Chapter 4 Chemical Reactions and Stoichiometry PDF

Chapter 4 Chemical Reactions and Stoichiometry 1  Learning Goals By the end of this module, students will be able to 1. Balance the chemical equations, determine the limiting reagent in a reaction, calculate the theoretical a...

Chapter 4 Chemical Reactions and Stoichiometry 1  Learning Goals By the end of this module, students will be able to 1. Balance the chemical equations, determine the limiting reagent in a reaction, calculate the theoretical and percent yields, and explain the concept of stoichiometry. 2. Determine solution concentrations and dilutions molarity, and understand the principles of solubility; 3. Recognize and balance different types of chemical reactions including: a) Precipitation: Predict solubility and write precipitation reactions b) Acid/base: Identify strong/weak acids and bases and write balanced neutralization reactions c) Gas-evolution: recognize and balance reactions that evolve gases d) Redox: Identify oxidations states and balance redox reaction in both acid and base solution.  Global Warming The average earth surface temperature has significantly increased since 1850. During the same period, atmospheric CO2 levels have risen 25%. Are the two trends causal? Global Average Surface Temperature www.climate.gov 3  The Greenhouse Effect The greenhouse gases in the atmosphere – allow sunlight to enter the atmosphere. – warm Earth’s surface. – prevent some of the – The balance between heat generated by the incoming and outgoing energy sunlight from escaping. from the sun determines Earth’s average temperature. Syukuro Manabe The Nobel Prize in Physics 2021 “for the physical modelling of Earth’s climate, quantifying variability and reliably predicting global warming” 4  The Sources and Quantity of Increase CO2 One source of CO2 is the combustion reactions of fossil fuels we use to get energy. C8H18 CO2 – How much CO2 is released? The balanced chemical equations for fossil-fuel combustion reactions provide the exact relationships between the amount of fossil fuel burned and the amount of carbon dioxide emitted. – 16 CO2 molecules are produced for every 2 molecules of octane burned. 5 Describing Reactions: Chemical Equations We use a balanced chemical equation to describe a chemical reaction. – Just like in math, there two sides and they are separated with an arrow 1. Reactants (on the left) 2. Products (on the right) – The number of atoms on each side must be the same as well as the overall charge For example: 2H2(g) + O2(g)  2H2O(l) 6 How to Write Balanced Chemical Equations 1. Write down a skeletal equation; 2. Count the atoms of each element, and overall charge on both sides of the equation; 3. Balance the numbers of atoms by placing coefficients in front of the compounds: (a) begin with atoms in more complex substances first; (b) lastly balance atoms in free elements; 4. Check the balance of each type of atom and overall charge on both sides. 5. May indicate the states of reactants and products. (g) Gas (l) Liquid (s) Solid (aq) Aqueous (water solution) 7  Q: Balance the following reaction: C3H8(g) + O2(g)  CO2(g) + H2O(g) Left: C (3); H(8); O(2)  Right: C (1), H (2), O (3) Begin with C: C3H8(g) + O2(g)  3 CO2(g) + H2O(g) Balance H: C3H8(g) + O2(g)  3 CO2(g) + 4 H2O(g) Balance O: Left: O(2)  Right: 3 X 2 + 4 X 1, O(10) C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(g) Check: Left: C(3); H(8); O (10); Right: C(3); H(8); O (10) 8  Quantities in Chemical Reactions The amount of every substance used and made in a chemical reaction is related to the amounts of all the other substances in the reaction (Mole-to-Mole Conversions). – Law of conservation of mass – Balancing equations by balancing atoms (Dalton’s atomic theory) and charges (conservation of # of electrons) The numerical relationship between chemical quantities in a chemical reaction is called stoichiometry. 9  Reaction Stoichiometry The coefficients in a chemical reaction specify the relative amounts (ratio) in moles of each of the substances involved in the reaction. 