Nitrogen monoxide and hydrogen react at high temperatures according to the following equation:
2NO(g)+2H2(g)→N2(g)+2H2O(g) 2\text{NO}(\text{g}) + 2\text{H}_2(\text{g}) \rightarrow \text{N}_2(\text{g}) + 2\text{H}_2\text{O}(\text{g}) 2NO(g)+2H2(g)→N2(g)+2H2O(g)Explain how increasing the temperature increases the rate of this reaction. In your answer, describe what a labelled sketch using Boltzmann distributions would show (stating the labels for both axes and how the distribution curve changes with a temperature increase).
The rate of this reaction is investigated by carrying out three experiments at a constant temperature. The results are shown in the table below:
| Experiment | [NO(g)] / mol dm−3[\text{NO}(\text{g})] \ / \ \text{mol dm}^{-3}[NO(g)] / mol dm−3 | [H2(g)] / mol dm−3[\text{H}_2(\text{g})] \ / \ \text{mol dm}^{-3}[H2(g)] / mol dm−3 | Initial rate / mol dm−3s−1\ / \ \text{mol dm}^{-3}\text{s}^{-1} / mol dm−3s−1 |
|---|---|---|---|
| 1 | 2.50 × 10-3 | 3.00 × 10-3 | 4.50 × 10-4 |
| 2 | 5.00 × 10-3 | 3.00 × 10-3 | 1.80 × 10-3 |
| 3 | 1.00 × 10-2 | 9.00 × 10-3 | 2.16 × 10-2 |
Determine the orders with respect to NO\text{NO}NO and H2\text{H}_2H2, the rate equation, and the value of the rate constant, kkk, including its units. Explain your reasoning.
Nitrogen dioxide and ozone react according to the following equation:
2NO2(g)+O3(g)→N2O5(g)+O2(g) 2\text{NO}_2(\text{g}) + \text{O}_3(\text{g}) \rightarrow \text{N}_2\text{O}_5(\text{g}) + \text{O}_2(\text{g}) 2NO2(g)+O3(g)→N2O5(g)+O2(g)The rate equation for this reaction is:
rate=k[NO2(g)][O3(g)] \text{rate} = k[\text{NO}_2(\text{g})][\text{O}_3(\text{g})] rate=k[NO2(g)][O3(g)]Suggest a possible two-step mechanism for this reaction, where the first step is much slower than the second step.
