Use the table data to work out the rate equation, including the order for each reactant. Show how you reach your answer.
Calculate the value of the rate constant, $k$, from the data in experiment 2. Give its units.
Explain the difference between heterogeneous and homogeneous catalysts.
Complete the table using ticks ($\checkmark$) to show whether the catalyst used in each reaction is heterogeneous or homogeneous: manufacture of ammonia in the Haber process; removal of nitrogen oxides from car exhausts; oxidation of sulfur dioxide in the atmosphere.
Some reactions are catalysed by one of the products formed. This is known as autocatalysis. One example is the reaction between acidified manganate(VII) ions, $\text{MnO}_4^-$, and ethanedioic acid, $(\text{CO}_2\text{H})_2$. $\text{Mn}^{2+}$ ions act as the catalyst for this reaction. Without a catalyst, the reaction is slow. Balance the equation for this reaction: $\ldots\text{MnO}_4^- + \ldots\text{H}^+ + \ldots(\text{CO}_2\text{H})_2 \rightarrow \ldots\text{Mn}^{2+} + \ldots\text{CO}_2 + \ldots\text{H}_2\text{O}$.
The graph shown is a concentration-time graph for a typical reaction. On the axes below, sketch the curve you would expect for the autocatalysed reaction in (i).
Describe, with the aid of a reaction pathway diagram, the effect of a catalyst on a reversible reaction. Suggest why catalysts are used in industrial processes.
The reaction used in the Haber process to make ammonia is shown here: $\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$, $\Delta H^{\circ} = -92\,\text{kJ mol}^{-1}$. At $500\,^{\circ}\text{C}$, when pressure is expressed in atmospheres, the numerical value of $K_p$ for this equilibrium is $1.45 \times 10^{-5}$. Write the expression for $K_p$ for this equilibrium. Calculate the partial pressure of $\text{NH}_3$ at equilibrium at $500\,^{\circ}\text{C}$, when the partial pressure of $\text{N}_2$ is $20\,\text{atm}$ and that of $\text{H}_2$ is $60\,\text{atm}$.