Fundamentals of Chemical Reactions
A chemical reaction is a process in which one or more substances (reactants) undergo chemical transformation to form new substances (products) with entirely different chemical properties and identities. Chemical changes involve the breaking and forming of chemical bonds between atoms.
Observational Indicators of a Chemical Reaction:
- Change in State: E.g., Combustion of wax or hydrogen and oxygen gases forming liquid water.
- Change in Color: E.g., Rusting of iron ($Fe$ turns reddish-brown) or reaction between copper sulfate solution and iron nail.
- Evolution of Gas: E.g., Zinc reacting with dilute sulfuric acid releases effervescent hydrogen gas ($H_2 \uparrow$).
- Change in Temperature: E.g., Reaction of quicklime ($CaO$) with water is vigorously exothermic, releasing intense heat.
- Formation of a Precipitate: E.g., Mixing lead nitrate and potassium iodide solutions produces a brilliant yellow precipitate of lead iodide ($PbI_2 \downarrow$).
In any chemical reaction, atoms of one element do not change into those of another element, nor do atoms disappear from the mixture. Only reorganization of chemical bonds takes place.
Balancing Chemical Equations
The Law of Conservation of Mass
According to Antoine Lavoisier's Law of Conservation of Mass, matter can neither be created nor destroyed in a chemical reaction. Therefore:
$$\sum \text{Mass of Reactants} = \sum \text{Mass of Products}$$
This implies that the total number of atoms of each element must remain exactly identical before and after a chemical reaction. An equation where atom counts are equal on both sides is called a Balanced Chemical Equation.
Step-by-Step Balancing Method (Hit-and-Trial):
Consider the reaction of iron with steam:
$$\text{Fe}(s) + \text{H}_2\text{O}(g) \rightarrow \text{Fe}_3\text{O}_4(s) + \text{H}_2(g)$$
- Step 1 (Count Atoms):
- Reactants: $\text{Fe} = 1$, $\text{H} = 2$, $\text{O} = 1$
- Products: $\text{Fe} = 3$, $\text{H} = 2$, $\text{O} = 4$
- Step 2 (Balance the element with greatest atoms - Oxygen):
Multiply $\text{H}_2\text{O}$ by 4:
$$\text{Fe} + 4\text{H}_2\text{O} \rightarrow \text{Fe}_3\text{O}_4 + \text{H}_2$$
- Step 3 (Balance Hydrogen):
Reactants now have $8\text{ H}$ atoms. Multiply product $\text{H}_2$ by 4:
$$\text{Fe} + 4\text{H}_2\text{O} \rightarrow \text{Fe}_3\text{O}_4 + 4\text{H}_2$$
- Step 4 (Balance Iron):
Products have $3\text{ Fe}$ atoms. Multiply reactant $\text{Fe}$ by 3:
$$3\text{Fe}(s) + 4\text{H}_2\text{O}(g) \rightarrow \text{Fe}_3\text{O}_4(s) + 4\text{H}_2(g)$$
- Step 5 (Verification): Both sides now possess $3\text{ Fe}$, $8\text{ H}$, and $4\text{ O}$ atoms. The equation is balanced!
Major Types of Chemical Reactions
| Reaction Type | Defining Mechanism | Representative Balanced Equation |
|---|---|---|
| Combination | Two or more reactants combine to form a single product | $\text{CaO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$ |
| Thermal Decomposition | Single reactant splits into simpler products upon heating | $2\text{FeSO}_4(s) \xrightarrow{\Delta} \text{Fe}_2\text{O}_3(s) + \text{SO}_2(g) + \text{SO}_3(g)$ |
| Electrolytic Decomposition | Decomposition driven by passage of electric current | $2\text{H}_2\text{O}(l) \xrightarrow{\text{Electricity}} 2\text{H}_2(g) + \text{O}_2(g)$ |
| Photochemical Decomposition | Decomposition triggered by absorption of light energy | $2\text{AgCl}(s) \xrightarrow{\text{Sunlight}} 2\text{Ag}(s) + \text{Cl}_2(g)$ |
| Displacement | More reactive element displaces a less reactive element | $\text{Fe}(s) + \text{CuSO}_4(aq) \rightarrow \text{FeSO}_4(aq) + \text{Cu}(s)$ |
| Double Displacement | Mutual exchange of ions between two aqueous compounds | $\text{Na}_2\text{SO}_4(aq) + \text{BaCl}_2(aq) \rightarrow \text{BaSO}_4(s)\downarrow + 2\text{NaCl}(aq)$ |
Exothermic vs Endothermic Reactions:
- Exothermic Reactions: Chemical processes that release energy into surroundings in the form of heat, light, or sound ($\Delta H < 0$).
