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This unit covers equilibrium constant, Le Chatelier's principle and ICE tables — essential concepts for AP Chemistry. Use our interactive study games to test your understanding, or review questions in traditional format below.
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This unit covers equilibrium constant, Le Chatelier's principle and ICE tables — essential concepts for AP Chemistry. Use our interactive study games to test your understanding, or review questions in traditional format below.
Key Concepts Breakdown
1 Equilibrium Constant
The equilibrium constant (Kc or Kp) expresses the ratio of product concentrations to reactant concentrations, each raised to the power of their stoichiometric coefficients, at equilibrium. A large K (>>1) favors products; a small K (<<1) favors reactants. K is temperature-dependent only—changing concentration, pressure, or adding a catalyst does NOT change K.
Key Points
- For aA + bB ⇌ cC + dD: Kc = [C]^c[D]^d / [A]^a[B]^b; pure solids and liquids are omitted
- Kp = Kc(RT)^Δn, where Δn = moles of gaseous products − moles of gaseous reactants
- If a reaction is reversed, K_new = 1/K; if multiplied by n, K_new = K^n
- Q < K: reaction proceeds forward; Q > K: reaction proceeds reverse; Q = K: at equilibrium
N2(g) + 3H2(g) ⇌ 2NH3(g), Kc = 6.0 × 10^−2 at 500°C. At a given moment, [N2] = 0.10 M, [H2] = 0.30 M, [NH3] = 0.020 M. Which direction does the reaction proceed?
Calculate Q = [NH3]^2 / ([N2][H2]^3) = (0.020)^2 / (0.10)(0.30)^3 = 4.0 × 10^−4 / 2.7 × 10^−3 ≈ 0.148. Since Q (0.148) > Kc (0.060), the system has too many products relative to equilibrium. Therefore the reaction proceeds in the reverse direction, consuming NH3 and forming N2 and H2 until Q = K.
2 Le Chatelier's Principle
Le Chatelier's principle states that when a system at equilibrium is subjected to a stress, the equilibrium shifts in the direction that partially relieves that stress. Stresses include changes in concentration, pressure/volume (for gases), and temperature. Only a temperature change alters the value of K.
Key Points
- Adding a reactant or removing a product shifts equilibrium forward (toward products); the reverse is also true
- Increasing pressure (decreasing volume) shifts equilibrium toward the side with fewer moles of gas; if Δn = 0, no shift occurs
- For an exothermic reaction, increasing temperature shifts equilibrium left (decreases K); for endothermic, increasing temperature shifts right (increases K)
- Adding an inert gas at constant volume has NO effect; adding it at constant pressure shifts toward more moles of gas
Consider: 2SO2(g) + O2(g) ⇌ 2SO3(g) ΔH = −198 kJ. Predict the effect on equilibrium position of (a) increasing temperature and (b) decreasing the volume of the container.
For (a): the forward reaction is exothermic, so heat is a product; increasing temperature adds stress on the product side, shifting equilibrium left toward reactants, decreasing [SO3] and increasing K^−1 (K decreases). For (b): decreasing volume increases pressure; the left side has 3 moles of gas (2 SO2 + 1 O2) while the right has 2 moles (2 SO3); the system shifts right toward fewer moles of gas to relieve the pressure increase, producing more SO3.
3 ICE Tables
ICE (Initial, Change, Equilibrium) tables are the primary algebraic tool for calculating equilibrium concentrations when given initial conditions and K. Set up the change row using stoichiometric ratios in terms of a single variable x, then substitute the equilibrium row expressions into the K expression and solve. On the AP exam, you must recognize when the 5% approximation (x << initial concentration) is valid to avoid the quadratic.
Key Points
- Change row signs: reactants lose (−), products gain (+), scaled by stoichiometric coefficients
- 5% approximation: if K is very small (K < 10^−3) and initial concentrations are reasonable, assume x is negligible; verify: x/[initial] × 100% < 5%
- If approximation fails, use the quadratic formula; for AP, the problem usually signals which approach to use
- For Kp ICE tables, use partial pressures (atm) instead of molar concentrations
H2(g) + I2(g) ⇌ 2HI(g), Kc = 49.0 at 458°C. If 1.00 mol H2 and 1.00 mol I2 are placed in a 1.00 L flask, find the equilibrium concentrations of all species.
Set up ICE: Initial [H2] = [I2] = 1.00 M, [HI] = 0. Change: −x, −x, +2x. Equilibrium: (1.00−x), (1.00−x), 2x. Substitute into K: Kc = (2x)^2 / (1.00−x)^2 = 49.0. Taking the square root of both sides: 2x/(1.00−x) = 7.0. Solving: 2x = 7.0 − 7.0x → 9.0x = 7.0 → x = 0.778. Therefore [H2] = [I2] = 0.222 M and [HI] = 1.556 M at equilibrium.
Questions, answered.
What is Equilibrium?
Equilibrium is Unit 7 of AP Chemistry, covering equilibrium constant, Le Chatelier's principle and ICE tables.
How to study for AP Chemistry Unit 7?
Start with the Quick Summary above, review the Key Concepts, then test yourself with our interactive study games. Aim for 80%+ accuracy before moving on.
How many questions are in this unit?
This unit has 30+ review questions across 5 different game modes.