Reaction Rates and Chemical Equilibrium Notes

Reaction Rates

• Particles must collide to react.
• Activation energy: minimum collision energy for a reaction to occur.
• Reaction rate depends on activation energy: higher activation energy, slower rate.

Energy Diagrams

• Endothermic: Energy increases over time
• Exothermic: Energy decreases over time

Factors Affecting Reaction Rates (4)

1. Temperature
   • Temperature increases, Reaction Rate increases
   • More collisions occur with greater force.
2. Concentration
   • Increased molarity of reactants produces more products because there is less distance between particles, leading to more collisions.
3. Particle Size
   • Smaller particle size increases surface area.
   • More collisions occur.
   • Example: Sugar dissolves faster in coffee when granulated compared to a sugar cube.
4. Type of Covalent Bond
   • Single: Fastest
   • Double:
   • Triple: Slowest (harder to break)

Catalysts and Inhibitors

• Catalysts

  • Increase reaction rate without being consumed.
  • Lower activation energy.
  • Biological catalysts = Enzymes

• Inhibitors

  • Remove or interfere with catalysts, decreasing the reaction rate.

Chemical Equilibrium

• Some reactions go to completion ($\longrightarrow$)
• Many reactions are reversible and reach equilibrium ($\rightleftharpoons$)
• Equilibrium = equal
• Favors product
• Favors reactants

Chemical Equilibrium Definition

• State where reactant and product concentrations are constant because the rates of formation and consumption are equal.
• Forward reaction rate = reverse reaction rate.
• Concentrations are constant (not necessarily equal).
• Dynamic State

Equilibrium Constant

• The equilibrium constant $(K_{eq})$ is the ratio of products to reactants at equilibrium.
• For the reaction: $aA + bB \rightleftharpoons cC + dD$
• The equilibrium constant expression is: $K_{eq} = \frac{[C]^c [D]^d}{[A]^a [B]^b}$
  • Coefficients become exponents in the $K_{eq}$ expression.
• $K_{eq} > 1$: Favors products
• $K_{eq} < 1$: Favors reactants

$K_{eq}$ Problem

• Reaction: $N<em>2 + 3H</em>2 \rightleftharpoons 2NH_3$
• At equilibrium in a 1L flask: $[H<em>2] = 0.15M$, $[N</em>2] = 0.25M$, $[NH_3] = 0.10M$
• $K_{eq} = \frac{[0.10]^2}{[0.15]^3 [0.25]} = 11.9$
• Since $K_{eq} > 1$, the products are favored.

Reverse Reaction $K_{eq}$

• Reaction: $2NH<em>3 \rightleftharpoons N</em>2 + 3H_2$
• $K_{eq} = \frac{[0.15]^3 [0.25]}{[0.10]^2} = 0.084375$
• Since $K_{eq} < 1$, the reactants are favored.
• The $K<em>{eq}$ for the reverse reaction is the inverse of the $K</em>{eq}$ for the forward reaction.

Le Chatelier’s Principle

• If stress is applied to a system at equilibrium, the system shifts to reduce the stress.
• Three types of stress:
  1. Change in Concentration
  2. Change in Pressure
  3. Change in Temperature

Concentration Stress

• Applies to gases and aqueous solutions.
• For the reaction: $A + B \rightleftharpoons C$
  • Add reactant: system shifts to favor products.
  • Add product: system shifts to favor reactants.
  • Remove reactant: system shifts to favor reactants.
  • Remove product: system shifts to favor products.

Temperature Stress

• Affects all states of matter.
• Exothermic Reaction:
  • $H<em>2 + I</em>2 \rightleftharpoons 2HI + heat$
  • Heat added: shift to favor reactants.
  • Heat removed: shift to favor products.
• Endothermic Reaction:
  • $2HI + heat \rightleftharpoons H<em>2 + I</em>2$
  • Heat added: shift to favor products.
  • Heat removed: shift to favor reactants.
• Treat heat as another reactant or product.

Pressure Stress

• Applies to gases only.
• Increase pressure: shift to the side with fewer gas molecules.
• Decrease pressure: shift to the side with more gas molecules.
• Example 1: $2NO<em>2(g) \rightleftharpoons N</em>2O_4(g)$
  • 2 gas molecules vs. 1 gas molecule, products favored with pressure increase
• Example 2: $2NH<em>3(g) \rightleftharpoons N</em>2(g) + 3H_2(g)$
  • 2 gas molecules vs. 4 gas molecules, reactants favored with pressure increase
• Example 3: $Mg(s) + 2HCl(aq) \rightleftharpoons MgCl<em>2(aq) + H</em>2(g)$
  • 0 gas molecules vs. 1 gas molecule, reactants favored with pressure increase

Spontaneous Reactions

• Reactions that occur naturally and favor product formation at equilibrium.
• Does not refer to reaction speed.

Driving Forces in Chemical Reactions

• Enthalpy ($\Delta H$): Minimum enthalpy is favored (exothermic, -$\Delta H$).

Entropy

• Entropy ($\Delta S$): Maximum entropy (disorder) is favored.
• $solid \longrightarrow liquid \longrightarrow gas$
• Least disorder to most disorder.
• Increase in # of particles = increase in disorder.
• Increase in temperature = increase in disorder.

Spontaneity

• A reaction is spontaneous when driving force(s) favor the products.
• Examples:
  • $heat + A(s) + B(s) \longrightarrow C(g) + D(aq)$
  • $A(g) + B(g) + heat \longrightarrow C(g)$
  • $A(g) \longrightarrow C(g) + D(g) + heat$
  • $A(g) + B(aq) \longrightarrow C(s) + D(aq) + heat$