Chemical Equilibrium Review

Chemical Equilibrium

Definition of Chemical Equilibrium

Chemical equilibrium is a state in which there are no observable changes as time passes. It occurs when:

• The rates of the forward and reverse reactions are equal.
• The concentrations of the reactants and products remain constant.

Types of Equilibrium

• Physical Equilibrium: Example: H2O (liquid)
• Chemical Equilibrium: Example: N2O4 (gas)

Representation of Reactions

• Reaction Example: N2O4 (g)

Equilibrium Concentrations

Table 15.1 - Initial and Equilibrium Concentrations at 25°C
| Initial Concentration
| Equilibrium Concentration
| Ratio of Concentrations at Equilibrium
| [NO2] | [N2O4] | K (Equilibrium Constant) |
| 0.000 M | 0.670 M | 4.63 × 10^-3 |
| 0.0500 M | 0.446 M | 4.66 × 10^-3 |
| 0.0300 M | 0.500 M | 4.60 × 10^-3 |
| 0.200 M | 0.000 M | 4.63 × 10^-3 |

Law of Mass Action

The equilibrium constant expression (K) is defined as follows:
K = rac{[NO₂]^2}{[N₂O₄]}
K Values:

• $K >> 1$ indicates that products are favored.
• $K << 1$ indicates that reactants are favored.

Types of Equilibrium

1. Homogeneous Equilibrium: All reacting species in the same phase. Example:

   • N2O4 (g)
   • Equilibrium Constants:
     • K_c = rac{[NO2]^2}{[N2O4]}
     • Kp = rac{P{NO2}^2}{P_{N2O4}}
     • Kp = Kc(RT)^{ riangle n} where $ riangle n$ is the difference in moles of gaseous products and reactants.

2. Heterogeneous Equilibrium: Reactants and products in different phases. Example:

   • CaCO3 (s)
   • Equilibrium constant expression:
     • K_c = [CO2]; only gases are included.

Calculation of Constants and Examples

• For reactions at equilibrium, K is calculated by determining concentrations or partial pressures.
• Example of Equilibrium Calculations:
  • Given equilibrium pressures and using equations to find unknowns.
  • Example: For 2NO(g) + O2(g)

Le Châtelier’s Principle

• When a system at equilibrium is disturbed, the equilibrium shifts to counteract the disturbance.

Effects of Changes

1. Change in Concentration:

   • Increase of reactant shifts right; increase of product shifts left.

2. Change in Volume/Pressure:

   • Increasing pressure shifts the reaction towards the side with fewer moles of gas.

3. Change in Temperature:

   • For exothermic reactions, an increase in temperature decreases K; for endothermic, it increases K.

Summary of Le Châtelier’s Principle

Importance of Catalysts

• Adding a catalyst lowers the activation energy for both the forward and reverse reactions, thus accelerating the rate at which equilibrium is reached but does not alter the equilibrium constant.

Conclusion

Understanding chemical equilibrium is essential to predict how a system reacts to changes in concentration, pressure, and temperature. The various types of equilibrium, constants, and the principles governing the reactions provide a comprehensive outlook on the behaviors of these systems in different conditions.