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C H E M I S T R Y Chapter 3 Mass Relationships in Chemical Reactions Chemical Reaction In a chemical reaction • a chemical change produces one or more new substances. • there is a change in the composition of one or more substances. Chemical Reaction In a chemical reaction • old bonds are broken and new bonds are formed. • atoms in the reactants are rearranged to form one or more different substances. Chemical Reaction In a chemical reaction • the reactants are Fe and O2. • the new product Fe2O3 is called rust. Chemical Equations A chemical equation gives • the formulas of the reactants on the left of the arrow. • the formulas of the products on the right of the arrow. Reactants O2 (g) Product CO2 (g) C(s) 5 Symbols Used in Equations Symbols in chemical equations show TABLE • the states of the reactants. • the states of the products. • the reaction conditions. 6 Chemical Equations are Balanced In a balanced chemical reaction • no atoms are lost or gained. • the number of reacting atoms is equal to the number of product atoms. Copyright © 2009 by Pearson Education, Inc. 7 Balancing Chemical Equations A balanced chemical equation shows that the law of conservation of mass is adhered to. In a balanced chemical equation, the numbers and kinds of atoms on both sides of the reaction arrow are identical. 2Na(s) + Cl2(g) 2NaCl(s) left side: right side: 2 Na 2 Cl 2 Na 2 Cl Balancing Chemical Equations 1. Write the unbalanced equation using the correct chemical formula for each reactant and product. H2(g) + O2(g) 2. H2O(l) Find suitable coefficients—the numbers placed before formulas to indicate how many formula units of each substance are required to balance the equation. 2H2(g) + O2(g) 2H2O(l) Balancing Chemical Equations 3. Reduce the coefficients to their smallest whole-number values, if necessary, by dividing them all by a common denominator. 4H2(g) + 2O2(g) 4H2O(l) divide all by 2 2H2(g) + O2(g) 2H2O(l) Chapter 3/10 Balancing Chemical Equations 4. Check your answer by making sure that the numbers and kinds of atoms are the same on both sides of the equation. 2H2(g) + O2(g) 2H2O(l) left side: right side: 4H 2O 4H 2O Copyright © 2008 Pearson Prentice Hall, Inc. Chapter 3/11 Balancing Chemical Equations Do not change subscripts when you balance a chemical equation. You are only allowed to change the coefficients. H2(g) + O2(g) 2H2(g) + O2(g) Balanced properly 2H2O(l) H2O(l) unbalanced H2(g) + O2(g) H2O2(l) Chemical equation changed! Examples Balance the coefficients from reactants to products. A. __Fe2O3(s) + __C(s) __Fe(s) + __CO2(g) B. __Co2O3(s) + __ C(s) __Co(s) + __CO2(g) Write a balanced equation for the reaction between solid silicon dioxide and solid carbon to produce solid silicon carbide and carbon monoxide gas 13 Chemical Symbols on Different Levels 2H2(g) + O2(g) microscopic: macroscopic: 2H2O(l) 2 molecules of hydrogen gas react with 1 molecule of oxygen gas to yield 2 molecules of liquid water. 0.56 kg of hydrogen gas react with 4.44 kg of oxygen gas to yield 5.00 kg of liquid water. How can we relate the two to each other? Avogadro’s Number and the Mole Molecular Mass: Sum of atomic masses of all atoms in a molecule. Formula Mass: Sum of atomic masses of all atoms in a formula unit of any compound, molecular or ionic. HCl: C 2H 4: 1.0 amu + 35.5 amu = 36.5 amu 2(12.0 amu) + 4(1.0 amu) = 28.0 amu Molar Mass of Na2SO4 Calculate the molar mass of K3PO4. Element Number of Moles Na Atomic Mass Total Mass in Na2SO4 S O 16 Stoichiometry: Chemical Arithmetic Stoichiometry: The relative proportions in which elements form compounds or in which substances react. aA + bB Grams of A Molar Mass of A Moles of A cC + dD Moles of B Mole Ratio Between A and B (Coefficients) Grams of B Molar Mass of B Stoichiometry: Chemical Arithmetic Aqueous solutions of sodium hypochlorite (NaOCl), best known as household bleach, are prepared by reaction of sodium hydroxide with chlorine gas: 2NaOH(aq) + Cl2(g) NaOCl(aq) + NaCl(aq) + H2O(l) How many grams of NaOH are needed to react with 25.0 g Cl2? Grams of Cl2 Molar Mass Moles of Cl2 Mole Ratio Moles of NaOH Grams of NaOH Molar Mass Stoichiometry: Chemical Arithmetic The commercial production of iron from iron ore involves the reaction of