🌐 Chemical Equilibria
🔁 1. Reversible Reactions and Dynamic Equilibrium
(a) Reversible Reactions
- A reversible reaction is a chemical reaction where the products can react to reform the reactants.
- It is represented by a double arrow (⇌).
Example: N2(g)+3H2(g)⇌2NH3(g)
(b) Dynamic Equilibrium
- Occurs in closed systems.
- The rate of forward reaction = rate of reverse reaction.
- The concentration of reactants and products remains constant, though the reactions are still occurring.
(c) Need for Closed System
- A closed system prevents the loss of reactants or products, ensuring conditions remain constant and equilibrium can be established.
⚖️ 2. Le Chatelier’s Principle
If a change is made to a system at dynamic equilibrium, the position of equilibrium will shift to oppose the change and restore equilibrium.
🔁 3. Application of Le Chatelier’s Principle
Changes that affect equilibrium:
| Change | Effect (with Example) |
|---|---|
| Temperature | Exothermic: Increase shifts left; Endothermic: Increase shifts right |
| Concentration | Increase in reactant → shift to product side |
| Pressure (gases only) | Increase favors side with fewer moles of gas |
| Catalyst | No effect on position, but speeds up attainment of equilibrium |
📐 4. Equilibrium Constant (Kc)
Expression for Kc: Kc=[products]stoichiometric coefficients / [reactants]stoichiometric coefficients
Example for: H2(g)+I2(g)⇌2HI(g)
🌬️ 5. Mole Fraction and Partial Pressure
- Mole Fraction (χ): χA=mol of A / total mols of all gases
- Partial Pressure (p): pA=χA×Total Pressurep
📊 6. Equilibrium Constant (Kp)
Expression for Kp (gaseous systems only): Kp=(pproducts)coefficients / (preactants)coefficients
Example for: N2(g)+3H2(g)⇌2NH3(g)
🧮 7–8. Kc and Kp Calculations
- Use ICE (Initial, Change, Equilibrium) tables to calculate equilibrium concentrations or pressures.
- Quadratic equations not required.
- Apply Kc or Kp formula to find unknowns.
🔺 9. Effects on the Value of K
| Change | Effect on Kc or Kp |
|---|---|
| Temperature | Yes: changes value depending on exo/endothermic nature |
| Concentration / Pressure / Catalyst | No change to K value |
🏭 10. Industry Examples
Haber Process – Manufacture of Ammonia: N2+3H2⇌2NH3 ΔH=−92 kJ mol−1
- Low temp favours product but slows rate → compromise at ~450°C
- High pressure (200 atm) favours product but expensive
- Iron catalyst speeds up reaction
Contact Process – Manufacture of Sulfuric Acid: 2SO2+O2⇌2SO3 ΔH=−196 kJ mol−1
- 450°C, 2 atm, vanadium(V) oxide catalyst
🔬 Brønsted–Lowry Theory of Acids and Bases
🧪 1–2. Common Acids and Alkalis
Acids:
- HCl (Hydrochloric Acid)
- H₂SO₄ (Sulfuric Acid)
- HNO₃ (Nitric Acid)
- CH₃COOH (Ethanoic Acid)
Alkalis:
- NaOH (Sodium Hydroxide)
- KOH (Potassium Hydroxide)
- NH₃ (Ammonia)
🔄 3. Brønsted–Lowry Theory
- Acid: proton (H⁺) donor
- Base: proton (H⁺) acceptor
Example:
NH3+HCl→NH4++Cl−
💪 4–6. Strong & Weak Acids and Bases
| Type | Description | Examples |
|---|---|---|
| Strong Acid/Base | Fully dissociates | HCl, NaOH |
| Weak Acid/Base | Partially dissociates | CH₃COOH, NH₃ |
Behavior Differences:
- Strong acids react faster with metals and have lower pH
- Weak acids produce fewer ions, so lower conductivity
💧 7–8. Neutralisation and Salt Formation
- Neutralisation: H++OH−→H2O
- Salt formation: occurs when acid reacts with base/alkali/carbonate/metal
📉 9. Titration Curves
Shape of pH curves depends on:
- Strong/weak acid and strong/weak base combinations
- Buffer regions (especially in weak acid/base titrations)
- Sharp pH changes at equivalence points
🧪 10. Choosing Indicators
Suitable indicators must change color near equivalence point pH:
| Titration Type | Indicator |
|---|---|
| Strong acid + strong base | Methyl orange or phenolphthalein |
| Weak acid + strong base | Phenolphthalein |
| Strong acid + weak base | Methyl orange |
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