뒤로Chapter 16: Chemical Equilibrium – Study Notes
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Chemical Equilibrium
Introduction to Chemical Equilibrium
Chemical equilibrium is a fundamental concept in chemistry describing the state in which the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products. This chapter explores the principles, calculations, and implications of equilibrium in chemical systems.
Hemoglobin Equilibrium System
Hemoglobin (Hb) is a protein in red blood cells that binds with oxygen (O2) to form oxyhemoglobin (HbO2) in a reversible reaction:
Reaction: Hb + O2 ↔ HbO2
The concentrations of Hb, O2, and HbO2 are interdependent.
The equilibrium constant (K) quantifies the relative concentrations at equilibrium.
A large K indicates a high concentration of products at equilibrium.
Changing the concentration of any component shifts the equilibrium to restore balance.
Oxygen Transport and Equilibrium Shifts
In the lungs: High O2 concentration shifts equilibrium to the right, forming more HbO2.
In the muscles: Low O2 concentration causes HbO2 to dissociate, increasing free O2.
Fetal Hemoglobin
Fetal hemoglobin has a larger equilibrium constant than adult hemoglobin.
It binds O2 more efficiently, facilitating oxygen transfer from mother to fetus in the placenta.
Reaction Dynamics and Reversibility
As a reaction proceeds, reactants are consumed and products are formed. The forward reaction rate decreases as reactant concentration drops, while the reverse reaction rate increases as product concentration rises. Reversible reactions can proceed in both directions.
Dynamic Equilibrium
At equilibrium, the rates of forward and reverse reactions are equal.
Concentrations of all chemicals remain constant.
Equilibrium does not mean equal concentrations of reactants and products.
The position of equilibrium may favor products or reactants depending on the value of K.
Law of Mass Action and Equilibrium Constant
The law of mass action relates the balanced chemical equation to the concentrations of reactants and products at equilibrium:
For a general reaction: aA + bB ↔ cC + dD
The equilibrium constant expression is:
K is unitless and always written as products over reactants.
Interpreting the Value of K
Large K (>1): Equilibrium favors products.
Small K (<1): Equilibrium favors reactants.
Relationships Between K and Chemical Equations
Reversing the reaction inverts K.
Multiplying coefficients by a factor raises K to that power.
Adding equations multiplies their equilibrium constants.
Equilibrium Constants for Gaseous Reactions
Equilibrium constant can be expressed in terms of partial pressures (Kp) or concentrations (Kc).
Relationship between Kp and Kc:
Where R is the gas constant, T is temperature, and Δn is the change in moles of gas.
Heterogeneous Equilibria
Pure solids and liquids are not included in the equilibrium constant expression.
Only substances in solution (aq) or gaseous state (g) are included.
Calculating Equilibrium Constants
Measure equilibrium concentrations of reactants and products.
K is independent of initial concentrations and only depends on temperature.
ICE Tables for Equilibrium Calculations
ICE tables (Initial, Change, Equilibrium) are used to organize and calculate equilibrium concentrations:
Set up initial concentrations.
Define changes based on stoichiometry.
Calculate equilibrium concentrations.
The Reaction Quotient (Q)
Q is calculated like K but with current concentrations.
Compare Q to K to predict the direction of the reaction:
Condition | Direction |
|---|---|
Q < K | Forward (toward products) |
Q > K | Reverse (toward reactants) |
Q = K | At equilibrium |
Finding Equilibrium Concentrations
Prepare ICE table.
Compare Q to K to determine direction.
Define changes in terms of x.
Solve for x using algebraic or quadratic methods.
Approximate if K is very small or large.
Le Châtelier’s Principle
Le Châtelier’s principle predicts how a system at equilibrium responds to disturbances:
If a system at equilibrium is disturbed, it shifts to minimize the disturbance.
Disturbances include changes in concentration, volume/pressure, and temperature.
Effect of Concentration Changes
Adding a reactant shifts equilibrium toward products.
Removing a reactant shifts equilibrium toward reactants.
Adding/removing solids or liquids does not affect equilibrium.
Effect of Volume and Pressure Changes
Decreasing volume (increasing pressure) shifts equilibrium toward the side with fewer gas molecules.
Increasing volume (decreasing pressure) shifts equilibrium toward the side with more gas molecules.
Adding an inert gas at constant volume has no effect.
Effect of Temperature Changes
Exothermic reactions: Heat is a product. Increasing temperature shifts equilibrium toward reactants and decreases K.
Endothermic reactions: Heat is a reactant. Increasing temperature shifts equilibrium toward products and increases K.
Summary Table: Effects on Equilibrium
Disturbance | Shift | K Value |
|---|---|---|
Add reactant | Toward products | Unchanged |
Remove reactant | Toward reactants | Unchanged |
Decrease volume | Toward fewer gas molecules | Unchanged |
Increase temperature (exothermic) | Toward reactants | Decreases |
Increase temperature (endothermic) | Toward products | Increases |
Key Equations
Equilibrium constant expression:
Relationship between Kp and Kc:
Example: Hemoglobin Equilibrium
In the lungs, high O2 concentration shifts equilibrium to form more HbO2. In the muscles, low O2 concentration causes HbO2 to dissociate, releasing O2 for cellular respiration.
