Overview
This lecture explains the law of multiple proportions, demonstrates how to apply it to compound data, and works through several examples with calculations.
The Law of Multiple Proportions
- The law states: when two elements form multiple compounds, the ratio of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.
- This law helps differentiate compounds made from the same elements by comparing mass ratios.
Example: Carbon Monoxide vs. Carbon Dioxide
- Carbon monoxide (CO): 12 g carbon + 16 g oxygen; dividing by 12 gives 1 g carbon + 1.33 g oxygen.
- Carbon dioxide (CO₂): 12 g carbon + 32 g oxygen; dividing by 12 gives 1 g carbon + 2.67 g oxygen.
- The ratio of 2.67 to 1.33 is approximately 2:1, matching the whole number ratio in the law.
Example: Nitrogen and Oxygen Compounds
- Three compounds: oxygen masses per 1 g nitrogen are 1.142 g, 2.284 g, and 2.855 g.
- 2.284 / 1.142 = 2; 2.855 / 1.142 = 2.5 (or 5:2 after multiplying by 2); 2.855 / 2.284 = 1.25 (or 5:4 after multiplying by 4).
- All ratios reduce to small whole numbers, demonstrating the law.
Example: Sulfur and Oxygen Compounds
- Compound A: 50 g sulfur, 50 g oxygen (100 g total); Compound B: 40 g sulfur, 60 g oxygen (100 g total).
- Per 1 g sulfur: Compound A has 1 g oxygen, Compound B has 1.5 g oxygen.
- 1.5 / 1 = 1.5; multiply by 2 to get a 3:2 whole number ratio, proving the law.
Applying the Law
- Always reduce to a comparison per 1 g of the first element.
- Find ratios of the second element between compounds and simplify to small whole numbers.
Key Terms & Definitions
- Law of Multiple Proportions — The ratio of masses of one element that combine with a fixed mass of another can be reduced to small whole numbers.
- Compound — A substance formed from two or more elements chemically bonded.
Action Items / Next Steps
- Practice applying the law by converting mass data to ratios and reducing them to whole numbers.
- Review similar example problems for practice.