A chemical reaction is represented by a chemical equation using the chemical formulas of reactants and products.
Example:
| Symbol | Meaning |
|---|---|
| → | Yields / produces |
| ⇌ | Reversible reaction |
| + | "And" or "combined with" |
| (s) | Solid state |
| (l) | Liquid state |
| (g) | Gaseous state |
| (aq) | Aqueous solution |
| Δ | Heat is applied |
| ↑ | Gas evolved |
| ↓ | Precipitate formed |
| hv | Light energy |
| catalyst | Catalyst is present |
A chemical equation must be balanced — the number of atoms of each element must be the same on both sides, consistent with the Law of Conservation of Mass.
Example — Balancing combustion of methane:
Unbalanced: CH₄ + O₂ → CO₂ + H₂O
Balanced:
| Element | Reactant Side | Product Side |
|---|---|---|
| C | 1 | 1 |
| H | 4 | 4 |
| O | 4 | 4 |
Definition: A reaction in which two or more substances combine to form a single product.
General Form:
Characteristics:
Examples:
a) Element + Element → Compound:
b) Compound + Element → Compound:
c) Compound + Compound → Compound:
d) Formation of Metal Oxides:
Significance:
Definition: A reaction in which a single compound breaks down into two or more simpler substances.
General Form:
Characteristics:
Types of Decomposition:
a) Thermal Decomposition (by heat):
b) Electrolytic Decomposition (by electricity):
c) Photolytic Decomposition (by light):
The decomposition of silver halides is the basis of photography.
Significance:
Definition: A reaction in which a more reactive element displaces a less reactive element from its compound.
General Form:
where A is more reactive than B.
a) Metal Displacement:
Note: A less reactive metal cannot displace a more reactive one.
The Activity Series of Metals (Most → Least Reactive):
Rule: Any metal higher in the series can displace any metal below it from its salt solution.
b) Non-Metal Displacement:
A more reactive non-metal displaces a less reactive non-metal from its compound.
Reactivity of Halogens:
c) Displacement of Hydrogen from Acids:
Note: Metals below hydrogen in the activity series (Cu, Ag, Au, Pt) do not displace hydrogen from dilute acids.
Definition: A reaction in which there is an exchange of ions between two compounds, typically in aqueous solution.
General Form:
a) Precipitation Reaction:
Solubility Rules (for predicting precipitates):
| Generally Soluble | Exceptions |
|---|---|
| All Na⁺, K⁺, NH₄⁺ salts | None |
| All nitrate (NO₃⁻) salts | None |
| Most chloride salts | AgCl, PbCl₂, Hg₂Cl₂ are insoluble |
| Most sulfate salts | BaSO₄, PbSO₄, CaSO₄ are insoluble |
| Most carbonate salts | Only Na₂CO₃, K₂CO₃, (NH₄)₂CO₃ are soluble |
| Most hydroxide salts | Only NaOH, KOH, Ca(OH)₂ (slightly) are soluble |
b) Neutralization Reaction (Acid-Base Reaction):
Net ionic equation (strong acid + strong base):
c) Gas-Forming Reactions:
Definition: Reactions involving the transfer of electrons between species.
#### Key Concepts:
| Term | Definition | Change in Oxidation Number |
|---|---|---|
| Oxidation | Loss of electrons | Increase in oxidation number |
| Reduction | Gain of electrons | Decrease in oxidation number |
| Oxidizing Agent | The species that is reduced (gains electrons) | — |
| Reducing Agent | The species that is oxidized (loses electrons) | — |
Mnemonic: OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons)
#### Example:
#### Types of Redox Reactions:
a) Direct Combination:
Mg: 0 → +2 (oxidized); O: 0 → −2 (reduced)
b) Combustion:
c) Corrosion:
d) Thermite Reaction:
Al: 0 → +3 (oxidized, reducing agent); Fe: +3 → 0 (reduced, oxidizing agent). Extremely exothermic — used in welding.
e) Disproportionation:
A single element is simultaneously oxidized and reduced in the same reaction.
f) Balancing Redox Reactions — Half-Reaction Method:
Example (Acidic solution):
Oxidation: Fe²⁺ → Fe³⁺ + e⁻
Reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
Final balanced:
Definition: Reactions that release energy (usually as heat) to the surroundings.
