Table of contents

1. Intro to General Chemistry3h 48m
2. Atoms & Elements4h 16m
3. Chemical Reactions4h 10m
4. BONUS: Lab Techniques and Procedures1h 36m
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 compound contains 15.94% boron, with the remainder being fluorine. What is the empirical formula of the compound?

B3F

B2F6

BF3

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Verified step by step guidance

Step 1: Begin by understanding that the empirical formula represents the simplest whole-number ratio of atoms in a compound. We need to determine the ratio of boron to fluorine atoms in the compound.

Step 2: Assume a 100 g sample of the compound for simplicity. This means the sample contains 15.94 g of boron and the remainder, 84.06 g, is fluorine.

Step 3: Convert the mass of each element to moles using their molar masses. The molar mass of boron (B) is approximately 10.81 g/mol, and the molar mass of fluorine (F) is approximately 19.00 g/mol. Use the formula: \( \text{moles} = \frac{\text{mass}}{\text{molar mass}} \).

Step 4: Calculate the moles of boron: \( \frac{15.94 \text{ g}}{10.81 \text{ g/mol}} \) and the moles of fluorine: \( \frac{84.06 \text{ g}}{19.00 \text{ g/mol}} \).

Step 5: Determine the simplest whole-number ratio of moles of boron to moles of fluorine by dividing each by the smallest number of moles calculated. This will give you the empirical formula of the compound.