Chapter 18 Electrochemistry Redox Reac)on One or more elements change oxida)on number all single displacement, and combus)on, some synthesis and decomposi)on Always have both oxida)on and reduc)on split reac)on into oxida)on half reac)on and a reduc)on half reac)on Oxidizing agent is reactant molecule that causes oxida)on (undergoes reduc)on) Reducing agent is reactant molecule that causes reduc)on (undergoes oxida)on) Oxida)on & Reduc)on Oxida/on is the process that occurs when oxida)on number of an element increases element loses electrons half reac)on has electrons as products Reduc/on is the process that occurs when oxida)on number of an element decreases element gains electrons half reac)ons have electrons as reactants 1
Recall: Rules for Assigning Oxida)on States 1. Free elements have an oxida)on state = 0 2. Monatomic ions have an oxida)on state equal to their charge 3. (a) the sum of the oxida)on states of all the atoms in a compound is 0 (b) the sum of the oxida)on states of all the atoms in a polyatomic ion equals the charge on the ion 4. (a) Group I metals have an oxida)on state of +1 in all their compounds 4. (b) Group II metals have an oxida)on state of +2 in all their compounds Recall: Rules for Assigning Oxida)on States 5. In their compounds, nonmetals have oxida)on states according to the table below Nonmetal Oxidation State Example F -1 CF 4 H +1 CH 4 O -2 CO 2 Group 7A -1 CCl 4 Group 6A -2 CS 2 Group 5A -3 NH 3 Oxida)on / Reduc)on Oxida)on and reduc)on 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
Redox: Examples H 2 S (g) + NO 3 (aq) S (s) + NO (g) (in acid) Redox: Examples MnO 2(s) + HBr(aq) MnBr 2(aq) + Br 2(l) (in acid) 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 )me When we discuss electric current, we are discussing the amount of electric charge that passes a point in a given period of )me whether as electrons flowing through a wire or ions flowing through a solu)on 3
Redox Reac)ons & Current Redox reac)ons involve the transfer of electrons from one substance to another Therefore, redox reac)ons have the poten)al to generate an electric current In order to use that current, we need to separate the place where oxida)on is occurring from the place that reduc)on is occurring Electric Current Flowing Directly Between Atoms Electric Current Flowing Indirectly Between Atoms 4
Electrochemical Cells Electrochemistry is the study of redox reac)ons that produce or require an electric current The conversion between chemical energy and electrical energy is carried out in an electrochemical cell Spontaneous redox reac)ons take place in a voltaic cell aka galvanic cells Nonspontaneous redox reac)ons can be made to occur in an electroly/c cell by the addi)on of electrical energy Electrochemical Cells Oxida)on and reduc)on reac)ons kept separate half cells Electron flow through a wire along with ion flow through a solu)on cons)tutes an electric circuit Requires a conduc)ve 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 Anode electrode where oxida)on occurs anions a^racted to it connected to posi)ve end of ba^ery in electroly)c cell loses weight in electroly)c cell Cathode electrode where reduc)on occurs ca)ons a^racted to it connected to nega)ve end of ba^ery in electroly)c cell gains weight in electroly)c cell electrode where pla)ng takes place in electropla)ng 5
Voltaic Cell Current and Voltage The number of electrons that flow through the system per second is the current (unit = Ampere) 1 A of current = 1 Coulomb of charge flowing by each second 1 A = 6.242 x 10 18 electrons/second Electrode surface area dictates the number of electrons that can flow The difference in poten)al energy between the reactants and products is the poten/al difference (unit = Volt) 1 V of force = 1 J of energy/coulomb of charge the voltage needed to drive electrons through the external circuit amount of force pushing the electrons through the wire is called the electromo/ve force, emf Cell Poten)al The difference in poten)al energy between the anode/the cathode in a voltaic cell is called the cell poten/al The cell poten)al depends on the rela)ve ease with which the oxidizing agent is reduced at the cathode and the reducing agent is oxidized at the anode The cell poten)al under standard condi)ons is called the standard emf, E cell 25 C, 1 atm for gases, 1 M concentra)on of solu)on sum of the cell poten)als for the half reac)ons 6
Cell Nota)on Shorthand descrip)on of Voltaic cell Electrode electrolyte electrolyte electrode Oxida)on half cell on leg, reduc)on half cell on the right Single = phase barrier if mul)ple electrolytes in same phase, a comma is used rather than ogen use an inert electrode Double line = salt bridge Fe(s) Fe 2+ (aq) MnO 4- (aq), Mn 2+ (aq), H + (aq) Pt(s) Standard Reduc)on Poten)al A half reac)on with a strong tendency to occur has a large + half cell poten)al When two half cells are connected, the electrons will flow so that the half reac)on with the stronger tendency will occur We cannot measure the absolute tendency of a half reac)on, we can only measure it rela)ve to another half reac)on We select as a standard half reac)on the reduc)on of H + to H 2 under standard condi)ons, which we assign a poten)al difference = 0 v standard hydrogen electrode, SHE 7
