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Chapter 10: Moles, Empirical & Molecular Formulas, Hydrates

Chapter 10: Moles, Empirical & Molecular Formulas, Hydrates. What is a Mole ?. SI unit for amount of substance is called mole A mole measures the number of particles/atoms/molecules within a substance

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Chapter 10: Moles, Empirical & Molecular Formulas, Hydrates

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  1. Chapter 10: Moles, Empirical & Molecular Formulas, Hydrates

  2. What is a Mole? SI unit for amount of substance is called mole A mole measures the number of particles/atoms/molecules within a substance A mole refers to a specific number of particles (it is a counting unit, similar to a dozen)

  3. 1 mole = 6.02 x 1023 particles/atoms/molecules 6.02 x 1023 is also known as Avogadro’s Number 1 mole magnesium = 6.02 x 1023Mg atoms 1 mole zinc = 6.02 x 1023Zn atoms 1 mole lead = 6.02 x 1023Pb atoms

  4. Molar Mass Although 1 mole always contains the same number of particles(6.02 x 1023), the mass of one mole varies depending on the substance Molar Mass: The mass of onemole of a substance. Mass of one mole of an element is EQUAL to its ATOMIC MASS,which isexpressed in grams. 1 mole of magnesium = 24.31 grams 1 mole of zinc = 65.39 grams 1 mole of lead = 207.2 grams

  5. Molar Mass of a Compound • To calculate the molar mass a compound, first determine the molar mass of each element in the compound • Then, add together each element’s molar mass • The sum is the molar mass of the compound • Example #1: Determine the molar mass of NaOH. • Example #2:Determine the molar mass of HCN.

  6. Mole Conversions • Many times in chemistry, you will need to express a measurement in a unit different from the one given or measured initially • Conversion Factor:Ratio of equivalent measurements used to convert from one unit to another

  7. Mole Grams Molar Mass Mole ⬄ Grams Conversions • To convert between grams and moles, use molar mass as the conversion factor Example #1:mol⬄ grams — How many grams are in 3.01 mol of iron (Fe)? Example #2:grams ⬄ mol — How many moles are in 75 g of Fe? Example #3: mol⬄ grams — What is the mass of 2.55 mol of the compound KMnO4? Example #4:grams ⬄ mol – How many moles are in 35.0 g of HCl?

  8. Mole ⬄ Grams Conversions: Example Problem Solutions Example #1: mol ⬄ grams — How many grams are in 3.01 mol of iron (Fe)? Example #2:grams ⬄ mol — How many moles are in 75 g of Fe? 75 g Fe x 1 mol Fe = 1.3 mol Fe 55.847 g Fe

  9. Mole ⬄ Grams Conversions: Example Problem Solutions Example #3: mol⬄ grams — What is the mass of 2.55 mol of the compound KMnO4? 2.55 mol x 158 g = 402.9 g x 1 mol Example #4:grams ⬄ mol – How many moles are in 35.0 g of HCl? 35.0 g x 1 mol = 0.959 mol x 36.5 g

  10. Mole ⬄ Particles/Atoms Conversions • To convert between particles/atoms and moles, use Avogadro’s Number (6.02 × 1023) Example #1:particles ⬄ mol — How many moles equal 2.41 × 1024 atoms of aluminum (Al)? Example #2:mol ⬄ particles — How many atoms are in 3.45 mol of Al? Particles Mole Avogadro’s #

  11. Mole ⬄ Particles/Atoms Conversions: Example Problem Solutions Example #1:particles ⬄ mol — How many moles equal 2.41 × 1024 atoms of aluminum (Al)? Example #2:mol ⬄ particles — How many atoms are in 3.45 mol of Al?