2 C8H18(l) + 25 O2(g) 16 CO2(g) + 18 H2O(g) – 2 molecules of C8H18 react with 25 molecules of O2 to form 16 molecules of CO2 and 18 molecules of H2O. NOTE: can replace “molecules” with “moles” and each statement is still correct! – 2 moles of C8H18 react with 25 moles of O2 to form 16 moles of CO2 and 18 moles of H2O. 1 mol C8H18 generates 8 mol CO2 10  Making Pizza The number of pizzas you can make depends on the amount of the ingredients you use. 1 crust + 5 oz tomato sauce + 2 cups cheese 1 pizza This relationship can be expressed mathematically. 1 crust : 5 oz sauce : 2 cups cheese : 1 pizza We can compare the maximum amount of pizza that can be made from 10 cups of cheese: – Since 2 cups cheese : 1 pizza, then, 11 Suppose That We Burn 22.0 Moles of C8H18; How Many Moles of CO2 Form? 2 C8H18(l) + 25 O2(g) 16 CO2(g) + 18 H2O(g) stoichiometric ratio: 2 moles C8H18 : 16 moles CO2 (Or, 1 mole C8H18 gives 8 moles CO2, thus 22.0 moles of C8H18 give 22 X 8 = 176 moles of CO2) The combustion of 22 moles of C8H18 adds 176 moles of CO2 to the atmosphere. 12  Limiting Reactant, Theoretical Yield Recall our pizza recipe: 1 crust + 5 oz tomato sauce + 2 cups cheese 1 pizza If we have 4 crusts, 10 cups of cheese, and 15 oz tomato sauce. How many pizzas can we make? We have enough crusts to make We have enough cheese to make We have enough tomato sauce to make 13  Limiting Reactant We have enough crusts for four pizzas, enough cheese for five pizzas, but enough tomato sauce for only three pizzas. – We can make only three pizzas. The tomato sauce limits how many pizzas we can make. 14  Theoretical Yield Tomato sauce is the limiting reactant, the reactant that makes the least amount of product. The limiting reactant is also known as the limiting reagent. The maximum number of pizzas we can make depends on this ingredient. In chemical reactions, we call this the theoretical yield. This is the amount of product that can be made in a chemical reaction based on the amount of limiting reactant. The ingredient that makes the least amount of pizza determines how many pizzas you can make (theoretical yield). 15  In a Chemical Reaction For reactions with multiple reactants, it is likely that one of the reactants will be completely used before the others. When this reactant is used up, the reaction stops and no more product is made. The reactant that limits the amount of product is called the limiting reactant. – It is sometimes called the limiting reagent. – The limiting reactant gets completely consumed. Reactants not completely consumed are called excess reactants. The amount of product that can be made from the limiting reactant is called the theoretical yield. 16  Actual Yield: More Making Pizzas Assume that while making pizzas, we may burn a pizza, drop one on the floor, or other uncontrollable events happen so that we only make two pizzas. The actual amount of product made in a chemical reaction is called the actual yield. We can determine the efficiency of making pizzas by calculating the percentage of the maximum number of pizzas we actually make. In chemical reactions, we call this the percent yield. 17 Apply These Concepts to a Chemical Reaction Consider the combustion of methane: CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) – Our balanced equation for the combustion of methane implies that every one molecule of CH4 reacts with two molecules of O2. 18 Practice: combustion of methane If we have five moles of CH4 and eight moles of O2, how much CO2 will be produced? CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) – Based on stoichiometry, one mole of CH4 needs two moles of O2 for combustion; – 5 moles of CH4 would needs 10 moles of O2 ; – Therefore, O2 is the limiting reactant. – 8 moles of O2 can burn only 4 moles of CH4, to produce 4 moles of CO2. 19 Practice: mass-to-mass conversions Consider the reaction of photosynthesis: 6CO2(g) + 6H2O(l)  6O2(g) + C6H12O6(aq) With 37.8 g of CO2 and excess of H2O, how many grams of C6H12O6 will be ideally produced? Plan: mA nB nA  S mB  nB M B MA nA 20  Solution Chemistry Many chemical reactions occur in solution – Increases chances of chemicals interacting and makes the reaction more efficient When table salt is mixed with water, it seems to disappear or become a liquid (dissolved), resulting a homogeneous mixture called solutions. The component of the solution that changes state is called the solute. The component that keeps its state is called the solvent. – If both components start in the same state, the major component is the solvent. 