General Chemistry
Chemical Equilibrium
Forward and reverse reactions at equal rates — dynamic balance
Chemical equilibrium is the state where forward and reverse reaction rates are equal — concentrations don't change but reactions continue (dynamic). Reached for reversible reactions A + B ⇌ C + D. Equilibrium constant K = [products]/[reactants] (with stoichiometric exponents). K > 1: products favored. K < 1: reactants favored. Le Chatelier's principle: system disturbed reverts to maintain equilibrium. Affected by: concentration, temperature, pressure (gases), catalysts (no — only kinetics, not equilibrium position). Foundation of chemical thermodynamics.
- Reversible⇌ symbol; both directions occur
- Equilibrium constantK = [products]/[reactants] (with exponents)
- Le ChatelierSystem opposes disturbance
- K vs QQ is reaction quotient at any time; equals K at equilibrium
- CatalystsDon't change K, only how fast equilibrium reached
- TemperatureChanges K (most important factor)
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Why equilibrium matters
- Industrial. Maximize product yield (Haber, etc.).
- Biology. Reactions in cells at equilibrium or close.
- Acid-base. Buffer chemistry.
- Solubility. Ionic compounds in solution.
- Atmospheric. Pollution, ozone chemistry.
- Thermodynamics. Equilibrium is thermodynamic minimum.
- Process engineering. Design conditions for product.
Common misconceptions
- Equilibrium means stopped. Dynamic; both directions ongoing.
- Equal concentrations means equilibrium. Only if K = 1.
- K depends on concentration. Only on T and reaction.
- Adding catalyst shifts equilibrium. No — equal effect on forward and reverse.
- Le Chatelier always solves problem. Qualitative; quantitative requires K.
- Equilibrium fast. Can be very slow without catalyst.
Frequently asked questions
What's chemical equilibrium?
State where forward and reverse rates are equal. For reversible reaction aA + bB ⇌ cC + dD: rate forward = rate reverse. Concentrations no longer change (macroscopically). But reactions still happening — dynamic equilibrium. Often, both directions are happening. Eventually, system reaches steady state.
What's the equilibrium constant K?
K = [C]^c[D]^d/([A]^a[B]^b). Concentrations at equilibrium with stoichiometric exponents. K > 1: products dominate. K < 1: reactants dominate. K = 1: roughly equal. K is unitless (technically, but often reported with units). Different K values: Kc (concentrations), Kp (partial pressures for gas), K (general).
What's Le Chatelier's principle?
System at equilibrium responds to disturbance to partially counteract it. Add reactant: forward shifts. Add product: reverse shifts. Increase pressure: shift toward fewer gas moles. Increase temperature: shift away from heat (toward cooler side). Helps predict response to changes. Doesn't change K (except temperature change).
How does temperature change equilibrium?
Different temperature gives different K (because thermodynamics differs). Endothermic reaction: heat is "reactant"; increasing T increases K (more products). Exothermic: heat is "product"; increasing T decreases K. Van't Hoff equation describes T dependence. Industrial: choose T for desired K.
Do catalysts shift equilibrium?
No. Catalysts speed up both forward and reverse rates equally. Equilibrium reached faster, but at same K. K is determined by thermodynamics (ΔG); catalysts don't change ΔG. Important industrial principle: use catalysts to reach equilibrium quickly, but K (and thus product yield) determined by ΔG of reaction.
What's reaction quotient Q?
Same form as K, but with current (not equilibrium) concentrations. If Q < K: reaction not at equilibrium; will shift forward. If Q > K: shift reverse. If Q = K: at equilibrium. Useful for predicting which direction reaction will go from any state.
How does pressure affect gaseous reactions?
Increase pressure (decrease volume): system shifts toward fewer gas moles. Example: N₂ + 3H₂ ⇌ 2NH₃. Forward: 4 → 2 moles of gas. Increase P: shift forward. Decrease P: shift reverse. Inert gases (no participation in reaction) don't affect equilibrium. Pressure changes only affect gas phase reactions.