Oxidation and Reduction

Introductory Chemistry, 3rd EditionNivaldo Tro : Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA Introductory Chemistry, 3rd EditionNivaldo Tro Chapter 16 Oxidation and Reduction 2009, Prentice Hall

Oxidation–Reduction Reactions : Tro's Introductory Chemistry, Chapter 16 2 Oxidation–Reduction Reactions Oxidation–reduction reactions are also called redox reactions. All redox reactions involve the transfer of electrons from one atom to another. Spontaneous redox reactions are generally exothermic, and we can use their released energy as a source of energy for other applications. Convert the heat of combustion into mechanical energy to move our cars. Use electrical energy in a car battery to start our car engine.

Combustion Reactions : Tro's Introductory Chemistry, Chapter 16 3 Combustion Reactions Combustion reactions are always exothermic. In combustion reactions, O2 combines with all the elements in another reactant to make the products. 4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) + energy CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) + energy

Reverse of Combustion Reactions : Tro's Introductory Chemistry, Chapter 16 4 Reverse of Combustion Reactions Since combustion reactions are exothermic, their reverse reactions are endothermic. The reverse of a combustion reaction involves the production of O2. energy + 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g) energy + CO2(g) + 2 H2O(g) → CH4(g) + 2 O2(g) Reactions in which O2 is gained or lost are redox reactions.

Oxidation and Reduction:One Definition : Tro's Introductory Chemistry, Chapter 16 5 Oxidation and Reduction:One Definition When an element attaches to an oxygen during the course of a reaction it is generally being oxidized. In CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g), C is being oxidized in this reaction, but H is not. When an element loses an attachment to oxygen during the course of a reaction, it is generally being reduced. In 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g), the Fe is being reduced. One definition of redox is the gain or loss of O, but it is not the best.

Another Oxidation–Reduction : Tro's Introductory Chemistry, Chapter 16 6 Another Oxidation–Reduction Consider the following reactions: 4 Na(s) + O2(g) → 2 Na2O(s) 2 Na(s) + Cl2(g) → 2 NaCl(s) The reaction involves 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+2O–(s) 2 Na(s) + Cl2(g) → 2 Na+Cl–(s)

Oxidation and Reduction:Another Definition : Tro's Introductory Chemistry, Chapter 16 7 Oxidation and Reduction:Another Definition In order to convert a free element into an ion, the atoms must gain or lose electrons. Of course, if one atom loses electrons, another must accept them. Reactions where electrons are transferred from one atom to another are redox reactions. Atoms that lose electrons are being oxidized, atoms that gain electrons are being reduced. 2 Na(s) + Cl2(g) → 2 Na+Cl–(s) Na → Na+ + 1 e– (oxidation) Cl2 + 2 e– → 2 Cl– (reduction)

Practice—Identify the Element Being Oxidized and the Element Being Reduced. : Tro's Introductory Chemistry, Chapter 16 8 Practice—Identify the Element Being Oxidized and the Element Being Reduced. 2 C(s) + O2(g) → 2 CO(g) Mg(s) + Cl2(g) → MgCl2(s) Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)

Practice—Identify the Element Being Oxidized and the Element Being Reduced, Continued. : Tro's Introductory Chemistry, Chapter 16 9 Practice—Identify the Element Being Oxidized and the Element Being Reduced, Continued. 2 C(s) + O2(g) → 2 CO(g) Mg(s) + Cl2(g) → MgCl2(s) Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s) C is oxidized because it is gaining an attachment to O. O is reduced; there has to be reduction and it’s the only other element. Mg is oxidized because it is becoming a cation by losing electrons. Cl is reduced because it is becoming an anion by gaining electrons. 0 0 2+ − Mg is oxidized because it is becoming a cation by losing electrons. Fe2+ is reduced because it is gaining electrons to become neutral.

Oxidation–Reduction : Tro's Introductory Chemistry, Chapter 16 10 Oxidation–Reduction 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, Cl is reduced. Na is the reducing agent, Cl2 is the oxidizing agent.

