Chemical Equilibrium Notes

Chemical Reactions and Equilibrium

Reversible Reactions

• Most chemical reactions do not go to completion and are reversible to some extent.
• Reversible reactions can occur in either direction:
  • Reactants ⇄ Products
  • Products ⇄ Reactants

Chemical Equilibrium

• When a chemical reaction occurs in a closed container, the quantities of the components change until the composition remains unchanged, as long as the system remains undisturbed. At this point, the system is said to be at equilibrium.

Definition of Chemical Equilibrium

• Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal.
• The concentrations of the reactants and products no longer change with time, but are constant.
• It may appear as though no change is occurring at equilibrium.

Characteristics of the Chemical Equilibrium State

• The rate of the forward reaction equals the rate of the reverse reaction.
• The equilibrium position can be approached from any direction.
• Macroscopic properties such as concentrations of reactants and products are constant.
• Microscopic properties are continuous; the forward and reverse reactions continue to occur.
• If the state is disturbed by an external stress, the system adjusts to minimize the stress.

The Law Of Mass Action

• The rate of a chemical reaction is directly proportional to the product of the masses (concentrations or activities) of the reacting substances.

Equilibrium Constant

• The "Equilibrium Constant" (K) describes an equilibrium reaction.
• It is defined as the ratio of the concentration of products (raised to a power) to the concentration of reactants (raised to a power).
• For a reaction: $aA(aq) + bB(aq) ⇄ cC(aq) + dD(aq)$
• The equilibrium constant (K) is: $K = \frac{[C(aq)]^c[D(aq)]^d}{[A(aq)]^a[B(aq)]^b}$

Kc and Kp

• $K_c$ is the equilibrium constant used for aqueous solutions.
• $K_p$ is the equilibrium constant used for gases.
• The relationship between $K<em>c$ and $K</em>p$ is: $K<em>p = K</em>c(RT)^{\Delta n}$

What Does the Equilibrium Constant Tell Us?

• The value of the equilibrium constant indicates the relative quantities of reactants and products formed at equilibrium, giving us an indication of the yield.
• When:
  • K > 1: There are more products than reactants.
  • K < 1: There are more reactants than products.
  • $K ≈ 1$: Products are almost equal to reactants.

Characteristics of Equilibrium Constant

• The value of K is constant at constant temperature (T) and pressure (P).
• The value of K does not depend on the concentrations of the various species involved in the reaction.
• For a reversible reaction, K for the backward reaction is the inverse (reciprocal) of K for the forward reaction.
• A catalyst has no effect on K.
• The numerical value of K depends on the stoichiometry of the chemical equation representing the reaction.

K Value & Stoichiometry

• Consider the reaction: $A<em>2 + B</em>2 ⇄ 2AB$ (1)
• It can also be expressed as: $\frac{1}{2}A<em>2 + \frac{1}{2}B</em>2 ⇄ AB$ (2)
• If $K<em>1$ is the equilibrium constant for reaction (1) and $K</em>2$ is the equilibrium constant for reaction (2), then: $K<em>2 = \sqrt{K</em>1}$

Le Chatelier’s Principle

• Reactions occur to attain some form of stability. When they attain this state (equilibrium), they want to remain there.
• If certain conditions of the system are changed, there would be a shift in the equilibrium position.

Definition of Le Chatelier’s Principle

• If the condition(s) of the system at equilibrium changes (or “stress is applied”), the system responds in such a way to absorb or reduce that stress and achieve a new equilibrium state.

Factors Affecting Equilibrium

• Factors that affect equilibrium include:
  1. Change in concentration
  2. Change in volume and pressure
  3. Change in temperature

Change in Concentration

• Changing the concentration of a substance will shift the equilibrium to the side that would reduce that change in concentration.
• Example: $CO + 2H<em>2 ⇄ CH</em>3OH$
• If we increase [CO], the system would move in such a way as to reduce [CO]; hence the reaction proceeds to the right and [$CH_3OH$] would increase.

Change in Volume and Pressure

• Changes in pressure only affect gas phase reactions.
• Changes in pressure are attributable to changes in volume.
• The equilibrium concentrations of the products and reactants do not directly depend on the external pressure.
• However, a change in pressure due to a change in volume of the system (internal pressure) will shift the equilibrium position.
• If we were to observe the following reaction: $N<em>2 + 3H</em>2 ⇄ 2NH_3$ (4 molecules ⇄ 2 molecules)
• When the volume of the system is changed, the partial pressures of the gases change.
• Because there are more moles of gas on the reactant side, this change is more significant on the reactant side, causing the equilibrium to be disturbed.
• Thus, an increase in pressure due to decreasing volume causes the reaction to proceed to the side with the fewer moles of gas so that pressure can be reduced, and vice versa.

Change in Temperature

• The effect of changing the temperature in the equilibrium can be easily viewed by incorporating heat as either a reactant or product.
• When $∆H$ is negative (-ve), we treat heat as a product.
• When $∆H$ is positive (+ve), we treat heat as a reactant.
• Hence, we can tell whether increasing or decreasing the temperature would favor the forward or reverse reaction by applying the same notion as in concentration changes.
• Consider the equilibrium system given by: $2NO<em>2(g) ⇄ N</em>2O_4(g) \quad ∆H = -58.0 \frac{kJ}{mol}$
  • (Brown) ⇄ (Colorless)
• In this reaction, heat is a product: $2NO<em>2(g) ⇄ N</em>2O_4(g) + heat$
• If we were to increase the temperature (add heat), the equilibrium would shift to the left to reduce heat (mixture is darker brown).
• If we were to decrease the temperature (remove heat), the equilibrium would shift to the right to produce heat (gas mixture appears a paler brown).
• A decrease in temperature favors the exothermic process.
• An increase in temperature favors the endothermic process.