Stoichiometry Mole-Mass Relationships in Chp 2

Molecular Formulas • The actual number of moles of each element in the smallest unit of the compound. •In water (H 2O), ammonia (NH 3), methane (CH 4), and ionic compounds, the empirical and molecular formulas are identical. •In some cases the molecular formula is a whole-number multiple of the empirical formula. • Using the empirical formula to obtain the molecular formula. Hydrogen peroxide has the empirical formula HO (17.01g/mol) Dividing its molar mass (34.02 g/mol) by the empirical formula mass gives the whole-number multiple: The molecular formula is H 2O2 . 19 Skill 3-5 : Lactic acid (M = 90.08 g/mol) contains 40.0 mass % C, 6.71 mass % H, and 53.3 mass % O. (a) Determine the empirical formula of lactic acid. (b) Determine the molecular formula. Plan: Assume 100 g lactic acid to express each mass % as grams. Convert grams to moles and find the empirical formula. Solution: Express mass % as grams; assume 100 g lactic acid: Similarly, there are 6.71 g H and 53.3 g O. Converting from grams of elements to moles: 20 Similarly, we have 6.66 mol H and 3.33 mol O. Constructing the preliminary formula: Converting to integer subscripts: Determining the molecular formula: Plan: Divide the molar mass by the empirical formula mass. Solution: Calculate the molecular formula: 21 Combustion Analysis • Combustion analysis gives the amounts of carbon and hydrogen in a sample of combustible compound. • The C forms CO 2 , which is absorbed in the first chamber. • The H forms H 2O, which is absorbed in the second chamber. • Weigh the chambers before and after combustion. • Calculate the masses of CO 2 and H 2O • Use them to calculate the masses of C and H in the compound. • From the masses, calculate the empirical formula. 22 Skill 3-5 contd: Vitamin C (M= 176. g/mol) contains C, H and O. A 1.000 g sample was placed in a combustion apparatus: What is the molecular formula of vitamin C? Plan: -Use changes in mass of CO 2 and H 2O absorbers to calculate the moles of C and H present in the sample. -Find the mass of C, using the mass fraction of C in CO 2 . Likewise, find the mass of H from the mass of H2O. -The mass of vitamin C minus the sum of the C and H masses gives the mass of O. -Construct the empirical formula and molecular formula. 23 Solution: Finding the masses of combustion products: Calculating mass fractions of the elements: 24 Calculate masses of C and H: Calculate mass of O by difference: Find the moles of elements: Divide mass of each by its molar mass gives 0.0341 mol C, 0.046 mol H, and 0.0341 mol O. Construct the preliminary formula: 25 Divide through by the smallest subscript: Determine the empirical formula. By inspection, find that 3 is the smallest multiple that makes all subscripts into integers: Determine the molecular formula: 26 Chemical Formulas and the Structures of Molecules • Chemical formulas represent real three-dimensional objects. • The molecular formula tells the number of each type of atom. • Molecular formulas may not be unique; the same types and numbers of atoms can bond to each other in more than one way •Can be different compounds or structural isomers of the same compound (same molecular formulas). 27 Structural Formulas •Consider C2H 6O; two very different compounds have this molecular formula: -ethanol, the intoxicating substance present in wine and beer -dimethyl ether, a colorless gas once used in refrigeration •Their radically different physical and chemical behaviors are the result of different molecular structures. ethanol dimethyl ether 28 Concept 6. The importance of balancing equations for the quantitative study of chemical reactions •For an equation to depict accurately the amounts of chemicals involved in a reaction, it must be balanced. • The same number of each type of atom must appear on both sides of the equation. • In a flashbulb Mg wire and oxygen gas react to yield magnesium oxide powder. 29 Translate the chemical statement into a “skeleton” equation: chemical formulas arranged in an equation format. reactants “yield” products Balancing the atoms •Each blank must contain a balancing coefficient, a numerical multiplier of all the atoms in the formula that follows it. •Stepwise, match the atoms on each side, element by element, beginning with the most complex substance. 31 Requires one Mg on the left. And one oxygen on the left. Adjusting the coefficients. •The smallest whole-number coefficients are required. •To remove the 1/2 O2, multiply the whole equation by 2: •A coefficient of 1 is implied by the presence of the substance and is not written: • Check the numbers of atoms: Reaction equation is balanced. 32 Specifying the states of matter. •A complete equation shows the physical state of each substance or whether it is dissolved in water. solid (s) liquid (l) gas (g) aqueous solution (aq) • From the statement, Mg “wire” is solid, O 2 is a gas, and “powdery” MgO is a solid: 2Mg(s) + O2(g) 2MgO(s) 33 Skill 3-6: Octane (C8 H 18) burns to form carbon dioxide and water vapor. Write a balanced equation for this reaction. Solution: Translate the statement into a skeleton equation (with coefficient blanks). Octane and oxygen are reactants; oxygen is molecular O 2 . Carbon dioxide and water vapor are products: Balance the atoms. Start with the more complex reactant, C 8H 18 , then balance O 2: 8 C on left requires 8 CO 2 on right: 34 18 H on left requires 9 H2O. 25 O on right requires 25/2 O2 Adjust coefficients. Check that the equation is balanced: reactants (16 C, 36 H, 50 O) = products (16 C, 36 H, 50 O) States of matter: C 8H 18 is liquid; O 2 , CO 2 , and H 2O vapor are gases: 35 Concept 7 The mole-mass-number information contained in a balanced equation • A balanced equation relates quantities of atoms, molecules, formula units, moles of substance, and masses. 36 Concept 8 The relation between amounts of reactants and products • Quantitative relationships in a balanced equation are expressed as stoichiometrically equivalent molar ratios. • Use stoichiometrically equivalent molar ratios to determine how much of one substance forms from (or reacts with) another. • 3 mol CO 2 is stoichiometrically equivalent to 4 mol H 2O, 5 mol O2 is stoichiometrically equivalent to 3 mol CO 2 , and so on. • In the combustion of propane, how many moles of O 2 are consumed when 10.0 mol H 2O are produced? starting amt. molar ratio equiv. amt. 37 Skill 3-7: Copper is obtained from sulfide ores by roasting the ore with O2 to form powdered copper(I) oxide and gaseous sulfur dioxide. (a) How many moles of oxygen are required to roast 10.0 mol copper(I) sulfide? Write the balanced equation first. The reactants are Cu 2S and O 2 , and the products are Cu 2O and SO 2 : Given the moles of Cu 2S; need to find the moles of O 2. (b) How many grams of sulfur dioxide are formed when 10.0 mol copper(I) sulfide are roasted? 38 Find mass of SO 2 that forms from the given mol of reactant (Cu2S). (c) How many kg of oxygen are needed to form 2.86 kg Cu 2O? Convert mass Cu2O to mol: Mol Cu2O to mol O2: Mol O2 to mass O2: 39