Practice—Identify the Oxidizing and Reducing Agents. : Tro's Introductory Chemistry, Chapter 16 11 Practice—Identify the Oxidizing and Reducing Agents. 2 C(s) + O2(g) → 2 CO(g) Mg(s) + Cl2(g) → MgCl2(s) Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s) C is oxidized because it is gaining attachment to O. O is reduced; there has to be reduction and it’s the only other element. Mg is oxidized because it is becoming a cation by losing electrons. Cl is reduced because it is becoming an anion by gaining electrons. 0 0 2+ − Mg is oxidized because it is becoming a cation by losing electrons. Fe2+ is reduced because it is gaining electrons to become neutral.

Practice—Identify the Oxidizing and Reducing Agents, Continued. : Tro's Introductory Chemistry, Chapter 16 12 Practice—Identify the Oxidizing and Reducing Agents, Continued. 2 C(s) + O2(g) → 2 CO(g) Mg(s) + Cl2(g) → MgCl2(s) Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s) C is the reducing agent because it contains the element that is oxidized. O is the oxidizing agent because it contains the element that is reduced. 0 0 2+ − Mg is the reducing agent because it contains the element that is oxidized. Cl2 is the oxidizing agent because it contains the element that is reduced. Mg is the reducing agent because it contains the element that is oxidized. Fe2+ is the oxidizing agent because it contains the element that is reduced.

Electron Bookkeeping : Tro's Introductory Chemistry, Chapter 16 13 Electron Bookkeeping For reactions that are not metal + nonmetal, or do not involve O2, we need a method for determining how the electrons are transferred. Chemists assign a number to each element in a reaction called an oxidation state that allows them to determine the electron flow in the reaction. Although they look like them, oxidation states are not ion charges! Oxidation states are imaginary charges assigned based on a set of rules. Ion charges are real, measurable charges.

Rules for Assigning Oxidation States : Tro's Introductory Chemistry, Chapter 16 14 Rules for Assigning Oxidation States Rules are in order of priority. Free elements have an oxidation state = 0. Na(s) = 0 and Cl2(g) = 0 in 2 Na(s) + Cl2(g) 2 NaCl(s). Monoatomic ions have an oxidation state equal to their charge. Na = +1 and Cl = -1 in NaCl(s). a. The sum of the oxidation states of all the atoms or ions in a compound is 0. Na = +1 and Cl = -1 in NaCl, and (+1) + (-1) = 0.

Rules for Assigning Oxidation States, Continued : Tro's Introductory Chemistry, Chapter 16 15 Rules for Assigning Oxidation States, Continued b. The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on the ion. N = +5 and O = -2 in NO3–, (+5) + 3(-2) = -1. a. Group I metals have an oxidation state of +1 in all their compounds. Na = +1 in NaCl. b. Group II metals have an oxidation state of +2 in all their compounds. Mg = +2 in MgCl2.

Rules for Assigning Oxidation States, Continued : Tro's Introductory Chemistry, Chapter 16 16 Rules for Assigning Oxidation States, Continued In their compounds, nonmetals have oxidation states according to the table below. Nonmetals higher on the table take priority.

Practice—Assign an Oxidation State to Each Element in the Following: : Tro's Introductory Chemistry, Chapter 16 17 Practice—Assign an Oxidation State to Each Element in the Following: F2 Mg2+ KCl SO2 PO43− BaO2

Practice—Assign an Oxidation State to Each Element in the Following, Continued: : Tro's Introductory Chemistry, Chapter 16 18 Practice—Assign an Oxidation State to Each Element in the Following, Continued: F2 F = 0 (Rule 1) Mg2+ Mg = +2 (Rule 2) KCl K = +1 (Rule 4a) and Cl = -1 (Rule 5) SO2 O = -2 (Rule 5) and S = +4 (Rule 3a) PO43− O = -2 (Rule 5) and P = +5 (Rule 3b) BaO Ba = +2 (Rule 4b) and O = -2

Oxidation and Reduction:A Better Definition : Tro's Introductory Chemistry, Chapter 16 19 Oxidation and Reduction:A Better Definition Oxidation occurs when an atom’s oxidation state increases during a reaction. Reduction occurs when an atom’s oxidation state decreases during a reaction. CH4 + 2 O2 → CO2 + 2 H2O -4 +1 0 +4 –2 +1 -2

Practice—Assign Oxidation States and Identify the Oxidizing and Reducing Agents in Each of the Following: : Tro's Introductory Chemistry, Chapter 16 20 Practice—Assign Oxidation States and Identify the Oxidizing and Reducing Agents in Each of the Following: 3 H2S + 2 NO3– + 2 H+ ® 3 S + 2 NO + 4 H2O MnO2 + 4 HBr ® MnBr2 + Br2 + 2 H2O