22  Solution Concentration: Molarity In solution reaction, solutes, not solvent often directly participate a chemical reaction and listed as reagents in chemical equation. Solutions can be described qualitatively, as dilute or concentrated; A common quantitative way to express solution concentration is molarity (M). – Molarity is the amount of solute (in moles) divided by the volume of solution (in liters). 23 Preparing 1 L of a 1.00 M NaCl Solution 24  Using Molarity in Calculations We can use the molarity of a solution as a conversion factor between moles of the solute and volumes (litres) of the solution. – How much NaCl in 1 L of a 0.500 M NaCl solution? – How many litres of a 0.500 M NaCl solution contains 0.25 mol of NaCl? 25  Solution Dilution Often, solutions are stored as concentrated stock solutions. To make solutions of lower concentrations from these stock solutions, more solvent is added. – The amount of solute doesn’t change, just the volume of solution: moles solute in solution 1 = moles solute in solution 2 The concentrations and volumes of the stock and new solutions are inversely proportional: M1∙V1 = M2∙V2 26  Solution Dilution n M V MV  n M 1V1  n  M 2V2 M 1V1  M 2V2 M 2V2 V1  M1 (0.500 M)(3.00 L) Vc  Preparing 3.00 L of 0.500 mol L−1 10.0 M CaCl2 from a 10.0 mol L−1 Stock Vc  0.150 L Solution. 27  Solution Stoichiometry Molarity can be used to convert between the amount of reactants and/or products in a chemical reaction. – What volume of B is required to react with a known volume of A given the concentrations of the solutions? nA  c AVA nB nB S VB  nA cB 28  What Happens When a Solute Dissolves? There are attractive forces between the solute particles, as well as between the solvent molecules. When we mix the solute with the solvent, there are attractive forces between the solute particles and the solvent molecules. If the attractions between solute and solvent are strong enough, the solute will dissolve. 29  Solute (NaCl) and Solvent (H2O) Interactions There is an uneven distribution of electrons within the water molecule. When NaCl is put into water, the attraction of Na+ and Cl– ions to water molecules (ion-dipole) competes with the attraction among the oppositely charged ions themselves. 30  Sodium Chloride Dissolving in Water Each ion is attracted to the surrounding water molecules and pulled off and away from the crystal. When it enters the solution, the ion is surrounded by water molecules, shielding it from other ions. The result is a solution with free moving charged particles able to conduct electricity. 31 Sugar Dissolved in Water 32  Electrolyte and Nonelectrolyte Solutions Electrolytes: materials that dissolve in water to form a solution that will conduct electricity. – Ionic substances, such as NaCl that completely dissociate into ions in H2O, are strong electrolytes. Nonelectrolytes: materials, mostly molecular compounds, that dissolve in water as intact neutral molecules, to form a solution that will not conduct electricity. A solution of salt (an electrolyte) conducts electrical current. A solution of sugar (a nonelectrolyte) does not. 33  Binary Acids (H-Y) Acids are molecular compounds, but they do ionize when dissolved in water. – When acids ionize, they form H+ cations and also anions. The percentage of molecules that ionize varies from one acid to another. Acids that ionize virtually 100% are called strong acids. HCl(aq) H+(aq) + Cl−(aq) Acids that only ionize a small percentage are called weak acids. HF(aq)  H+(aq) + F−(aq) 34  Strong and Weak Electrolytes Strong electrolytes are materials that dissolve completely as ions. – Ionic compounds and strong acids – Solutions conduct electricity well Weak electrolytes are materials that dissolve mostly as molecules, but partially as ions. – Weak acids – Solutions conduct electricity, but not well