Half Cell Poten)als SHE reduc)on poten)al is defined to be exactly 0 v Half reac)ons with a stronger tendency toward reduc)on than the SHE have a + value for E red Half reac)ons with a stronger tendency toward oxida)on than the SHE have a value for E red E cell = E oxida)on + E reduc)on E oxida)on = E reduc)on when adding E values for the half cells, do not mul/ply the half cell E values, even if you need to mul)ply the half reac)ons to balance the equa)on 8
Example Calculate E cell for the reac)on at 25 C Al (s) + NO 3 (aq) + 4 H + (aq) Al3+ (aq) + NO (g) + 2 H 2 O (l) Example Sketch and Label the Voltaic Cell Fe(s) Fe 2+ (aq) Pb 2+ (aq) Pb(s), Write the Half Reac)ons and Overall Reac)on, and Determine the Cell Poten)al under Standard Condi)ons. 9
Example ox: Fe(s) Fe 2+ (aq) + 2 e E = +0.45 V red: Pb 2+ (aq) + 2 e Pb(s) E = 0.13 V tot: Pb 2+ (aq) + Fe(s) Fe 2+ (aq) + Pb(s) E = +0.32 V Predic)ng Whether a Metal Will Dissolve in an Acid Acids dissolve in metals if the reduc)on of the metal ion is easier than the reduc)on of H + (aq) Metals whose ion reduc)on reac)on lies below H + reduc)on on the table will dissolve in acid E cell, ΔG and K For a spontaneous reac)on one the proceeds in the forward direc)on with the chemicals in their standard states ΔG < 1 (nega)ve) E > 1 (posi)ve) K > 1 ΔG = RTlnK = nfe cell n is the number of electrons F = Faraday s Constant = 96,485 C/ mol e 10
Example Calculate ΔG for the reac)on I 2(s) + 2 Br (aq) Br 2(l) + 2 I (aq) Nonstandard Condi)ons the Nernst Equa)on ΔG = ΔG + RT ln Q E = E (0.0592/n) log Q at 25 C when Q = K, E = 0 use to calculate E when concentra)ons not 1 M E at Nonstandard Condi)ons 11
Example Calculate E cell at 25 C for the reac)on 3Cu (s) + 2MnO 4 (aq) + 8H + (aq) 2MnO 2(s) + Cu2+ (aq) +4H 2 O (l) Concentra)on Cells It is possible to get a spontaneous reac)on when the oxida)on and reduc)on reac)ons are the same, as long as the electrolyte concentra)ons are different The difference in energy is due to the entropic difference in the solu)ons the more concentrated solu)on has lower entropy than the less concentrated Electrons will flow from the electrode in the less concentrated solu)on to the electrode in the more concentrated solu)on LeClanche Acidic Dry Cell electrolyte in paste form ZnCl 2 + NH 4 Cl or MgBr 2 anode = Zn (or Mg) Zn(s) Zn 2+ (aq) + 2 e cathode = graphite rod MnO 2 is reduced 2 MnO 2 (s) + 2 NH 4+ (aq) + 2 H 2 O(l) + 2 e 2 NH 4 OH(aq) + 2 Mn(O)OH(s) cell voltage = 1.5 v expensive, nonrechargeable, heavy, easily corroded 12
Alkaline Dry Cell same basic cell as acidic dry cell, except electrolyte is alkaline KOH paste anode = Zn (or Mg) Zn(s) Zn 2+ (aq) + 2 e cathode = brass rod MnO 2 is reduced 2 MnO 2 (s) + 2 NH 4+ (aq) + 2 H 2 O(l) + 2 e 2 NH 4 OH(aq) + 2 Mn(O)OH(s) cell voltage = 1.54 v longer shelf life than acidic dry cells and rechargeable, li^le corrosion of zinc Lead Storage Ba^ery 6 cells in series electrolyte = 30% H 2 SO 4 anode = Pb Pb(s) + SO 4 2 (aq) PbSO 4 (s) + 2 e cathode = Pb coated with PbO 2 PbO 2 is reduced PbO 2 (s) + 4 H + (aq) + SO 4 2 (aq) + 2 e PbSO 4 (s) + 2 H 2 O(l) cell voltage = 2.09 v rechargeable, heavy Fuel Cells Like ba^eries in which reactants are constantly being added so it never runs down! Anode and Cathode both Pt coated metal Electrolyte is OH solu)on Anode Reac)on: 2 H 2 + 4 OH 4 H 2 O(l) + 4 e Cathode Reac)on: O 2 + 4 H 2 O + 4 e 4 OH 13
Electrochemical Cells In all electrochemical cells, oxida)on occurs at the anode, reduc)on occurs at the cathode In voltaic cells, anode is the source of electrons and has a ( ) charge cathode draws electrons and has a (+) charge In electroly)c cells electrons are drawn off the anode, so it must have a place to release the electrons, the + terminal of the ba^ery electrons are forced toward the anode, so it must have a source of electrons, the terminal of the ba^ery Electrolysis Electrolysis is the process of using electricity to break a compound apart Electrolysis is done in an electroly)c cell Electroly)c cells can be used to separate elements from their compounds generate H 2 from water for fuel cells recover metals from their ores 14
Electrolysis of Water Electrolysis of Pure Compounds Must be in molten (liquid) state Electrodes normally graphite Ca)ons are reduced at the cathode to metal element Anions oxidized at anode to nonmetal element Electrolysis of NaCl (l) 15
Corrosion Corrosion is the spontaneous oxida)on of a metal by chemicals in the environment Since many materials we use are ac)ve metals, corrosion can be a very big problem Rus)ng Rust is hydrated iron (III) oxide Moisture must be present water is a reactant required for flow between cathode and anode Electrolytes promote rus)ng enhances current flow Acids promote rus)ng lower ph = lower E red 16