Practice—Assign Oxidation States and Identify the Oxidizing and Reducing Agents in Each of the Following, Continued: : Tro's Introductory Chemistry, Chapter 16 21 3 H2S + 2 NO3– + 2 H+ ® 3 S + 2 NO + 4 H2O MnO2 + 4 HBr ® MnBr2 + Br2 + 2 H2O +1 -2 +5 -2 +1 0 +2 -2 +1 -2 oxidizing agent reducing agent +4 -2 +1 -1 +2 -1 0 +1 -2 reducing agent Oxidizing agent Practice—Assign Oxidation States and Identify the Oxidizing and Reducing Agents in Each of the Following, Continued:

Balancing Redox Reactions : Tro's Introductory Chemistry, Chapter 16 22 Balancing Redox Reactions Some redox reactions can be balanced by the method we previously used, but many are hard to balance using that method. Many are written as net ionic equations. Many have elements in multiple compounds. The main principle is that electrons are transferred, so if we can find a method to keep track of the electrons, it will allow us to balance the equation.

Balancing Redox Reactions by the Half-Reaction Method : Tro's Introductory Chemistry, Chapter 16 23 Balancing Redox Reactions by the Half-Reaction Method In this method, the reaction is broken down into two half-reactions, one for oxidation and another for reduction. Each half-reaction includes electrons. Electrons go on the product side of the oxidation half-reaction—loss of electrons. Electrons go on the reactant side of the reduction half-reaction—gain of electrons. Each half-reaction is balanced for its atoms. Then the two half-reactions are adjusted so that the electrons lost and gained will be equal when added.

Slide 24 : Tro's Introductory Chemistry, Chapter 16 24 Assign oxidation states and determine element oxidized and element reduced. Separate into oxidation and reduction half-reactions. Balancing Redox Reactions in Acidic Solution

Slide 25 : 25 3. Balance half-reactions by mass. First balance atoms other than O and H. Balance O by adding H2O to side that lacks O. Balance H by adding H+ to side that lacks H. Finished if in acidic solution. If in basic solution, add enough OH− to neutralize the H+, rewrite H+ + OH− as H2O. Add to both sides. Then cancel H2O on both sides. Fe2+ → Fe3+ Balancing Redox Reactions in Acidic Solution, Continued Tro's Introductory Chemistry, Chapter 16

Slide 26 : Tro's Introductory Chemistry, Chapter 16 26 4. Balance each half-reaction, with respect to charge, by adjusting the numbers of electrons. Electrons on product side for oxidation. Electrons on reactant side for reduction. 5. Balance electrons between half-reactions. 6. Add half-reactions, canceling electrons and common species. 7. Check. Fe2+ → Fe3+ + 1 e- MnO4– + 8H+ + 5 e- → Mn2+ + 4H2O } x 5 5 Fe2+ + MnO4– + 8H+ → Mn2+ + 4H2O + 5 Fe3+ Balancing Redox Reactions in Acidic Solution, Continued

Practice—Balance the Following Equation:Cu+ + I2 → Cu2+ + I– : Tro's Introductory Chemistry, Chapter 16 27 Practice—Balance the Following Equation:Cu+ + I2 → Cu2+ + I–

Practice—Balance the Following Equation, Continued:Cu+ + I2 → Cu2+ + I– : Tro's Introductory Chemistry, Chapter 16 28 Practice—Balance the Following Equation, Continued:Cu+ + I2 → Cu2+ + I– +1 0 +2 -1 oxid red oxid: Cu+ → Cu2+ red: I2 → I– oxid: Cu+ → Cu2+ red: I2 → 2 I– oxid: Cu+ → Cu2+ + 1 e- red: I2 + 2 e- → 2 I– oxid: Cu+ → Cu2+ + 1 e- } x 2 red: I2 + 2 e- → 2 I– 2 Cu+ + I2 → 2 Cu2+ + I2

Practice—Balance the Following Equationin Acidic Solution: I– + Cr2O72- → Cr3+ + I2 : Tro's Introductory Chemistry, Chapter 16 29 Practice—Balance the Following Equationin Acidic Solution: I– + Cr2O72- → Cr3+ + I2