When compounds, such as acetic acid, containing a polyatomic ion dissolve, the polyatomic ion stays together. CH3CO2H (aq)  H+(aq) + CH3CO2−(aq) 35 Classes of Dissolved Materials 36  The Solubility of Ionic Compounds When an ionic compound dissolves in water, the ions are dissociated and solvated. However, not all ionic compounds dissolve in water. – If we add AgCl to water, it remains solid and appears as a white powder at the bottom of the water. In general, a compound is termed soluble if it dissolves in water and insoluble if it does not. 37  Solubility of Salts If we mix solid AgNO3 with water, it dissolves and forms a strong electrolyte solution. Silver chloride, on the other hand, is almost completely insoluble. – If we mix solid AgCl with water, virtually all of it remains as a solid within the liquid water. 38  When Will a Salt Dissolve? Whether a particular compound is soluble or insoluble depends on several factors. Predicting whether a compound will dissolve in water is not easy. The best way to do it is to do some experiments to test whether a compound will dissolve in water, and then develop some rules based on those experimental results. – We call this method the empirical method. 39  Solubility Rules Table 4.2 General Solubility Rules for Ionic Compounds in Water 1. All salts containing cations of group 1 metals (alkali metals, Li+, Na+, K+, etc.) and ammonium ions (NH4+) are soluble. 2. All nitrates (NO3−), ethanoates (acetates, CH3COO−), chlorates (ClO3−), and perchlorates (ClO4−) are soluble. 3. Salts containing Ag+, Pb2+, and Hg22+ are insoluble. 4. Most chlorides (Cl−), bromides (Br−), and iodides (I−) are soluble. 5. Sulfates (SO42−) are soluble, except those containing Ca2+, Sr2+, Ba2+. 6. Carbonates (CO32−), hydroxides (OH−), oxides (O2−), phosphates (PO43−), and sulfides (S2−) are generally insoluble. Empirical rules with exceptions! 40  Precipitation Reactions Precipitation reactions are reactions in which a solid forms when we mix two solutions. – Reactions between aqueous solutions of ionic compounds produce an ionic compound that is insoluble in water. – The insoluble product is called a precipitate. Predicting precipitation reaction? Original Salts KI Pb(NO3)2 41 Precipitation of Lead(II) Iodide 42  Predicting Precipitation Reactions 1. Determine what ions each aqueous reactant has. 2. Determine formulas of all possible products. – Exchange ions. (+) ion from one reactant with (–) ion from other – Balance charges of combined ions to get the formula of each product. 3. Determine solubility of each product in water. – Use the solubility rules. – If product is insoluble or slightly soluble, it will precipitate. 4. If neither product will precipitate, write no reaction after the arrow. 43  Predicting Precipitation Reactions 5. If any of the possible products are insoluble, write their formulas as the products of the reaction using (s) after the formula to indicate solid. Write any soluble products with (aq) after the formula to indicate aqueous. 6. Balance the equation. – Remember to only change coefficients, not subscripts. 44  Representing Aqueous Reactions An equation showing the complete neutral formulas for each compound in the aqueous reaction as if they existed as molecules is called a molecular equation. Pb(NO3 ) 2 (aq )  2 KI(aq)  PbI2 ( s )  KNO3 (aq ) In actual solutions of soluble ionic compounds, dissolved substances are present as ions. Equations that describe the material’s structure when dissolved are called complete ionic equations. 45  Complete Ionic Equation Rules of writing the complete ionic equation: – Aqueous strong electrolytes are written as ions. Soluble salts, strong acids, strong bases – Insoluble substances, weak electrolytes, and nonelectrolytes are written in molecule form. Solids, liquids, and gases are not dissolved, hence molecule form Pb 2 (aq)  2 NO3– (aq)  2 K  (aq )  2 I – ( aq )  PbI2 ( s )  2 NO3– (aq)  2 K  (aq) Some of the ions that appear unchanged on both sides of the equation are called spectator ions, because they do not participate in the reaction. 