The Mole Chemistry – Chapter 11

Measuring Matter  What measurements do we use?  Pair  Dozen  Gross  Ream  Counting Particles  Atoms and molecules are extremely small, so it’s virtually impossible to count them  Mole (mol) – SI base unit used to measure the amount of a substance  Number of representative particles, carbon atoms, in exactly 12 g of pure C-12

 A mole of anything contains 6.02 x representative particles  Representative particles – any kind of particle, such as atoms, molecules, formula units, electrons, or ions  Ex: representative particle in a mole of water is the water molecule; representative particle in a mole of copper is the copper atom; representative particle in a mole of sodium chloride is the formula unit

 Converting Moles to Particles and Particles to Moles  If you buy 3 ½ dozen roses, how many roses did you buy? Conversion factor: 12 roses/1 dozen 3.5 doz x 12 roses/1 doz = 42 roses  Determine how many representative particles of sucrose are in 3.5 moles of sucrose (for sucrose, the representative particle is a molecule) Conversion factor: 6.02 x molecules/1 mol 3.5 mol x 6.02 x molecules/1 mol = 2.11 x molecules

Mass and the Mole  Atomic masses of all elements are established relative to C-12  Ex: An atom of H-1 has a mass of 1 amu; the mass of an atom of He-4 is 4 amu  The mass of 1 mol of C-12 atoms is 12 g  The mass of 1 mol of H-1 is 1 g  The mass of 1 mol of He-4 is 4 g  Molar mass – mass in grams of one mole of any pure substance  Molar mass of any element is equal to its atomic mass and has the units g/mol  Ex: An atom of Mn has an atomic mass of amu; therefore the molar mass of Mn is g/mol

 Suppose that you need 3 mol of Mn for a chemical reaction 3 mol Mn x 54.9 g Mn/1 mol Mn = 165 g Mn  The mole is at the CENTER of all chemistry mass conversions  Mass must be converted to moles before being converted to atoms and atoms must be converted to moles before calculating their mass  Ex: How many molecules of water are in 1 g of water? 1 g H 2 O x 1 mol H 2 O/18.02 g H 2 O x 6.02 x molecules H 2 O/1 mole H 2 O = 3.34 x molecules H 2 O

Moles of Compounds (Refer to textbook for example problems)

Empirical and Molecular Formulas  Percent Composition % by mass = mass of element/mass of compound x 100  Percent composition – percent by mass of each element in a compound  Percent composition from the chemical formula – use the chemical formula to calculate the molar mass of water (18.02 g/mol) and assume you have an g sample  B/c the % composition is always the same regardless of the size of the sample, assume that the sample is the size of one mole g H/18.02 g H 2 O x 100 = 11.2% H 16 g O/18.02 g H 2 O x 100 = 88.8% O

 Empirical Formula – formula w/ the smallest whole- number mole ratio of the elements  Ex: The empirical formula for hydrogen peroxide is HO; the molecular formula is H 2 O 2  Molecular Formula – specifies the actual number of atoms of each element in one molecule or formula unit of the substance