Characteristics:
Examples:
| Reaction | ΔH (kJ/mol) |
|---|---|
| Combustion of methane: CH₄ + 2O₂ → CO₂ + 2H₂O | −890 |
| Neutralization: HCl + NaOH → NaCl + H₂O | −57.1 |
| Formation of water: 2H₂ + O₂ → 2H₂O | −572 |
| Rusting of iron: 4Fe + 3O₂ → 2Fe₂O₃ | −1648 |
| Respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O | −2870 |
Definition: Reactions that absorb energy from the surroundings.
Characteristics:
Examples:
| Reaction | ΔH (kJ/mol) |
|---|---|
| Decomposition of CaCO₃: CaCO₃ → CaO + CO₂ | +178 |
| Photosynthesis: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂ | +2870 |
| Dissolution of NH₄NO₃ | +25.7 |
| Melting of ice | +6.01 |
Application: Instant cold packs use the endothermic dissolution of NH₄NO₃ in water.
Brønsted-Lowry Definition:
Lewis Definition:
Strength of Acids and Bases:
| Strong Acids | Strong Bases |
|---|---|
| HCl, HBr, HI | NaOH, KOH, LiOH |
| HNO₃ | Ca(OH)₂, Ba(OH)₂ |
| H₂SO₄, HClO₄ | Sr(OH)₂ |
Strong acids/bases completely dissociate in water; weak acids/bases partially dissociate.
pH Scale:
| pH Range | Nature |
|---|---|
| 0–6.9 | Acidic |
| 7.0 | Neutral |
| 7.1–14 | Basic |
Two soluble ionic compounds in aqueous solution react to form an insoluble product (precipitate) that settles out of solution.
(Spectator ions like Na⁺ and NO₃⁻ are omitted from the net ionic equation.)
Reactions that produce a gaseous product, often observed as bubbling.
Definition: Reactions in which all reactants and products are in the same phase.
Definition: Reactions in which reactants and products exist in two or more phases.
Definition: Reactions that proceed in one direction only.
Representation: Single arrow (→)
Characteristics:
Definition: Reactions that can proceed in both the forward and reverse directions simultaneously.
Representation: Double arrow (⇌)
Characteristics:
Le Chatelier's Principle — Effects:
| Stress | Shift Direction |
|---|---|
| Increase concentration of reactant | Toward products |
| Increase concentration of product | Toward reactants |
| Increase pressure | Toward fewer moles of gas |
| Increase temperature (exothermic reaction) | Toward reactants |
| Increase temperature (endothermic reaction) | Toward products |
| Add catalyst | No shift (reaches equilibrium faster) |
Definition: A rapid reaction between a substance and oxygen, producing heat and light.
General Form:
Complete Combustion (excess O₂):
Incomplete Combustion (limited O₂):
Incomplete combustion of hydrocarbons produces carbon monoxide (CO), a toxic gas, and/or carbon (soot).
Definition: Reactions in which a molecule undergoes a structural rearrangement to form an isomer — same molecular formula, different structure.
Definition: Reactions in which small molecules (monomers) join together to form a very large molecule (polymer).
a) Addition Polymerization:
b) Condensation Polymerization:
Monomers join with the elimination of a small molecule (usually water).
n(H₂N−(CH₂)₆−NH₂) + n(HOOC−(CH₂)₄−COOH) → Nylon 6,6 + nH₂O
Definition: A reaction in which a compound is broken down by reaction with water.
Ester hydrolysis
Sucrose → Glucose + Fructose
Definition: A reaction between a carboxylic acid and an alcohol to form an ester and water.