46  Net Ionic Equation Pb 2 (aq)  2 NO3– (aq)  2 K  (aq )  2 I – (aq )  PbI 2 ( s )  2 NO3– (aq )  2 K  (aq )  An ionic equation in which the spectator ions are removed is called a net ionic equation. Pb 2 (aq)  2 I  (aq )  PbI 2 ( s ) 47  Q: Write the ionic and net ionic equation for the following reaction. Molecular equation: 2 KOH(aq) + Mg(NO3)2(aq) 2 KNO3(aq) + Mg(OH)2(s) Complete ionic equation: 2 K+(aq) + 2 OH−(aq) + Mg2+(aq) + 2 NO3−(aq) K+(aq) + 2 NO3−(aq) + Mg(OH)2(s) 2 K+(aq) + 2 OH−(aq) + Mg2+(aq) + 2 NO3−(aq) K+(aq) + 2 NO3−(aq) + Mg(OH)2(s) Net ionic equation: 2 OH−(aq) + Mg2+(aq) Mg(OH)2(s) 48  Acid–Base Reactions Arrhenius Definitions: Acid: Substance that produces H+ in aqueous solution HCl(aq) H+(aq) + Cl–(aq) – Some acids—called polyprotic acids These acids contain more than one ionizable proton and release them sequentially. For example, sulfuric acid, H2SO4 is a diprotic acid. Strong in its first ionizable proton, but weak in its second. Base: Substance that produces OH– ions in aqueous solution NaOH(aq) Na+(aq) + OH–(aq) 49  Acid–Base Reactions Also called neutralization reactions because the acid and base neutralize each other’s properties: 2 HNO3(aq) + Ca(OH)2(aq) Ca(NO3)2(aq) + 2 H2O(l) The net ionic equation for an acid–base reaction is: H+(aq) + OH(aq) H2O(l) – as long as the salt that forms is soluble in water. 50  Acids in Solution: Hydronium Ion Acids ionize in water to form H+ ions. – More precisely, the H from the acid molecule is donated to a water molecule to form hydronium ion, H3O+. Chemists may use H+ and H3O+ interchangeably. 51  Acids and Bases in Solution Bases dissociate in water to form OH ions. – Bases, such as NH3, that do not contain OH ions, produce OH by pulling H off water molecules. In the reaction of an acid with a base, the H+ from the acid combines with the OH from the base to make water. The cation from the base combines with the anion from the acid to make the salt. 52  Acid–Base Reaction HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) 53  Some Common Acids and Bases Table 4.3 Common Acids and Bases Name of Acid Formula Name of Base Formula Hydrochloric acid HCl Sodium hydroxide NaOH Hydrobromic acid HBr Lithium hydroxide LiOH Hydroiodic acid HI Potassium hydroxide KOH Nitric acid HNO3 Calcium hydroxide Ca(OH)2 Sulfuric acid H2SO4 Barium hydroxide Ba(OH)2 Perchloric acid HClO4 Ammonia* NH3 (weak base) Acetic acid CH3COOH (weak acid) Blank Blank Hydrofluoric acid HF (weak acid) Blank Blank *Ammonia does not contain OH−, but it produces OH− in a reaction with water that occurs only to a small extent: NH3 (aq) + H2O (l) NH4+ (aq) + HO– (aq) 54 Q: Predict the Product of the Reactions 1. HCl(aq) + Ba(OH)2(aq) (H+ + Cl−) + (Ba2+ + OH−) → (H+ + OH−) + (Ba2+ + Cl−) HCl(aq) + Ba(OH)2(aq) → H2O(l) + BaCl2 2 HCl(aq) + Ba(OH)2(aq) 2 H2O(l) + BaCl2(aq) 2. H2SO4(aq) + LiOH(aq) (H+ + SO42−) + (Li+ + OH−) → (H+ + OH−) + (Li+ + SO42−) H2SO4(aq) + LiOH(aq) → H2O(l) + Li2SO4 H2SO4(aq) + 2 LiOH(aq) 2 H2O(l) + Li2SO4(aq) 55  Gas-Evolving Reactions Some reactions form a gas directly from the ion exchange. K2S(aq) + H2SO4(aq) K2SO4(aq) + H2S(g) Other reactions form a gas by the decomposition of one of the ion exchange products into a gas and water. NaHCO3(aq) + HCl(aq) NaCl(aq) + H2CO3(aq) H2CO3(aq) H2O(l) + CO2(g) 56 Gas-Evolution Reaction 57  Types of Compounds That Undergo Gas- Evolution Reactions Table 4.4 Reactant Intermediate Gas Type Product Evolved Example Sulfides None H2S 2 HCl(aq) + K2S(aq) → H2S(g) + 2 KCl(aq) Carbonates and bicarbonates H2CO3 CO2 2 HCl(aq) + K2CO3(aq) → H2O(l) + CO2(g) + 2 KCl(aq) Sulfites and bisulfites H2SO3 SO2 2 HCl(aq) + K2SO3(aq) → H2O(l) + SO2(g) + 2 KCl(aq) Ammonium NH4OH NH3 NH4Cl(aq) + KOH(aq) → H2O(l) + NH3(g) + KCl(aq) 58  Oxidation–Reduction Reactions The reactions in which electrons are transferred from one reactant to the other are called oxidation-reduction reactions. – These are also called redox reactions. – Many redox reactions involve the reaction of a substance with oxygen, but not always. – 4 Fe(s) + 3 O2(g) 2 Fe2O3(s) (rusting) – 2 C8H18(l) + 25 O2(g) 16 CO2(g) + 18 H2O(g) (combustion) – 2 H2(g) + O2(g) 2 H2O(g) (combustion) 59  Combustion as Redox 2 H2(g) + O2(g) 2 H2O(g) Insert