Practice—Balancing Redox Reactions : Tro's Introductory Chemistry, Chapter 16 30 Practice—Balancing Redox Reactions Assign oxidation states and determine element oxidized and element reduced. Separate into oxidation and reduction half-reactions. oxid: I− → I2 red: Cr2O72– → Cr3+

Practice—Balancing Redox Reactions, Continued : Tro's Introductory Chemistry, Chapter 16 31 Practice—Balancing Redox Reactions, Continued 3. Balance half-reactions by mass. First balance atoms other than O and H. Balance O by adding H2O to side that lacks O. Balance H by adding H+ to side that lacks H. Finished if in acidic solution. If in basic solution, add enough OH− to neutralize the H+, rewrite H+ + OH− as H2O. Add to both sides. Then cancel H2O on both sides. oxid: I− → I2 oxid: 2 I− → I2 red: Cr2O72– → Cr3+ red: Cr2O72– → 2 Cr3+ red: Cr2O72– → 2Cr3+ +7H2O Cr2O72– +14H+ → 2Cr3+ +7H2O

Practice—Balancing Redox Reactions, Continued : Cr2O72– +14H+ + 6e−→ 2Cr3+ +7H2O Tro's Introductory Chemistry, Chapter 16 32 Practice—Balancing Redox Reactions, Continued 4. Balance each half-reaction, with respect to charge, by adjusting the numbers of electrons. Electrons on product side for oxidation. Electrons on reactant side for reduction. 5. Balance electrons between half-reactions. 6. Add half-reactions, canceling electrons and common species. 7. Check. } x 3 2 I− → I2 2 I− → I2 + 2e− Cr2O72– + 14H+ → 2Cr3+ + 7H2O Cr2O72– +14H+ + 6e−→ 2Cr3+ +7H2O 2 I− → I2 + 2e− 6 I− → 3 I2 + 6e− Cr2O72– +14H+ + 6 I−→ 2Cr3+ +7H2O + 3 I2

Balancing Redox Reactions in Basic Solution : Tro's Introductory Chemistry, Chapter 16 33 Balancing Redox Reactions in Basic Solution Assign oxidation states and determine element oxidized and element reduced. Separate into oxidation and reduction half-reactions.

Slide 34 : MnO4– + 4H+ + 4OH− → MnO2 + 2H2O + 8OH− MnO4– + 4H2O → MnO2 + 2H2O + 8OH− 3. Balance half-reactions by mass. First balance atoms other than O and H. Balance O by adding H2O to side that lacks O. Balance H by adding H+ to side that lacks H. Finished if in acidic solution. If in basic solution, add enough OH− to neutralize the H+, rewrite H+ + OH− as H2O. Add to both sides. Then cancel H2O on both sides. CN− → CNO− CN− + H2O → CNO− CN− + H2O → CNO− + 2 H+ MnO4– → MnO2 MnO4– → MnO2 + 2H2O MnO4– + 4H+ → MnO2 + 2H2O CN− + H2O + 2OH− → CNO− + 2H+ + 2OH− CN− + H2O + 2OH− → CNO− + 2H2O CN− + 2OH− → CNO− + H2O MnO4– + 2H2O → MnO2 + 8OH− Balancing Redox Reactions in Basic Solution, Continued

Slide 35 : 3CN− + 6OH− → 3CNO− + 3H2O + 6e− Tro's Introductory Chemistry, Chapter 16 35 4. Balance each half-reaction with respect to charge by adjusting the numbers of electrons. Electrons on product side for oxidation. Electrons on reactant side for reduction. 5. Balance electrons between half-reactions. 6. Add half-reactions, canceling electrons and common species. 7. Check. MnO4– + 2H2O → MnO2 + 4OH− } x 3 3CN– + 2MnO4– + H2O → 3CNO– + 2MnO2 + 2OH– CN− + 2OH− → CNO− + H2O CN− + 2OH− → CNO− + H2O + 2e− MnO4– + 2H2O + 3e− → MnO2 + 4OH− CN− + 2OH− → CNO− + H2O + 2e− MnO4– + 2H2O + 3e− → MnO2 + 4OH− } x 2 2MnO4– + 4H2O + 6e− → 2MnO2 + 8OH− Balancing Redox Reactions in Basic Solution, Continued

