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Chapter 10 – Chemical Quantities I II III IV What is the Mole? A unit of measurement used in chemistry. A counting number like – a dozen eggs, a ream of paper, a bushel of apples, etc. A mole describes an amount of matter. Definition: A mole of any element is defined as the number of atoms of that element equal to the number of atoms in exactly 12.0 grams of carbon-12. Atomic Mass Mass of a single atom 10-23 grams Atomic Mass Unit (amu) is used to describe the atomic mass of an atom. A scale was devised to express the mass of all the different atoms. Based on the mass of an atom relative to carbon-12. Atomic Masses are listed in the periodic table. What it means: 1 atom of 12C has a mass of 12.01 amu 1 mole of 12C has a mass of 12.01 grams What it means: 1 atom of 35Cl has a mass of 35.45 amu 1 mole of 35Cl has a mass of 35.45 grams Formula Mass Is calculated using the atomic mass of each element in the compound, times how many atoms of each element are in the compound. Example: Find the formula mass of H2O H = 2 x 1.0 = 2.0 grams O = 1 x 16.0 = 16.0 grams 18.0 grams/mole Example: Find the formula mass of NaOH Na = 1 x 23.0 = 23.0 grams O = 1 x 16.0 = 16.0 grams H = 1 x 1.0 = 1.0 grams 40.0 grams/mole The Atomic Mass or Formula Mass Tells the mass in grams of 1 mole of a substance There are 6.02 x 1023 objects in one mole. 602,000,000,000,000,000,000,000 A. What is the Mole? 1 mole of hockey pucks would equal the mass of the moon! 1 mole of basketballs would fill a bag the size of the earth! 1 mole of pennies would cover the Earth 1/4 mile deep! A mole of water A mole of sugar A mole of copper Example: 16.0 g of oxygen = 1 mole of oxygen = 6.02 x 1023 atoms of oxygen Example: 18.0 g of water = 1 mole of water = 6.02 x 1023 molecules of water Example: 40.0 g of NaOH dissolved in water = 1 mole of Na+ ions and = 1 mole of OH ions 6.02 x 1023 Na+ ions and 6.02 x 1023 OH- ions Molar Mass Mass in grams of 1 mole of a substance. Molar Mass Examples water H2O 2(1.01) + 16.00 = 18.02 g/mol sodium chloride NaCl 22.99 + 35.45 = 58.44 g/mol Molar Mass Examples sodium bicarbonate NaHCO3 22.99 + 1.01 + 12.01 + 3(16.00) = 84.01 g/mol sucrose C12H22O11 12(12.01) + 22(1.01) + 11(16.00) = 342.34 g/mol Molar Conversions MASS molar mass 6.02 1023 ÷ x MOLES IN GRAMS NUMBER OF x ÷ (g/mol) (particles/mol) PARTICLES Molar Conversion Examples How many moles of carbon are in 26 g of carbon? 26 g C 1 mol C 12.01 g C = 2.2 mol C Molar Conversion Examples How many molecules are in 2.50 moles of C12H22O11? 6.02 1023 2.50 mol molecules 1 mol = 1.51 1024 molecules C12H22O11 Molar Conversion Examples the mass of 2.1 1024 molecules of NaHCO3. Find 2.1 1024 molecules 1 mol 84.01 g 6.02 1023 1 mol molecules = 293 g NaHCO3 Chapter 10 – % Composition, Empirical and Molecular Formulas Suggested Reading: Pages 305 - 313 I II III IV Percentage Composition the percentage by mass of each element in a compound mass of element % composition 100 total mass Percentage Composition Find %Cu = %S = the % composition of Cu2S. 128g Cu 160 g Cu2S 32 g S 160 g Cu2S 100 = 80% Cu 100 = 20% S Percentage Composition Find the percentage composition of a sample that is 28 g Fe and 8.0 g O. 28 g 100 = 78% Fe %Fe = 36 g %O = 8.0 g 36 g 100 = 22% O Percentage Composition How many grams of copper are in a 38.0-gram sample of Cu2S? Cu2S is 80% Cu (38.0 g Cu2S)(0.80) = 30.4 g Cu Percentage Composition Find the mass percentage of water in calcium chloride dihydrate, CaCl2•2H2O? %H2O = 36 g 146 g 100 = 25% H2O Empirical Formula Smallest whole number ratio of atoms in a compound C 2H 6 reduce subscripts CH3 Empirical Formula 1. Find mass (or %) of each element. 2. Find moles of each element. 3. Divide moles by the smallest # to find subscripts. 4. When necessary, multiply subscripts by 2, 3, or 4 to get whole #’s. Empirical Formula Song (to help you remember steps) 1. Percent to Mass 2. Mass to Mole 3. Divide by Small 4. Multiply ‘til Whole Empirical Formula Find the empirical formula for a sample of 25.9% N and 74.1% O. 25.9 g 1 mol 14 g 74.1 g 1 mol 16 g = 1.85 mol N =1N 1.85 mol = 4.63 mol O = 2.5 O 1.85 mol Empirical Formula N1O2.5 Need to make the subscripts whole numbers multiply by 2 N2O5 Molecular Formula “True Formula” - the actual number of atoms in a compound empirical formula CH3 ? molecular formula C2H6 Molecular Formula 1. Find the empirical formula. 2. Find the empirical formula mass. 3. Divide the molecular mass by the empirical mass. 4. Multiply each subscript by the answer from step 3. MF mass n EF mass EF n Molecular Formula The empirical formula for ethylene is CH2. Find the molecular formula if the molecular mass is 28.1 g/mol? empirical mass = 14.03 g/mol 28.1 g/mol 14 g/mol = 2.00 (CH2)2 C2H4