According to Dalton's work with gases, the total pressure of a gaseous mixture is the sum of the partial pressures exerted by the different types of gases. Each gas would exert a pressure if it were in the container by itself, and it would exert the same pressure if it were in the container along with other gases (assuming no interaction between the particles). Based upon the relationship between P and n in the Ideal Gas Law, the mole ratio in a mixture of gases determines the partial pressure of each gas.
HW: Read "Dalton's Law of Partial Pressures" article; finish Worksheet 1 (Unit 8), journal entry due Tuesday - complete the analysis and conclusion questions from the "Chemistry in a Bag" handout using the data provided.
Unit 8 Worksheet 1 Answers:
(assume all zeros to be significant)
1. 3.9 atm
2. 95.1 mL
3. 14 mL
4. 750 mmHg
5. %nCO2 = 33.3%; %nN2 = 50.0%; %nHe = 16.7%
6. PN2 = 600 mmHg; PO2 = 150 mmHg; PAr = 8 mmHg
7. 20%
8. 710 mmHg; 97%
9. 713 mmHg; 23 mL
Unit 5 – Determining an Empirical Formula (Journal Entry)
Introduction
In 1778, Lavoisier concluded that combustion was a reaction of oxygen in the air with a sample of matter. He realized that as the substance burned gained mass, the same mass was lost from the surrounding air. A great deal of chemical knowledge has been amassed by using simple combustion experiments conducted with crucibles, burners, and balances.
An empirical formula gives the simplest whole number ratio of the different atoms in a compound. The empirical formula does not necessarily indicate the exact number of atoms in a single molecule. This information is given by the molecular formula, which is always a simple multiple of the empirical formula.
In this experiment, you will determine the empirical formula of a magnesium-oxygen product, a compound that is formed when magnesium metal reacts with oxygen gas. According to the law of conservation of mass, the total mass of the products must equal the total mass of the reactants in a chemical reaction. Therefore,
mass Mg + mass O2 = mass MgxOy
Since you will measure the mass of magnesium and the magnesium-oxygen product, you will be able to calculate the mass of oxygen consumed during the reaction.
Data Table
Items Mass (g)
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Magnesium 2.54
Magnesium oxide 3.95
Calculations:
1. Assuming all of the magnesium has reacted, what is the mass of magnesium reacted?
2. What is the mass of the magnesium oxide produced?
3. Determine the mass of oxide in the magnesium oxide.
4. Determine the number of moles of magnesium, then the number of moles of oxide.
5. Determine the ratio:
Conclusion:
1. Since you believe that atoms combine in simple, whole-number ratios, what do you think is the likely ratio?
2. How does your value compare to the accepted value?
3. What is the empirical formula of magnesium oxide?
4. Write a balanced equation for the reaction.
5. What are some possible errors in your experiment?
6. What was the limiting reactant in the experiment?
7. What was the % yield of magnesium oxide?
