CFR 13: Chemical Equilibria and Physiological Consequences
Chemical Equilibria and Physiological Consequences
Learning Outcomes
- Define chemical equilibrium.
- Write equilibrium constant expressions (K) for reactions involving solids, liquids, and gases for both forward and reverse reactions.
- Understand the implications of K values:
- K << 1 indicates reactants are favored.
- K >> 1 indicates products are favored.
- Given a balanced chemical equation, determine the value of K using equilibrium concentrations of reactants and products.
- Given initial concentrations of reactants, find equilibrium concentrations using K.
- State Le Chatelier’s principle and apply it in various scenarios affecting equilibrium:
- Changes in concentrations,
- Temperature adjustments,
- Pressure or volume changes,
- Addition of catalysts.
- Relate to the bicarbonate blood buffer system in terms of response to hypoventilation or hyperventilation and blood acidity.
- Discuss hemoglobin and oxygen equilibria related to altitude sickness.
Equilibrium Concepts
- Reversible Reactions: Reactions that can proceed in both directions.
- Dynamic Equilibrium: State achieved when the rates of the forward and reverse reactions are equal, despite ongoing reactions.
- Physical Equilibrium: Example: Liquid water and vapor reaching equilibrium in a closed system.
Chemical Equilibrium
- Occurs when forward and reverse reaction rates are equal, leading to constant concentrations of reactants and products:
- Example: N2O4 (g) ⇌ 2NO2 (g)
- Concentrations remain steady irrespective of changes in the system until disturbed.
Equilibrium Constant Expressions (K)
- General Formula:
[ K_{eq} = \frac{[products]}{[reactants]}] - For a chemical reaction:
- aA + bB ⇌ gG + hH
--> [ K_{eq} = \frac{[G]^g[H]^h}{[A]^a[B]^b} ] - Values of K are specific to temperature and reaction type.
Special Considerations for Equilibrium Constants
- Inclusion of Phases:
- Gases: Include partial pressures in Kp expressions.
- Solids/Pure Liquids: Concentrations can be omitted from equilibrium expressions as they remain constant.
Analyzing K Values
- K >> 1: Reaction favors formation of products.
- Example: 2 H2(g) + O2(g) → 2 H2O(g)
- K << 1: Reaction favors formation of reactants.
- Example: CH3CO2H(aq) + H2O(l) ⇌ CH3COO⁻(aq) + H3O⁺(aq)
Solving Equilibrium Problems
- Determine K from equilibrium concentrations.
- Calculate equilibrium concentrations from initial concentrations using K.
Example Problems
- Finding K:
- Given equilibrium concentrations from a balanced chemical equation, calculate K.
- Example: Given H2(g) + I2(g) ⇌ 2 HI(g) and concentrations, find K.
- Using K to find equilibrium concentrations:
- Example: For PCl3(g) + Cl2(g) ⇌ PCl5(g), with initial moles in a vessel, calculate Kc.
Le Chatelier’s Principle
- If a stress is applied to a system at equilibrium, the system adjusts to counteract that stress:
- Concentration Changes:
- Increasing reactants shifts toward products;
- Decreasing reactants shifts toward reactants.
- Temperature Changes: Treat heat as a reactant/product:
- For exothermic reactions, increase in temperature shifts toward reactants.
- Pressure/Volume Changes:
- Increasing pressure shifts towards the side with fewer gas moles.
Bicarbonate Blood Buffer System
- CO2 + H2O ⇌ H+ + H2CO3 + HCO3⁻
- Changes in pH during exercise affect this equilibrium.
- Hypoventilation vs. hyperventilation effects on this system.
Altitude Sickness and Hemoglobin
- At higher altitudes, lower O2 pressure results in hypoxia:
- Hb + 4O2 ⇌ Hb(O2)4
- The equilibrium shifts to respond to reduced available O2.
Haber Process for Ammonia Production
- Continuously remove NH3 as it forms.
- Keep reactant concentrations high.
- Increase pressure to shift equilibrium right.
- Lower temperature where possible favors the exothermic reaction.
Conclusion
- Understanding chemical equilibria and Le Chatelier’s principle is crucial for predicting system responses to external changes, particularly in physiological systems like blood buffering and respiratory health.