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AP Chemistry
Chapter 20 Outline
20
Electrochemistry
20.1 Oxidation States
20.1.1 Summary of rules for assigning oxidation numbers:
20.1.1.1
Uncombined elements have an oxidation number of 0.
20.1.1.2
For monatomic ions, the ion charge is the oxidation number.
20.1.1.3
In compounds, hydrogen usually has an oxidation number of +1
20.1.1.3.1 In metal hydrides, hydrogen has an oxidation number of -1.
20.1.1.4
In compounds, oxygen usually has an oxidation number of -2.
20.1.1.4.1 In peroxides, oxygen has an oxidation number of -1.
20.1.1.4.2 In a binary compound with fluorine, oxygen has an oxidation number of +2.
20.1.1.5
In compounds, fluorine always has an oxidation number of -1.
20.1.1.6
The sum of the oxidation numbers is 0 for a neutral compound
20.1.1.7
For a polyatomic ion, the sum of the oxidation numbers is the charge of the ion.
20.1.2 Redox reaction = reaction in which oxidation numbers change
20.1.2.1
If one substance loses electrons, another substance must gain electrons
20.1.3 Oxidation = loss of electrons
20.1.3.1
When a substance is oxidized, its oxidation number increases
20.1.4 Reduction = gain of electrons
20.1.4.1
When a substance is reduced, its oxidation number decreases.
20.1.4.2
LEO the lion says GER, or OIL RIG
20.1.5 The species that is oxidized is the REDUCING AGENT
20.1.6 The species that is reduced is the OXIDIZING AGENT
20.2 Balancing oxidation-reduction reactions
20.2.1 Both mass and charge must be conserved.
20.2.1.1
Use “half reactions”, i.e., write out the oxidation step and the reduction step
separately
20.2.1.2
The number of electrons lost must equal the number of electrons gained.
20.2.2 Balancing redox reactions
20.2.2.1
Write skeletons for the oxidation and reduction half reactions.
20.2.2.2
For 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.
20.2.2.3
Multiply the half-reactions by integers to equal the numbers of electrons in both
half reactions.
20.2.2.4
Add the two half-reactions and simplify.
20.2.2.5
If balancing in basic conditions, then add OH- to neutralize any H+ and simplify.
20.3
Voltaic Cells (aka galvanic cells)
20.3.1 The energy released in a spontaneous redox reaction can be used to perform electrical
work.
20.3.1.1
Physically separate the reduction half from the oxidation half to create a flow of
electrons through an external circuit.
20.3.2 electrode = strip of solid metal, connected to external circuit
20.3.2.1
anode = electrode where oxidation occurs
20.3.2.1.1 negative electrode (by convention)
20.3.2.1.2 during reaction, anode will lose mass (as metal turns into ions in solution)
20.3.2.2
cathode = electrode where reduction occurs
20.3.2.2.1 positive electrode (by convention)
20.3.2.2.2 during reduction, cathode will gain mass (as ions gaining electrons deposit on
electrode)
20.3.2.3
AN OX
RED CAT
20.3.3 half cell = a (metal) electrode immersed in a solution of its own ions
20.3.3.1
anode solution will become more concentrated during reaction
20.3.3.2
cathode solution will become less concentrated during reaction
20.3.3.3
Electrons travel from the anode through the external wire to the cathode.
20.3.3.4
Salt 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
20.3.3.4.3 You need to be able to generate correctly labeled sketches of the components of a
voltaic cell
20.4 Cell EMF under standard conditions
20.4.1 1 volt = 1 J/1 C
(a couloumb is a mole of electrons)
20.4.2 Electrons flow from the anode to the cathode because of a difference in potential energy.
20.4.2.1
Potential energy of electrons is higher in the anode than in the cathode.
20.4.2.2
Electromotive force (EMF) = the potential difference that pushes electrons
through the external circuit
20.4.2.3
Cell potential = the EMF of a voltaic cell = cell voltage = Ecell
20.4.2.4
For spontaneous reactions (i.e. voltaic cells), Ecell >0
20.4.2.5
Standard conditions = 1 M concentration, 1 atm (for gases), 25oC
20.4.3 Standard reduction potential = a measure of the tendency of a reduction half-reaction to
occur, relative to a standard= Eo
20.4.3.1
Standard hydrogen electrode
Eo = 0 by convention
20.4.3.2
The more positive the value of Eo, the greater the tendency of the reactant of the
half-reaction to be reduced.
20.4.3.2.1 Higher Eo  stronger oxidizing agent
20.4.3.3
The 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!
o
20.4.3.3.2
Lower E  stronger reducing agent
20.4.3.4
Tabulated for many reduction half-reactions
20.4.3.5
Intensive property! Multiplying a half-reaction by a constant value does not
change the value of Eo
20.4.4 To find Ecell,
20.4.4.1
Find half-reactions on table of standard reduction potentials
20.4.4.2
The reaction that is higher: keep as written (i.e., reduction)
20.4.4.2.1
This reaction occurs at the CATHODE
20.4.4.3
The 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
20.4.5.1
The sum of the reactions (after equalizing the number of e- lost and gained) gives
the overall reaction for the cell
20.5 Free energy and redox reactions
20.5.1 Any reaction that can occur in a voltaic cell to produce a positive EMF must be
spontaneous.
20.5.1.1
A positive EMF value indicates a spontaneous process.
20.5.1.2
A negative EMF value indicates a non-spontaneous process.
20.5.2
KNOW THIS EQUATION!
G o   nFE o
20.5.2.1
n = the number of electrons transferred in the reaction
20.5.2.2
F = Faraday’s constant = the quantity of electrical charge on one mole of
electrons = 96485 C/mol
20.6 Cell EMF under nonstandard conditions
20.6.1 As a battery runs, the reactant and product concentrations change. Eventually, the battery
is “dead.”
20.6.1.1
A dead battery = a system at equilibrium
20.6.1.2
Cell EMF depends on reactant and product concentrations
20.6.2 Nernst equation
20.6.2.1
In general, if reactants increase relative to products, EMF increases
20.6.2.2
If products increase relative to reactants, EMF decreases
0.0592
0.0592
[ products]
20.6.2.3
KNOW THIS
Ecell  Eo 
log Q  Eo 
log
n
n
[reac tan ts]
20.6.3 Concentration cell = cell based solely on the EMF generated because of a difference in
concentration
20.6.3.1
Basis 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
20.9 Electrolysis
20.9.1 It is possible to use electrical energy to cause non-spontaneous redox reactions to occur.
20.9.1.1
Electrolysis reactions = reactions driven by an outside source of electrical energy
20.9.1.2
Ecell is < 0
20.9.2 Electrolytic cells consist of two electrodes in a molten salt or solution
20.9.2.1
Oxidation occurs at the anode
20.9.2.2
In electrolytic cells, the anode is the positive electrode
20.9.2.3
Reduction occurs at the cathode
20.9.2.3.1
In electrolytic cells, the cathode is the negative electrode
20.9.2.4
For a more detailed discussion of electrolysis reactions, go to the packet on
electrolysis from the Ultimate Chemical Equations book.
20.9.3 Quantitative aspects of electrolyis
20.9.3.1
Couloumb = amperes x seconds
20.9.3.2
To calculate the quantities of substances involved in an electrolytic process:
 1mole e   1mole metal  gfm of metal 



20.9.3.2.1
grams metal  (amperes)  (sec onds)
96485
C
n
e

1
mole




n = number of electrons needed to go from cation to neutral atom
for that metal
Be able to do these calculations!