Learning Module 32: Compositional Stoichiometry

Definition of Stoichiometry

• Stoichiometry pertains to the relationship between quantities of substances involved in chemical reactions or in forming compounds.

• Due to the Law of Simple Proportions, these ratios can be expressed as whole numbers (integers).

• Compositional Stoichiometry Problems: Focus on stoichiometric ratios confined to a single compound.

• Reaction Stoichiometry Calculations: Involve stoichiometric ratios associated with balanced chemical reactions.

Dalton's Atomic Theory

• The foundation of stoichiometry statements is based on principles laid out by Dalton:

  1. Elements consist of small, indivisible particles called atoms.

  2. All atoms of a given element have identical properties, which differ from those of other elements.

  3. Atoms cannot be created, destroyed, or transformed into different elements.

  4. Molecules form when atoms of different elements combine in simple, whole number ratios.

  5. The relative number and type of atoms in a molecule remain constant.

• While Dalton's laws have limitations, the fundamental principles are applicable in the context of the module.

Importance of Ratios in Molecules

• Variations in the ratios of atoms within a molecule can significantly affect its properties.

• Stoichiometry is used to calculate the ratios of elements in a compound, aiding in determining the molecular formula.

Key Examples in Compositional Stoichiometry

Example 1: Mass of Magnesium

• Task: Calculate the mass of a magnesium (Mg) atom in grams.

• Expression: (\text{mass of Mg atom} = (\text{valence factor} )(\text{conversion factor} ))

Example 2: Atoms in a Small Mass

• Task: Calculate the number of atoms in one-millionth of a gram (10^{-6} \text{ g}) of magnesium (Mg).

• Methods involve using conversion calculations linked to Avogadro's number.

Example 3: Atoms in Moles

• Task: Determine how many atoms are in 1.67 \text{ moles} of magnesium (Mg).

• Formula: (\text{number of atoms} = (1.67 \text{ moles}) (6.022 \times 10^{23} \text{ atoms/mole}))

Example 4: Moles from Mass

• Task: Find out how many moles of magnesium are in 73.4 \text{ g} of magnesium.

• Formula: (\text{moles of Mg} = \frac{73.4 \text{ g}}{24.3 \text{ g/mole}}) \rightarrow \text{resulting in } 3.0 \text{ moles}

Key Concepts in Molecular Stoichiometry

• Understanding molar masses is integral to calculating quantities in stoichiometry.

• Example 6: Calculate the number of propane molecules in 74.6 \text{ g} of propane (C₃H₈).

• Formula: (\text{propane molecules} = \frac{74.6 \text{ g}}{\text{formula weight of } C{3}H{8}})

Mass Contributions

• Hydrofluoric Acid (HF): Calculation of mass percentage contributions from each constituent, especially from Hydrogen (H).

  • Molar mass of HF is approximately 20 \text{ g/mol}; contributions are calculated from individual atom weights.

Understanding Diatomic Elements

• Diatomic elements are essential in stoichiometry, represented by molecules containing two atoms (e.g., O₂, N₂)

• Students should recognize the distinction between moles of atoms vs moles of molecules, leading to potential confusion in calculations.

• Example of Ozone (O_{3}): How many moles, molecules, and atoms are contained in a 60\text{-gram} sample?

  • Moles = (\frac{60 \text{ g}}{48 \text{ g/mole}}) = 1.25 \text{ moles}

  • Molecules = (1.25 \text{ moles}) (6.022 \times 10^{23} \text{ molecules/mole})

  • Atoms = \text{molecules} (3 \text{ atoms/molecule})

PROBLEM SET 1

1. Using the concept of valence and conversion factors, calculate the mass of a single Magnesium (Mg) atom in grams.

2. Determine the number of atoms present in a sample of Magnesium weighing 5.0×10^-6g

3. If a chemist has 2.45 moles of Magnesium, how many individual atoms do they have? (Assume Avogadro's number is 6.022×1023 atoms/mole).

4. How many moles of Magnesium are found in a 121.5 g sample? (Note: Mg atomic mass is approximately 24.3 g/mol).

5. Calculate the formula weight of propane (C₃H₈) given atomic weights for Carbon (12.011) and Hydrogen (1.01).

6. Determine the total number of propane molecules in a 200g sample.

7. In Hydrofluoric Acid (HF), calculate the mass percentage contribution of Hydrogen (H) and Fluorine (F).

8. For an 80-gram sample of Ozone (O₃), how many moles of ozone molecules are present?

9. For an 80-gram sample of Ozone (O₃), how many individual ozone molecules does this represent?

10. For an 80-gram sample of Ozone (O₃), what is the total count of oxygen atoms in this sample?

PROBLEM SET 2

11. Using Avogadro's number, calculate the mass of a single atom of Gold (Au) if its atomic mass is approximately 197 \text{ g/mol}.

12. How many atoms are contained in a 15\text{-mg} (0.015 \text{ g}) sample of Carbon (C)?

13. A flask contains 4.2 \text{ moles} of Helium (He). Calculate the total number of Helium atoms.

14. Calculate the number of moles in a 500\text{-gram} block of Iron (Fe). (Atomic mass: 55.85 \text{ g/mol}).

15. Determine the formula weight of Glucose

16. Calculate the total number of Ethanol molecules (C₂H₅OH) in a 46 gram sample.

17. Find the mass percentage of Oxygen (O) in Water (H_{2}O).

18. For a 140\text{-gram} sample of Nitrogen gas (N_{2}), calculate the number of moles of molecules present.

19. For a 160\text{-gram} sample of Oxygen gas (O_{2}), calculate the total number of individual oxygen atoms.

PROBLEM SET 3

20. Using Avogadro’s number, calculate the mass of a single atom of silver (Ag).
    (Atomic mass = 107.87 g/mol)

    1. 1.79×10^-22 g of Ag

21. How many atoms of phosphorus (P) are contained in 0.032 g of phosphorus?
    (Atomic mass = 30.97 g/mol)

    1. 6.2×10²⁰ atoms of P

22. A container holds 5.75 moles of neon (Ne).
    Calculate the total number of neon atoms present.

    1. 3.45×10²⁴ atoms of Ne

23. Determine the number of moles of copper (Cu) in a 127 g copper sample.
    (Atomic mass = 63.55 g/mol)

    1. 2 moles of Cu

24. Calculate the molar mass of calcium nitrate, Ca(NO₃)₂.

    1. 164g/mol

25. How many molecules of carbon dioxide (CO₂) are present in 88.0 g of CO₂?

    1. 12×10²³ CO₂ molecules

26. A sample contains 3.0 × 10²³ molecules of ammonia (NH₃).
    What is the mass of this sample?

    1. 10g of NH₃

27. Find the mass percent of sulfur (S) in sulfuric acid (H₂SO₄).

    1. 33% of S

28. How many moles of nitrogen gas (N₂) are contained in 84.0 g of N₂?

    1. 3 mol of N₂

29. A 96.0 g sample of oxygen gas (O₂) is given.
    Calculate the total number of oxygen atoms in the sample.

    1. 3.6×10²⁴ atoms of O

30. A sample contains 2.50 moles of ozone (O₃).
    Determine the total number of oxygen atoms present.

    1. 4.5×10²⁴ O atoms

31. How many grams of hydrogen are present in 54.0 g of water (H₂O)?

    1. 6g of H

32. A carbon monoxide (CO) sample contains 28.0 g of oxygen.
    How many grams of carbon are present?

    1. 21g of C

33. Explain in one sentence why it is important to distinguish between moles of atoms and moles of molecules when solving stoichiometry problems.