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Chapter 3

Chapter 3. General Stoichiometry. Stoichiometry. Chemical Stoichiometry- Studies quantities of materials consumed and produced in chemical reactions. Essentially, we do this by using the average atomic masses of each element.

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Chapter 3

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  1. Chapter 3 General Stoichiometry

  2. Stoichiometry • Chemical Stoichiometry- • Studies quantities of materials consumed and produced in chemical reactions. • Essentially, we do this by using the average atomic masses of each element. • Average atomic mass the sum of the atomic masses of naturally occurring isotopes of an element. • Formula: Σ (AMU Isotope1 x percent occurrence)1→∞

  3. Average Mass of an Element • For example, Carbon exists as 12C, 13C, 14C • If these three exist naturally as 98.89% 12C, 1.11% 13C, and a negligible amount of 14C, what is the average atomic mass of Carbon? • 12C = 12u, 13C = 13.00335u

  4. Average Mass of an Element • Let’s try another example: • Copper (Cu) exists naturally in two isotopes, 63Cu (62.93u) and 65Cu (64.93u). • 63Cu = 69.09%, 65Cu = 30.91% • What is the average atomic mass?

  5. The Mole • Avogadro’s Number : 6.022 x 1023 • Don’t you dare ever forget that number….EVER!!!1 • 6.022 x 1023 amu = 1 gram • Atomic mass from periodic table indicates mass of 1 atom of element in amu, also indicates the mass of 1 mole of atoms of element in grams.

  6. Percent Composition • Useful for describing the composition of a compound. • Obtain mass percents of the elements from the formula of the compound by comparing the mass of each element present in one mole of the compound to the total mass of one mole of the compound.

  7. Percent Composition • In other words: % composition = (Mass of element in 1 mol of compound) molecular mass of compound x 100

  8. Lets try a few examples • A chemist determines that 1.26 g of iron reacts with 0.54 g of oxygen to form rust. What is the percent composition of each element in the new compound? • In the lab, a chemist analyzed a sample of methanol and found that it was made of 6.2 g of carbon, 4.1 g of hydrogen, and 15.9 g of oxygen. What is the percent composition of each element?

  9. The Mole • How grams of oxygen are present in a 30.0 gram sample of potassium chlorate? • How many atoms are in a 14g sample of Carbon? • How many moles are in a sample of Potassium with 1.533 x 1064 atoms? • How many grams are in 1.66 x 109 molecules of Magnesium Phosphate?

  10. A few more examples • What is the percent composition by mass for each element in sodium phosphate, Na₃PO₄? • What is the percent composition by mass for each element in hydrogen peroxide, H₂O₂?

  11. Chapter 3 Part II General Stoichiometry

  12. Determining Empirical Formulas • Empirical Formula: the simplest whole number ratio of atoms in a compound. • First determine the molar mass if not given, THIS IS USUALLY PROVIDED. • Next, determine the number of moles of each element based on the mass percents in 100 grams of the compound. • Divide the moles of each element by the lowest number of moles.

  13. Determining Empirical Formulas:Examples • A molecule that is 75% carbon, 25% hydrogen, what is the empirical formula? • A molecule that is 22.1% aluminum, 25.4% phosphorus, 52.5% oxygen, what is the empirical formula? • A molecule that is 25.3% copper, 12.9% sulfur, 25.7% oxygen, 36.1 % water, what is the empirical formula?

  14. Determining Empirical Formulas and Molecular Formula: • A white powder is analyzed and found to contain 43.64% phosphorus and 56.36% Oxygen. The compound has a molar mass of 283.88 g/mol. What are the compounds empirical and molecular formula? • Caffeine contains 49.48% carbon, 5.15 % hydrogen, 28.87% nitrogen, and 16.49 % oxygen by mass. The Molar mass is 194.2 g/mol, what is the molecular formula.

  15. Stoichiometry: Amounts of Reactants and Products. • Steps for solving Stoichiometry questions • 1. Balance the equation • 2. Convert known masses to moles • 3. Use balanced equation to set up mole ratios • 4. Use mole ratios to calculate number of moles of the desired variable • 5. Convert moles back to grams if needed. The best way to learn this is to do it. I cant explain it any better than that.

  16. Stoichiometry Examples • 18.0 grams of carbon is burned in 55.0 grams of oxygen. How many grams of carbon dioxide are formed? • When 13.5 grams of methane (CH4) burns in 40.0 grams of oxygen, how many grams of water are formed?

  17. Lets set up tomorrow's lab

  18. Stoichiometry Examples

  19. Week 2 Homework • Start reading Chapter 3 • Stoichiometry and limiting reagent worksheet for homework.

  20. AP Chem • Start your labs, use some athletic tape to attach your capillary tube to your thermometer. • I have powdered Alum in the hood for everyone.

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