Figure 4.22 on Pg. 176 60  Redox without O2 2 Na(s) + Cl2(g) 2 NaCl(s) Insert Figure 4.24 on Pg. 177 61  Reactions of Metals with Nonmetals Consider the following reactions: 4 Na(s) + O2(g) → 2 Na2O(s) 2 Na(s) + Cl2(g) → 2 NaCl(s) The reactions involve a metal reacting with a nonmetal. In addition, both reactions involve the conversion of free elements into ions. 4 Na(s) + O2(g) → 2 Na+2O2– (s) 2 Na(s) + Cl2(g) → 2 Na+Cl–(s) 62  Oxidation and Reduction To convert a free element into an ion, the atoms must gain or lose electrons. Reactions where electrons are transferred from one atom to another are redox reactions. Oxidation is the process of a substance losing electron(s); Reduction is the process of a substance gaining electron(s): OIL RIG: Oxidation Is Loss; Reduction Is Gain 63  Redox Reaction The transfer of electrons does not need to be a complete transfer (as occurs in the formation of an ionic compound) for the reaction to qualify as oxidation–reduction. – For example, consider the reaction between hydrogen gas and chlorine gas: H2(g) + Cl2(g) 2 HCl(g) When hydrogen bonds to chlorine, the electrons are unevenly shared, resulting in an increase of electron density (reduction) for chlorine and a decrease in electron density (oxidation) for hydrogen. 64  Oxidation States An oxidation state is an imaginary “charge” number assigned to each element in a compound to help us determine the electron flow in the reaction. – Based on the assumption that electrons were completely transferred to the more electronegative atom. – Oxidation states are not ion charges! Ion charges are real, measurable charges. The terms “oxidation state” and “oxidation number” mean the same thing in literature. The difference in electronegativity (EN) between atoms determines the direction of electron transfer, which in turn determine the oxidation state. 65  Rules for Assigning Oxidation States The following are general rules: 1. Free elements have an oxidation state = 0. – Na = 0 and Cl2 = 0 in 2 Na(s) + Cl2(g) 2. Monatomic ions have an oxidation state equal to their charge. – Na = +1 and Cl = −1 in NaCl 3. The sum of the oxidation states of all atoms in a neutral molecule is always zero – Na = +1 and Cl = −1 in NaCl, (+1) + (−1) = 0 a polyatomic ion is equal to the charge of the ion – N = +5 and O = −2 in NO3–, (+5) + 3(−2) = −1 66  Rules for Assigning Oxidation States Typical oxidation states of specific elements: 4. Hydrogen (H): Typically +1 when bonded to nonmetals (except B), -1 when bonded to metals. – H : +1 in H2O; -1 in NaH and BH3. 5. Oxygen (O): Usually -2, except in peroxides (-1) and in compounds with fluorine. – O : -2 in H2O; -1 in H2O2; +2 in OF 6. Fluorine (F): Always -1 in compounds. – F : -1 in CF4 7. (a) Alkali metals (Group 1): +1 in compounds (b) Alkaline earth metals (Group 2): +2 in compounds – Na: +1 in NaCl; Mg: +2 in MgCl2 67  Rules for Assigning Oxidation States 8. Identify the known oxidation states using the rules above for specific atoms, then use the overall charge of the compound to deduce the unknown oxidation states. 9. When necessary, use electronegativity (EN) to determines how the electrons are transferred. 68  Identifying Redox Reactions Oxidation: An increase in oxidation state Reduction: A decrease in oxidation state – Carbon changes from an oxidation state of 0 to an oxidation state of +4. Carbon loses electrons and is oxidized. – Sulfur changes from an oxidation state of 0 to an oxidation state of –2. Sulfur gains electrons and is reduced. 69  Redox Reactions Oxidation and reduction must occur simultaneously. – If an atom loses electrons another atom must take them. The reactant that reduces an element in another reactant is called the reducing agent. – The reducing agent contains the element that is oxidized. The reactant that oxidizes an element in another reactant is called the oxidizing agent. – The oxidizing agent contains the element that is reduced. 2 Na(s) + Cl2(g) → 2 Na+Cl–(s) Na is oxidized, while Cl is reduced. Na is the reducing agent, and Cl2 is the oxidizing agent. 