When 10 grams of reactant A reacts with 20 grams of reactant B, only 15 grams of product C is produced. What is the percentage yield of this reaction?

A. 50%

B. 75%

C. 100%

D. 150%

Which of the following statements best describes a limiting reactant?

A. The reactant that is completely consumed in a reaction.

B. The reactant that is present in excess.

C. The reactant that determines the amount of product formed.

D. Both A and C.

In a balanced chemical equation, the number of atoms of each element on the reactant side is:

A. Greater than on the product side.

B. Lesser than on the product side.

C. Equal to the number on the product side.

D. Not related to the product side.

If the actual yield of a reaction is 80 grams and the theoretical yield is 100 grams, what is the percentage yield?

A. 60%

B. 70%

C. 80%

D. 90%

Which factor does NOT affect the percentage yield in real-world reactions?

A. Purity of reactants.

B. Temperature.

C. Catalysts.

D. Color of the reactants.

a) Define the concept of a limiting reactant and explain its significance in predicting product yields. [3]

b) Given the reaction: 2H₂ + O₂ → 2H₂O, if 4 moles of hydrogen gas react with 2 moles of oxygen gas, identify the limiting reactant and the number of moles of water produced. [3]

a) Describe the importance of balancing chemical equations in stoichiometric calculations. [3]

b) For the reaction: N₂ + 3H₂ → 2NH₃, if 28 grams of nitrogen gas reacts with an excess of hydrogen gas, calculate the mass of ammonia produced. (Given: Molar mass of N₂ = 28 g/mol, NH₃ = 17 g/mol) [3]

a) Define the term 'percentage yield' and explain factors that might affect it in real-world reactions. [4]

b) If a reaction has a theoretical yield of 50 grams, but only 40 grams of product is obtained, calculate the percentage yield. [2]

a) Define Avogadro's number and explain its significance in relating the macroscopic world to the atomic/molecular level. [3]

b) How many atoms are present in 2 moles of carbon? [2]

c) If you have 12 grams of carbon, how many moles of carbon atoms do you have? (Given: Molar mass of Carbon = 12 g/mol) [2]

a) Describe the difference between an element, a compound, and a mixture. [3]

b) Given a sample containing iron and sulphur, how would you determine if it's a mixture or a compound? [3]

c) Why is the mole concept crucial for stoichiometric calculations? [2]

In a reaction where 5 moles of A react with 3 moles of B to produce C, if only 2 moles of B are available, which is the limiting reactant?

A. A

B. B

C. C

D. Both A and B

The balanced equation for the combustion of methane is CH₄ + 2O₂ → CO₂ + 2H₂O. If 3 moles of O₂ are available for the reaction, how many moles of CO₂ will be produced?

A. 1.5 moles

B. 2 moles

C. 3 moles

D. 6 moles

Which of the following is NOT a reason for a reaction not achieving 100% yield?

A. Side reactions.

B. Loss of product during purification.

C. Complete conversion of reactants to products.

D. Reactants not being pure.

In the reaction 2A + 3B → C, if 10 moles of A and 15 moles of B are reacted, which reactant will be left in excess?

A. A

B. B

C. Both A and B

D. Neither A nor B

A reaction has a theoretical yield of 50 grams, but in a laboratory setting only produces 40 grams. What is the percentage yield?

A. 60%

B. 70%

C. 80%

D. 90%

a) Explain the concept of excess reactants in a chemical reaction. [3]

b) For the reaction: 2H₂ + O₂ → 2H₂O, if 5 moles of hydrogen gas and 3 moles of oxygen gas are mixed, which is the excess reactant? [2]

c) How many moles of the excess reactant will remain unreacted? [2]

a) What factors can affect the percentage yield in real-world reactions? [3]

b) If a reaction has a theoretical yield of 60 grams, but only 45 grams of product is obtained, calculate the percentage yield. [2]

c) Why might the actual yield be less than the theoretical yield in many chemical reactions? [3]

a) Define Avogadro's number and explain its significance in the context of the mole concept. [3]

b) Calculate the number of molecules in 0.5 moles of water. [2]

c) If 9 grams of water is decomposed into its constituent elements, how many moles of hydrogen gas would be produced? (Given: Molar mass of Water = 18 g/mol) [3]

d) Why is it essential to balance chemical equations before performing stoichiometric calculations? [2]

a) Describe the difference between limiting and excess reactants in a chemical reaction. [3]

b) For the reaction: N₂ + 3H₂ → 2NH₃, if 2 moles of nitrogen gas and 5 moles of hydrogen gas are mixed, which is the limiting reactant? [3]

c) How many moles of ammonia will be produced from the given reactants? [2]

d) Explain why, in industrial processes, it might be beneficial to use an excess of one reactant. [2]

a) Define the term "percentage yield" in the context of chemical reactions. [2]

b) If a reaction has a theoretical yield of 80 grams, but only 60 grams of product is obtained, calculate the percentage yield. [2]

c) List two factors that can lead to a reduced percentage yield in a chemical reaction. [2]

d) Why is achieving a high percentage yield important in industrial chemical processes? [3]