1. Chemical reactions can be reversible or irreversible.
2. Irreversible reactions are shown with an arrow →.
3. Example of irreversible reaction: methane combustion → CO₂ + H₂O.
4. Reversible reactions don’t finish completely, especially in a closed vessel.
5. In reversible reactions, reactant and product concentrations become constant.
6. This constant state is called chemical equilibrium.
7. Chemical equilibrium occurs when forward and reverse reactions proceed at the same rate.
8. Reactants and products continuously interconvert at equilibrium.
9. Thesystem reaches a dynamic balance with time.
10. Forward reaction converts reactants → products.
11. Forward reaction starts fast, then slows down as reactants deplete.
12. Reverse reaction converts products → reactants.
13. Reverse reaction starts slow, then speeds up until equilibrium is attained.
14. At equilibrium, concentrations of reactants/products don’t change.
15. Equilibrium constant (Kc) is constant at a given temperature.
16. Kc = [C]^c [D]^d / [A]^a [B]^b for reaction aA + bB ⇌ cC + dD.
17. Kc changes with temperature.
18. Large Kc means reaction goes nearly to completion.
19. Small Kc means reaction proceeds little in the forward direction.
20. Reaction quotient (Qc) represents the ratio of product to reactant concentrations at any moment.
21. At equilibrium, Qc = Kc.
22. Qc < Kc → reaction moves forward.
23. Qc > Kc → reaction moves backward.
24. Le Chatelier’s principle: equilibrium shifts to reduce external stress.
25. Changing concentration, pressure, or temperature disturbs equilibrium.
26. Increasing pressure shifts equilibrium toward fewer moles of gas.
27. Exothermic reactions shift left with increased temperature.
28. Endothermic reactions shift right with increased temperature.
29. Haber process: N₂ + 3H₂ ⇌ 2NH₃ (exothermic, ΔH = -ve).
30. Optimum conditions for Haber process: 200-300 atm, 400-500°C.
31. Catalyst reduces time to reach equilibrium but doesn’t change equilibrium position.
32. Solubility product (Ksp) is the product of molar concentrations of ions of a sparingly soluble salt.
33. Ksp remains constant for a given salt at a specific temperature.
34. Common ion effect reduces solubility of a sparingly soluble salt.
35. Applications of common ion effect: buffering, purification, soap precipitation, and analysis.