Practice—Balance the Following Equationin Basic Solution: I– + Cr2O72- → Cr3+ + I2 : Tro's Introductory Chemistry, Chapter 16 36 Practice—Balance the Following Equationin Basic Solution: I– + Cr2O72- → Cr3+ + I2

Balancing Redox Reactions : Tro's Introductory Chemistry, Chapter 16 37 Balancing Redox Reactions Assign oxidation states and determine element oxidized and element reduced. Separate into oxidation and reduction half-reactions. oxid: I− → I2 red: Cr2O72– → Cr3+

Balancing Redox Reactions,Continued : 38 Balancing Redox Reactions,Continued 3. Balance half-reactions by mass. First balance atoms other than O and H. Balance O by adding H2O to side that lacks O. Balance H by adding H+ to side that lacks H. Finished if in acidic solution. If in basic solution, add enough OH− to neutralize the H+, rewrite H+ + OH− as H2O. Add to both sides. Then cancel H2O on both sides. oxid: I− → I2 oxid: 2 I− → I2 red: Cr2O72– → Cr3+ red: Cr2O72– → 2 Cr3+ red: Cr2O72– → 2Cr3+ +7H2O Cr2O72– +14H+ → 2Cr3+ +7H2O Cr2O72– +14H+ +14OH− → 2Cr3+ +7H2O +14OH− Cr2O72– +14H2O → 2Cr3+ +7H2O +14OH− Cr2O72– +7H2O → 2Cr3+ +14OH−

Balancing Redox Reactions,Continued : Tro's Introductory Chemistry, Chapter 16 39 Balancing Redox Reactions,Continued 4. Balance each half-reaction with respect to charge by adjusting the numbers of electrons. Electrons on product side for oxidation. Electrons on reactant side for reduction. 5. Balance electrons between half-reactions . 6. Add half-reactions, canceling electrons and common species. 7. Check. } x 3 2 I− → I2 2 I− → I2 + 2e− Cr2O72– +7H2O → 2Cr3+ +14OH− Cr2O72– +7H2O + 6e−→ 2Cr3+ +14OH− 2 I− → I2 + 2e− 6 I− → 3 I2 + 6e− Cr2O72– +7H2O + 6e−→ 2Cr3+ +14OH− Cr2O72– +7H2O + 6 I−→ 2Cr3+ +14OH− + 3 I2

Will a Reaction Take Place? : Tro's Introductory Chemistry, Chapter 16 40 Will a Reaction Take Place? Reactions that are energetically favorable are said to be spontaneous. They can happen, but the activation energy may be so large that the rate is very slow. The relative reactivity of metals can be used to determine if some redox reactions are spontaneous.

Single Displacement Reactions : Tro's Introductory Chemistry, Chapter 16 41 Single Displacement Reactions Also known as single replacement reactions. A more active free element displaces a less active element in a compound. Metals displace metals or H. Cu + 2 AgNO3 ® Cu(NO3)2 + 2 Ag Mg + 2 HCl ® MgCl2 + H2 Nonmetals displace nonmetals. 2 KI + Br2 ® 2 KBr + I2 Carbon displaces metals from oxides. 3 C + Fe2O3 ® 3 CO + 2 Fe Always redox.

Tendency to Lose Electrons : Tendency to Lose Electrons Some metals have a greater tendency to lose electrons than others. Metallic-free elements are always oxidized. The greater the tendency of a metal to lose electrons, the easier it is to oxidize. The greater the tendency of a metal to lose electrons, the harder it is to reduce its cations. If Metal A has a greater tendency to lose electrons than Metal B, then: A(s) + B+(aq)  A+(aq) + B(s), but: A+(aq) + B(s)  no reaction. Tro's Introductory Chemistry, Chapter 16 42

Activity Series of Metals : 43 Zn + Fe2+ ® Fe + Zn2+ Activity Series of Metals Listing of metals by reactivity. Free metal higher on the list displaces metal cation lower on the list. Metals above H will dissolve in acid: Cu + Fe2+ ® no reaction

Slide 44 : Tro's Introductory Chemistry, Chapter 16 44 Mg is above Cu on the activity series.

