Chapter 3 Mathematics of Formulas and Equations PAGES (46-56)

Aim: To Apply Simple Math Relations to the Solving of Chemical Problems

Gram Formula Mass Compounds are made up of more than one element Gram formula mass- sum of the atomic masses of all of the atoms in the compound To find gram formula mass (GFM): List the elements present Count how many of each element Find each element’s atomic mass Multiple the atomic mass by the number of each element Add it together

Examples: Find each of the GFM NaBr H2O Ca(OH)2 Mg(NO3) 2 (NH4)3PO4

Open your book to p. 47 (1-13) 4 4 = 254g/mol 3 4 1 2

Percent Composition Composition of a substance within a compound, expressed as a percent Table T To find the percent composition: Find the GFM of the compound(s) Find the mass of the part in the question Plug each part into the formula Hydrates- Compound plus a water part CuSO4 * 5H2O Many times, they will ask for how much water there is within the hydrate

Examples: What is the percent composition in each compound? Hydrogen in H2O Carbon in CO2 Oxygen in Ca(OH)2 Nitrogen in Fe(NO3)3 Chlorine in KClO3

Examples: Find the amount of water within each hydrate Na2CO3 * 10H2O CaSO4 * 2H2O CH2O * 3H2O CuSO4 * 5H2O Mg3(PO4) * 4H2O

Open your book to p. 48-49 (14-26)

The MOLE A collective noun Number of atoms present 6.02 x 1023 Dozen, gross, pair, etc. Number of atoms present Actually, it’s the number of atoms present in 12 grams of Carbon-12 6.02 x 1023 Avagadro’s number GFM = 1 mol of a substance

Converting Between Grams and Moles Set up a proportion: 1 mol = # of moles GFM # of grams Converting Grams to Moles: Moles = number of grams x 1/GFM Converting Moles to Grams: Grams = number of moles x GFM/1

Examples: Convert the following grams to moles Converting Grams to Moles: Moles = number of grams/GFM 36g of H2O 185g of KCl 80.0g of C2H5Cl

Examples: Convert the following moles to grams Converting Moles to Grams: Grams = number of moles x GFM 3mol of CO2 2.1mol of NaCl 0.12mol of C2H6 1mol of H2O

Finding Molecular Formulas from Empirical Formulas When the molecular mass and the empirical formula are known, you can find the molecular formula. If the empirical formula is CH2, then you know that the empirical formula’s mass must be 14g. If the molecular mass is 56g, it must be a multiple of the empirical. So, divide the given mass by the GFM of the empirical formula 56 / 14 = The answer will ALWAYS be a whole number. Now, multiply each element by the number to find the molecular formula. What is it?

Examples: Find the Molecular Formula Molecular Mass = 36g Empirical Formula = H2O Molecular Mass = 140g Empirical Formula = CH2 Molecular Mass = 180g Empirical Formula = CH2O Molecular Mass = 60g Empirical Formula = NO

Mole Relations in Balanced Equations The coefficients in a balanced equation equal the number of moles of each element/compound. So, the ratios of reactants and products will ALWAYS stay the same. For example, 2H2 + O2 => 2H2O For every 2 moles of H2 used, there are 2 moles of H2O produced. For every 2 moles of H2 used, there is 1 mole of O2 also used. Now, if 4 moles of H2 used, how many moles of H2O are produced? Remember that it’s the SAME ratio! Also, if 3 moles of H2 used, how many moles of O2 are used?

Examples: CH4 + 2O2 => CO2 + 2H2O 2CO + O2 => 2CO2 What is the mole ratio of CH4 to O2? So, how many moles of O2 are need for 3mol of CH4? 2CO + O2 => 2CO2 What is the mole ratio of CO to CO2? So, how many moles of CO are needed to create 1mol of CO2? Mg + 2HCl => MgCl2 + H2 What is the mole ratio of HCl to MgCl2? So, if 3mol HCl were used, how many MgCl2 were created?