Fe2O3 with CO to yield iron metal plus carbon dioxide: Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g) Predict how many grams of CO will react with 0.500 moles Fe2O3? Yields of Chemical Reactions Actual Yield: The amount actually formed in a reaction. Theoretical Yield: The amount predicted by calculations from the limiting reactant. Percent Yield = actual yield theoretical yield x 100% Reactions with Limiting Amounts of Reactants Limiting Reactant: The reactant that is present in limiting amount. The extent to which a chemical reaction takes place depends on the limiting reactant. Excess Reactant: Any of the other reactants still present after determination of the limiting reactant. Reactions with Limiting Amounts of Reactants At a high temperature, ethylene oxide reacts with water to form ethylene glycol which is an automobile antifreeze and a starting material in the preparation of polyester polymers: C2H4O(aq) + H2O(l) C2H6O2(l) Examples How many sandwiches can be made from 10 slices of bread and 8 slices of cheese? Which one is the limiting reactant? The balanced chemical equation is 2Bd + Ch Bd2Ch Example Ammonia, NH3, can be synthesized by the following reaction: 2 NO(g) + 5 H2(g) 2 NH3(g) + 2 H2O(g) Starting with 86.3 g NO and 25.6 g H2 a. Determine the limiting reactant b. Find the theoretical yield of NH3 in grams Reactions with Limiting Amounts of Reactants Lithium oxide is used aboard the space shuttle to remove water from the air supply according to the equation: Li2O(s) + H2O(g) 2LiOH(s) If 80.0 g of water are to be removed and 65.0 g of Li2O are available, which reactant is limiting? How many grams of excess reactant remain? How many grams of LiOH are produced? Percent Composition and Empirical Formulas Percent Composition: Expressed by identifying the elements present and giving the mass percent of each. Empirical Formula: It tells only the ratios of the atoms in a compound. Molecular Formula: It tells the actual numbers of atoms in a compound. It can be either the empirical formula or a multiple of it. multiple = molecular mass empirical formula mass Steps in determine the Empirical formula Step 1: Obtain the mass of each element (in grams) Step 2: Determine the numbers of atom present Step 3: Divide the smallest moles by numbers of each atoms to obtain the closet integer as possible. Step 4: If the result ended with 0.5, 0.33, 1.125, 1.50 etc… then multiply with a factor to get the nearest integer as possible. Example An unknown sample gives the following mass percent: 17.5% Na, 39.7% Cr and 42.8% O. What is the empirical formula? Mass percents Molar masses Moles Mole ratios Relative mole ratios Subscripts Example A compound containing nitrogen and oxygen is decomposed in the laboratory and produces 24.5 g nitrogen and 70.0 g oxygen. Calculate the empirical formula of the compound. Examples Ibuprofen, an aspirin substitute, has the following mass percent composition: C 75.69%, H 8.80% an O 15.51%. What is the empirical formula of ibuprofen? Example Butanedione – a main component in the smell and taste of butter and cheese – contains the elements carbon, hydrogen, and oxygen. The empirical formula of butanedione is C2H3O and its molar mass is 86.09 g/mol. Find its molecular formula Determining Empirical Formulas: Elemental Analysis Combustion Analysis: A compound of unknown composition (containing a combination of carbon, hydrogen, and possibly oxygen) is burned with oxygen to produce the volatile combustion products CO2 and H2O, which are separated and weighed by an automated instrument called a gas chromatograph. hydrocarbon + O2(g) carbon hydrogen xCO2(g) + yH2O(g) Percent Composition and Empirical Formulas A colorless liquid has a composition of 84.1 % carbon and 15.9 % hydrogen by mass. Determine the empirical formula. Also, assuming the molar mass of this compound is 114.2 g/mol, determine the molecular formula of this compound. Example Upon combustion, a 0.8233 g sample of a compound containing only carbon, hydrogen, and oxygen produced 2.445 g CO2 and 0.6003 g H2O. Find the empirical formula of the compound Example A compound contains only carbon, hydrogen, and oxygen. Combustion of 10.68 mg of compound yields 16.01 mg CO2 and 4.37 mg H2O. The molar mass of the compound is 176.1 g/mol.