Table of Oxidation Half-Reactions : Table of Oxidation Half-Reactions 45

Table of Oxidation Half-Reactions, Continued : Table of Oxidation Half-Reactions, Continued Any oxidation half-reaction that is higher on the list will give a spontaneous reaction when combined with the reverse of a half-reaction that is lower on the list. The reverse of an oxidation half-reaction is a reduction half-reaction. Metals will dissolve in acid if their oxidation half-reaction is above H2  2H++ 2e−. Tro's Introductory Chemistry, Chapter 16 46

Example—Complete and Balance the Following Reaction:Al(s) + NiCl2(aq) ® : Example—Complete and Balance the Following Reaction:Al(s) + NiCl2(aq) ® Check the activity series. Al more reactive. If more reactive metal uncombined, then you will get reaction. Reaction. Determine what ion the uncombined metal will form. Al ® Al+3 Determine formula of new compound. AlCl3 Complete and balance; check solubility of salt. 2 Al(s) + 3 NiCl2(aq) ® 3 Ni(s) + 2 AlCl3(aq) Tro's Introductory Chemistry, Chapter 16 47

Practice—Predict the Products and Balance the Equation. : Tro's Introductory Chemistry, Chapter 16 48 Practice—Predict the Products and Balance the Equation. Mg + H3PO4 ® Cu + H2SO4 ® Al + Fe2+ ®

Slide 49 : Tro's Introductory Chemistry, Chapter 16 49 3 Mg + 2 H3PO4 ® Mg3(PO4)2 + 3 H2 Cu + H2SO4 ® No reaction. 2 Al + 3 Fe2+ ® 2 Al3+ + 3 Fe Practice—Predict the Products and Balance the Equation, Continued.

Electrical Current : Electrical Current When we talk about the current of a liquid in a stream, we are discussing the amount of water that passes by in a given period of time. When we discuss electric current, we are discussing the amount of electric charge that passes a point in a given period of time. Whether as electrons flowing through a wire or ions flowing through a solution. Tro's Introductory Chemistry, Chapter 16 50

Redox Reactions and Current : Redox Reactions and Current Redox reactions involve the transfer of electrons from one substance to another. Therefore, redox reactions have the potential to generate an electric current. In order to use that current, we need to separate the place where oxidation is occurring from the place that reduction is occurring. Tro's Introductory Chemistry, Chapter 16 51

Electric Current Flowing Directly Between Atoms : Electric Current Flowing Directly Between Atoms 52

Electric Current Flowing Indirectly Between Atoms : Electric Current Flowing Indirectly Between Atoms Tro's Introductory Chemistry, Chapter 16 53

Electrochemical Cells : Tro's Introductory Chemistry, Chapter 16 54 Electrochemical Cells Electrochemistry is the study of redox reactions that produce or require an electric current. The conversion between chemical energy and electrical energy is carried out in an electrochemical cell. Spontaneous redox reactions take place in a voltaic cell. Also known as galvanic cells. Batteries are voltaic cells. Nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy.

Electrochemical Cells, Continued : Tro's Introductory Chemistry, Chapter 16 55 Electrochemical Cells, Continued Oxidation and reduction reactions kept separate. Half-cells. Electron flow through a wire, along with ion flow through a solution, constitutes an electric circuit. Requires a conductive solid (metal or graphite) electrode to allow the transfer of electrons. Through external circuit. Ion exchange between the two halves of the system. Electrolyte.

Electrodes : Tro's Introductory Chemistry, Chapter 16 56 Electrodes Anode Electrode where oxidation occurs. Anions attracted to it. Connected to positive end of battery in electrolytic cell. Loses weight in electrolytic cell. Cathode Electrode where reduction occurs. Cations attracted to it. Connected to negative end of battery in electrolytic cell. Gains weight in electrolytic cell. Electrode where plating takes place in electroplating.

Voltaic Cell : 57 Voltaic Cell

Current and Voltage : Tro's Introductory Chemistry, Chapter 16 58 Current and Voltage The number of electrons that flow through the system per second is the current. Electrode surface area dictates the number of electrons that can flow. The amount of force pushing the electrons through the wire is the voltage. The farther the metals are separated on the activity series, the larger the voltage will be.

Current : Tro's Introductory Chemistry, Chapter 16 59 Current The amount of water that passes a point each second is called the current of the river. The number of electrons that pass a point each second is called the current of the electricity.

Voltage : Tro's Introductory Chemistry, Chapter 16 60 Voltage Gravity is the force pulling the water down the river. Voltage is the force pushing the electrons down the wire.

