Use the following chemical equation to solve these problems:
SO3(g) + H2O(g) --> H2SO4(g)
1) You have 11.4 mol SO3 and an excess of water. How much sulfuric acid does the reaction theoretically produce? If the reaction only produces 9.5 mol SO3, what is the percent yield?
Step 1: What is the molar ratio of SO3 to H2SO4?
Step 2: Multiply 11.4 mol SO3 times this molar ratio, making sure to place SO3 on the bottom so that it cancels out and your answer is in moles of H2SO4.
Step 3: Put a box around your answer using the correct number of significant figures and the correct units.
Step 4: Divide your actual yield (given in the problem) by your theoretical yield and multiply by 100% to find the percent yield.
2) Given 5 mol SO3 and 4.5 mol H2O, which is the limiting reactant? What is the theoretical yield of this reaction given these starting amounts? If the reaction only produces 3 mol H2SO4, what is the percent yield?
Step 1: Find the molar ratio of H2O to SO3. Multiply the number of moles of SO3 by this molar ratio, making sure that the SO3 appears in the denominator of the ratio.
Step 2: Compare your answer to the number of moles of water. If your answer is greater than the number of moles of water, then water is the limiting reactant. If it is less than the number of moles of water, then SO3 is the limiting reactant. Repeat this calculation starting with H2O. Indicate the limiting reactant and put a box around it.
Step 3: Use the number of moles of the limiting reactant to find the number of moles the reaction will theoretically produce by multiplying that number by the molar ratio between that reactant and the product. Remember where the reactant belongs in the expression of that ratio? Put a box around this theoretical yield being sure to express it using the correct number of s.f. and the correct units.
Step 4: Divide the actual yield given in the problem by the theoretical yield you just found and multiply by 100%. Put a box around this, the final answer required by the problem.
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More Complex: Using Grams
Use the following chemical equation to solve this problem:
4NH3(g) + 5O2(g) --> 4NO(g) + 6H2O(g)
3) Starting with 67.42 g of ammonia gas (NH3) and an excess of oxygen, how much NO is produced theoretically? How much water? If an experiment shows that 71.33 g of water are produced under these conditions, what is the percent yield?
Step 1: Convert grams of ammonia to moles of ammonia since stoichiometry only works if all quantities are expressed as moles. The molar mass of NH3 is 17.03 g/mol).
Step 2: Find the molar ratio of NH3 to NO. Multiply the number of moles of NH3 by this ratio. I won’t say this time which one belongs in the denominator. Put a box around this theoretical yield being sure to express it using the correct number of s.f. and the correct units.
Step 3: Repeat step two substituting water for NO. Box in your answer.
Step 4: Find the number of moles of water represented by 71.33 g. Divide this number by the theoretical yield you found in the previous step and multiply by 100%. What should you do with the answer? Box it in.
More Complex: Using Density
Use the following chemical equation to solve thise problem:
2Al(s) + 3CuCl2(aq) --> 2AlCl3(aq) + 3Cu(s)
4) You are given 5.793 mL of solid aluminum and 15 g CuCl2. Which is the limiting reagent? What is the theoretical yield of copper? What is the volume of copper? Aluminum D = 2.702 g/mL. Copper D = 8.92 g/mL.
Step 1: Calculate grams of aluminum: Multiply the volume by the density to find the mass. Can you show, using the conversion method we have used this whole semester, why this is the correct procedure?
Step 2: Find the number of moles of copper chloride by finding its molar mass and then multiplying the mass given in the problem by the proper conversion factor.
Step 3: Find the molar ratio of Al to CuCl2 given in the equation. Multiply the number of moles of Al by the ratio, written with the proper reactant in the denominator. Can you see which is the limiting reactant? Box in your answer.
Step 4: Use the limiting reactant to calculate the amount of copper metal that the reaction would theoretically produce. Convert this answer to grams and then to mL.
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More Complex: Using Gases
Use the following chemical equation to solve these problems:
3H2(g) + N2(g) --> 2NH3(g)
5) How many liters of ammonia gas will be produced when 2.00 L of hydrogen react with 4.00 L of nitrogen? If 1.00 L of ammonia are actually collected, what is the percent yield?
Step 1: It is necessary to explain here that stoichiometry can be used for chemical equations involving gases without converting to moles. Avogadro’s Law states that equal amounts (numbers of moles) of any gas occupy equal volumes at the same temperature and pressure. So step 1 is just to find the proper molar ratio for determining the limiting reactant.
Step 2: Multiply the number of liters of nitrogen by the ratio found in step 1. How many liters of hydrogen would be needed to completely react with the nitrogen? Indicate which reactant is the limiting reactant and box in your answer.
Step 3: Use the number of liters of the limiting reactant to find out how many liters of ammonia will be produced. Box in your answer.
Step 4: Find the percent yield using the same procedure used in previous problems. It’s OK to calculate percent yield using liters!
6) A reaction of 15.78 g N2 with excess H2 will produce how many liters of ammonia? The reaction will take place at standard temperature and pressure (0°C and 1 atm).
Step 1: First, find out how many moles of N2 are being reacted.
Step 2: Next, find out how many moles of ammonia would result.
Step 3: Now use V = nRT/P to find the volume of ammonia gas. (Hint: use the temperature and pressure given in the problem. R = 0.0821 L·atm/K·mol)
7) What mass of hydrogen would be needed to react with excess nitrogen to produce enough ammonia to fill a 2.35 L container at 25°C to a pressure of 98.56 kPa?
Step 1: First, find out how many moles of NH3 are needed to fulfill the conditions given in the problem. Use the equation n = PV/RT. R = 8.314 L·kPa/K·mol.
Step 2: Use this number of moles to find the number of moles of hydrogen needed using an appropriate molar ratio.
Step 3: Now find the molar mass of hydrogen and use it to convert the number of moles you just found to grams.