Question
With respect to the previous question, one way to reduce the formation of acid rain involves trapping the $\mathrm{SO}_{2}$ by passing smokestack gases through a spray of calcium oxide and $\mathrm{O}_{2} .$ The product of this reaction is calcium sulfate.a. Write a balanced chemical equation describing this reaction.b. How many metric tons of calcium sulfate would be produced from each ton of $\mathrm{SO}_{2}$ that is trapped?
Step 1
The reaction is between sulfur dioxide (SO2), calcium oxide (CaO), and oxygen (O2) to produce calcium sulfate (CaSO4). The balanced chemical equation is: \[CaO + SO_{2} + 1/2O_{2} \rightarrow CaSO_{4}\] Show more…
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One of the results of acid rain is the reaction of sulfuric acid, H2SO4, with calcium carbonate, CaCO3, to produce calcium sulfate, CaSO4, and carbonic acid, H2CO3. The chemical equation for this reaction is: H2SO4 + CaCO3 = CaSO4 + H2CO3 A student demonstrates the reaction by mixing 25.00 g of H2SO4 with 25.00 g of CaCO3. a. How many moles of each reactant does the student mix? b. Identify the limiting reactant. Explain how you know this. c. Calculate the theoretical mass of the excess reactant. d. Soon after mixing the reactants, the student determines that 33.17 g of CaSO4 was produced by the reaction. Is this a reasonable result? Show calculations to defend your answer.
Reducing $\mathrm{SO}_{2}$ Emissions One way in which $\mathrm{SO}_{2}$ is removed from the "stack" gases of coal-burning power plants is by spraying the gases with fine particles of solid calcium oxide suspended in $\mathrm{O}_{2}$ gas. The product of the reaction of $\mathrm{SO}_{2}$ $\mathrm{CaO},$ and $\mathrm{O}_{2}$ is calcium sulfate. a. Write a balanced chemical equation for this reaction. b. How many metric tons of calcium sulfate would be produced from each ton of $\mathrm{SO}_{2}$ that is trapped?
One method of reducing acid rain is "scrubbing" the combustion products before they are emitted from power plant smoke stacks. The process involves addition of an aqueous suspension of lime $(\mathrm{CaO})$ to the combustion chamber and stack, where the lime reacts with $\mathrm{SO}_{2}$ to give calcium sulfite $\left(\mathrm{CaSO}_{3}\right): \mathrm{CaO}(a q)+\mathrm{SO}_{2}(g) \longrightarrow \mathrm{CaSO}_{3}(a q)$ (a) How much lime (in grams) is needed to remove $1 \mathrm{~mol}$ of $\mathrm{SO}_{2} ?$ (b) How much lime (in kilograms) is needed to remove $1 \mathrm{~kg}$ of $\mathrm{SO}_{2} ?$
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