AP Chemistry
Chapter 20 Outline
20Electrochemistry
20.1Oxidation States
20.1.1Summary of rules for assigning oxidation numbers:
20.1.1.1Uncombined elements have an oxidation number of 0.
20.1.1.2For monatomic ions, the ion charge is the oxidation number.
20.1.1.3In compounds, hydrogen usually has an oxidation number of +1
20.1.1.3.1In metal hydrides, hydrogen has an oxidation number of -1.
20.1.1.4In compounds, oxygen usually has an oxidation number of -2.
20.1.1.4.1In peroxides, oxygen has an oxidation number of -1.
20.1.1.4.2In a binary compound with fluorine, oxygen has an oxidation number of +2.
20.1.1.5In compounds, fluorine always has an oxidation number of -1.
20.1.1.6The sum of the oxidation numbers is 0 for a neutral compound
20.1.1.7For a polyatomic ion, the sum of the oxidation numbers is the charge of the ion.
20.1.2Redox reaction = reaction in which oxidation numbers change
20.1.2.1If one substance loses electrons, another substance must gain electrons
20.1.3Oxidation = loss of electrons
20.1.3.1When a substance is oxidized, its oxidation number increases
20.1.4Reduction = gain of electrons
20.1.4.1When a substance is reduced, its oxidation number decreases.
20.1.4.2LEO the lion says GER, or OILRIG
20.1.5The species that is oxidized is the REDUCING AGENT
20.1.6The species that is reduced is the OXIDIZING AGENT
20.2Balancing oxidation-reduction reactions
20.2.1Both mass and charge must be conserved.
20.2.1.1Use “half reactions”, i.e., write out the oxidation step and the reduction step separately
20.2.1.2The number of electrons lost must equal the number of electrons gained.
20.2.2Balancing redox reactions
20.2.2.1Write skeletons for the oxidation and reduction half reactions.
20.2.2.2For each half reaction, BE SURE YOU CAN DO THIS!
20.2.2.2.1 Balance the elements other than H and O.
20.2.2.2.2 Add H2O to balance O atoms.
20.2.2.2.3 Add H+ to balance H atoms.
20.2.2.2.4 Add e- to balance charge; the sum of the charges should be the same on both sides.
20.2.2.3Multiply the half-reactions by integers to equal the numbers of electrons in both half reactions.
20.2.2.4Add the two half-reactions and simplify.
20.2.2.5If balancing in basic conditions, then add OH- to neutralize any H+ and simplify.
20.3Voltaic Cells (aka galvanic cells)
20.3.1The energy released in a spontaneous redox reaction can be used to perform electrical work.
20.3.1.1Physically separate the reduction half from the oxidation half to create a flow of electrons through an external circuit.
20.3.2electrode = strip of solid metal, connected to external circuit
20.3.2.1anode = electrode where oxidation occurs
20.3.2.1.1negative electrode (by convention)
20.3.2.1.2during reaction, anode will lose mass (as metal turns into ions in solution)
20.3.2.2cathode = electrode where reduction occurs
20.3.2.2.1positive electrode (by convention)
20.3.2.2.2during reduction, cathode will gain mass (as ions gaining electrons deposit on electrode)
20.3.2.3AN OX RED CAT
20.3.3half cell = a (metal) electrode immersed in a solution of its own ions
20.3.3.1anode solution will become more concentrated during reaction
20.3.3.2cathode solution will become less concentrated during reaction
20.3.3.3Electrons travel from the anode through the external wire to the cathode.
20.3.3.4Salt bridge = allows ions to move to maintain charge neutrality in both half-cells
20.3.3.4.1 Anions travel toward the anode
20.3.3.4.2 Cations travel toward the cathode
20.3.3.4.3You need to be able to generate correctly labeled sketches of the components of a voltaic cell
20.4Cell EMF under standard conditions
20.4.11 volt = 1 J/1 C (a couloumb is a mole of electrons)
20.4.2Electrons flow from the anode to the cathode because of a difference in potential energy.
