The reaction between hydrogen iodide (HI) and oxygen (O₂) occurs in two steps:
HI + O2 → HIO2 (fast)
HIO2 + HI → H2O + I2 (slow)
Identify the molecularity of each elementary step.
Write the rate law for the overall reaction based on the rate-determining step.
Justify why the second step determines the rate of the overall reaction.
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Consider the reaction between carbon monoxide (CO) and chlorine (Cl2):
CO (g) + Cl2 (g) → COCl2 (g)
Explain why increasing the temperature increases the rate of reaction.
Describe how the orientation of CO and Cl2 molecules during collision affects the reaction rate.
Define what is meant by an “effective collision.”
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The graph below shows the Maxwell-Boltzmann distribution curve for a gas at 300 K.
Sketch the distribution curve for the same gas at 350 K on the same graph. Ensure your curve shows the shift in the fraction of particles with energy ≥ Ea.
Explain how the shift in the fraction of particles with energy ≥ Ea affects the reaction rate at 350 K compared to 300 K.
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The reaction between methane (CH4) and chlorine (Cl) is initiated by light:
CH4 + Cl2 → CH3Cl + HCl
Explain the role of light in initiating the reaction.
Draw a reaction energy profile for the exothermic reaction described in part (a), labelling the activation energy and ΔH.
Explain what a transition state is in a reaction.
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The decomposition of hydrogen peroxide (H2O2) is commonly catalysed by manganese(IV) oxide (MnO2):
2H2O2 → 2H2O + O2
Explain how the steric factor (p) influences the rate constant (k) by affecting molecular collisions in the decomposition of H2O2.
Use the collision model to explain why the decomposition of H2O2 is slow without a catalyst.
Predict how the addition of MnO2 changes the activation energy and the reaction pathway, including its effect on intermediates.
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The following data were collected for the rate constant (kkk) of a reaction at different temperatures:
Temperature (K) | Rate Constant, k (s−1) |
|---|---|
300 | 0.0045 |
320 | 0.0092 |
340 | 0.0178 |
The relationship between the rate constant and temperature is described by the Arrhenius equation:
k = Ae−Ea/RT
Using the Maxwell-Boltzmann distribution, explain why the reaction rate increases as the temperature rises from 300 K to 320 K.
Predict whether the sensitivity of the rate constant (k) to temperature would increase, decrease, or remain the same if the activation energy (Ea) were higher. Justify your answer.
Explain how the data provided can be used to calculate Ea.
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Consider the following reaction mechanism:
NO + O3 → NO2 + O2 (fast)
NO + NO2 → N2O3 (slow)
The overall reaction releases energy as products form. Write the overall balanced chemical equation for this reaction.
Draw the energy profile for this mechanism, clearly labelling the activation energy (Ea) for each step, the intermediate, and ΔH.
Justify why the first step has a lower activation energy than the second step.
Explain how the addition of a catalyst changes the energy profile and the reaction rate.
Explain how increasing the concentration of NO2 would affect the overall rate of the reaction. Justify your answer.
A student proposes using temperature to speed up the reaction. Predict how increasing the temperature affects the fraction of collisions with energy ≥ Ea and explain the molecular basis of this change.
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The decomposition of hydrogen peroxide (H2O2) occurs via the following mechanism:
H2O2 → 2HO. (slow)
HO. + H2O2 → H2O + HO2. (fast)
HO2. + HO. → H2O + O2
Write the rate law for the overall reaction based on the rate-determining step.
Identify the intermediate(s) in this reaction mechanism and explain your reasoning.
Explain why increasing the temperature increases the rate of this reaction, referring to the Maxwell-Boltzmann distribution.
The decomposition of H2O2 involves bond breaking in the first step. Explain how molecular collisions must have sufficient energy and correct orientation for this step to occur.
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