them. The arrowhead points towards the products and shows the direction of the reaction.
Example: Magnesium + Oxygen → Magnesium oxide
Skeletal Chemical Reaction
A chemical reaction in which, the reactants and products are identified with their chemical formula but their quantity and proportion are not identified.
Example: Mg + O2 → MgO
The physical state of reactants and the products are mentioned to make the chemical reaction more informative. e.g. we use (g) for gas, (l) for liquid, (s) for solid and (aq) for aqueous.
The insoluble substance formed in a chemical reaction is called precipitate.
- The formation of precipitate in a chemical reaction is represented by the Downward Facing arrow (↓), immediately after the precipitate.
- Precipitation Reaction: Any reaction that produces a precipitate can be called a precipitation reaction.
Evolution of a gas in a chemical reaction is represented by an Upward Facing Arrow (↑), immediately after the chemical formula of that gas.
A balanced equation is the one in which the number of atoms on the reactant and product sides is equal.
Balancing a chemical equation by Trial and Error method
Most chemical reactions can be balanced by trial and error, using three simple principles.
- The number of atoms of each element must be the same on each side of the reaction arrow.
- You can manipulate only the coefficients and not the subscripts.
- The final set of coefficients should be whole numbers and should be the smallest whole numbers that will do the job.
Example: Balancing a given chemical equation (Ferric Oxide with Carbon)
Step 1. Write the word-equation
Ferric Oxide + Carbon → Carbon Dioxide + Iron
Step 2. Represent the word-equation as a Chemical Equation
Fe2O3 + C → CO2 + Fe (Unbalanced)
Step 3. List the number of atoms of different elements present in the unbalanced equation.
|Element||Number of atoms in reactants (LHS)||Number of atoms in products (RHS)|
Step 4. Let’s begin with balancing CO2 molecule (as in the reaction, the maximum number of atoms is of Oxygen)
Here we observe that there are 3 atoms of oxygen in LHS and 2 atoms in the RHS. So, to make them equal we’ll multiply both CO2 and Fe2O3 with a number so that the resultant number must be the LCM of 2 and 3 i.e. 6, thus we’ll multiply Fe2O3 with 2 and CO2 with 3, hence
2Fe2O3 + C → 3CO2 + Fe (Partially balanced)
Step 5. Now the number of atoms of Oxygen in both RHS and LHS is equal. Now, we’ll work on iron. We observe that there are 4 atoms of Fe in LHS and 1 in RHS. To equate them, we just have to multiply Fe with 4 in RHS (as LCM of 1 and 4 is 4). Thus, we get
2Fe2O3 + C → 3CO2 + 4Fe (Partially balanced)
Step 6. After Fe, now we’ll work on Carbon. We observe that there is just 1 carbon atom in LHS, whereas there are 2 Carbon atoms in RHS. So, we need to multiply Carbon in LHS with 2 (LCM of 1 and 3 is 3), and thus we get the final balanced chemical equation in which all the elements are balanced (LHS = RHS).
2Fe2O3 + 3C → 3CO2 + 4Fe (Balanced)
Step 7. The last step would be to write the symbols of the physical states of the compounds. Here, as we know that Ferric Oxide, Carbon and Iron are solids, whereas CO2 is a gas. Thus, we get our final chemical equation
2Fe2O3(s) + 3C(s) → 3CO2(g) + 4Fe(s)
Types of Reaction (Based on Absorption/Desorption of Heat)
1. Exothermic Reaction
Reactions in which heat energy is released along with the formation of products are called exothermic chemical reactions.
Examples of Exothermic Reaction:
- Reaction of Quick Lime (Calcium oxide) with Water: CaO(s) + H2O(l) → Ca(OH)2(aq) + Heat
- Reaction of Methane with Oxygen: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + Heat
- Respiration (Reaction of Glucose with Oxygen): C6H12O6(aq) + 6O2(aq) → 6CO2(aq) + 6H2O(l) + Heat
- The decomposition of vegetable matter into compost.
2. Endothermic Reaction
Reactions which require or absorb heat energy from its surroundings are called exothermic chemical reactions.
Examples of Endothermic Reaction:
Ba(OH)2 + 2NH4Cl + Heat → BaCl2 + 2NH3 + 2H2O
Types of Reaction (Based on the behaviour of the reactants)
I. Combination Reaction
The reaction in which two or more substances (elements or compounds) combine to form a new single substance.
A + B → AB
Examples of Combination Reaction:
- Reaction of Quick Lime and Water: CaO(s) + H2O(l) → Ca(OH)2(aq)
- Burning of Coal: C(s) + O2(g) → CO2 (g)
- Formation of water from the combination of Hydrogen and Oxygen: 2H2(g) + O2(g) → 2H2O(l)
II. Decomposition Reaction
The reaction in which a single reactant decomposes to give two or more products.
AB → A + B
Decomposition reactions can be divided into three types:
When a decomposition reaction is carried out by heating.
- Thermal Decomposition of Ferrous Sulphate: 2FeSO4(s) + Heat → Fe2O3(s) + SO2(g) + SO3(g)
- Thermal Decomposition of Limestone: CaCO3(s) + Heat → CaO(s) + CO2(g)
- Thermal Decomposition of Lead Nitrate: 2Pb(NO3)2(s) + Heat → 2PbO(s) + 4NO2(g) + O2(g)
When a decomposition reaction is carried out by electric current.
