Chemical Equilibrium- Principles, Constants, and Calculations
What Chemical Equilibrium Actually Is
Most textbooks describe chemical equilibrium as "when forward and reverse reactions occur at the same rate." That's technically correct but useless for problem-solving. Here's what you actually need to know:
Equilibrium is the point where reactants convert to products at the exact same speed as products convert back to reactants. The amounts of each substance stop changing—not because the reactions stop, but because they cancel each other out.
This is dynamic. Molecules are still reacting. The concentrations just aren't shifting anymore.
The Equilibrium Constant (K)
The equilibrium constant tells you where the balance sits. High K? Reaction favors products. Low K? Reactants dominate. K = 1? Roughly equal amounts of everything.
Writing the Expression
For a general reaction:
aA + bB ⇌ cC + dD
The equilibrium expression is:
K = [C]c[D]d / [A]a[B]b
Notice: products over reactants. Always. Coefficients become exponents.
Critical rule: Solids and liquids don't appear in the expression. Only gases and aqueous solutions count.
Kc vs Kp
Kc uses concentrations in mol/L. Kp uses partial pressures. They relate through the ideal gas equation, but most problems give you one and expect you to work with it directly.
If you need to convert:
Kp = Kc(RT)Δn
Where Δn = (moles of gaseous products) − (moles of gaseous reactants).
Le Chatelier's Principle: The Disturbance Rule
Le Chatelier's principle predicts what happens when you mess with a system at equilibrium. It states that the system will shift to counteract the change. That's it. Nothing mystical.
What Actually Shifts the Equilibrium
- Concentration changes: Adding reactant pushes toward products. Adding product pushes toward reactants. Removing stuff pulls the reaction toward whatever you removed.
- Temperature changes: This one's different. You need to know if the reaction is exothermic or endothermic. Raising temperature favors the endothermic direction. Lowering temperature favors the exothermic direction.
- Pressure changes: Only matters if gases are involved. Increasing pressure pushes toward the side with fewer gas molecules. Decreasing pressure pushes toward more gas molecules.
- Catalysts: They speed up both directions equally. Equilibrium position doesn't change. Only the time to reach equilibrium changes.
Common Misconception
Students often think changing pressure alters K. It doesn't. K is temperature-dependent only. Pressure changes shift the equilibrium, but the constant value stays the same at constant temperature.
Equilibrium Constants Comparison Table
| K Value | What It Means | Practical Implication |
|---|---|---|
| K > 103 | Reaction strongly favors products | Assume essentially complete conversion |
| 10-3 < K < 103 | Significant amounts of both | Must calculate equilibrium concentrations |
| K < 10-3 | Reaction strongly favors reactants | Assume essentially no product formation |
How to Solve Equilibrium Problems
Here's the method that actually works. No guessing, no shortcuts—just systematic algebra.
Step 1: Write the Balanced Equation
Can't solve what you haven't set up. Make sure coefficients are correct before touching anything else.
Step 2: Set Up an ICE Table
ICE stands for Initial, Change, Equilibrium. It's not optional—it's how you organize what you know.
For each species, you'll track:
- Initial: Starting concentration or partial pressure (often zero for products if starting with reactants only)
- Change: How much forms or disappears (use stoichiometry—multiply by the coefficient ratio)
- Equilibrium: Initial plus change
Step 3: Plug Into the Equilibrium Expression
Substitute your equilibrium concentrations into K = [products]/[reactants]. Set it equal to the given K value.
Step 4: Solve for x
Most equilibrium problems give you a quadratic equation. You have two choices:
- Use the quadratic formula (reliable, works every time)
- Check if x is negligible (if K is very small or very large, x might be tiny enough to ignore in addition/subtraction)
If you assume x ≈ 0 and the calculated x is less than 5% of the initial value, your assumption holds. If not, solve the quadratic.
Step 5: Calculate Equilibrium Concentrations
Back-substitute your x value to find actual concentrations. Verify by plugging back into the equilibrium expression—you should get the original K value.
Getting Started: Worked Example
Problem: For the reaction N2(g) + 3H2(g) ⇌ 2NH3(g), Kc = 0.040. If you start with [N2] = 1.0 M and [H2] = 1.0 M with no ammonia, find equilibrium concentrations.
Solution:
Set up the ICE table:
| N2 | H2 | NH3 | |
|---|---|---|---|
| Initial | 1.0 | 1.0 | 0 |
| Change | −x | −3x | +2x |
| Equilibrium | 1.0−x | 1.0−3x | 2x |
Write the expression:
K = [NH3]2 / [N2][H2]3 = (2x)2 / (1.0−x)(1.0−3x)3 = 0.040
Solving this gives x ≈ 0.11 M.
Final concentrations:
- [N2] = 0.89 M
- [H2] = 0.67 M
- [NH3] = 0.22 M
The Reaction Quotient (Q)
Q uses the same formula as K, but you use current concentrations instead of equilibrium values. It tells you which direction the reaction needs to shift to reach equilibrium.
- Q < K: Reaction shifts right (toward products)
- Q > K: Reaction shifts left (toward reactants)
- Q = K: Already at equilibrium
This is useful for predicting what happens when you mix things or add/subtract species mid-reaction.
What Determines K
K depends only on temperature. Not on initial concentrations. Not on pressure. Not on catalysts. Temperature changes K. Everything else doesn't.
This is why equilibrium problems always specify temperature—it's the only variable that actually changes the constant.
Common Mistakes to Avoid
- Forgetting to include coefficients as exponents in the equilibrium expression
- Including solids and liquids in the expression (don't)
- Confusing initial concentrations with equilibrium concentrations
- Assuming x is negligible without checking the 5% rule
- Forgetting that pressure changes don't affect K
- Mixing up Kc and Kp units
Quick Reference
Before you start any equilibrium problem:
- Check what K value is given and whether it's Kc or Kp
- Identify all species—eliminate solids and liquids from the expression
- Set up the ICE table completely before writing any equations
- Know if you're expected to solve exactly or use approximations
That's the whole game. Set up correctly, solve systematically, verify your answer.