10: Decomposition of Hydrogen Peroxide
- Page ID
- 516594
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- To conduct the catalyzed decomposition of hydrogen peroxide under various conditions.
- To determine the rate law expression for the reaction.
- To calculate the rate constant for the reaction.
- To calculate the activation energy for the reaction.
INTRODUCTION
The decomposition of hydrogen peroxide, \(\ce{H2O2}\), in an aqueous solution proceeds very slowly. For instance, a commercial bottle of 3 % hydrogen peroxide is stable for a long duration. The decomposition reaction is described by the equation:
\[ 2\,\ce{H2O2}\,(aq) \rightarrow 2\,\ce{H2O}\,(l) + \ce{O2}\,(g) \]
To accelerate this naturally slow process, several catalysts can be utilized, including potassium iodide (KI), manganese(IV) oxide (\(\ce{MnO2}\)), or the enzyme catalase.
In this experiment, the catalyzed decomposition of hydrogen peroxide is carried out in a closed vessel. By doing so, the reaction rate can be determined by monitoring the pressure increase that occurs as oxygen gas (\(\ce{O2}\)) is produced.
This approach allows for several calculations and determinations:
- The rate constant for the reaction can be calculated.
- By varying the initial molar concentration of the hydrogen peroxide solution (as detailed in Parts A, B, and C), the rate law expression for the reaction can be determined.
- By conducting the reaction at different temperatures (such as approximately 20 °C and 30 °C, as detailed in Part D), the activation energy (\(E_a\)) for the reaction can be calculated.
As the reaction proceeds, the concentration of \(\ce{H2O2}\) decreases, causing the reaction to slow down. This means your Pressure vs. Time graph will eventually curve.
- The Goal: We want the rate at the very beginning, where concentrations are known (initial mixing).
- The Method: You must fit a linear trendline only to the first 60 seconds (or the straightest early region). If you include the curved part later in the run, your slope (rate) will be artificially low.
Rate Law:
\[ rate = k[\ce{H2O2}]^m[\ce{I-}]^n \]
Initial Rate:
\[ rate = \frac{\Delta[\text{reactant}]}{\Delta t} \]
Determining Order:
\[ \frac{rate_1}{rate_2} = \left( \frac{[\text{reactant}]_2}{[\text{reactant}]_1} \right)^n \]
- 10.1: Decomposition of Hydrogen Peroxide - Experiment
- This page details an experimental procedure for decomposing hydrogen peroxide with potassium iodide, emphasizing safety precautions regarding gas pressure. It is structured in four parts that vary by concentration and temperature to measure reaction rates, including specific equipment needs and chemical preparations. Instructions cover data collection, temperature recording, and equipment cleaning, with a focus on thorough data analysis and \(R^2\) validation for reliability.
- 10.2: Decomposition of Hydrogen Peroxide - Pre-lab
- This page explores the role of catalysts in chemical reactions, focusing on how they lower activation energy and enhance reaction rates. It examines the decomposition of hydrogen peroxide and related pressure buildup in closed vessels, along with calculating molarity of hydrogen peroxide at various concentrations. The page also touches on necessary unit conversions for rate calculations, highlighting the connection between pressure, volume, and molarity in evaluating reaction rates.
- 10.3: Decomposition of Hydrogen Peroxide - Data and Report
- This page details an experimental procedure for investigating the kinetics of hydrogen peroxide decomposition catalyzed by potassium iodide. It covers data collection, initial rate determination, concentration calculation, and deriving the rate law. Additionally, the reaction order is calculated using logarithmic ratios, and the Arrhenius equation is applied to determine activation energy. Proper unit usage and rounding reaction orders are emphasized throughout the procedure.