Acetic acid + Ethanol → Ethyl acetate + Water. This is reversible; the reverse is hydrolysis.
Definition: Hydrolysis of an ester (fat/oil) by a strong base (NaOH or KOH) to produce soap and glycerol.
Definition: An anaerobic biochemical process in which sugars are converted to alcohol and CO₂ by enzymes.
Glucose → Ethanol + Carbon dioxide
Lactic acid fermentation:
Glucose → Lactic acid
| Acid | Base | Salt | Water |
|---|---|---|---|
| HCl | NaOH | NaCl | H₂O |
| H₂SO₄ | 2KOH | K₂SO₄ | 2H₂O |
| HNO₃ | Ca(OH)₂ | Ca(NO₃)₂ | 2H₂O |
| CH₃COOH | NaOH | CH₃COONa | H₂O |
Special cases:
Definition: Reactions initiated or driven by light energy (photons).
Definition: Reactions involving the conversion of chemical energy to electrical energy (or vice versa) through redox processes.
a) Galvanic (Voltaic) Cell — Spontaneous reaction producing electricity:
b) Electrolytic Cell — Non-spontaneous reaction driven by external electricity:
Definition: Reactions involving changes in the nucleus of an atom, transforming one element into another.
Types:
| Type | Description | Example |
|---|---|---|
| Alpha decay (α) | Emission of a ⁴He nucleus | Ra-226 → Rn-222 + He-4 |
| Beta decay (β⁻) | Neutron → proton + electron emission | C-14 → N-14 + e⁻ |
| Gamma emission (γ) | Release of high-energy photon | Co-60* → Co-60 + γ |
| Fission | Heavy nucleus splits into lighter nuclei | U-235 + n → Ba-141 + Kr-92 + 3n |
| Fusion | Light nuclei combine to form heavier | H-2 + H-3 → He-4 + n |
| Type | General Form | Key Feature | Example |
|---|---|---|---|
| Combination | A + B → AB | Two or more → one product | 2H₂ + O₂ → 2H₂O |
| Decomposition | AB → A + B | One compound → simpler substances | CaCO₃ → CaO + CO₂ |
| Single Displacement | A + BC → AC + B | More reactive displaces less reactive | Zn + CuSO₄ → ZnSO₄ + Cu |
| Double Displacement | AB + CD → AD + CB | Ion exchange | AgNO₃ + NaCl → AgCl↓ + NaNO₃ |
| Combustion | Fuel + O₂ → Oxides | Rapid reaction with O₂ | CH₄ + 2O₂ → CO₂ + 2H₂O |
| Acid-Base | Acid + Base → Salt + H₂O | Proton transfer | HCl + NaOH → NaCl + H₂O |
| Redox | Electron transfer | OIL RIG | Zn + Cu²⁺ → Zn²⁺ + Cu |
| Precipitation | Soluble ions → Insoluble solid | Precipitate forms | Ag⁺ + Cl⁻ → AgCl↓ |
| Esterification | Acid + Alcohol ⇌ Ester + H₂O | Ester formation | CH₃COOH + EtOH ⇌ CH₃COOEt + H₂O |
| Hydrolysis | Compound + H₂O → Products | Bond breaking by water | CH₃COOEt + H₂O → CH₃COOH + EtOH |
| Polymerization | n(Monomer) → Polymer | Monomers join | n(CH₂=CH₂) → polyethylene |
| Fermentation | Sugar → Alcohol + CO₂ | Anaerobic process | C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂ |
| Photochemical | Reactants →(hν) Products | Light-driven | 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂ |
| Nuclear | Nucleus transformation | Change in atomic number/mass | U-235 → fission products + energy |
Most reactions proceed through a series of elementary steps, each involving a small number of molecules.
Example — Decomposition of N₂O₅:
Mechanism:
Every reaction requires a minimum energy input to initiate bond breaking — this is the activation energy.
