Chemical Reaction
There are two types of changes happening in chemistry :
1) Physical change (Water to Ice or water to steam)
2) Chemical change (Two or more reactants form a different element )
Definition of chemical change:
A chemical reaction is a process where substances (called reactants) are
transformed into different substances (called products) through the breaking
and forming of chemical bonds.
Types of chemical reactions:
1) Directional Classification:
a. Reversible Reactions (⇌)
N₂ + 3H₂ ⇌ 2NH₃ (Haber process - making ammonia)
Can proceed in both directions
CO₂ + H₂O ⇌ H₂CO₃ (Carbon dioxide in water)
b. Irreversible Reactions (→)
Proceed only in one direction
2KClO₃ → 2KCl + 3O₂ (Potassium chlorate decomposition)
C + O₂ → CO₂ (Burning carbon)
NaOH + HCl → NaCl + H₂O (Acid-base neutralization)
2) Temperature Classification
a. Exothermic Reactions (-ΔH)
Release heat energy
C + O₂ → CO₂ + heat (Combustion of carbon)
2Na + 2H₂O → 2NaOH + H₂ + heat (Sodium in water)
CaO + H₂O → Ca(OH)₂ + heat (Lime + water)
b. Endothermic Reactions (+ΔH)
Absorb heat energy
CaCO₃ + heat → CaO + CO₂ (Limestone decomposition)
2HgO + heat → 2Hg + O₂ (Mercury oxide decomposition)
N₂ + O₂ + heat → 2NO (Nitrogen oxide formation)
3) Electron Transfer Classification
a. Oxidation-Reduction (Redox) Reactions
Involve electron transfer
Types:
Oxidation: Loss of electrons
Reduction: Gain of electrons
Examples:
Zn + Cu²⁺ → Zn²⁺ + Cu
o Zinc loses electrons (oxidized)
o Copper gains electrons (reduced)
2Fe + 3Cl₂ → 2FeCl₃
o Iron loses electrons (oxidized)
o Chlorine gains electrons (reduced)
b. Non-Redox Reactions
No electron transfer occurs
Examples:
AgNO₃ + NaCl → AgCl + NaNO₃ (Precipitation reaction)
HCl + NaOH → NaCl + H₂O (Acid-base neutralization)
BaCl₂ + Na₂SO₄ → BaSO₄ + 2NaCl (Double displacement)
Chemical Equilibrium
Chemical equilibrium occurs when the rate of the forward reaction equals the
rate of the reverse reaction in a reversible chemical reaction.
Factors of Chemical Reactions
Chemical reactions are influenced by several key factors that determine how
fast they occur, whether they happen at all, and how far they proceed. Here
are the main factors:
1. Concentration
Effect: Higher concentration of reactants increases reaction rate
Why: More particles in each space means more frequent collisions between
reactant molecules
Examples:
Hydrochloric acid + Zinc: 2M HCl reacts faster with zinc than 0.1M
HCl
Combustion: Pure oxygen makes things burn much faster than air
(21% oxygen)
Cooking: Concentrated lemon juice works faster to "cook" fish in
ceviche
2. Temperature
Effect: Higher temperature increases reaction rate exponentially
Why:
Particles move faster and collide more frequently
More particles have enough energy to overcome activation energy
barrier
Examples:
Food spoilage: Milk spoils faster at room temperature than
refrigerated
Cooking: Higher oven temperature cooks food faster
Rule of thumb: Every 10°C increase roughly doubles reaction rate
3. Surface Area
Effect: Larger surface area increases reaction rate
Why: More surface exposed means more contact points for reaction
Examples:
Antacid tablets: Chewable tablets work faster than whole tablets
Sugar: Powdered sugar dissolves faster than sugar cubes
Firewood: Kindling burns faster than logs
Digestion: Chewing food increases surface area for enzymes
4. Catalyst
Effect: Speeds up reaction without being consumed
Why: Provides alternative reaction pathway with lower activation energy
Types:
Homogeneous: Same phase as reactants (acid catalysis)
Heterogeneous: Different phase (solid catalyst with gas/liquid
reactants)
Examples:
Enzymes: Speed up biological reactions (pepsin for protein digestion)
Car catalytic converter: Converts harmful gases to less harmful ones
Haber process: Iron catalyst for ammonia production
5. Pressure (for gas reactions)
Effect: Higher pressure increases reaction rate and can shift equilibrium
Why: Compresses gas molecules closer together, increasing collision
frequency
Examples:
Haber process: High pressure favors ammonia formation
Carbonated drinks: High pressure keeps CO₂ dissolved
Pressure cooker: Higher pressure allows higher temperature, faster
cooking
Le Chatelier's Principle
What is Le Chatelier's Principle?
