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Chapter 9 Chemical Quantities:

Georgia Performance Standards: SC2 (c, d & e): Students will relate how the Law of Conservation of Matter is used to determine chemical composition in compounds and chemical reactions.

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Chapter 9 Chemical Quantities:

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  1. Georgia Performance Standards: SC2 (c, d & e): Students will relate how the Law of Conservation of Matter is used to determine chemical composition in compounds and chemical reactions. c. Apply concepts of the mole and Avogadro’s number to conceptualize and calculate the empirical & molecular formulas, Mass, moles and molecules relationships, and molar volumes of gases. Essential Questions: -How is the molar mass calculated for various compounds? How do you convert between moles, mass, and number of atoms? How is the mass % of an element in a compound calculated? Chapter 9 Chemical Quantities:

  2. Mass, Moles, Number of atoms relationships

  3. Performance Tasks: • Calculate the molar mass for various compounds. • Use the mole concept, Avogadro’s number, & molar mass to convert among moles, mass, and number of atoms • Calculate the mass percent of an element in a compound

  4. Law of Conservation of Mass • We balance chemical equations to satisfy the law of conservations of mass • Mass is neither created nor destroyed, only transformed

  5. Chemical Equations Chemical equations are concise representations of chemical reactions.

  6. CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g) Anatomy of a Chemical Equation

  7. Reactants appear on the left side of the equation. CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g) Anatomy of a Chemical Equation

  8. Products appear on the right side of the equation. CH4 (g) + 2 O2 (g) CO2(g) + 2 H2O(g) Anatomy of a Chemical Equation

  9. The states of the reactants and products are written in parentheses to the right of each compound. CH4 (g) + 2 O2 (g) CO2(g) + 2 H2O(g) Anatomy of a Chemical Equation

  10. Coefficients are inserted to balance the equation. CH4 (g) + 2 O2 (g) CO2(g) + 2 H2O(g) Anatomy of a Chemical Equation

  11. Subscripts and Coefficients Give Different Information • Subscripts tell the number of atoms of each element in a molecule.

  12. Subscripts and Coefficients Give Different Information • Subscripts tell the number of atoms of each element in a molecule • Coefficients tell the number of molecules.

  13. Notice • The number of atoms of each type of element must be the same on both sides of a balanced equation. • Subscripts must not be changed to balance an equation. • A balanced equation tells us the ratio of the number of molecules which react and are produced in a chemical reaction. • Coefficients can be fractions, although they are usually given as lowest integer multiples. • Trial and error is a valid method to balance a chemical equation.

  14. Balancing Chemical Equations

  15. Chemical Equation • C2H5OH+3O22CO2 + 3H2O • The equation is balanced. • 1 mole of ethanol reacts with 3 moles of oxygen to produce 2 moles of carbon dioxide and 3 moles of water

  16. Formula Weights

  17. Formula Weight (FW) • A formula weight is the sum of the atomic weights for the atoms in a chemical formula. • So, the formula weight of calcium chloride, CaCl2, would be Ca: 1(40.1 amu) + Cl: 2(35.5 amu) 111.1 amu • Formula weights are generally reported for ionic compounds.

  18. + H: 6(1.0 amu) Molecular Weight (MW) • A molecular weight is the sum of the atomic weights of the atoms in a molecule. • For the molecule ethane, C2H6, the molecular weight would be C: 2(12.0 amu) 30.0 amu

  19. Atomic Masses • Elements occur in nature as mixtures of isotopes • Carbon = 98.89% 12C 1.11% 13C <0.01% 14C Carbon atomic mass = 12.01 amu

  20. (number of atoms)(atomic weight) x 100 % element = (FW of the compound) Percent Composition One can find the percentage of the mass of a compound that comes from each of the elements in the compound by using this equation:

  21. (2)(12.0 amu) %C = (30.0 amu) x 100 24.0 amu = 30.0 amu Percent Composition So the percentage of carbon in ethane is… = 80.0%

  22. Percent Composition • Mass percent of an element: • For iron in iron (III) oxide, (Fe2O3)

  23. Moles

  24. The Mole • The number equal to the number of carbon atoms in exactly 12 grams of pure 12C. • 1 mole of anything = 6.022 * 1023 units of that thing

  25. Avogadro’s Number • 6.02 x 1023 • 1 mole of 12C has a mass of 12 g.

  26. Molar Mass • By definition, a molar mass is the mass of 1 mol of a substance (i.e., g/mol). • The molar mass of an element is the mass number for the element that we find on the periodic table. • The formula weight (in amu’s) will be the same number as the molar mass (in g/mol).

  27. Using Moles Moles provide a bridge from the molecular scale to the real-world scale.

  28. Mole Relationships • One mole of atoms, ions, or molecules contains Avogadro’s number of those particles. • One mole of molecules or formula units contains Avogadro’s number times the number of atoms or ions of each element in the compound.

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