Sulfur dioxide gas, SO2\text{SO}_2SO2, reacts with oxygen gas, O2\text{O}_2O2, to produce sulfur trioxide gas, SO3\text{SO}_3SO3. The reaction is reversible. Write the balanced symbol equation for this reaction.
The conditions used to make sulfur trioxide in this industrial process are:
- a pressure of 2 atmospheres
- a temperature of 450 ∘C450\ ^\circ\text{C}450 ∘C.
The forward reaction is exothermic.
A chemical plant operator decreases the temperature used to 350 ∘C350\ ^\circ\text{C}350 ∘C.
What happens to the rate of the reaction when the temperature is decreased?
The plant operator thinks that the decrease in temperature will increase the yield of sulfur trioxide. Is the operator correct? Explain your answer.
The plant wants to reduce the cost of operating the process and decides to decrease the pressure used to 0.5 atmospheres. Write about two disadvantages of using a lower pressure to make sulfur trioxide.
Potassium sulfate can be prepared in the laboratory by neutralizing potassium hydroxide solution with dilute sulfuric acid using a titration method, as shown in the diagram.
The student adds an indicator to the potassium hydroxide solution in a conical flask, then adds dilute sulfuric acid from a burette until the indicator changes color. The student then crystallises the solution, but is left with impure potassium sulfate crystals.
What should the student have done to obtain pure potassium sulfate crystals?
Give one reason why this laboratory method to make potassium sulfate is not used in industry.