Le Chatelier's Principle states that when a system at equilibrium is
disturbed, it will shift to counteract the disturbance and restore
equilibrium.
Think of it like this: If you push a balanced system, it will push back to
maintain balance!
The Basic Idea
Imagine you're on a seesaw that's perfectly balanced. If someone adds
weight to one side, the seesaw will tip. But if the seesaw could adjust itself, it
would try to restore balance by adding weight to the other side.
Chemical equilibrium works similarly - when you change conditions, the
reaction shifts to oppose that change
Effect of Concentration
Rule:
Add more reactants → equilibrium shifts RIGHT (toward products)
Add more products → equilibrium shifts LEFT (toward reactants)
Remove reactants → equilibrium shifts LEFT
Remove products → equilibrium shifts RIGHT
Example: Making Ammonia
N₂ + 3H₂ ⇌ 2NH₃
What happens:
Add more N₂ → More NH₃ is produced (shifts right)
Remove NH₃ → More NH₃ is produced to replace it (shifts right)
Add more NH₃ → Some NH₃ breaks down back to N₂ and H₂ (shifts left)
Real-life analogy: If you keep eating cookies from a jar, your mom might
bake more cookies to refill it!
Effect of Temperature
Rule:
Increase temperature → favors the endothermic direction (absorbs
heat)
Decrease temperature → favors the exothermic direction (releases
heat)
Example:
N2 + O2 ⇌ 2NO (-180 KJ) This is EXOTHERMIC.
Increase temperature → shifts LEFT (backward)
Decrease temperature → shifts RIGHT (forward)
3 H2 + N2 ⇌ 2NH3 (+92KJ) This is ENDOTHERMIC.
Increase temperature → shifts RIGHT (forward)
Decrease temperature → shifts LEFT (backward)
Real-life analogy: When you're cold, you want to do activities that warm
you up. When you're hot, you want to do things that cool you down.
Effect of Pressure (for gas reactions only)
Rule:
Increase pressure → shifts toward the side with fewer gas
molecules
Decrease pressure → shifts toward the side with more gas
molecules
Example: Haber Process
N₂ + 3H₂ ⇌ 2NH₃
Count the molecules:
Left side: 1 + 3 = 4 gas molecules
Right side: 2 gas molecules
What happens:
Increase pressure → Shifts right (4 molecules → 2 molecules)
Decrease pressure → Shifts left (2 molecules → 4 molecules)
Real-life analogy: When a crowded room gets more crowded, people try to
move to a less crowded area.
Graphs & Chemical Equilibrium
Graphs of Exothermic & Endothermic Reactions
Intro to Redox Reaction
What is a Redox Reaction?
Redox stands for Reduction-Oxidation reactions. These are chemical
reactions where electrons are transferred between substances.
Oxidation
Definition: Loss of electrons
Memory trick: "OIL" = Oxidation Is Loss (of electrons)
What happens:
o Oxidation number increases
o Substance becomes more positive
Reduction
Definition: Gain of electrons
Memory trick: "RIG" = Reduction Is Gain (of electrons)
What happens:
o Oxidation number decreases
o Substance becomes more negative
Oxidizing Agent (Oxidant)
Causes oxidation in other substances
Gets reduced itself (gains electrons)
Examples: O₂, Cl₂, KMnO₄
Reducing Agent (Reductant)
Causes reduction in other substances
Gets oxidized itself (loses electrons)
Examples: H₂, metals like Zn, Mg
Rules for Assigning Oxidation Numbers:
1. Pure elements = 0 (e.g., O₂, N₂, Fe)
2. Group 1 metals = +1 (e.g., Na⁺, K⁺)
3. Group 2 metals = +2 (e.g., Mg²⁺, Ca²⁺)
4. Oxygen = -2 (except in peroxides)
5. Hydrogen = +1 (except in metal hydrides)
6. Sum of oxidation numbers = charge on ion/molecule
Metals and non metals:
Oxidation Numbers of Different Types of Elements
Group 1: Alkali Metals
Elements: Li, Na, K, Rb, Cs, Fr Oxidation Number: Always +1
Examples: Na⁺, K⁺, Li⁺
Never varies because they easily lose 1 electron
Group 2: Alkaline Earth Metals
Elements: Be, Mg, Ca, Sr, Ba, Ra Oxidation Number: Always +2
Examples: Mg²⁺, Ca²⁺, Ba²⁺
Never varies because they easily lose 2 electrons
Group 17: Halogens
Elements: F, Cl, Br, I, At Oxidation Numbers: -1, +1, +3, +5, +7
Most common: -1 (F is always -1)
Examples:
o Cl⁻ (-1), ClO⁻ (+1), ClO₂⁻ (+3), ClO₃⁻ (+5), ClO₄⁻ (+7)
Group 18: Noble Gases
Elements: He, Ne, Ar, Kr, Xe, Rn Oxidation Number: Usually 0
Rarely form compounds (except some Xe compounds)
Transition Metals (Groups 3-12)
Variable oxidation states - can have multiple values Examples:
Iron (Fe): +2, +3 (most common)
Copper (Cu): +1, +2
Manganese (Mn): +2, +3, +4, +6, +7
Chromium (Cr): +2, +3, +6
Common Non-Metals
Oxygen (O)
Most common: -2
In peroxides: -1 (H₂O₂)
In OF₂: +2
In compounds with F: positive values
Hydrogen (H)
Most common: +1
In metal hydrides: -1 (NaH, CaH₂)
Nitrogen (N)
Range: -3 to +5
Common: -3 (NH₃), 0 (N₂), +5 (HNO₃)
Carbon (C)
Range: -4 to +4
Examples: -4 (CH₄), 0 (C), +4 (CO₂)
Sulfur (S)
Range: -2 to +6
Common: -2 (H₂S), +4 (SO₂), +6 (SO₃)
Phosphorus (P)
Range: -3 to +5
Common: -3 (PH₃), +3 (PCl₃), +5 (PCl₅)
Post-Transition Metals
Aluminum (Al)
Always: +3
Tin (Sn)
Common: +2, +4
Lead (Pb)
Common: +2, +4
Bismuth (Bi)
Common: +3, +5
Quick Reference Rules
1. Pure Elements
Always: 0
Examples: O₂, N₂, Fe, Cu
2. Simple Ions
Oxidation number = charge
Examples: Na⁺ (+1), Cl⁻ (-1), Ca²⁺ (+2)
3. Compounds
Sum of all oxidation numbers = 0
For polyatomic ions = charge on ion
4 Basic Types of Redox Reactions
1. Combination (Synthesis) Reactions
Pattern: A + B → AB
What happens:
Two elements combine to form one compound
Both elements start with oxidation number 0
They change to positive and negative oxidation numbers
Examples:
2Mg + O₂ → 2MgO
o Mg: 0 → +2 (loses electrons = oxidized)
o O: 0 → -2 (gains electrons = reduced)
H₂ + Cl₂ → 2HCl
o H: 0 → +1 (oxidized)
o Cl: 0 → -1 (reduced)
Real-life example:
Formation of rust starts when iron combines with oxygen
2. Decomposition Reactions
Pattern: AB → A + B
What happens:
One compound breaks down into elements
Elements go from positive/negative oxidation numbers to 0
Examples:
2H₂O → 2H₂ + O₂ (using electricity)
o H: +1 → 0 (gains electrons = reduced)
o O: -2 → 0 (loses electrons = oxidized)
2HgO → 2Hg + O₂ (heating)
o Hg: +2 → 0 (reduced)
o O: -2 → 0 (oxidized)
Real-life example:
Electrolysis of water to produce hydrogen and oxygen gases
3. Single Displacement Reactions
Pattern: A + BC → AC + B
What happens:
One element replaces another in a compound
The free element gets oxidized
The element in the compound gets reduced
Examples:
Zn + CuSO₄ → ZnSO₄ + Cu
o Zn: 0 → +2 (loses electrons = oxidized)
o Cu: +2 → 0 (gains electrons = reduced)
Fe + CuSO₄ → FeSO₄ + Cu
o Fe: 0 → +2 (oxidized)
o Cu: +2 → 0 (reduced)
Real-life example:
Putting an iron nail in copper sulfate solution - copper metal forms on the nail
4. Combustion Reactions
Pattern: Fuel + O₂ → CO₂ + H₂O + energy
What happens:
A substance burns in oxygen
The fuel gets oxidized
Oxygen gets reduced
Energy is released as heat and light
Examples:
CH₄ + 2O₂ → CO₂ + 2H₂O (burning methane)
o C: -4 → +4 (oxidized)
o O: 0 → -2 (reduced)
C + O₂ → CO₂ (burning carbon)
o C: 0 → +4 (oxidized)
o O: 0 → -2 (reduced)
Real-life example:
Burning wood, gas stoves, car engines, candles
