Chemical Equilibrium and Equilibrium Shifts

• N₂O₄(g) → 2 NO₂(g) has an equilibrium constant of 217 with [N₂O₄] at 1.50 x 10⁻³ M and [NO₂] at 0.571 M.
• For 2 SO₂(g) + O₂(g) ⇌ 2 SO₃(g), the [O₂] at equilibrium is 0.0313 M when [SO₂] is 2.00 M, [SO₃] is 10.0 M, and K is 800.0.

Equilibrium Shifts

• N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g) + heat:
  • Removing NH₃ shifts the equilibrium to the right.
  • Decreasing pressure shifts the equilibrium to the left.
  • Decreasing temperature shifts the equilibrium to the right.
  • Adding a catalyst does not shift the equilibrium.
• CO₂(g) + C(s) + heat ⇌ 2 CO(g):
  • Increasing temperature shifts the equilibrium to the right.
  • Adding CO shifts the equilibrium to the left.
  • Adding C does not shift the equilibrium.
  • Increasing pressure shifts the equilibrium to the left.
• CO₂(g) + H₂(g) + heat ⇌ CO(g) + H₂O(l):
  • Decreasing temperature shifts the equilibrium to the left.
  • Adding H₂ shifts the equilibrium to the right.
  • Increasing pressure shifts the equilibrium to the right.
• For SO₂ (g) + Cl₂ (g) ⇌ SO₂Cl₂ (g):
  • Adding Cl₂ shifts the equilibrium to the right.
  • Removing SO₂Cl₂ shifts the equilibrium to the right.
  • Removing SO₂ shifts the equilibrium to the left.
  • If the Kc value decreases when temperature increases, the reaction is exothermic.
  • Increasing pressure by decreasing volume shifts the equilibrium to fewer gaseous moles, decreasing the moles of Cl₂.
• For 2 HI(g) ⇌ H₂(g) + I₂(g), at 520°C, the equilibrium concentrations with K = 0.016 and an initial [HI] of 0.50 M are [HI] = 0.396 M, [H₂] = 0.052 M, and [I₂] = 0.052 M.
• For 2NOCl (g) ⇌ 2NO (g) + Cl₂ (g), if 2.50 mol of NOCl is initially in a 1.50L chamber and 28.0% dissociates, Keq is 0.0354.

Ammonia Synthesis

• For N₂ (g) + 3H₂(g) ⇌ 2NH₃ (g) with Kc = 1.2 at 375°C, starting with [H₂] = 0.76 M, [N₂] = 0.60 M, and [NH₃] = 0.48 M.:
  • The reaction shifts forward.
  • [N₂] and [H₂] decrease.
  • [NH₃] increases.
• For H₂ (g) + Br₂ (g) ⇌ 2HBr (g) with Kc = 2.18x10¹⁰ at 730°C, starting with 3.20 moles HBr in a 12.0L vessel, [H₂] = [Br₂] = 1.81x10⁻⁴ M and [HBr] = 0.267 M at equilibrium.
• BaCO₃ has a solubility of 7.00 x 10⁻⁵ mol/L; its Ksp is 4.90 x 10⁻⁹.
• CuS has a Ksp of 6.31 x 10⁻³⁶; its solubility is 2.40 x 10⁻¹⁶ g/L.
• Al(OH)₃ has a Ksp of 1.26 x 10⁻³³; its solubility is 2.61 x 10⁻⁹ mol/L.
• When 105 mL of 0.10 M AgNO₃ is added to 125 mL of 0.35 M K₂CrO₄, a precipitate of Ag₂CrO₄ will form because Qsp = 4.0 x 10⁻⁴ > Ksp = 1.12 x 10⁻¹².
• A precipitate will form when enough Ag⁺ is added to tap water with [Cl⁻] = 2 x 10⁻⁴ M to make [Ag⁺] = 1 x 10⁻⁵ M because Qsp = 2 x 10⁻⁹ > Ksp.
• The solubility of BaSO₄ in a solution already 0.1 M in SO₄²⁻ is 1 x 10⁻⁵ mol/L.

Solubility Product

• PbBr₂ has a solubility product of 8.9x10⁻⁶.
  • Molar solubility in pure water is 0.0131 M.
  • Molar solubility in 0.20 M KBr is 2.23x10⁻⁴ M.
  • Molar solubility in 0.20 M Pb(NO₃)₂ is 3.33x10⁻³ M.