Where: k = rate constant; A = frequency factor; Eₐ = activation energy; R = gas constant; T = temperature.
Lower Eₐ → faster reaction
A catalyst increases the rate of a reaction by providing an alternative pathway with a lower activation energy, without being consumed.
Types of Catalysis:
| Type | Description | Example |
|---|---|---|
| Homogeneous | Catalyst in same phase as reactants | Acid catalysis of ester hydrolysis |
| Heterogeneous | Catalyst in different phase from reactants | Pt in catalytic converters; Fe in Haber process |
| Enzymatic | Biological catalysts (proteins) | Amylase in digestion |
| Factor | Effect | Explanation |
|---|---|---|
| Concentration | ↑ Concentration → ↑ Rate | More molecules → more collisions |
| Temperature | ↑ Temperature → ↑ Rate | More molecules exceed Eₐ |
| Surface Area | ↑ Surface Area → ↑ Rate | More exposed particles |
| Catalyst | Presence → ↑ Rate | Lowers Eₐ |
| Nature of Reactants | Varies | Ionic > covalent reactions |
| Pressure (gases) | ↑ Pressure → ↑ Rate | Higher concentration |
Where: k = rate constant; [A], [B] = concentrations; m, n = reaction orders (determined experimentally, not from stoichiometry).
| Application | Reaction Type | Details |
|---|---|---|
| Haber Process | Combination, Redox, Reversible | N₂ + 3H₂ ⇌ 2NH₃ (fertilizer production) |
| Contact Process | Combination, Redox | 2SO₂ + O₂ → 2SO₃ → H₂SO₄ |
| Downs Process | Electrolytic decomposition | 2NaCl(l) → 2Na + Cl₂ |
| Hall-Héroult Process | Electrolytic decomposition | 2Al₂O₃ → 4Al + 3O₂ |
| Cement Manufacturing | Thermal decomposition | CaCO₃ → CaO + CO₂ |
| Steel Production | Redox | Fe₂O₃ + 3CO → 2Fe + 3CO₂ (Blast furnace) |
| Process | Reaction Type | Details |
|---|---|---|
| Photosynthesis | Endothermic, Combination | 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂ |
| Respiration | Exothermic, Combustion | C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + ATP |
| Digestion | Hydrolysis | Proteins → amino acids; starch → glucose |
| Fermentation | Decomposition | Glucose → ethanol + CO₂ |
| Application | Reaction Type |
|---|---|
| Cooking food | Thermal decomposition, Maillard reaction |
| Burning fuel | Combustion |
| Rusting of iron | Oxidation (redox) |
| Batteries | Electrochemical (redox) |
| Bleaching | Oxidation |
| Baking (baking soda) | Acid-base, gas evolution |
| Photography | Photochemical decomposition |
| Antacids | Neutralization |
Chemical reactions are the fundamental processes by which matter transforms. Understanding the types of chemical reactions is essential for:
Key Takeaways:
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Advanced Types of Chemical Reactions: Redox, Photochemical, Nuclear & More is a fundamental concept in inorganic chemistry. Understanding the mechanisms, reaction conditions, and stereo-chemical outcomes is crucial for mastering organic chemistry. Our curated resources provide step-by-step visualizations to help you excel.
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What happens in a disproportionation reaction?
LEARNING SUPPORT
It is a specific type of redox reaction where a single substance acts as both the oxidizing and reducing agent. This means the same element in a single reactant is simultaneously oxidized to a higher oxidation state and reduced to a lower oxidation state.
A catalyst speeds up a reaction by providing an alternative reaction pathway that has a lower activation energy barrier. Crucially, the catalyst itself is not permanently consumed in the reaction.
Chemical reactions only involve the breaking and making of bonds between atoms via valence electrons; the identity of the elements remains unchanged. Nuclear reactions involve changes within the atomic nucleus, transforming one element into another and releasing vastly more energy.