20.4.2.1Potential energy of electrons is higher in the anode than in the cathode.
20.4.2.2Electromotive force (EMF) = the potential difference that pushes electrons through the external circuit
20.4.2.3Cell potential = the EMF of a voltaic cell = cell voltage = Ecell
20.4.2.4For spontaneous reactions (i.e. voltaic cells), Ecell >0
20.4.2.5Standard conditions = 1 M concentration, 1 atm (for gases), 25oC
20.4.3Standard reduction potential = a measure of the tendency of a reduction half-reaction to occur, relative to a standard= Eo
20.4.3.1Standard hydrogen electrode Eo = 0 by convention
20.4.3.2The more positive the value of Eo, the greater the tendency of the reactant of the half-reaction to be reduced.
20.4.3.2.1Higher Eo stronger oxidizing agent
20.4.3.3The more negative the value of Eo, the less tendency for this reduction reaction to occur
20.4.3.3.1 i.e, the reverse, oxidation half-reaction becomes more likely!
20.4.3.3.2 Lower Eo stronger reducing agent
20.4.3.4Tabulated for many reduction half-reactions
20.4.3.5Intensive property! Multiplying a half-reaction by a constant value does not change the value of Eo
20.4.4To find Ecell,
20.4.4.1Find half-reactions on table of standard reduction potentials
20.4.4.2The reaction that is higher: keep as written (i.e., reduction)
20.4.4.2.1 This reaction occurs at the CATHODE
20.4.4.3The reaction that is lower: reverse, and change the sign of Eo
20.4.4.3.1 This reaction occurs at the ANODE
20.4.5 The sum of the Eo values gives Ecell
20.4.5.1The sum of the reactions (after equalizing the number of e- lost and gained) gives the overall reaction for the cell
20.5Free energy and redox reactions
20.5.1Any reaction that can occur in a voltaic cell to produce a positive EMF must be spontaneous.
20.5.1.1A positive EMF value indicates a spontaneous process.
20.5.1.2A negative EMF value indicates a non-spontaneous process.
20.5.2 KNOW THIS EQUATION!
20.5.2.1n = the number of electrons transferred in the reaction
20.5.2.2F = Faraday’s constant = the quantity of electrical charge on one mole of
electrons = 96485 C/mol
20.6Cell EMF under nonstandard conditions
20.6.1As a battery runs, the reactant and product concentrations change. Eventually, the battery is “dead.”
20.6.1.1A dead battery = a system at equilibrium
20.6.1.2Cell EMF depends on reactant and product concentrations
20.6.2Nernst equation
20.6.2.1In general, if reactants increase relative to products, EMF increases
20.6.2.2If products increase relative to reactants, EMF decreases
20.6.2.3 KNOW THIS
20.6.3Concentration cell = cell based solely on the EMF generated because of a difference in concentration
20.6.3.1Basis for pH meters & function of nerve cells
You are not expected to know sections 20.7 and 20.8 in any detail for the AP exam
20.9Electrolysis
20.9.1It is possible to use electrical energy to cause non-spontaneous redox reactions to occur.
20.9.1.1Electrolysis reactions = reactions driven by an outside source of electrical energy
20.9.1.2Ecell is < 0
20.9.2Electrolytic cells consist of two electrodes in a molten salt or solution
20.9.2.1Oxidation occurs at the anode
20.9.2.2In electrolytic cells, the anode is the positive electrode
20.9.2.3Reduction occurs at the cathode
20.9.2.3.1 In electrolytic cells, the cathode is the negative electrode
20.9.2.4For a more detailed discussion of electrolysis reactions, go to the packet on electrolysis from the Ultimate Chemical Equations book.
20.9.3Quantitative aspects of electrolyis
20.9.3.1Couloumb = amperes x seconds
20.9.3.2To calculate the quantities of substances involved in an electrolytic process:
20.9.3.2.1
n = number of electrons needed to go from cation to neutral atom for that metal
Be able to do these calculations!