Percentage yield

• What does percentage yield indicate in an experiment?


  It indicates the overall success of an experiment.
• How is percentage yield calculated?


  Percentage yield is calculated by comparing actual yield to theoretical yield.
• What is the actual yield in a chemical reaction?


  The actual yield is the amount of product you get from the experiment.
• Does the percentage yield depend on the amount of reactants used?


  No, percentage yield is a percentage and does not depend on the amount of reactants.
• What is the formula for calculating percentage yield?


  Percentage yield = (mass of product actually made / maximum theoretical mass of product) × 100
• What does a 100% yield indicate?


  A 100% yield means that all the expected product was obtained.
• What does a 0% yield indicate?


  A 0% yield means that no product was formed from the reactants.
• Why is it important for industrial processes to have a high percentage yield?


  To reduce waste and lower costs.
• Why is it impossible to achieve a 100% yield in real-life reactions?


  Some product or reactant is always lost during the process.
• What are three common reasons for not achieving 100% yield in reactions?

  1. Not all reactants react to make a product (reversible reactions).
  2. Side reactions occur, forming unwanted products.
  3. Product loss during separation from the reaction mixture.
• What happens in reversible reactions that affects yield?


  Products can turn back into reactants, preventing 100% yield.
• In the Haber process, what reactions occur simultaneously?


  N₂ + 3H₂ → 2NH₃ and 2NH₃ → N₂ + 3H₂.
• What is the maximum theoretical yield of zinc chloride in the given example?


  1. 72 g.
• What is the actual yield of zinc chloride after purification in the example?


  1. 31 g.
• How do you calculate the percentage yield of zinc chloride in the example?


  Percentage yield = (2.31 g / 2.72 g) × 100.
• What is the percentage yield of zinc chloride calculated from the example?


  Percentage yield = $\frac{2.31}{2.72} \times 100 \approx 85.04\%$
• What is the topic of the study material?

  Topic 3 — Quantitative Chemistry
• What is the formula for percentage yield?


  Percentage yield = (actual yield / theoretical yield) x 100
• What is actual yield in a chemical reaction?


  Actual yield is the amount we actually get when we carry out the reaction.
• What is theoretical yield in a chemical reaction?

  Theoretical yield is the yield we expect to get based on our calculations.
• Why might we not achieve the theoretical yield in a chemical reaction?


  We might not achieve the theoretical yield because reactants might not all react, side reactions may occur, or some product may be lost during the process.
• What are the reasons for not achieving the expected amount of product in a chemical reaction?

  1. Reactants might not all react:

  • Slow reactions or reversible reactions (e.g., N₂ + 3H₂ ↔ 2NH₃)

  1. Side reactions:

  • Reactants produce different products (e.g., N₂ + 3H₂ → NH₃ and N₂ + 2O₂ → 2NO₂)

  1. Loss of product during the process:

  • Gaseous products can escape or solid may remain on equipment.
• Given the reaction 2H₂ + O₂ → 2H₂O, what is the theoretical yield if 2g of H₂ and 16g of O₂ are used?


  Theoretical yield is 18g.
• If the actual yield from the reaction 2H₂ + O₂ → 2H₂O is 15g, what is the percentage yield?


  Percentage yield is 83.3%.
• What does a percentage yield of 83.3% indicate about a chemical reaction?


  It indicates that we obtained 83.3% of the expected theoretical yield.
• In the example given, what is the actual yield of water produced?


  The actual yield is 15g.
• What is the theoretical yield of water produced from 2g of H₂ and 16g of O₂?


  The theoretical yield is 18g.
• How does a reversible reaction affect the actual yield of a product?


  A reversible reaction may not allow all reactants to convert into products, leading to a lower actual yield.
• What impact do side reactions have on the yield of a desired product?


  Side reactions can divert reactants to produce unwanted products, reducing the yield of the desired product.