Chapter 5 Chemical Reactions 5.1 Chemical Equations A chemical equation shows the chemical change taking place. The state of each substance is written in parentheses after the formula: s for solids, l for liquids, g for gases, and aq for aqueous. Reactants H 2 (g) + O 2 (g) H 2 O (l) Products Balancing Chemical Equations Balanced chemical equation the number of atoms of each element in the reactants is equal to the number of atoms of that same element in the products. Reactions must be balanced to obey the law of conservation of mass. Coefficients are written to the left of each reactant or product in order to achieve balance. 2 H 2 (g) + O 2 (g) 2 H 2 O (l) 1
Examples of Balancing Chemical Equations Methane, CH 4,is the main ingredient in natural gas. It combines with oxygen, O 2, when it burns to form carbon dioxide, CO 2, and water, H 2 O. Write a balanced chemical equation for this reaction. Balance the following chemical equations: Mg(s) + O 2 (g) MgO(s) Na 3 PO 4 (aq) + Ca(NO 3 ) 2 (aq) NaNO 3 (aq) + Ca 3 (PO 4 ) 2 (s) C 6 H 14 (l) + O 2 (g) CO 2 (g) + H 2 O(g) More Examples An aqueous solution of potassium carbonate is mixed with an aqueous solution of iron(iii) nitrate resulting in a precipitate of iron(iii) carbonate. Write a balanced chemical equation for this reaction. Ammonia gas reacts with oxygen gas to form nitrogen monoxide gas and steam. Write a balancd chemical equation for this reaction. 5.2 Types of Reactions Chemical reactions are often classified into categories according to characteristics of the reactions. The following is a useful classification scheme: 2
Oxidation-reduction (redox) reactions involve one substance losing electrons (being oxidized) and another substance gaining electrons (being reduced). 5.3 Redox Reactions LEO say GER Lose Electrons Oxidation, Gain Electrons Reduction OIL RIG Oxidation Is Loss, Reduction Is Gain Oxidation Numbers Oxidation numbers (also called oxidation states) are positive or negative numbers assigned to elements in chemical formulas according to a specific set of rules Rule 1: The O.N. of an element is 0. Examples: Fe (0), Cl 2 (0), and Ca(0) Rule 2: The O.N. of a monatomic ion is equal to the charge of the ion. This includes when the monatomic ion is in a compound. Examples: Mg 2+ (+2), O 2- (-2), Cl - (-1), Na 2 S (Na = +1) and Ca(NO 3 ) 2 (Ca = +2) Oxidation Numbers (continued) Rule 3: The O.N. of hydrogen is +1. Examples: HBr (H = +1) and H 2 SO 4 (H = +1) Rule 4: The O.N. of oxygen is -2 except in peroxides it is -1. Examples: MgO & HBrO 3 (O = -2), and H 2 O 2 (O = -1) Rule 5: The algebraic sum of the oxidation numbers of all atoms in a compound or polyatomic ion equals the charge of the compound or ion. For the compound, the charge always equals zero. Example: MgSO 4 (O.N. of Mg = +2 by rule 2, O.N. of O = - 2 by rule 4, and O.N. of S = +6 by algebra and rule 5) 3
Oxidation Number Practice Determine the oxidation number for each atom in the following compounds or ions. CaCl 2 Na 3 PO 4 ClO 3 1- H 2 SO 3 Oxidizing and Reducing Agents Reducing agent the substance that is oxidized. It caused the other substance to be reduced. Oxidizing agent the substance that is reduced. It caused the other substance to be oxidized. Oxidizing and Reducing Agent Examples Determine the oxidizing and reducing agents as well as what is being reduced and what is being oxidized in the following reactions: 2Na(s) +2H 2 O(l) H 2 (g) + 2NaOH(aq) S 2 O 8 2- (aq) + 2 I - (aq) 4NO 2 (g) + O 2 (g) I 2 (aq) + 2SO 4 2- (aq) 2N 2 O 5 (g) Hint: First, assign oxidation numbers. 4
5.4 Decomposition Reactions In decomposition reactions, one substance is broken down into two or more simpler substances. Decomposition reactions may be either redox or nonredox reactions. The general form of the equation for a decomposition reaction is: A B + C. An example of a redox decomposition reaction is: 2HI(g) H 2 (g) + I 2 (g) An example of a nonredox decomposition reaction is: H 2 CO 3 (aq) CO 2 (g) + H 2 O(l) 5.5 Combination Reactions In combination reactions two or more substances react to form a single substance. Combination reactions may be either redox or nonredox reactions. The general form of the equation for a combination reaction (also known as addition or synthesis reactions) is: A + B C An example of a redox combination reaction is: 2Fe(s) + 3O 2 (g) 3Fe 2 O 3 (s) An example of a nonredox combination reaction is: N 2 O 5 (g) + H 2 O(l) 2HNO 3 (aq) 5.6 Replacement Reactions Single-replacement reactions one element replaces another in a compound. They are always redox reactions because an element turns into an ion and an ion turns into an element The general form of the equation for a single replacement reaction (also known as substitution reactions) is: A + BX B + AX In this equation, A and B represent elements and AX and BX are compounds. An example of a single replacement reaction is: Fe(s) + Cu(NO 3 ) 2 (aq) Cu(s) + Fe(NO 3 ) 2 (aq) 5