  12. Mole Conversion Practice, (cont.) 2.) A chemist produced 11.25 g of magnesium (Mg). How many moles of Mg were produced? 3.) Determine the number of silver (Ag), atoms that are contained in 0.650 mol of Ag. 11.25 g x 1 mol = 0.463 mol x 24.3 g 0.65 mol x 6.02 x 1023 atoms = 3.91 x 1023 atoms x 1 mol

  13. Particle Mole Avogadro’s # Molar Mass Gram Particles ⬄ Grams Conversions • To convert between particles and grams, use BOTHmolar mass ANDAvogadro’s Number (cannot go directly from particles to grams and vice versa; must first find MOLES) Example #1:particles ⬄ grams — How many grams are in 7.00 × 1033 molecules of H2O? Example #2:grams ⬄ particles — How many molecules are in 8.25 g of H2?

  14. Particles ⬄ Grams Conversions: Example Problem Solutions Example #1:particles ⬄ grams — How many grams are in 7.00 × 1033 molecules of H2O? Example #2:grams ⬄ particles — How many molecules are in 8.25 g of H2?

  15. Empirical Formulas

  16. Percent Composition • The percent by mass of an element in a compound • Percent by mass (element) = mass of element x 100 • mass of compound • For example, in H2O, you have 2 g of hydrogen and 16 g of oxygen. The total mass of the compound is 18 g. • Hydrogen: 2 g x 100 = 11% hydrogen • 18 g • Oxygen:16 g x 100 = 89% oxygen • 18 g • The percentages should always equal 100%

  17. Percent Composition • Practice Problem • Sodium bicarbonate (NaHCO3), commonly known as baking soda, is an active ingredient in some antacids used for the relief of indigestion. Determine the percent composition of each element in NaHCO3. • NaHCO3 formula mass = 84 g • Elements = Na, H, C, O

  18. Determining the Empirical Formula of a Compound If we know how much of each element is present in a compound, we can determine the formula of the compound Amounts may be expressed in: percentages grams moles If composition is given in terms of percentages, assume you have a 100 g sample of the compound

  19. The empirical formula, or simple formula, is a formula in which the elements are in their lowest whole number ratio. The empirical formula may or may not be the actual formula. Example: The formula for glucose is C6H12O6 . Its empirical formula is CH2O. The subscripts in a chemical formula represent a MOLE ratio of elements in the compound. If given mass (grams), must convert to moles

  20. DETERMINING THE EMIRICAL FORMULA OF A COMPOUND Example #1: Analysis shows a compound consists of 32.4% Na, 22.6% S, and 45.0% O. Determine the empirical formula for this compound. Example #2: A compound is found to contain 4.43 g phosphorus and 5.72 g oxygen. Determine the empirical formula for this compound.

  21. Molecular Formulas • Molecular Formula: Actual formula of a compound; may or may not be the same as the empirical formula • Actual formula is a multiple of the empirical formula. • C6H12O6 = 6 (CH2O) • In order to find the molecular formula, you must be given the formula mass of the compound. • Divide the formula mass by the empirical mass. This will give you the multiple. • CH2O = 30 g • C6H12O6 = 180 g In this case, the multiple is 6.

  22. DETERMINING THE MOLECULAR FORMULA OF A COMPOUND • Example #1: Determine the molecular formula of a compound with an empirical formula of CH and a formula mass of 78 amu. • Example #2: A compound with a formula mass of 60.0 amu is found to be 39.9% carbon, 6.7% hydrogen, and 53.4% oxygen by mass. Find its molecular formula.

  23. Hydrates

  24. Hydrates • Hydrate: A compound that has a specific number of water molecules bound to its atoms • Formula for a hydrate is written as follows: • Formula for ionic compound x H2O • x = number of water molecules

  25. Molar Mass of Hydrates • When given the formula of a hydrate, the water molecules are included in the molar mass along with the ionic compound • Example #1: CuSO4 5 H2O • CuSO4 = 159.6 g • H2O = (5) 18.0 = 90 g • 159.6 g + 90 g = 249.6 g • Example #2:BaCl2 2 H2O • BaCl2 = 208.3 g • H2O = 2 (18.0) = 36 g • 208. 3 g + 36 g = 244.3 g

  26. Hydrates

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