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Use the following chemical equation to solve these problems:

SO_{3}(g) + H_{2}O(g) --> H_{2}SO_{4}(g)

SO

1) You have 11.4 mol SO_{3} and an excess of water. How much sulfuric acid does the reaction theoretically produce? If the reaction only produces 9.5 mol SO_{3}, what is the percent yield?

Step 2: Multiply 11.4 mol SO

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 SO_{3} and 4.5 mol H_{2}O, which is the limiting reactant? What is the theoretical yield of this reaction given these starting amounts? If the reaction only produces 3 mol H_{2}SO_{4}, what is the percent yield?

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 SO

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.

Use the following chemical equation to solve this problem:

4NH_{3}(g) + 5O_{2}(g) --> 4NO(g) + 6H_{2}O(g)

4NH

3) Starting with 67.42 g of ammonia gas (NH_{3}) 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 2: Find the molar ratio of NH

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.

Use the following chemical equation to solve thise problem:

2Al(s) + 3CuCl_{2}(aq) --> 2AlCl_{3}(aq) + 3Cu(s)

2Al(s) + 3CuCl

4) You are given 5.793 mL of solid aluminum and 15 g CuCl_{2}. 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 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 CuCl

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.

Use the following chemical equation to solve these problems:

3H_{2}(g) + N_{2}(g) --> 2NH_{3}(g)

3H

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 gasesStep 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 N_{2} with excess H_{2} will produce how many liters of ammonia? The reaction will take place at standard temperature and pressure (0°C and 1 atm).

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 NHStep 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.