Equilibria: dynamic equilibrium, Le Chatelier’s principle, equilibrium constants
Equilibria in Physical Chemistry
Dynamic Equilibrium
Imagine a busy intersection 🚗 where cars (reactants) move into the intersection and then leave it as new cars (products). In a chemical system, the forward reaction (reactants → products) and the reverse reaction (products → reactants) happen at the same rate. The system looks steady because the number of molecules on each side stays the same, but molecules are still moving back and forth. This is dynamic equilibrium – a constant dance with no net change.
- Forward and reverse rates are equal: $r_f = r_r$
- No net change in concentration of reactants or products.
- Changing the system (temperature, pressure, concentration) will disturb the balance.
Le Chatelier’s Principle ⚖️
When you add a stress to a system at equilibrium, the system will shift to counteract that stress. Think of it like a seesaw: if you put a heavy weight on one side, the other side will rise.
- Increase concentration of a reactant or product: the system shifts to consume the added species.
- Decrease concentration: the system shifts to produce more of that species.
- Change temperature: for an exothermic reaction, adding heat shifts left; for endothermic, it shifts right.
- Change pressure (for gases): increasing pressure favors the side with fewer moles of gas.
Use the “counteracting stress” rule to predict the direction of shift.
Equilibrium Constants (K)
The equilibrium constant tells us how far the reaction has gone at equilibrium.
For a general reaction: $$aA + bB \rightleftharpoons cC + dD$$
The equilibrium constant expression is:
$$K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$$
There are two common forms:
- Kc – uses concentrations (mol L⁻¹).
- Kp – uses partial pressures (atm).
Large K (≫1) means products dominate; small K (≪1) means reactants dominate.
Sample Equilibrium Table
| Reaction | Kc | Kp |
|---|---|---|
| $$\ce{N2(g) + 3H2(g) <=> 2NH3(g)}$$ | $$1.0 \times 10^{-5}$$ | $$1.0 \times 10^{-5}$$ |
| $$\ce{2SO2(g) + O2(g) <=> 2SO3(g)}$$ | $$1.2 \times 10^{3}$$ | $$2.5 \times 10^{3}$$ |
Exam Tips 📚
- Always write the equilibrium expression before plugging in numbers.
- Check units: Kc uses concentrations, Kp uses pressures.
- When applying Le Chatelier, remember the “stress” and the system’s response.
- For multi-step reactions, consider the net change in moles of gas to predict pressure effects.
- Practice converting between Kc and Kp using the relation: $$K_p = K_c (RT)^{\Delta n}$$ where Δn = moles of gaseous products – moles of gaseous reactants.
Revision
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