RRB ALP Science: Chemical Reactions and Equations

Chemical Reactions and Equations form the foundation for understanding how substances change and interact in chemistry. This topic is highly important from an exam perspective, with at least one question commonly asked. It helps students learn how reactants convert into products and how these changes are represented symbolically. Mastering this concept improves clarity in balancing equations, identifying reaction types, and understanding the principles behind chemical transformations.

Also Read: 

• RRB ALP Selection Process
• RRB ALP Syllabus

What is a Chemical Reaction?

A chemical reaction occurs when two or more reactants interact under specific conditions (e.g., heating, cooling, or catalyst addition) to form a new product. This new product possesses distinct physical and chemical properties compared to the original reactants. Essentially, if a given reactant transforms into a new substance, a chemical reaction has taken place. These reactions represent chemical changes, which are transformations leading to the formation of one or more new compounds.

Characteristics of Chemical Changes & Permanent Changes

During a chemical reaction, new substances are invariably formed. These reactions can involve a change in energy, either absorbing or releasing heat. Although there might appear to be a change in mass, equations must be balanced to uphold the principle of mass conservation. Many chemical changes are permanent and irreversible. 

For instance, an unripe mango ripening cannot revert to its unripe state (Memory Tip: Remember an unripe fruit changing state is a permanent change). Other examples include cooking of food, rusting of iron, heating of lead nitrate, souring of milk, and ripening of fruits.

Example: Rusting of Iron

Rusting of iron is a chemical reaction forming a new substance: iron oxide. This process is irreversible under normal conditions. A frequently asked exam question addresses whether the mass increases or decreases when iron rusts. The mass of the iron object increases when it rusts (Memory Tip: If an iron object initially weighs 1 kg, after rusting, its weight might increase to about 1.25 kg due to the addition of oxygen). This mass change is a consequence of the chemical change, involving the addition of oxygen.

Exothermic Reactions

An exothermic reaction is a chemical reaction that releases energy (typically heat) into its surroundings. The occurrence of the reaction is accompanied by the release of energy.

Examples:

1. Burning of Carbon: C(s) + O₂(g) → CO₂(g) + Energy

2. Formation of Ammonia (Haber Process): N₂(g) + 3H₂(g) → 2NH₃(g) + Energy. This synthesis is specifically known as the Haber Process, which is an exothermic reaction.

Endothermic Reactions

An endothermic reaction is a chemical reaction that absorbs energy (typically heat) from its surroundings. Energy must be supplied to the reactants for the reaction to proceed.

Examples:

1. Formation of Nitrogen Monoxide: N₂(g) + O₂(g) + Energy → 2NO(g)

2. Decomposition of Mercuric Oxide: 2HgO(s) + Energy → 2Hg(l) + O₂(g)

Exothermic vs. Endothermic Reactions

Representing Chemical Equations

A chemical equation provides a symbolic representation of a chemical change, using the symbols and formulas of the reactants and products. Each chemical element has a fixed symbol (e.g., H, Cl, K, O). Understanding positive and negative ionic components is crucial.

The products formed are fixed by the reactants and reaction conditions. For example, potassium permanganate (KMnO₄) reacting with hydrochloric acid (HCl) yields potassium chloride (KCl), manganese chloride (MnCl₂), water (H₂O), and chlorine gas (Cl₂).

Balancing Chemical Equations & Law of Mass Conservation

Chemical equations are typically balanced using the hit-and-trial method. This is necessary to adhere to the Law of Mass Conservation, which states that mass can neither be created nor destroyed in an isolated chemical reaction. Therefore, the total mass of reactants must equal the total mass of products. This means the number of atoms for each element must be identical on both sides of a balanced chemical equation.

Balancing Chemical Equations: Example 1

Reaction of Zinc (Zn) with Sulfuric Acid (H₂SO₄):

Unbalanced Equation: Zn + H₂SO₄ → ZnSO₄ + H₂

This equation is already balanced as each element (Zn, H, S, O) has the same number of atoms on both sides.

Balancing Chemical Equations: Example 2

Reaction of Iron (Fe) with Water (H₂O):

Unbalanced Equation: Fe + H₂O → Fe₃O₄ + H₂

1. Balance Iron (Fe): 3Fe + H₂O → Fe₃O₄ + H₂

2. Balance Oxygen (O): 3Fe + 4H₂O → Fe₃O₄ + H₂

3. Balance Hydrogen (H): 3Fe + 4H₂O → Fe₃O₄ + 4H₂
   Balanced Equation: 3Fe + 4H₂O → Fe₃O₄ + 4H₂

Types of Chemical Reactions

Chemical reactions are classified based on how reactants transform into products. Understanding these types helps identify reaction patterns, predict products, and balance equations correctly. The most common types include displacement, double displacement, oxidation, and reduction reactions.

Combination Reaction

A combination reaction occurs when two or more reactants combine to form a single new product (A + B → AB).

