Chemical Equilibrium In-Depth Notes

Initial Reaction Phase:
Reactants are consumed and products are formed.
Concentrations of reactants decrease, while product concentrations increase.
Forward reaction rate decreases as reactant concentration decreases.
Equilibrium Phase:
Products can re-form reactants if not allowed to escape, leading to an increase in the reverse reaction rate as product concentration rises.
Reversible Reactions:
Reactions that can proceed in both forward and reverse directions, represented as:
reactants ⇌ products

Equilibrium Condition:
Forward and reverse reaction rates equal out; reaction is in a state of balance.
Concentrations of all substances remain constant at equilibrium.
Key Definitions:
Equilibrium ≠ Equal:
- Some reactions favor products (almost all reactants consumed), while others favor reactants (only a small percentage of reactants converted).

General Reaction Notation:
For a reaction: aA + bB ⇌ cC + dD.
The Law of Mass Action describes the relationship between concentrations at equilibrium.
Equilibrium Constant (K):
Defined for a reaction at equilibrium:
K = Kc = ([C]^c * [D]^d) / ([A]^a * [B]^b)
Rate Law Dynamics:
Forward reaction: N2O4 (g) → 2 NO2 (g); Rate = kf[N2O4]
Reverse reaction: 2 NO2 (g) → N2O4 (g); Rate = kr[NO2]^2
At equilibrium,
kf[N2O4] = kr[NO2]^2 → Relation to K:
K = kf/kr

Interpreting K:
If K >> 1: reaction is product-favored (dominates at equilibrium).
If K << 1: reaction is reactant-favored.
For moderate K (10^-3 < K < 10^3): significant amounts of both reactants and products.
Example Calculation:
For the reaction N2(g) + 3 H2(g) ⇌ 2 NH3(g), calculate Kc from concentrations:
- [N2]eq = 0.082 M, [H2]eq = 0.66 M, [NH3]eq = 3.8 M.

Reciprocal K:
For reverse reaction, Kreversed = 1/Kforward.
K Changes with Stoichiometry:
If the reaction coefficients are doubled or halved, the equilibrium constant is raised to the respective power (K^n).
Example:
Given Kc of 0.212 for
N2O4(g) ⇌ 2 NO2(g), determine K for the reverse reaction and adjusted scales.

Overall K Calculation:
The equilibrium constant of a net reaction made up of two or more steps is the product of the equilibrium constants for the individual steps.
Example Calculation:
Perform K calculations for composed reactions involving the principles learned.

Gas Concentration and Partial Pressure:
Kc relates to Kp through
Δn = (moles of gaseous products) - (moles of gaseous reactants);
Kp = Kc(RT)^Δn.
Unitless K:
Concentrations and partial pressures are relative to standard conditions.
Example Calculation:
Convert Kc to Kp using calculated parameters.

Homogeneous vs Heterogeneous:
Homogeneous equilibria: all substances in the same phase.
Heterogeneous equilibria: include different phases; solids and pure liquids excluded from K expressions.
Example: Reaction of PbCl2(s) converting to Pb2+(aq).

Prepare ICE (Initial, Change, Equilibrium) table.
Fill in initial concentrations.
Determine changes in concentrations.
Calculate the equilibrium values.
Determine K based on equilibrium concentrations using established relationships.

Practical Example:
Use real concentration values for reactions to showcase the K calculation process.

Effect of Changing Conditions:
If disturbed by temperature, pressure, or concentration changes, the system shifts to restore equilibrium.
Concentration Changes:
- Adding reactants shifts right, removing reactants shifts left.
Volume Changes:
- Decreasing volume increases pressure, shifting towards fewer gas molecules; vice versa.
Temperature Changes:
- Exothermic: raising temperature shifts left (favoring reactants); lowering shifts right (favoring products).
- Endothermic: raising shifts right; lowering shifts left.

Kc and Kp are fundamental in predicting the behavior of chemical systems under varied conditions. Understanding their calculations and relationships is crucial for mastering chemical equilibria, as is predicting shifts in response to disturbances based on Le Châtelier’s Principle.
Constant observations of real concentration shifts provide deep insight into chemical dynamics.