Q1. Balance the following chemical equation:

Background

Topic: Chemical Equation Balancing

This question tests your ability to balance chemical equations, which is fundamental to stoichiometry. A balanced equation must have equal numbers of each type of atom on both sides, following the law of conservation of mass.

Key concepts:

• Law of Conservation of Mass: Atoms cannot be created or destroyed

• Coefficients: Numbers placed before formulas to balance atoms

• Subscripts: Numbers in formulas that cannot be changed

Step-by-Step Guidance

1. List all atoms present on each side:

   Start by identifying all the different types of atoms in the equation. Count how many of each atom appear on the reactant side (left) and the product side (right). This inventory helps you see what needs to be balanced.

2. Count atoms on each side:

   Count the total number of each type of atom on both sides. Remember that subscripts indicate how many atoms of that element are in the compound, and coefficients (which you'll add) multiply everything in the formula. Start with the current unbalanced equation.

   Count Fe and O atoms on both sides.

3. Start with the most complex compound:

   Begin balancing with the compound that has the most atoms or the most different elements. In this case, is more complex. Place a coefficient in front of it if needed, but often you'll start by balancing other elements first.

4. Balance one element at a time:

   Choose one element to balance first. It's often easiest to start with elements that appear in only one compound on each side, or to save oxygen/hydrogen for last if they appear in multiple compounds. Add coefficients (whole numbers) in front of formulas to balance atoms, but never change subscripts within formulas.

   Add coefficients to balance the Fe atoms first, then work on O atoms.

5. Check and adjust:

   After adding coefficients, recount all atoms on both sides. If they don't match, adjust your coefficients. You may need to use the least common multiple to find appropriate coefficients. Continue adjusting until all atoms are balanced.

   Verify that Fe and O atoms are equal on both sides.

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Q2. How many moles of are produced when 2.0 moles of completely combust?

Background

Topic: Stoichiometry and Chemical Reactions

This question tests your ability to use balanced chemical equations to determine the quantitative relationships between reactants and products. You'll need to use mole ratios from the balanced equation to convert between moles of different substances.

Key formula:

Where:

• Mole ratio = ratio of coefficients from balanced equation

• Combustion reaction:

Step-by-Step Guidance

1. Write and verify the balanced equation:

   First, ensure you have the correct balanced equation for the combustion of ethane (). Combustion reactions involve a hydrocarbon reacting with oxygen to produce carbon dioxide and water. Make sure the equation is balanced before proceeding with stoichiometric calculations.

2. Identify the mole ratio:

   From the balanced equation, identify the ratio between and . The coefficients tell you that 2 moles of produce 4 moles of . This gives you a conversion factor to use in your calculation.

3. Set up the conversion:

   Use the mole ratio as a conversion factor. Multiply the given moles of reactant by the ratio of product to reactant coefficients. This dimensional analysis approach ensures your units cancel correctly.

4. Substitute the values:

   Plug in the given number of moles of and the coefficients from the balanced equation. Make sure to use the correct ratio with the product coefficient in the numerator and reactant coefficient in the denominator.

   Substitute:

5. Calculate the result:

   Perform the multiplication to find the number of moles of produced. Simplify the ratio first if possible, then multiply by the given moles of reactant.

   Complete the calculation to find the final answer.

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Q3. What is the percent yield if 45.0 g of product is obtained from a reaction with a theoretical yield of 60.0 g?

Background

Topic: Percent Yield

This question tests your understanding of percent yield, which compares the actual amount of product obtained in an experiment to the theoretical maximum amount that could be produced. Percent yield is important in assessing the efficiency of chemical reactions and identifying sources of product loss.

Key formula:

Where:

• Actual Yield = amount of product actually obtained (in grams or moles)

• Theoretical Yield = maximum amount of product possible (calculated from stoichiometry)

• Percent Yield = efficiency of the reaction (expressed as a percentage)

Step-by-Step Guidance

1. Identify the actual and theoretical yields:

   Determine which value is the actual yield (what was actually obtained in the experiment) and which is the theoretical yield (what should have been produced according to calculations). The actual yield is typically less than or equal to the theoretical yield due to experimental losses.

2. Understand the percent yield formula:

   The percent yield formula compares the actual yield to the theoretical yield and expresses this as a percentage. This tells you how efficient your reaction was - a 100% yield means you got exactly what was expected, while lower percentages indicate product loss.

3. Substitute the values into the formula:

   Replace the actual yield and theoretical yield in the formula with the given values. Make sure both values are in the same units (both in grams, or both in moles) so the units cancel properly.

4. Calculate the ratio:

   First, divide the actual yield by the theoretical yield. This gives you a decimal value between 0 and 1 (or possibly 1 if the yield is perfect). This ratio represents what fraction of the theoretical yield was actually obtained.

   Calculate to find the decimal ratio.

5. Convert to percentage:

   Multiply the decimal ratio by 100 to convert it to a percentage. This final step gives you the percent yield, which is typically reported with appropriate significant figures based on your input values.

   Multiply your calculated ratio by 100 to get the final percent yield.

Try solving on your own before revealing the answer!