Equilibrium - Class 11 Chemistry

Chemical Equilibrium

Overview: Chemical and Ionic Equilibrium, Equilibrium Constants, Le Chatelier's Principle, pH, and Buffer Solutions.

1. Chemical Equilibrium

State where rate of forward rxn equals rate of backward rxn. Dynamic nature.

K_c = [C]^c[D]^d / [A]^a[B]^b

K_p = K_c(RT)^Δn_g

For Solid and Pure Liquid, active mass is 1.
Applications: Predicting extent of rxn, direction of rxn.

2. Le Chatelier's Principle

If a system at equilibrium is subjected to change in T, P, or C, equilibrium shifts to counteract the change.

Conc: Increase reactant -> Shift Forward.
Pressure: Increase P -> Shift to side with fewer moles of gas.
Temp: Increase T -> Shift to endothermic direction.
Catalyst: No effect on K, only time to reach eq.

3. Ionic Equilibrium

Acids and Bases

Arrhenius: H+ donor, OH- donor.
Bronsted-Lowry: Proton donor, Proton acceptor (Conjugate pairs).
Lewis: Electron pair acceptor, Electron pair donor.

Ionization of Water

K_w = [H+][OH-] = 10^-14 (at 298K)

pH = -log[H+]

4. Buffer & Hydrolysis

Buffer Solution

Resists change in pH. Henderson-Hasselbalch Eq.

pH = pKa + log([Salt]/[Acid])

Solubility Product (Ksp)

Product of ion conc raised to power stoichiometry in saturated solution.

IP < Ksp: Unsaturated.
IP = Ksp: Saturated.
IP > Ksp: Precipitation occurs.

Numericals - Equilibrium

Numericals

Equilibrium Constant (Kc)

Q1. PCl5 -> PCl3 + Cl2. At eq, concs are [PCl5]=1.41, [PCl3]=1.59, [Cl2]=1.59 M. Calc Kc.

Kc = [PCl3][Cl2] / [PCl5]

Kc = (1.59 × 1.59) / 1.41

Kc = 2.528 / 1.41 = 1.79.

Kp from Kc

Q2. Find Kp for Q1 at 500 K.

Kp = Kc (RT)^Δng

Δng = 2 - 1 = 1.

Kp = 1.79 × (0.0831 × 500)^1.

Kp = 1.79 × 41.55 = 74.37.

pH Calculation

Q3. Calculate pH of 0.001 M HCl.

[H+] = 10^-3 M.

pH = -log(10^-3)

pH = 3.

Weak Acid pH

Q4. Calculate pH of 0.1 M CH3COOH (Ka = 1.8 × 10^-5).

[H+] = √(Ka × C)

[H+] = √(1.8 × 10^-5 × 0.1)

[H+] = √(1.8 × 10^-6) = 1.34 × 10^-3.

pH = -log(1.34 × 10^-3) = 3 - 0.12 = 2.88.

Buffer Solution

Q5. pH of buffer with 0.1M CH3COOH and 0.1M CH3COONa. pKa = 4.74.

Henderson Eq: pH = pKa + log(Salt/Acid)

pH = 4.74 + log(0.1/0.1)

pH = 4.74 + log(1) = 4.74 + 0 = 4.74.

Solubility Product

Q6. Solubility of AgCl is 10^-5 mol/L. Calculate Ksp.

AgCl -> Ag+ + Cl-

Ksp = s × s = s².

Ksp = (10^-5)² = 10^-10.

Common Ion Effect

Q7. Calculate solubility of AgCl in 0.1 M NaCl. Ksp=10^-10.

Ksp = [Ag+][Cl-]. [Cl-] = 0.1 (from NaCl).

10^-10 = s' × 0.1.

s' = 10^-9 M.

(Solubility decreased from 10^-5 to 10^-9).

Hydrolysis of Salt

Q8. pH of 0.1 M NH4Cl. pKb(NH3) = 4.75.

Salt of SA + WB. Acidic.

pH = 7 - 1/2(pKb + log C)

pH = 7 - 1/2(4.75 + log 0.1)

pH = 7 - 1/2(4.75 - 1) = 7 - 1.875 = 5.125.

Predict Direction

Q9. Qc = 10, Kc = 5. Direction?

Qc > Kc.

Product conc is too high.

Reaction shifts BACKWARD.

Degree of Dissociation

Q10. At eq, HI is 20% dissociated. 2HI -> H2 + I2. Calc Kc.

Start: 1 mol HI. Eq: 0.8 HI, 0.1 H2, 0.1 I2.

Kc = (0.1 × 0.1) / (0.8)²

Kc = 0.01 / 0.64 = 0.0156.

