CHEM 123 Chapter 14 Flashcards

Reaction Rates and Concentration Changes

• Reaction rates are analyzed by measuring the concentration of reactants and products over time.
• Plotting concentration vs. time reveals characteristic curves for reactants and products, indicating their behavior during the reaction.
• For the equation Y ⟶ 2X, the concentration of Y decreases while X increases, demonstrating that Y is a reactant and X is a product.

Collision Theory

• The collision theory states that chemical reactions occur when reactant particles collide effectively.
• Increasing reactant concentrations enhances the frequency of collisions, thereby speeding up the reaction rate.
• A higher number of reactants leads to faster reactions compared to fewer reactants.

Reaction Order and Integration

• A reaction order can be determined by plotting concentration data in various forms ([A] vs. t, ln[A] vs. t, or 1/[A] vs. t).
• A linear plot of 1/[A] vs. t indicates a second-order reaction, with a specific equation relating concentration changes over time.
• A first-order reaction shows a direct relationship between half-life and concentration changes.

Rate Constants and Enthalpy

• The rate constant (k) can vary based on the order of the reaction.
• Endothermic and exothermic reactions have different enthalpy values, affecting activation energies for forward and reverse reactions. For instance, with ΔH = 24 kJ and Ea = 56 kJ, the activation energy for the reverse reaction is 32 kJ.

Transition States and Mechanisms

• Reaction mechanisms can be assessed using energy diagrams, revealing transition states that correspond to maxima in energy levels.
• A mechanism with faster steps can be inferred from lower activation energies in the diagram.

Average and Instantaneous Rates

• Average rates can be calculated from concentration changes over time intervals and are affected by stoichiometric relationships.
• Instantaneous rates are determined by the slope of the tangent line on a concentration vs. time graph at any given moment.

Reaction Mechanisms Consistency

• Only mechanisms that align with the observed rate laws and additional catalytic effects (such as light in specific reactions) are deemed valid.
• Mechanisms must eliminate intermediates from rate laws to maintain consistency.

Factors Affecting Rates

• Increasing temperature or adding a catalyst boosts reaction rates by enhancing molecular speeds and collision frequencies.
• Decreasing reactant concentration or surface area reduces reaction rates, as fewer particles are available for collision.

Classification of Rate Laws

• The order of reaction toward a reactant is indicated by its exponent in the rate law (e.g., for rate = k[B]^2, order in B = 2, order in A = 0).
• The overall order of the reaction is determined by summing the individual orders, guiding predictions of rate changes with concentration adjustments.

Reaction Rate Calculations

• Reaction rates respond predictably to concentration changes, and adjustments can be calculated based on exponents in the rate laws.
• Realistic initial rates can be calculated from modifications in the concentrations of reactants.

Summary of Chemical Process

• Understanding reaction dynamics, mechanisms, and the factors influencing rates helps predict product formation and optimize chemical reactions.### Reaction Rates and Orders
• Doubling the concentration of a second-order reactant increases the reaction rate by a factor of four (2²).
• Reducing concentration of a second-order reactant by half decreases the reaction rate by a factor of four (1/4).
• Calculation example: If the initial rate is 0.0890 M/s, halving concentration gives a rate of 0.0668 M/s.

Determining Reaction Order

• In rate laws, the exponent ( x ) indicates the order of the reaction with respect to a reactant ( A ).
• If the rate doubles when ( [A] ) is doubled, ( x = 1 ) (first-order relationship).
• If the rate quadruples when ( [A] ) is doubled, ( x = 2 ) (second-order relationship).
• Zero-order means the rate is independent of concentration.

Rate Constant Calculations

• For the reaction ( C_2H_5Cl \rightarrow C_2H_4 + HCl ):
  • Initial rates and concentrations lead to a rate constant ( k = 6.67 \times 10^{-30} ) s⁻¹.
• For the reaction ( A + 2B \rightarrow C + D ):
  • Understanding the initial rates shows the reaction is second-order with respect to ( A ) and zero-order with respect to ( B ), leading to ( k = 0.141 ) M⁻¹⋅s⁻¹.
• For the reaction involving ( 2NO_2 + O_3 ):
  • Calculation results in a rate constant ( k = 45000 ) M⁻¹⋅s⁻¹.

First-Order Reactions

• A first-order reaction has a rate constant.
• For ( A \rightarrow \text{products} ) with ( k = 0.550 ) s⁻¹, the time to decrease concentration from 0.790 M to 0.280 M is calculated to be 1.89 s.
• The mass of ( A ) remaining after 1.75 min with ( k = 0.0203 ) s⁻¹ is found to be 2.1 g.

Half-Life in Reactions

• The half-life of a first-order reaction is constant and can be used to determine ( k ).
• For a half-life of 202 min calculated from 14% decomposition over 44 min, ( k ) is deduced from the average lifetime.
• For a second-order reaction, decreased half-life indicates dependency on the initial concentration.

Concentration Change Over Time

• For a second-order reaction with ( k = 0.160 ) M⁻¹⋅s⁻¹, time for concentration to change from 0.990 M to 0.300 M is 14.5 s.
• For a zero-order reaction, with ( k = 0.0270 ) M⋅s⁻¹, time to decrease from 0.950 M to 0.280 M is calculated to be 24.8 s.

Plotting Concentration Data

• Different types of concentration-time plots can indicate the reaction order.
• For reaction data showing a linear plot of ( 1/[A] ) versus ( t ), the reaction order is determined as second-order with ( k = 0.0333 ) M⁻¹⋅s⁻¹.

Rate Constant Calculation Using Arrhenius Equation

• The Arrhenius equation ( k = Ae^{-Ea/(RT)} ) is essential for calculations.
• Activation energy must be converted to J/mol and temperature to Kelvin for accurate calculations.
• Live example results in ( k = 0.0332 ) s⁻¹ for a reaction at 55.0 °C with ( Ea = 84.0 ) kJ/mol and ( A = 7.83 \times 10^{11} ) s⁻¹.