Q6. For vanadinite, Pb5(VO4)3Cl, determine:

• Number of atoms in one formula unit

• Number ratio of Pb to V atoms

• The ratio, by mass, of Pb to O

• The mass of compound that contains 1.00 g Cl

Background

Topic: Chemical formulas, stoichiometry, and mass ratios

This question tests your understanding of interpreting chemical formulas, calculating atom ratios, and using molar masses to determine mass relationships in compounds.

Key Terms and Formulas

• Formula unit: The simplest ratio of ions represented in an ionic compound.

• Molar mass: The mass of one mole of a substance (g/mol).

• Mass ratio:

• Number ratio: The ratio of the number of atoms of each element in the formula unit.

Step-by-Step Guidance

1. Write out the chemical formula for vanadinite: .

2. Count the number of each type of atom in one formula unit. For example, how many Pb, V, O, and Cl atoms are present?

3. Determine the number ratio of Pb to V atoms by comparing their counts from the formula.

4. Calculate the mass of Pb and O in one formula unit using their atomic masses (, ). Set up the mass ratio of Pb to O.

5. To find the mass of compound containing 1.00 g Cl, first determine the mass percent of Cl in the compound using the molar mass of vanadinite. Then, set up a proportion to find the total mass of compound that contains 1.00 g Cl.

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Q7. Balance the following chemical equations using the smallest whole number coefficients.

Background

Topic: Balancing chemical equations

This question tests your ability to apply the law of conservation of mass by ensuring the same number of each type of atom appears on both sides of the equation.

Key Terms and Formulas

• Coefficient: The number placed in front of a chemical formula to indicate how many units are involved.

• Balancing: Adjusting coefficients so that the number of atoms of each element is equal on both sides.

Step-by-Step Guidance

1. For each equation, list the number of atoms of each element on both sides.

2. Start by balancing the element that appears in the fewest compounds first (often metals or complex ions).

3. Balance polyatomic ions as a unit if they appear unchanged on both sides.

4. Adjust coefficients as needed, checking after each change to ensure all elements remain balanced.

5. Repeat the process for each equation, but stop before writing the final set of coefficients.

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Q8a. Hydnocarpic acid contains 76.14% carbon, 11.18% hydrogen, and 12.68% oxygen by mass. Determine the empirical formula.

Background

Topic: Empirical formula determination from percent composition

This question tests your ability to convert mass percentages to a simplest whole-number ratio of atoms (empirical formula).

Key Terms and Formulas

• Empirical formula: The simplest whole-number ratio of atoms in a compound.

• Percent composition: The percent by mass of each element in a compound.

• Steps:

  1. Assume 100 g of compound (so percentages become grams).

  2. Convert grams to moles for each element using atomic masses.

  3. Divide each by the smallest number of moles to get the simplest ratio.

Step-by-Step Guidance

1. Assume you have 100 g of hydnocarpic acid: 76.14 g C, 11.18 g H, 12.68 g O.

2. Convert each mass to moles using atomic masses: , , .

3. Calculate the number of moles for each element.

4. Divide each mole value by the smallest number of moles calculated to get the simplest whole-number ratio.

5. If necessary, multiply all ratios by a small integer to get whole numbers for the empirical formula.

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Q8b. The experimental molar mass of hydnocarpic acid is 252.4 g/mol. Determine the molecular formula.

Background

Topic: Determining molecular formula from empirical formula and molar mass

This question tests your ability to relate the empirical formula mass to the actual molar mass to find the molecular formula.

Key Terms and Formulas

• Molecular formula: The actual number of atoms of each element in a molecule.

• Empirical formula mass: The molar mass of the empirical formula.

• Relationship: , where

Step-by-Step Guidance

1. Calculate the molar mass of the empirical formula () using atomic masses.

2. Divide the experimental molar mass (252.4 g/mol) by the empirical formula mass to find .

3. Multiply the subscripts in the empirical formula by to get the molecular formula.

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