70  Balancing Redox Reactions Balancing redox reactions requires you to balance both the atoms and the charge of both sides of the reaction. Redox reactions in aqueous solutions can be balanced by half-reaction method of balancing. The overall equation is split into two half-reactions: oxidation and reduction. Both are then balanced and added together. Redox reactions occur in either acidic or basic solutions – If the solution is acidic, you will need to add H+ and H2O at some point during the balancing process. – If the solution is basic, you will need to add OH- at some point during the balancing process. 71  Half-Reaction Method of Balancing 1. Assign oxidation states to all atoms, identify which elements are oxidized and reduced. 2. Separate the overall reaction into two half- reactions. 3. Balance each half-reaction with respect to mass First balance all elements other than H and O Balance O by adding H2O where needed Balance H by adding H+ where needed If reaction done in base, neutralize H+ with OH−. 4. Balance each half-reaction with respect to charge by adding electrons 72  Half-Reaction Method of Balancing 5. Make the number of electrons in each half-reaction equal by multiplying one or both reactions by a small whole number. 6. Add the two half reactions, cancelling electrons and other species as necessary. 7. Verify that the reaction is balanced with respect to both charge and mass. 73  Example: Balancing Redox in Acidic Solution Balance this redox reaction in acidic solution: Fe2+ + MnO4– → Fe3+ + Mn2+ 1.Assign oxidation states and determine what is oxidized and what is reduced: – Left side: Fe2+ is +2 – Right side: Fe3+ is +3 (oxidized) – Left side: Mn is +7 in MnO4- – Left side: Oxygen is -2 in MnO4- – Right side: Mn2+ is 2+ (reduced). 2.Separate into oxidation & reduction half-reactions: – ox: Fe2+ → Fe3+ (Fe2+ is the reducing agent) – red: MnO4– → Mn2+ (MnO4- is the oxidizing agent) 74  Example: Balancing Redox in Acidic Solution 3. Balance half-reactions by mass: First balance all elements other than H and O Balance O by adding H2O where needed Balance H by adding H+ where needed – Oxidation Half Reaction: Fe2+ → Fe3+ – Reduction Half Reaction: MnO4– → Mn2+ MnO4– → Mn2+ + 4H2O MnO4– + 8H+ → Mn2+ + 4H2O 75  Example: Balancing Redox in Acidic Solution 4. Balance each half-reaction with respect to charge by adding electrons (es) – # of electrons on product side for oxidation Fe2+ → Fe3+ + 1 e− – # of electrons on reactant side for reduction: MnO4– + 8H+ → Mn2+ + 4H2O LS: (-1) +(8) = +7 RS: +2 MnO4– + 8H+ + 5 e− → Mn2+ + 4H2O 76  Example: Balancing Redox in Acidic Solution 5. Balance es Fe2+ → Fe3+ + 1 e− } x 5; (since the other between half- half reaction involves 5 e-) reactions. 6. Add half- MnO4– + 8H+ + 5 e− → Mn2+ + 4H2O reactions, 5Fe2+ → 5 Fe3+ + 5 e− canceling es and common 5 Fe2+ + MnO – + 8H+ → Mn2+ + 4H O + 5 Fe3+ 4 2 species. reactant side Element product side 7. Check that 5 Fe 5 numbers of 1 Mn 1 atoms and total 4 O 4 charge are equal. 8 H 8 +17 charge +17 77  Balancing Redox in Basic Solution Balance this redox reaction in basic solution: Fe2+ + MnO4– → Fe3+ + Mn2+ When balancing redox reactions in basic solution, use the exact same steps for acidic solution except you will have to neutralize all H+ ions by adding OH- to both sides, at the end of the process (1 extra step only!) Since we already did this in acidic solution: 5 Fe2+ + MnO4– + 8H+ → Mn2+ + 4H2O + 5 Fe3+ So add 8 eq. of OH- on both sides to neutralize the 8H+: 5 Fe2+ + MnO4– + 8H+ + 8OH- → Mn2+ + 4H2O + 5 Fe3+ + 8OH- 5 Fe2+ + MnO4– + 8H2O → Mn2+ + 4H2O + 5 Fe3+ + 8OH- 5 Fe2+ + MnO4– + 4H2O → Mn2+ + 5 Fe3+ + 8OH- Check to make sure it is balanced (charge and atoms!) 78  Balancing Disproportionation Reactions Disproportionation is a special type of redox reaction in which a reactant is both oxidized and reduced. – The same species is the oxidant and the reductant ClO-(aq)  Cl-(aq) + ClO3-(aq) Balancing procedure is similar to balancing normal redox reactions. Two half reactions are identified that both start from the same starting compound that disproportionates. Balance as though in acidic solution. Add OH− to neutralize H+ if in basic solution in the same way as normal redox reactions. 79