1 / 8

Rate Laws

Rate Laws. Chemical reactions are reversible. 2NO 2 (g) → 2NO(g) + O 2 (g). O 2 (g) + 2NO(g) → 2NO 2 (g). The reverse reaction becomes important when enough products accumulate. Now Δ [NO 2 ] depends on the difference in the rates of the forward and reverse reactions.

juana
Télécharger la présentation

Rate Laws

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Rate Laws • Chemical reactions are reversible. 2NO2(g) → 2NO(g) + O2(g) O2(g) + 2NO(g) → 2NO2(g) • The reverse reaction becomes important when enough products accumulate. • Now Δ[NO2] depends on the difference in the rates of the forward and reverse reactions. • To avoid this complication reactions are studied at a point soon after the reactants are mixed, before products build up to significant levels. • Therefore, the reaction rate will depend only on the concentrations of the reactants.

  2. 2NO2(g) → 2NO(g) + O2(g) • Neglect reverse reaction and the rate can be expressed as: Rate Law k = rate constant n = order of the reactant Both must be determined by experiment. n can be an integer (including zero) or a fraction Concentrations of products do not appear because the rate is being studied under conditions where reverse reaction does not contribute to overall rate. Value of n must be determined by experiment; it cannot be written from balanced equation.

  3. Types of Rate Laws • Differential rate law or rate law: rate depends on concentration (from previous slide). • Integrated rate law: concentrations depend on time. • Both must be determined by experiment. • If one rate law is determined, we also know the other one. • Which one is determined by experiment depends on what types of data are easiest to collect.

  4. Method of Initial Rates • Common method for determining the form of the rate law. NH4+(aq) + NO2-(aq) → N2(g) + 2H2O(l) Values of n and m can be determined by observing how the initial rate depends on the initial concentration of NH4+ and NO2-. error on chart

  5. Rate order for NO2-: use #1 and 2, [NH4+] stays the same and [NO2-] changes Ratio of rates: Value of m = 1 Reaction is first order for NO2-

  6. Rate order for NH4+: use #2 and 3, [NO2-] stays the same and [NH4+] changes Ratio of rates: The value of n is also 1.

  7. The values of m and n are both 1 and the rate law is written: • Rate = k[NH4+][NO2-] • The rate law is first order in both NO2- and NH4+. • The overall reaction order is the sum of n and m. • For this reaction, n + m = 2. • The reaction is second order overall.

  8. The value of the rate constant, k, can be calculated using the results of any of the experiments. • Using data for #1, • Rate = k[NH4+][NO2-] • 1.35 x 10-7 mol/L∙s = k(0.100 mol/L)(0.0050 mol/L) • Then Note: the units for kare dependent upon the rate orders for the reactants.

More Related