Nitrogen Dioxide

• Nitrogen dioxide, dimerizes to form nitrogen tetroxide according to 2 NO₂ (g) ⇌ N₂O₄ (g).
  • Standard molar Gibbs free energy of reaction, ΔG°, is -5.31 kJ/mol.
  • The equilibrium constant, K, is 8.53 at 298 K.
• When 20.0mL of 1.0x10⁻⁵ M Ba(NO₃)₂ is added to 50.0mL of 1.0x10⁻⁵ M Na₂CO₃, BaCO₃ will not precipitate because Qsp = 2.04x 10⁻¹¹ < Ksp=5.10x10⁻⁹.
• For A (g) + B (g) ⇌ 2C (g) + heat, with 4.5 moles of A, B, and C initially in a 1.5 L flask at 27°C.:
  • The concentrations at equilibrium = K=55.1, Q=1 < Kč so rxn shifts forward, x=2.04, at EQ: [A]=[B]=0.96, [C]=7.1.
  • The Kc ratio decreases when heated because the reaction shifts reverse.
  • The Kc ratio does not change with the addition of an inert gas; only temperature changes K values.
• A saturated solution of lead contains 4.50 g of the compound per Litre. The Ksp is 1.6 x 10⁻⁵.
• For 2 Cl₂ (g) + 2 H₂O (g) ⇌ 4 HCl (g) + O₂ (g) with Keq= 752 at 480°C., with each reactant initially at 1.20 M and each product at 0.720 M, the reaction shifts right to reach equilibrium.
• For Heat + 4 CuO (s) + CH₄ (g) ⇌ CO₂ (g) + 4 Cu (s) + 2 H₂O (g) :
  • decreasing [CH₄] shifts the equilibrium to the left.
  • decreasing [CO₂] shifts the equilibrium to the right.
  • decreasing temperature shifts the equilibrium to the left.
  • decreasing volume shifts the equilibrium to the left.
  • decreasing pressure shifts the equilibrium to the right.
  • decreasing the mass of Cu (s) has no effect on the equilibrium.
• For CH₄ (g) + 2 H₂S (g) ⇌ CS₂ (g) + 4 H₂ (g), with initial [CH₄] = 0.23 M and [H₂S] = 0.34 M in a flask at 700°C. The Keq= 152 when [H₂S]eq = 0.04 M.
• For H₂O(g) + CO(g) ⇌ H₂(g) + CO₂(g) + 41.8 kJ, an equilibrium constant, K, at a temperature somewhat higher than 1000°C will be less than the constant for 1000°C.
• For 2 NO₂(g) ⇌ N₂O₄(g) + 59.0 kJ.(exothermic) :
  • Decreasing the external pressure would increase [N₂O₄(g)].
• If Q exceeds K at a given temperature, the rate of the reverse reaction is greater than the rate of the forward reaction until equilibrium is achieved.
• In a reversible reaction at equilibrium, the forward and reverse reactions do not stop.

Solubility Constant

• The Ksp for lead(II) chloride is best expressed as [Pb²⁺][Cl⁻]².

Nitrogen and Oxygen

• If starting with 0.10 M N₂ and 0.10 M O₂ in a 10.0 L container at 2000°C, with K at 0.10, 0.028 M of NO or 0.28 mol would be present at equilibrium.
  • The value of ΔG° under these conditions is 43000 J/mol.
  • Since both ΔH and ΔS are positive, increasing the temperature makes the forward process more spontaneous by making ΔG less positive.
• When adding 20.0 mL of 0.010 M Ba(NO₃)₂ to 50.0 mL of 0.0030 M Na₂CO₃, a precipitate of BaCO₃ will form because Qsp = 6.2 x 10⁻⁶ which is greater than the Ksp.
• A solution contains 1.0 x 10⁻⁴ M Cu⁺ and 2.0 x 10⁻³ M Pb²⁺, requires the least amount of I to exceed the Ksp so it will precipitate first as I is slowly added to the mixture.
• To separate Ag+, Ba²+, Cu²⁺ from a mixed solution:
  • First, add Cl⁻ to precipitate Ag+
  • Next, add SO₄²⁻ to precipitate Ba²⁺.
  • Finally, add anything that is insoluble with Cu²⁺ such as OH⁻, S²⁻, CO₃²⁻, etc.
• Successive precipitations are generally not "clean" because each involves an equilibrium rather than a process which goes to completion.
• The following reactions include:
• Sodium metal + chlorine gas → sodium chloride :
  • Na (s) + Cl₂ (g) → 2 NaCl (s) is a synthesis/redox reaction.
• Iron (II) sulfate + potassium hydroxide → iron (II) hydroxide + potassium sulfate:
  • FeSO₄ (aq) + 2 KOH (aq) → Fe(OH)₂ (s) + K₂SO₄ (aq)
  • Fe²⁺ (aq) + 2 OH⁻ (aq) → Fe(OH)₂ (s) is a precipitation reaction.
• Zinc metal + silver nitrate → zinc nitrate + silver metal:
  • Zn (s) + 2 AgNO₃ (aq) → Zn(NO₃)₂ (aq) + 2 Ag (s)
  • Zn(s) + 2 Ag⁺(aq) → Zn²⁺(aq) + 2 Ag(s) is a metal single replacement/redox reaction.
• Cobalt (III) bromide + magnesium metal will not occur because both products are aqueous.