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

Previous problemNext problem

Struggling with General Chemistry?

Join thousands of students who trust us to help them ace their exams!Watch the first video

Multiple Choice

A 4.386 g compound of hydrogen, carbon, and oxygen was burned in 9.389 g of oxygen, producing 6.184 g of carbon dioxide and 3.789 g of water. What is the empirical formula of the original sample?

C2H6O

CH2O

C3H6O2

C2H4O

Previous problemNext problem

Verified step by step guidance

Determine the mass of carbon in the carbon dioxide produced. Use the molar mass of CO2 (44.01 g/mol) and the fact that each mole of CO2 contains one mole of carbon (12.01 g/mol). Calculate the moles of CO2 and then the mass of carbon: \( \text{mass of carbon} = \frac{6.184 \text{ g CO}_2}{44.01 \text{ g/mol}} \times 12.01 \text{ g/mol} \).

Determine the mass of hydrogen in the water produced. Use the molar mass of H2O (18.02 g/mol) and the fact that each mole of H2O contains two moles of hydrogen (2.02 g/mol). Calculate the moles of H2O and then the mass of hydrogen: \( \text{mass of hydrogen} = \frac{3.789 \text{ g H}_2\text{O}}{18.02 \text{ g/mol}} \times 2.02 \text{ g/mol} \).

Calculate the mass of oxygen in the original compound. Subtract the mass of carbon and hydrogen from the total mass of the compound: \( \text{mass of oxygen} = 4.386 \text{ g} - (\text{mass of carbon} + \text{mass of hydrogen}) \).

Convert the masses of carbon, hydrogen, and oxygen to moles by dividing each by their respective atomic masses: \( \text{moles of C} = \frac{\text{mass of carbon}}{12.01 \text{ g/mol}} \), \( \text{moles of H} = \frac{\text{mass of hydrogen}}{1.008 \text{ g/mol}} \), \( \text{moles of O} = \frac{\text{mass of oxygen}}{16.00 \text{ g/mol}} \).

Determine the empirical formula by finding the simplest whole number ratio of moles of carbon, hydrogen, and oxygen. Divide each mole value by the smallest number of moles calculated in the previous step. Adjust the ratios to the nearest whole numbers to find the empirical formula.