Chem 2 Rate Law

Announcements for Thursday, 06FEB2025

  • Important dates and requests:

    • Exam 1 Conflict Exam Requests and Location Requests are due by Friday, 07FEB2025, 11:59 PM (EST).

    • This applies to students with classes or Rutgers sanctioned activities during the Exam 1 period (Wednesday, 19FEB2025, 7:45 PM – 9:05 PM).

    • Also applies to students whose classes end at 7:00 PM or later on the evening of the exam.

    • Refer to the Canvas announcement from 01FEB2025 for detailed instructions.

  • For any general questions, feel free to see me after class!

Try These On Your Own

  • A series of calculations regarding rate constants was provided:

    • Activation Energy Calculation:

      • Rate constants collected at several temperatures yielded a plot of ln k vs. 1/T with a slope of -1.25 × 10^4 K.

      • The activation energy (Ea) calculated: 104 kJ/mol.

    • First Order Rate Constant:

      • Given rate constants at 25 °C (4.82 × 10^–5 s–1) and at 70 °C (1.41 × 10^–2 s–1).

      • Calculated rate constant at 90 °C: 0.11 s–1.

    • Rate Constant Factor Change:

      • For a reaction with an activation energy of 28.9 kJ/mol, determine change from 115 °C to 230 °C.

      • It increases by a factor of: 7.75.

    • Half-Life Calculation:

      • For reaction A → 2B + C with activation energy 85.6 kJ/mol and a half-life of 17.5 h at 225 °C.

      • The half-life at 166 °C calculated: 282 h.

Summarizing Collision Model Requirements for a Reaction

  • Effective collisions are necessary for reactions:

    • Particles must be properly oriented for effective reactions to occur.

    • The specific orientation for each reaction may vary.

    • Sufficient energy must be present to overcome the activation energy barrier.

    • Absence of meeting these criteria results in no reaction.

Chemical Reactions on the Molecular Level

  • Collision theory stipulates that the reactants must collide to initiate a reaction:

    • Simple reactions can proceed in one step (e.g., I– + CH3Br → CH3I + Br–).

    • More complex reactions (e.g., N2(g) + 3 H2(g) → 2 NH3(g)) often occur in multiple steps due to the improbability of all reactants colliding simultaneously with the correct orientation.

    • The individual steps contribute to what is known as the reaction mechanism.

Reaction Mechanisms

  • Mechanisms are composed of a series of elementary reactions:

    • An elementary reaction cannot be simplified further—direct particle interaction occurs.

  • Reactions may consist of rearranging or decomposing particles (e.g., NO2 → NO radical + O radical + hn).

Reaction Mechanisms (continued)

  • Validation of proposed mechanisms requires experimental evidence:

    • A mechanism can be hypothesized but can never be fully proven.

    • It can, however, be discarded if it does not align with experimental data.

  • Requirements for a plausible mechanism:

    1. The sum of individual steps must equal the overall balanced reaction.

    2. The mechanism must align with the experimentally determined rate law.

Elementary Reactions and Molecularity

  • Molecularity refers to the number of reactant particles involved in an elementary reaction:

    • Unimolecular: Contains one particle.

    • Bimolecular: Requires two particles.

    • Termolecular: Requires three particles (rare).

Determining the Rate Law of an Elementary Reaction

  • The rate law can be established directly from the stoichiometry of an elementary reaction.

Requirement 1 – Adding Up Reactions

  • Example: 2 NO(g) + 2 H2(g) → N2(g) + 2 H2O(g)

    • Proposed mechanism involves intermediate species:

      1. 2 NO ⇌ N2O2

      2. N2O2 + H2 → N2O + H2O

      3. N2O + H2 → N2 + H2O

    • Overall reaction is reconcilable with the mechanism.

Requirement 2 – the Mechanism’s Rate Law

  • Each elementary step has its own activation energy, rate constant, and associated rate law:

  • The rate of the entire mechanism is limited by the rate-determining step (slowest step) with the highest activation energy.

  • The rate law of this step must correspond with the experimentally determined rate law for viability.

Determining the Rate Law from a Mechanism

  • For the overall reaction H2(g) + 2 IBr(g) → I2(g) + 2 HBr(g), the experimentally derived rate law is:

    • rate = k [H2][IBr]

    • Proposed mechanism includes:

      1. Step 1: H2 + IBr → HI + HBr (slow)

      2. Step 2: HI + IBr → I2 + HBr (fast)

  • The rate-determining step is step 1 matching the determined rate law for relativity.

Mechanisms with a Fast Equilibrium Step

  • Example reaction: 2 NO(g) + O2(g) → 2 NO2(g)

    • Step 1: 2 NO ⇌ N2O2 (fast)

    • Step 2: N2O2 + O2 → 2 NO2 (slow)

  • The rate-determining step is step 2, but includes the reactive intermediate [N2O2], which complicates the rate law.

Mechanisms with a Fast Equilibrium Step (continued)

  • Solution to rate law challenge involves substitution using concentration relations from step 1's equilibrium, culminating in a comprehensible overall rate law:

    • Rearranging yields the overall reaction involving measurable concentrations

    • Result is: rate = k [NO]^2 [O2].