BackChemical Equilibrium and Acids & Bases: Study Notes for General Chemistry
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Chapter 16: Chemical Equilibrium
Dynamic Equilibrium
Dynamic equilibrium occurs when the rate of the forward reaction equals the rate of the reverse reaction in a closed system. At this point, the concentrations of reactants and products remain constant over time, although both reactions continue to occur.
Definition: A state in which the forward and reverse reactions proceed at the same rate, resulting in no net change in the concentrations of reactants and products.
Example: In the reaction N2(g) + 3H2(g) \rightleftharpoons 2NH3(g), equilibrium is reached when the rate of ammonia formation equals the rate of its decomposition.
Interpreting Concentration vs. Time Graphs
Concentration vs. time graphs show how the concentrations of reactants and products change as a reaction proceeds. At equilibrium, the lines for all species become horizontal, indicating constant concentrations.
Key Point: The point where all concentration curves level off indicates equilibrium has been reached.
Reaction Quotient (Q) and Equilibrium Constant (K)
The reaction quotient (Q) and equilibrium constant (K) are mathematical expressions that describe the ratio of product and reactant concentrations at any point (Q) or at equilibrium (K).
General Expression: For a reaction aA + bB \rightleftharpoons cC + dD:
Kp: For gases, partial pressures are used:
Q vs. K: Q is calculated the same way as K, but with current (not necessarily equilibrium) concentrations or pressures.
Magnitude of the Equilibrium Constant
If K >> 1: Products are favored at equilibrium; the reaction proceeds nearly to completion.
If K << 1: Reactants are favored; very little product is formed at equilibrium.
Relating Equilibrium Constants for Related Reactions
Reversing a reaction:
Multiplying a reaction by n:
Adding reactions: Multiply their K values:
Converting Between Kp and Kc
Equation:
Where R is the gas constant (0.08206 L·atm·mol−1·K−1), T is temperature in Kelvin, and Δn is the change in moles of gas (moles of gaseous products minus moles of gaseous reactants).
Predicting Reaction Direction Using Q and K
If Q < K: The reaction proceeds forward (toward products) to reach equilibrium.
If Q > K: The reaction proceeds in reverse (toward reactants) to reach equilibrium.
If Q = K: The system is at equilibrium.
Calculating K from Equilibrium Concentrations or Pressures
Insert the equilibrium concentrations or partial pressures into the K expression for the reaction.
Example: For H2 + I2 \rightleftharpoons 2HI, if [H2] = 0.2 M, [I2] = 0.2 M, [HI] = 0.6 M at equilibrium:
ICE Tables for Equilibrium Calculations
ICE stands for Initial, Change, Equilibrium. ICE tables help organize data to solve for unknown equilibrium concentrations.
Steps:
Write the balanced equation.
Set up the ICE table with initial concentrations, changes (using variables), and equilibrium values.
Substitute equilibrium values into the K expression and solve for unknowns.
Le Chatelier’s Principle
Le Chatelier’s Principle states that if a system at equilibrium is disturbed, it will shift in the direction that minimizes the disturbance.
Concentration: Adding reactant shifts equilibrium toward products; removing reactant shifts toward reactants.
Volume/Pressure (for gases): Decreasing volume (increasing pressure) shifts equilibrium toward the side with fewer moles of gas.
Temperature: Increasing temperature favors the endothermic direction; decreasing temperature favors the exothermic direction.
Temperature Dependence of K
The value of K changes with temperature. For endothermic reactions, K increases with temperature; for exothermic reactions, K decreases with temperature.
Chapter 17 (Sections 1–5): Acids and Bases
Properties of Acids and Bases
Acids: Taste sour, turn blue litmus paper red, react with metals to produce H2 gas, conduct electricity in solution.
Bases: Taste bitter, feel slippery, turn red litmus paper blue, conduct electricity in solution.
Definitions of Acids and Bases
Arrhenius Definition:
Acid: Produces H+ ions in aqueous solution.
Base: Produces OH− ions in aqueous solution.
Brønsted-Lowry Definition:
Acid: Proton (H+) donor.
Base: Proton (H+) acceptor.
Strong vs. Weak Acids and Bases
Strong acids/bases: Completely ionize in solution.
Weak acids/bases: Partially ionize; exist in equilibrium with their ions.
Examples of strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4
Examples of strong bases: NaOH, KOH, Ba(OH)2
Identifying Strong and Weak Acids/Bases
Strong acids and bases are typically listed in tables; all others are considered weak.
Weak acids include acetic acid (CH3COOH), formic acid (HCOOH), etc.
Ionizable Protons in Carboxylic Acids
In carboxylic acids (R–COOH), the hydrogen attached to the –COOH group is the ionizable (acidic) proton.
Ranking Acid Strengths Using Ka
Ka (acid dissociation constant): Measures the strength of a weak acid.
Larger Ka: Stronger acid.
Smaller Ka: Weaker acid.
Ka Expression for a Weak Acid
For HA(aq) \rightleftharpoons H+(aq) + A−(aq):
Conjugate Acid-Base Pairs
Conjugate acid: The species formed when a base gains a proton.
Conjugate base: The species formed when an acid loses a proton.
Example: In NH3 + H2O \rightleftharpoons NH4+ + OH−:
NH3 = base, NH4+ = conjugate acid
H2O = acid, OH− = conjugate base
Writing Acid-Base Reactions
Write the reactants and products, showing the transfer of a proton from the acid to the base.
Example: CH3COOH + OH− \rightarrow CH3COO− + H2O
pH, pOH, and the Ion Product of Water (Kw)
pH:
pOH:
Relationship: (at 25°C)
Kw: (at 25°C)
pKw:
Acidic solution: pH < 7
Neutral solution: pH = 7
Basic solution: pH > 7
Calculating [H3O+], [OH−], pH, or pOH
If one value is known, use the relationships above to find the others.
Example: If pH = 3, then [H3O+] = M; [OH−] = M; pOH = 11.