Dead Battery : Tro's Introductory Chemistry, Chapter 16 61 Dead Battery As the reaction proceeds, the reactants get consumed and the voltaic cell “dies.” The current decreases until electrons can no longer flow through the wire.

LeClanché’s Acidic Dry Cell : Tro's Introductory Chemistry, Chapter 16 62 LeClanché’s Acidic Dry Cell Electrolyte in paste form. ZnCl2 + NH4Cl. Or MgBr2. Anode = Zn (or Mg). Zn(s) ® Zn2+(aq) + 2 e- Cathode = graphite rod. MnO2 is reduced. 2 MnO2(s) + 2 NH4+(aq) + 2 H2O(l) + 2 e- ® 2 NH4OH(aq) + 2 Mn(O)OH(s) Cell voltage = 1.5 v. Expensive, nonrechargeable, heavy, easily corroded.

Alkaline Dry Cell : Tro's Introductory Chemistry, Chapter 16 63 Alkaline Dry Cell Same basic cell as acidic dry cell, except electrolyte is alkaline KOH paste. Anode = Zn (or Mg). Zn(s) ® Zn2+(aq) + 2 e- Cathode = brass rod. MnO2 is reduced. 2 MnO2(s) + 2 NH4+(aq) + 2 H2O(l) + 2 e- ® 2 NH4OH(aq) + 2 Mn(O)OH(s) Cell voltage = 1.54 v. Longer shelf life than acidic dry cells and rechargeable; little corrosion of zinc.

Lead Storage Battery : Tro's Introductory Chemistry, Chapter 16 64 Lead Storage Battery Six cells in series. Electrolyte = 6 M H2SO4. Anode = Pb. Pb(s) + SO42-(aq) ® PbSO4(s) + 2 e- Cathode = Pb coated with PbO2. PbO2 is reduced. PbO2(s) + 4 H+(aq) + SO42-(aq) + 2 e- ® PbSO4(s) + 2 H2O(l) Cell voltage = 2.09 v. Rechargeable, heavy.

Fuel Cells : 65 Fuel Cells Like batteries in which reactants are constantly being added. So it never runs down! Anode and cathode both Pt-coated metal. Electrolyte is OH– solution. Anode reaction: 2 H2 + 4 OH– → 4 H2O(l) + 4 e-. Cathode reaction: O2 + 4 H2O + 4 e- → 4 OH–.

Nonspontaneous Redox Reaction : Tro's Introductory Chemistry, Chapter 16 66 Nonspontaneous Redox Reaction The reverse of a spontaneous reaction is nonspontaneous. To get it to run, an outside energy source must be supplied. Nonspontaneous redox reactions can be made to work by using a battery to force the electrons to flow in the nonspontaneous direction.

Electrolysis : Tro's Introductory Chemistry, Chapter 16 67 Electrolysis Electrolysis is the process of using electricity to break a compound apart. Electrolysis is done in an electrolytic cell. Electrolytic cells can be used to separate elements from their compounds. Generate H2 from water for fuel cells. Recover metals from their ores.

Electrolytic Cell : Tro's Introductory Chemistry, Chapter 16 68 Electrolytic Cell The + terminal of the battery = anode. The - terminal of the battery = cathode. Cations attracted to the cathode; anions attracted to the anode. Cations pick up electrons from the cathode and are reduced; anions release electrons to the anode and are oxidized. In electroplating, the work piece is the cathode. Cations are reduced at the cathode and plate onto the surface. The anode is made of the plate metal, the anode oxidizes and replaces the metal cations lost from the solution.

Electrolytic Cell—Electroplating : Tro's Introductory Chemistry, Chapter 16 69 Electrolytic Cell—Electroplating

Corrosion : Tro's Introductory Chemistry, Chapter 16 70 Corrosion Corrosion is the spontaneous oxidation of a metal by chemicals in the environment. Since many materials we use are active metals, corrosion can be a very big problem.

Preventing Corrosion : Tro's Introductory Chemistry, Chapter 16 71 Preventing Corrosion One way to reduce or slow corrosion is to coat the metal surface to keep it from contacting corrosive chemicals in the environment. Paint. Some metals, like Al, form an oxide that strongly attaches to the metal surface, preventing the rest from corroding. Another method to protect one metal is to attach it to a more reactive metal that is cheap. Sacrificial electrode.