- Electrolytic Decomposition of Sodium Chloride: 2NaCl (Electrolysis) → 2Na (at Cathode)+ Cl2 (at anode)
- Electrolytic Decomposition of Aluminium Chloride: 2AlCl3 (Electrolysis) → 2Al (at Cathode) + 3 Cl2 (at anode)
When a decomposition reaction is carried out by sunlight.
- Photo Decomposition of Silver Chloride: 2AgCl(s) + Sunlight → 2Ag(s) + Cl2(g)
- Photo Decomposition of Silver Bromide: 2AgBr(s) + Sunlight → 2Ag(s) + Br2(g)
III. Displacement Reaction
The chemical reaction in which an element displaces another element from its solution.
A + BC → AC + B
- A reaction of Iron and Copper Sulphate: Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
- A reaction of Zinc and Copper Sulphate: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
- A reaction of Lead and Copper Chloride: Pb(s) + CuCl2(aq) → PbCl2(aq) + Cu(s)
Note: In the above reactions, Iron, Zinc and Lead are more reactive than Copper. Thus they displace copper from its compounds.
IV. Double Displacement Reaction
The reaction in which two different atoms or group of atoms is mutually exchanged.
AB + CD → AD + CB
- Reaction of Sodium Sulphate and Barium Chloride: Na2SO4(aq) + BaCl2(aq) → BaSO4(s)↓ + 2NaCl(aq)
V. Redox Reaction
The reaction in which one reactant gets oxidised while the other gets reduced.
AO + B → A + BO
- Oxidation: Oxidation is the gain of oxygen or loss of hydrogen.
- Reduction: Reduction is the loss of oxygen or gain of hydrogen.
- The endothermic Reaction of Copper Oxide and Hydrogen: CuO + H2 → Cu + H2O
- The reaction between Zinc Oxide and Carbon: ZnO + C → Zn + CO
- The reaction between Magnesium Oxide and Hydrogen Chloride: MnO2 + 4HCl → MnCl2 + 2H2O + Cl2
The effects of Oxidation Reactions in everyday life
When metallic substances come in contact with acid or moisture then the substances corrode and the process is called corrosion.
- Rusting is the term used for corrosion of iron.
- Corrosion causes damage to car bodies, bridges, iron railings, ships and to all objects made of metals, especially those of iron.
When the food materials containing fats or oils are left for a long time, they are oxidised i.e., they become rancid and their smell and taste changes. The phenomenon is known as rancidity.
- Usually, substances which prevent oxidation (antioxidants) are added to foods containing fats and oil.
- Keeping food in airtight containers helps to slow down oxidation.
- Chips manufacturers usually flush bags of chips with a gas such as nitrogen to prevent the chips from getting oxidised.
Important Chemical Compounds Name and Formula
|Chemical Name||Formula||Chemical Name||Formula|
|Magnesium oxide||MgO||Sulphuric acid||H2SO4|
|Hydrochloric acid||HCl||Nitric acid||HNO3|
|Potassium sulphate||K2SO4||Sodium sulphate||Na2SO4|
|Zinc sulphate||ZnSO4||Barium sulphate||BaSO4|
|Copper sulphate||CuSO4||Ferrous sulphate||FeSO4|
|Ferrous oxide||Fe2O3||Ferric oxide||Fe3O4|
|Aluminium oxide||Al2O3||Lead oxide||PbO|
|Copper oxide||CuO||Lead nitrate||Pb(NO3)2|
|Calcium nitrate||Ca(NO3)2||Sodium nitrate||NaNO3|
|Aluminium chloride||AlCl3||Barium chloride||BaCl2|
|Lead chloride||PbCl2||Copper chloride||CuCl2|
|Manganese chloride||MnCl2||Magnesium chloride||MgCl2|
|Silver chloride||AgCl||Silver bromide||AgBr|
|Barium bromide||BaBr2||Sodium Chloride||NaCl|
Special Chemical Reactions
Plants and other organisms (that have chlorophyll) takes carbon dioxide and water to produce food (glucose) for themselves and oxygen as a byproduct. The reaction can only take place under sunlight which is tapped by the green colour pigments (chlorophyll).
6CO2 + 6H2O → C6H12O6 + 6O2
Aerobic respiration involves taking oxygen from the atmosphere to break glucose molecules for the production of energy along with byproducts (water and carbon dioxide).
6O2 + C6H12O6 → 6H2O + 6CO2 + Energy
Reactions of Lime
Exothermic reaction of Quick Lime with Water
The quick lime (CaO) reacts with water (H2O) to produce slaked lime (Ca(OH2)) and large amount of heat.
CaO + H2O → Ca(OH2) + Heat
Reaction of Slaked Lime with Carbon dioxide
The slaked lime (Ca(OH2)) reacts with carbon dioxide (CO2) to produce limestone and water.
Ca(OH2) + CO2 → CaCO3 + H2O
Endothermic reaction of Limestone
The Limestone on heating produces quick lime and carbon dioxide.
CaCO3 + Heat → CaO + CO2