STOICHIOMETRY Objective The purpose of this exercise is to give you some practice on some Stoichiometry calculations. Discussion The molecular mass of a compound is the sum of the atomic masses of all the atoms present in one molecule of the compound. It represents the mass of one molecule of that compound in atomic mass units (amu). The formula mass of an ionic compound is calculated by adding up the atomic masses of all the atoms present in one formula unit of the compound. A chemical reaction is a shorthand way of showing the changes occurring in any chemical reaction: reactant 1 + reactant 2 +...! product 1 + product 2 +... Every chemical equation must be balanced, which means that, for every element shown in the equation, the number of atoms on the left must equal the number on the right: Example: 3H2 + N2! 2 NH3 : 6 H atom & 2 N atoms (left) = 2 N atoms & 6 H atoms (right) The 3 before H2, the (implicit) 1 before N2, and the 2 before NH3 are the coefficients of the equation. The coefficients in chemical equations represent moles as well as atoms and molecules. The equation: 3H2 + N2! 2 NH3, for instance, tells us that three moles of H2 combines with one mole of N2 to form two moles of NH3. Knowing the number of moles, we can calculate the amounts of reactants and products: Molar mass of H2 is 2 g / mol; 1 mol has a mass of 2 g, 3 mol has a mass of 6 g. Molar mass of N2 is 28 g / mol; 1 mol has a mass of 28 g. Molar mass of N2 is 17 g / mol; 1 mol has a mass of 17 g, 2 mol has a mass of 34 g. 6 gram H2 + 28 gram N2! 34 gram NH3 Stoichiometry Relationship for molecular, liquid, solid, gas and aqueous s. Vol ( L) Liquid Density (g / cc) # of molecules / atoms N Av (6.02 10 23 ) particle (atomic) # of molecules / atoms N Av (6.02 10 23 ) particle (atomic) Liquid (l) Density (g / cc) Mass (g) Conc. (mol / L) Pressure ( atm) Temperature (K) Volume (L) Molar Mass (g / mol) Solid Aqueous Gas Balance equation Moles A!# ## ## ## # " Moles B Stoic. coefficient. R (.0821 atm L mol K ) Molar Mass (g / mol) Solid (g) Gas R (.0821 atm L mol K ) Aqueous Mass (g) Conc. (mol / L) Pressure ( atm ) Temperature ( K) Volume ( L)
Stoichiometry Name Show calculations setups and answers for all problems. 1. Use the equation given to solve the following problem: Na3PO4 + 3AgNO3! Ag3PO4 + 3NaNO3 a) How many moles of AgNO3 would be required to react with 1.0 mol of Na3PO4. b) How many moles of NaNO3 can be produced from 0.30 mol of Na3PO4? c) What weight of Ag3PO4 can be produced from 4.00 g of AgNO3? d) If you have 8.44 g of Na3PO4, what weight of AgNO3 will be needed for complete reaction? e) When 25.0 g of AgNO3 is combined with excess Na3PO4, 17.7 g Ag3PO4 is produced. What is the % yield? 2. Use the equation given to solve the following problem: 2 KMnO4 + 16HCl! 5 Cl2 + 2 KCl + 2MnCl2 + 8H2O a) How many moles of HCl are required to react with 45 g of KMnO4? b) How many Cl2 molecules will be produced using 5.0 mol of KMnO4?
c) To produce 55.0 g of MnCl2, what mass of HCl is needed? d) How many moles of water will be produced when 7.0 mol of KMnO4 are consumed? e) What is the maximum weight of Cl2 that can be produced by reacting 35.0 g of KMnO4 with 45.0 g of HCl? 3. Use the equation given to solve the following problem: 6KI + 8HNO3! 6KNO3 + 2NO + 3I2 + 4H2O [Note that 1 mole of any gas measured at STP (standard temp and pressure) will occupy 22.4 L volume.] a) If 38 g of KI are reacted, what mass of KNO3 will be formed? b) What volume of NO gas, measured at STP, will be produced if 37.0 g of HNO3 are consumed? c) If 0.50 mol of KI is to be reacted, what volume (ml) of 6.00 M HNO3 will be required? d) When the reaction produces 6.0 mol of NO, how many molecules of I2 will be produced? e) How many grams of iodine can be obtained by reacting 45.0 ml of 0.350 M KI solution?
4. Use the equation given to solve the following problem. Since all substances are in the gas at STP, use the conversion factor of 1 mole = 22.4 L for any gasous substance N2 (g) + H2 (g)! 2NH3 (g) a) If 3.0 mol of N2 react, how many moles of NH3 will form? b) When 3.50 mol of N2 react, what volume of NH3, measured at STP, will form? c) If 9.0 L of H2 are reacted, what volume of NH3 will form at STP? d) How many molecules of NH3 will form when 20.0 L of N2 react with excess hydrogen at STP? e) If 8.0 L of N2 and 20.0 L of H2 are mixed, what volume of NH3 can be produced under STP conditions. 5. Solution Stoichiometry a) What will be the percent composition by weight of a solution made by dissolving 12.0 g of zinc nitrate, Zn(NO3)2, in 45.0 g of water? Assume the densty of water is 1.00 g/ml b) How many moles of sodium hydroxide, NaOH, are required to prepare 2.00 L of 0.380 M solution?
c) What will be the molarity of a solution if 3.50 g of potassium hydroxide, KOH, are dissolved in water to make 150 ml of solution? d) What volume (ml) of 0.400 M solution can be prepared by dissolving 5.00 g of KOH of KOH in water? e) What weight of potassium bromide, KBr, could be recovered by evaporating 650 g of 15 percent KBr solution to dryness? f) What volume of 1.00 M HCl is needed to prepare 300 ml of 0.250 M HCl solution? g) A sulfuric acid solution has a density of 1.73 g/ml and contains 80 percent H2SO4 by weight. What is the molarity of this solution?