Introduction to Chemical Reactions

What are Chemical Reactions?

• Processes where atoms rearrange to form new substances.
• New substances have different properties from original ones.

How to Know if a Reaction Occurred?

• Change in state (solid, liquid, gas)
• Change in color
• Gas evolution (bubbles)
• Temperature change (gets hot or cold)
• Formation of a precipitate (solid settling in liquid)

Key Terms:

• Reactants: Starting substances.
• Products: Substances formed.

Chemical Equations - The Language of Chemistry

What is a Chemical Equation?

• A symbolic way to represent a chemical reaction.
• Uses chemical formulas for reactants and products.
• Arrow ($\rightarrow$) separates reactants from products.

Types of Equations:

• Word Equation: Names of substances (e.g., Magnesium + Oxygen $\rightarrow$ Magnesium Oxide)
• Symbolic Equation: Chemical formulas (e.g., $Mg + O_2 \rightarrow MgO$)

Slide 3: Balancing Chemical Equations

Why Balance?

• Based on the Law of Conservation of Mass: Mass is neither created nor destroyed.
• Total mass of reactants = Total mass of products.
• Number of atoms of each element must be equal on both sides.

Method:

• Hit and Trial Method (Commonly used)

Steps for Balancing:

1. Write unbalanced equation.
2. Count atoms of each element on both sides.
3. Start with elements appearing least or having most atoms.
4. General order: Metals $\rightarrow$ Non-metals $\rightarrow$ Oxygen $\rightarrow$ Hydrogen.
5. Verify the balanced equation.

Making Equations More Informative

Adding Physical States:

• (s) = solid
• (l) = liquid
• (g) = gas
• (aq) = aqueous solution (dissolved in water)

Indicating Reaction Conditions:

• Temperature ($\Delta$ for heat, or specific temp like $340 \, atm$)
• Pressure
• Catalyst (written above/below arrow)

Example: $CO(g) + 2H_2(g) \xrightarrow{340 \, atm} CH_3OH(l)$

Type 1: Combination Reactions

Definition:

• Two or more reactants combine to form a single product.

General Form:

$A + B \rightarrow AB$

Examples:

• Burning of coal: $C(s) + O_2(g) \rightarrow CO_2(g)$
• Formation of water: $2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$
• Quicklime + Water $\rightarrow$ Slaked lime: $CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq)$

Type 2: Decomposition Reactions

Definition:

• A single reactant breaks down into two or more simpler products.

General Form:

$AB \rightarrow A + B$

Requires Energy:

• Heat, light, or electricity.

Sub-types:

• Thermal Decomposition (Heat): $CaCO_3(s) \xrightarrow{Heat} CaO(s) + CO_2(g)$
• Electrolytic Decomposition (Electricity): $2H_2O(l) \xrightarrow{Electricity} 2H_2(g) + O_2(g)$
• Photolytic Decomposition (Light): $2AgCl(s) \xrightarrow{Sunlight} 2Ag(s) + Cl_2(g)$ (Used in photography)

Type 3: Displacement Reactions

Definition:

• A more reactive element displaces a less reactive element from its compound.

General Form:

$A + BC \rightarrow AC + B$

Examples:

• Iron displaces copper: $Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$
• Zinc displaces copper: $Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$

Type 4: Double Displacement Reactions

Definition:

• Exchange of ions between two compounds.

General Form:

$AB + CD \rightarrow AD + CB$

Sub-types:

• Precipitation Reaction: Forms an insoluble solid (precipitate $\downarrow$).

  Example: $BaCl_2(aq) + Na_2SO_4(aq) \rightarrow BaSO_4(s) \downarrow + 2NaCl(aq)$

• Neutralization Reaction: Acid + Base $\rightarrow$ Salt + Water.

  Example: $HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)$

Type 5: Oxidation and Reduction (Redox) Reactions

Oxidation:

• Addition of Oxygen
• Removal of Hydrogen
• Loss of Electrons (LEO)

Reduction:

• Removal of Oxygen
• Addition of Hydrogen
• Gain of Electrons (GER)

Redox Reaction:

• Oxidation and Reduction happen simultaneously.

Agents:

• Oxidizing Agent: Causes oxidation, gets reduced.
• Reducing Agent: Causes reduction, gets oxidized.

Example: $CuO(s) + H_2(g) \xrightarrow{Heat} Cu(s) + H_2O(l)$

• $CuO$ is reduced to $Cu$.
• $H_2$ is oxidized to $H_2O$.
• $CuO$ is the oxidizing agent.
• $H_2$ is the reducing agent.

Effects of Oxidation in Everyday Life - Corrosion

Definition:

• Metals slowly eaten away by air, moisture, or chemicals.

Examples:

• Rusting of Iron: Iron + Oxygen + Moisture $\rightarrow$ Hydrated Iron(III) Oxide (Rust).

  $4Fe(s) + 3O_2(g) + xH_2O(l) \rightarrow 2Fe_2O_3 \cdot xH_2O(s)$

• Black coating on silver ($Ag_2S$)
• Green coating on copper (basic copper carbonate)

Prevention:

• Painting, Oiling, Greasing
• Galvanizing (zinc coating)
• Electroplating
• Alloying

Effects of Oxidation in Everyday Life - Rancidity

Definition:

• Oxidation of fats and oils in food, leading to bad smell and taste.

Causes:

• Exposure to oxygen (air), light, moisture, bacteria.

Prevention:

• Adding antioxidants (prevent oxidation).
• Storing in airtight containers.
• Flushing with nitrogen gas (e.g., chips packets).
• Refrigeration.
• Storing away from light.