Examples:

1. Formation of Slaked Lime: CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat (Memory Tip: Slaked lime is used in paan; unslaked lime can burn).

2. Burning of Coal: C(s) + O₂(g) → CO₂(g)

3. Formation of Water: 2H₂(g) + O₂(g) → 2H₂O(l)

4. Formation of Carbon Dioxide: 2CO(g) + O₂(g) → 2CO₂(g)

5. Formation of Ammonium Chloride: NH₃(g) + HCl(g) → NH₄Cl(s) (Here, two compounds combine).

6. Decomposition of Vegetable Matter into Compost.

Decomposition Reaction

A decomposition reaction is when a single complex substance breaks down into two or more simpler components. These are the inverse of combination reactions.

Types:

1. Thermal Decomposition: Breakdown due to heat.

2. Electrolytic Decomposition (Electrolysis): Breakdown due to electricity.

3. Photolytic Decomposition (Photolysis): Breakdown due to light.

Electrolysis & Photolysis

• Electrolysis: An electric current passed through a substance (molten or solution) causes it to break down.

• Photolysis: A substance breaks down when exposed to light.

Examples of Decomposition Reactions

• Thermal Decomposition:

• 2HgO(s) --(Heat)--> 2Hg(l) + O₂(g)

• 2Cu(NO₃)₂(s) --(Heat)--> 2CuO(s) + 4NO₂(g) + O₂(g)

• 2FeSO₄(s) --(Heat)--> Fe₂O₃(s) + SO₂(g) + SO₃(g)

• CaCO₃(s) --(Heat)--> CaO(s) + CO₂(g)

• 2Pb(NO₃)₂(s) --(Heat)--> 2PbO(s) + 4NO₂(g) + O₂(g) (Ensure coefficients for two moles of lead nitrate are correct).

• Electrolytic Decomposition:

• Electrolysis of Water: 2H₂O(l) --(Electricity)--> 2H₂(g) + O₂(g)

• Electrolysis of Molten Lead Bromide: PbBr₂(l) --(Electricity)--> Pb(l) + Br₂(g)

• Photolytic Decomposition:

• Silver Chloride: 2AgCl(s) --(Sunlight)--> 2Ag(s) + Cl₂(g)

• Silver Bromide: 2AgBr(s) --(Sunlight)--> 2Ag(s) + Br₂(g)

Displacement Reaction

A displacement reaction occurs when one element displaces another element from its compound (a more reactive element replaces a less reactive one).

Examples:

1. Iron in Copper Sulfate: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

2. Zinc in Copper Sulfate: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

3. Lead in Copper Chloride: Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s)

Double Displacement Reaction

A double displacement reaction occurs when two compounds exchange their ions to form two new compounds. (Contrast with single displacement, where one element replaces another; here, two compounds exchange ionic pairs).

Example: Reaction of Sodium Sulfate (Na₂SO₄) with Barium Chloride (BaCl₂):

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

Oxidation Reaction

An oxidation reaction is a chemical reaction where an element or compound gains oxygen (addition of oxygen).

Example: Formation of Copper Oxide: 2Cu(s) + O₂(g) --(Heat)--> 2CuO(s)

Reduction Reaction

A reduction reaction is a chemical reaction in which an element or compound loses oxygen (removal of oxygen). (Memory Tip: Visualize oxygen "leaving" or "running away" from the compound).

Oxidizing and Reducing Agents

This concept is often a source of confusion. Remember that these agents function opposite to the process they facilitate.

• Oxidizing Agent: A substance that causes oxidation in another substance while being itself reduced.

• Reducing Agent: A substance that causes reduction in another substance while being itself oxidized.

Redox Reaction

A redox reaction (oxidation-reduction reaction) is a chemical reaction where both oxidation and reduction processes occur simultaneously within the same reaction.

Example: Reaction of Copper Oxide (CuO) with Hydrogen (H₂):

CuO(s) + H₂(g) → Cu(s) + H₂O(l)

Here, CuO is reduced (loses oxygen), and H₂ is oxidized (gains oxygen).

Real-world examples of oxidation include the rusting of iron and the spoilage of food.

Reactivity Series

The reactivity series lists metals in decreasing order of reactivity, predicting which element can displace another.

A mnemonic for a portion of the series: कनक मांगे आलू जो आयरन शीशा है

• क (Kanak) → Potassium (K)

• न (N) → Sodium (Na)

• क (K) → Calcium (Ca)

• मांगे (Maange) → Magnesium (Mg)

• आलू (Aaloo) → Aluminum (Al)

• जो (Jo) → Zinc (Zn)

• आयरन (Iron) → Iron (Fe)

• शीशा (Sheesha) → Lead (Pb)

• है (Hai) → Hydrogen (H)
  An element higher in the series can displace one lower in the series. For example, aluminum (Al) will displace zinc (Zn). Noble metals like gold (Au) and silver (Ag) are highly unreactive. Copper (Cu), platinum (Pt), and mercury (Hg) are among the least reactive metals.