Double-Replacement Reactions Double-replacement reactions two ionic compounds switch partners. They are never redox reactions. The following general form of the equation for double replacement reactions (also known as metathesis reactions) shows the partner-swapping characteristic of the reactions: AX + BY BX + AY An example of a double-replacement reaction is: AgNO 3 (aq) + NaCl(aq) AgCl(s) + NaNO 3 (aq) Reaction Classification Classify each of the following equations as redox or nonredox reactions and as combination, decomposition, singlereplacement, or double-replacement reactions. NaOH(aq) + HCl(aq) NaCl(aq) + H 2 O(l) 2 Mg(s) + O 2 (g) 2 MgO(s) Na 2 CO 3 (s) Na 2 O(s) + CO 2 (g) Mg(s) + 2 HCl(aq) MgCl 2 (aq) + H 2 (g) 5.7 Ionic Equations Ionic compounds and some polar covalent compounds break apart (dissociate) when they dissolve in water and form ions. Reactions between these dissolved substances can be shown using molecular, total ionic, and net ionic equations. 6
Ionic Equations (Continued) Molecular Equations: Each compound is represented by its formula. NaCl (aq) + AgNO 3 (aq) AgCl (s) + NaNO 3 (aq) Total Ionic Equations: All soluble ionic substances are represented by the ions they form in solution. Substances represented by s, l, or g (do not dissolve or do not dissociate) are represented by their formulas. Na + (aq) + Cl - (aq) + Ag + (aq) + NO 3- (aq) AgCl(s) + Na + (aq) + NO 3- (aq) Net Ionic Equations: Shows only the change that is actually taking place. All spectator ions (those that are the same on both sides of the equation) are not included in the net ionic equation. Ag + (aq) + Cl - (aq) AgCl(s) Ionic Equations Practice An aqueous solution of barium chloride is mixed with an aqueous solution of sodium sulfide. Write balanced molecular, total ionic, and net ionic equations for this reaction. Write a total ionic and a net ionic equation for the following reaction. BaCO 3 (s) + 2HNO 3 (aq) Ba(NO 3 ) 2 (aq) + CO 2 (g) + H 2 O(l) 5.8 Energy and Reactions Exothermic reactions chemical reactions that release heat. Ordinary combustion of a log in a fireplace is an example of an exothermic reaction. Endothermic reactions chemical reactions that absorb heat. Example of an endothermic process is the melting of ordinary ice (physical not chemical change). As the ice melts, heat is absorbed from the air surrounding the ice. 7
5.9 The Mole and Chemical Equations Stoichiometry the study of mass relationships in chemical reactions. Stoichiometry allows us to calculate: The amount of a substance in a rxn given the amount of another substance in the rxn Must start with a balanced chemical equation! The following represents the process used in stoichiometry: grams A Molar mass moles A Equation coefficient moles B Molar mass grams B Stoichiometry Examples Calculate the number of moles of H 2 S that would need to react with excess O 2 in order to produce 115 g of SO 2. Water is the other product. First you will need to write a balanced chemical equation. Then follow the process as outlined previously: grams of A to moles of A to moles of B. Calculate the mass of SO 2 that would be produced if 56.2 grams of O 2 reacts with excess H 2 S. Calculate the mass of H 2 S that would be needed to react with 107 g of O 2. 5.10 The Limiting Reactant Limiting reactant the reactant that will be used up first, and thus determines the amount of product that can be produced. How to approach limiting reactant problems: Calculate the amount of product that could be produced by each of the reactants available. The reactant that gives the least amount of product is then the limiting reactant. 8
Limiting Reagent in Kitchen A BLT sandwich requires 2 slices of bread, 2 leaves of lettuce, 4 pieces of bacon, and 2 tomato slices. How many sandwiches could you make if you were given: 7 slices of bread 5 leaves of lettuce 14 slices of bacon 9 tomatoes Limiting Reactant Examples Calculate the maximum amount (in grams) of SO 2 that could be produced by reacting 55.2 g of O 2 with 50.8 g of H 2 S. If 23.4 g of copper(ii) nitrate reacts with 6.25 g of sodium hydroxide, then how much of the precipitate (solid) will form? 9
5.11 Reaction Yields A reaction will not always produce the calculated amount of product. The maximum amount of product that could be produced is called the theoretical yield. The amount of product actually produced is called the actual yield. These two quantities are used to calculate the percentage yield using the following equation: actual yield % yield = 100 theoretical yield Reaction Yields Example A 2.00g sample of magnesium reacts with 2.00g of pure oxygen gas to form 2.96g of magnesium oxide. What is the percent yield for this reaction? 10