Formulas & Facts - Equilibrium

Equations & Formulas

Concept	Formula
Kp vs Kc	K_p = K_c(RT)^Δn_g
Reaction Quotient	Q = P/R (at any time)
pH	-log[H+]
Kw	[H+][OH-] = 10^-14
pH + pOH	14
Henderson (Acid)	pH = pKa + log([S]/[A])
Henderson (Base)	pOH = pKb + log([S]/[B])
Ksp (AB type)	s²
Ksp (AB2 type)	4s³
Ostwald Law	α = √(Ka/C)

50 NEET Facts

Key points for Equilibrium.

1. Law of Mass Action Rate is proportional to product of active masses of reactants. (Guldberg & Waage).

2. Equilibrium Constant (K) Depends only on Temperature.

3. Unit of K Depends on stoichiometry. (Conc)^delta_n.

4. Active Mass of Solid Taken as Unity (1).

5. Catalyst Does not change K. Helps reach equilibrium faster.

6. Inert Gas (Const Vol) No effect on Equilibrium.

7. Inert Gas (Const P) Volume increases. Equilibrium shifts to side with more moles.

8. Le Chatelier (Endothermic) High Temp favors forward reaction. K increases.

9. Le Chatelier (Exothermic) Low Temp favors forward reaction. K decreases with Temp.

10. Physical Equilibrium Ice-Water: Pressure increase favors side with less volume (Water). MP of ice decreases.

11. Strong Electrolyte Completely ionized (alpha ~ 1). NaCl, HCl.

12. Weak Electrolyte Partially ionized (alpha << 1). CH3COOH, NH4OH.

13. Common Ion Effect Suppresses dissociation of weak electrolyte. Decreases solubility.

14. Conjugate Acid-Base Differ by one Proton (H+). Strong acid has weak conjugate base.

15. Amphoteric Solvent Water acts as acid with NH3 and base with HCl.

16. Leveling Effect All strong acids appear equally strong in water. Differentiate in Glacial Acetic Acid.

17. pH scale Sorenson. 0 to 14. Can be negative for Conc > 1M.

18. Temperature on pH Kw increases with T. Neutral pH (7) decreases as T increases. (Water is neutral at any T, but pH < 7 at high T).

19. Acidic Buffer Weak Acid + Salt of WA-SB. pH < 7.

20. Basic Buffer Weak Base + Salt of WB-SA. pH > 7.

21. Buffer Capacity Max when [Salt] = [Acid]. pH = pKa.

22. Blood Buffer H2CO3 / HCO3- system. Maintains pH ~ 7.4.

23. Salt Hydrolysis SA+SB Neutral. No hydrolysis. pH = 7. (NaCl).

24. Salt Hydrolysis WA+SB Anionic hydrolysis. Basic solution. pH > 7. (CH3COONa).

25. Salt Hydrolysis SA+WB Cationic hydrolysis. Acidic solution. pH < 7. (NH4Cl).

26. Salt Hydrolysis WA+WB Both hydrolyze. pH depends on Ka vs Kb.

27. Indicator Theory Ostwald's theory (WA/WB). Quinonoid theory.

28. Methyl Orange Red (Acid) to Yellow (Base). Range 3.1-4.4.

29. Phenolphthalein Colorless (Acid) to Pink (Base). Range 8.3-10.0.

30. Solubility Product Limit Applicable only to sparingly soluble salts.

31. Simultaneous Equilibrium Equilibria sharing a common species satisfy all eq constants simultaneously.

32. Coupling If K1 and K2 coupled, K_new = K1 x K2.

33. Reversing Rxn K_new = 1/K.

34. Multiplying Rxn by n K_new = K^n.

35. Vapor Density Method Degree of dissociation x = (D-d)/((n-1)d).

36. Reaction Quotient Q At equilibrium, Q = K.

37. Q < K Reactants -> Products (Forward).

38. Lewis Acid Electron deficient. BF3, AlCl3, H+.

39. Lewis Base Electron rich. NH3, H2O, OH-.

40. Soft Water Lathers easily with soap.

41. Ksp and Precipitation Precipitation starts when Ionic Product > Ksp.

42. Group II Analysis H2S in HCl. Low S2- conc precipitates only low Ksp sulfides (CuS, PbS).

43. Group IV Analysis H2S in NH4OH. High S2- conc precipitates high Ksp sulfides (ZnS, MnS).

44. Solubility of Gas in Liquid Decreases with increase in Temp. (Exothermic). Henry's Law.

45. pKa value Smaller pKa, Stronger Acid.

46. Degree of Hydrolysis (h) h = √(Kw / (Ka x C)) for WA-SB.

47. pH of Rain Water ~5.6. Acid Rain < 5.6.

48. Solubility of AgCl in NH3 Increases due to complex formation [Ag(NH3)2]+.

49. Odd Electron Species NO, NO2. Paramagnetic.

50. Triple Point Three phases coexist. Kinetic energy same in all.

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