Table of contents

1. Intro to General Chemistry3h 48m
2. Atoms & Elements4h 16m
3. Chemical Reactions4h 10m
4. BONUS: Lab Techniques and Procedures1h 38m
5. BONUS: Mathematical Operations and Functions48m
6. Chemical Quantities & Aqueous Reactions3h 53m
7. Gases3h 49m
8. Thermochemistry2h 37m
9. Quantum Mechanics2h 58m
10. Periodic Properties of the Elements3h 9m
11. Bonding & Molecular Structure3h 29m
12. Molecular Shapes & Valence Bond Theory1h 57m
13. Liquids, Solids & Intermolecular Forces2h 23m
14. Solutions3h 1m
15. Chemical Kinetics2h 53m
16. Chemical Equilibrium2h 29m
17. Acid and Base Equilibrium5h 1m
18. Aqueous Equilibrium4h 47m
19. Chemical Thermodynamics1h 50m
20. Electrochemistry2h 42m
21. Nuclear Chemistry2h 36m
22. Organic Chemistry5h 7m
23. Chemistry of the Nonmetals2h 39m
24. Transition Metals and Coordination Compounds3h 16m

3. Chemical Reactions

Empirical Formula

General Chemistry

3. Chemical Reactions

Empirical Formula

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Multiple Choice

A 48.3 g sample of an aluminum iodine compound contains 3.2 g of aluminum. What is the empirical formula for the compound?

AlI3

Al2I3

Al3I2

AlI

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Verified step by step guidance

Step 1: Determine the mass of iodine in the compound by subtracting the mass of aluminum from the total mass of the compound. This is calculated as: 48.3 g (total mass) - 3.2 g (mass of aluminum) = 45.1 g (mass of iodine).

Step 2: Convert the mass of aluminum to moles using its molar mass. The molar mass of aluminum (Al) is approximately 26.98 g/mol. Use the formula: \( \text{moles of Al} = \frac{3.2 \text{ g}}{26.98 \text{ g/mol}} \).

Step 3: Convert the mass of iodine to moles using its molar mass. The molar mass of iodine (I) is approximately 126.90 g/mol. Use the formula: \( \text{moles of I} = \frac{45.1 \text{ g}}{126.90 \text{ g/mol}} \).

Step 4: Determine the simplest whole number ratio of moles of aluminum to moles of iodine. Divide the moles of each element by the smallest number of moles calculated in the previous steps.

Step 5: Use the ratio from Step 4 to write the empirical formula of the compound. The empirical formula is based on the simplest whole number ratio of atoms in the compound.