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Chapter 10

Chapter 10. Chemical Quantities. Chapter 10 Section 1. The Mole. Measuring matter. We measure by three different methods Counting- How many there are. Volume- How much space they take up. Mass- How much they weigh. We have to know conversion factors to relate these different measurements.

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Chapter 10

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  1. Chapter 10 Chemical Quantities

  2. Chapter 10 Section 1 The Mole

  3. Measuring matter • We measure by three different methods • Counting- How many there are. • Volume- How much space they take up. • Mass- How much they weigh. • We have to know conversion factors to relate these different measurements.

  4. Moles • A mole of a substance is 6.02 x 1023 particles. The mole is the SI unit for measuring amount of matter. • Avagadro’s number is 6.02 x 1023. • This is the number of particles for one mole of any subtance. • The particles making up a mole of substance can be atoms, molecules, or formula units.

  5. Converting • How many particles are in three moles • 3 moles x 6.02 x 1023 particles 1 mole 3 x 6.02 x 1023 = 1.81 x 1024 particles

  6. Atomic Mass • The atomic mass of an element is the mass of a single atom. The units are amu or atomic mass units. • The atomic mass of Carbon is 12.0 amu, the atomic mass of Hydrogen is 1.0 amu, so Carbon is 12x heavier than Hydrogen.

  7. Molar mass • Molar mass is the mass in grams of one mole of an element. • One mole of H has a mass of 1.0g. One mole of C is 12.0g • The molar mass of any element will contain 6.02x1023 atoms. • Always round molar masses to one decimal place.

  8. Compounds • To find molar mass of a compound, add the masses of each element in the compound. • Example CO3= 1 atom C, 3 atoms O • C 12.0amu x1=12.0 O 16.0amu x3 = 48.0 • 12.0 + 48.0 = 60.0 amu • Since we have exactly one mole, substitute grams for amu and we will have the molar mass of 60g for one mole of CO3.

  9. Chapter 10 Section 2 Mole relationships

  10. Mole to mass relationships • We can convert between moles and mass by using molar mass. • Mass = moles x molar mass • Moles = mass  molar mass

  11. Examples • How many grams are in 4 moles of H2O. The molar mass of H2O is 18g/mol. • Take 4 moles times 18 g/mol = 72g. • How many moles are in 132grams of CO2. • The molar mass of CO2 is 44g/mol. • Divide 132 grams by 44g/mol. = 3 moles

  12. Mole – volume relationships • Avagadro’s hypothesis- equal volumes of gases at the same temperature and pressure have equal number of particles. • Standard Temperature and Pressure STP- 0º C and 1 atm pressure. • AT STP 1 mole of any gas occupies a volume of exactly 22.4 L. • 22.4 L is called the molar volume of a gas.

  13. Mole- volume calculations • We will use 22.4L per one mole as a conversion factor between moles and vol. • Volume of gas= moles x 22.4 L/mol • Moles = volume of gas  22.4 L/mol

  14. Density and molar mass conversions • We can calculate molar mass from density by the following equation. • Molar mass = density at STP x 22.4 L/mol • Density at STP = molar mass  22.4 L/mol

  15. Chapter 10 Section 3 Percent Composition

  16. Percent composition • The relative amounts of the elements in a compound. • Also called percent by mass • % of element= mass of element x 100 mass of compound • % of element = mass in 1 mole x 100 molar mass of compound

  17. Example • Water (H2O) Molar mass is 18 • % Hydrogen = (2/18) x 100 = 11.1% H • % Oxygen = (16/18) x 100 = 88.9% O • Water is 11% Hydrogen and 89% oxygen • The individual percentages of any compound should add up to 100%

  18. Converting with % composition • You can use % composition to find the number of grams of any element when given the mass of a compound. • The number of grams of any element in a compound will be the total grams of the compound times the percent composition of that element  100.

  19. Example • Example – how many grams of Oxygen would be in 50 grams of water. • 50g H2O x 88.9g O = 44.45 grams O 100 g H2O

  20. Empirical formulas • The lowest whole number ratio of the atoms in a compound. • Example-The molecular formula of ethane is C2H6. • The empirical formula would be CH3. • Molecular formulas will always be the same as or a whole number multiple of the empirical formula.

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