- Respiration: $\text{C}_6\text{H}_{12}\text{O}_6(aq) + 6\text{O}_2(aq) \rightarrow 6\text{CO}_2(aq) + 6\text{H}_2\text{O}(l) + \text{Energy (38 ATP)}$
- Combustion of Natural Gas: $\text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(g) + \text{Heat}$
- Endothermic Reactions: Processes that absorb energy from surroundings ($\Delta H > 0$).
- Photosynthesis: $6\text{CO}_2 + 12\text{H}_2\text{O} \xrightarrow[\text{Chlorophyll}]{\text{Sunlight}} \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 + 6\text{H}_2\text{O}$
Oxidation, Reduction & Redox Reactions
In classical electronic and oxygen-hydrogen definitions:
- Oxidation: Addition of oxygen OR removal of hydrogen OR loss of electrons.
- Reduction: Addition of hydrogen OR removal of oxygen OR gain of electrons.
- Redox Reaction: Simultaneous oxidation of one reactant and reduction of another.
$$\text{CuO}(s) + \text{H}_2(g) \xrightarrow{\Delta} \text{Cu}(s) + \text{H}_2\text{O}(l)$$
- $\text{CuO}$ loses oxygen $\rightarrow$ Reduced to $\text{Cu}$ (Acts as the Oxidising Agent).
- $\text{H}_2$ gains oxygen $\rightarrow$ Oxidised to $\text{H}_2\text{O}$ (Acts as the Reducing Agent).
Everyday Effects of Oxidation
1. Corrosion of Metals
The slow destruction and degradation of metals caused by spontaneous chemical reactions with atmospheric gases like moisture ($H_2O$), oxygen ($O_2$), and carbon dioxide ($CO_2$).
- Rusting of Iron: Requires both air and water. Hydrated ferric oxide: $4\text{Fe} + 3\text{O}_2 + 2x\text{H}_2\text{O} \rightarrow 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$ (Reddish-brown flake).
- Corrosion of Copper: Forms a protective basic copper carbonate green coating: $2\text{Cu} + \text{CO}_2 + \text{O}_2 + \text{H}_2\text{O} \rightarrow \text{CuCO}_3 \cdot \text{Cu(OH)}_2$.
- Tarnishing of Silver: Reacts with atmospheric trace $H_2S$ forming black silver sulfide: $2\text{Ag} + \text{H}_2\text{S} \rightarrow \text{Ag}_2\text{S}(s) + \text{H}_2(g)$.
2. Rancidity
The aerobic oxidation of unsaturated fats and oils present in food items, resulting in foul odor, disagreeable taste, and degradation of nutritional value.
- Prevention Techniques:
- Packaging food with unreactive inert gases like Nitrogen ($N_2$).
- Adding artificial antioxidants such as BHA (Butylated Hydroxyanisole) and BHT.
- Storing foodstuffs in airtight containers away from direct sunlight and refrigerating.
High-Yield CBSE Board Exam Questions with Marking Scheme
- Q1 (3 Marks): A shiny brown element 'X' on heating in air becomes black in color. Name element 'X' and the black compound formed. Write the balanced chemical equation and state how the original color can be restored.
- Marking Scheme & Model Answer:
- Element 'X' is Copper ($\text{Cu}$). (1 Mark)
- Black compound formed is Copper(II) Oxide ($\text{CuO}$). (1 Mark)
- Equation: $2\text{Cu}(s) + \text{O}_2(g) \xrightarrow{\Delta} 2\text{CuO}(s)$ (1/2 Mark)
- Passing hydrogen gas over heated $\text{CuO}$ reduces it back to reddish-brown copper: $\text{CuO}(s) + \text{H}_2(g) \xrightarrow{\Delta} \text{Cu}(s) + \text{H}_2\text{O}(g)$. (1/2 Mark)
- Q2 (2 Marks): Why is the amount of gas collected in one of the test tubes during electrolysis of water double the amount collected in the other? Name this gas.
- Model Answer: Water ($\text{H}_2\text{O}$) consists of two parts hydrogen and one part oxygen by volume ($2:1$ ratio). During electrolysis: $2\text{H}_2\text{O}(l) \rightarrow 2\text{H}_2(g) + \text{O}_2(g)$. Since $2$ moles of hydrogen are liberated at the cathode for every $1$ mole of oxygen liberated at the anode, the volume of hydrogen gas collected is exactly double that of oxygen.
