Electrochemistry
Chapter 20
Introduction
• Electrochemistry involves looking at how
spontaneous redox reactions can be used to produce
electricity and how electricity can cause
nonspontaneous reactions to occur.
• Ex. Batteries use spontaneous redox reactions and
convert chemical energy into electrical energy
• Review of redox reactions:
• A species is oxidized when it loses electrons.
– Zinc loses two electrons, forming the Zn2+ ion.
• A species is reduced when it gains electrons.
– H+ gains an electron, forming H2.
20.2 – Balancing Redox Equations
• Half-Reactions = either the oxidation or the reduction process
shown alone with the # of electrons lost or gained.
• Ex. Cu(s) → Cu2+(aq) + 2e- (oxidation)
• Ex. Ag+(aq) + e- → Ag(s) (reduction)
• The # of electrons lost must equal the # gained, so the second
half-reaction becomes -- 2Ag+(aq) + 2e- → 2Ag(s)
• When you add them together, you get an oxidation-reduction
reaction:
• Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)
• For acidic solutions, follow the steps on the bottom of pg 860
and the top of page 861
• Sample exercise 20.2
• Balancing in Basic Solution:
• A reaction that occurs in basic solution can be
balanced as if it occurred in acid.
• Once the equation is balanced, add OH– to each side
to “neutralize” the H+ in the equation and create
water in its place.
• If this produces water on both sides,
subtract water from each side so it appears on only
one side of the equation.
• Sample exercise 20.3
20.3 – Voltaic Cells
• In spontaneous redox
reactions, electrons are
transferred and energy
is released.
• That energy can do
work if the electrons
flow through an
external device.
• This is a voltaic cell.
• Components of a Voltaic Cell:
 electrode = a strip of metal used in an electrochemical
experiment. It can be made of a metal that’s in the reaction,
or another conducting material such as platinum.
a) anode = the electrode at which oxidation occurs
b) cathode = the electrode at which reduction occurs
 half-cell = one of the compartments of a voltaic cell
 Salt bridge = a U-shaped tube containing an electrolyte
solution that neutralizes any charge buildup in the half-cells.
 Electrons flow through the
wire from the anode to
the cathode, producing an
electrical current.
• See sample exercise 20.4
• Important other points about a voltaic cell:
 If the electrodes are made of materials that participate in the
reaction (i.e. a zinc electrode in a solution of ZnSO4), than the
anode will lose mass as the neutral atoms of the electrode
oxidize into ions, and the cathode will gain mass as its ions
deposit as neutral atoms.
 If the electrodes are made of a conducting material that does
not participate in the reaction, their masses don’t change.
 The cations in the salt bridge will migrate to the cathode to
neutralize the negative charge that will build up at the + ions
become neutral, leaving the −ions (spectators) in excess.
 The anions in the salt bridge will migrate to the anode to
neutralize the excess + charge from the ions that are forming.
20.4 – Cell Potentials Under Standard
Conditions (E0cell)
• In a Voltaic cell, electrons flow spontaneously from the anode
to the cathode because of a difference in potential energy
(they flow from high to low potential energy).
• The difference in potential energy between the 2 electrodes
(called the cell potential, or Ecell) is measured in volts. Note:
your book also calls it emf
• One volt is one joule per coulomb (1 V = 1 J/C). An electron
has a charge of 1.60 x 10-19 C.
• E°cell is the standard cell potential [occurs at standard
conditions – 1M (or 1atm for gaseous) reactants and
products].
• Standard reduction potential (E°red):
• It’s a measurement of the tendency for a reduction process to
occur at an electrode. The value is determined by comparison
to a standard hydrogen electrode (H+→H2), which is given a
value of 0.
–
–
Substances that are easily reduced have a large + value for their E°red.
Substances that are easily oxidized have a large – value for their E°red.
• To calculate E°cell:
1) write the reduction half-equation and look up its E°red
2) write the oxidation half-equation and write the opposite
sign of its E°red (which is a E°oxid)
3) combine the half-equations and add the E°s. Note: E°
values are unaffected by coefficients.
• Voltaic cells will have a + E°cell
• See sample exercises 20.5 – 20.7
• The larger the + value for the
E°red, the more easily the
substance is reduced
• The more negative the value is for
E°red, the more likely the
substance will be oxidized.
• **metals at the top of the activity
series have the most negative
E°red
20.5 – Free Energy and Redox Reactions
•
•
•
•
•
•
•
Relationship of Ecell to spontaneity:
If Ecell is +, the rxn is spontaneous in the forward direction
If Ecell is -, the rxn is spontaneous in the reverse direction
If Ecell is 0, the rxn is at equilibrium
See sample exercise 20.9
Relationship of Ecell to free energy:
Go = -nFEocell , where n = the # of moles of electrons
transferred according to the balanced equation, and F = the
electric charge per mole of e-, called the Faraday constant
(96,485 C/mole e- = 96485J/V-mol). The units for Go are
J/mol
• Relationship between ΔG, Eocell , and K:
• ΔG° = –nFE° = –RT ln K
• See sample exercise 20.10
20.6 – Cell Potentials Under Nonstandard
Conditions
• Effect of Concentrations on Cell Voltage:
• Most cell measurements involve nonstandard conditions (i.e.
not 1M solute concentrations and/or not 1 atm)
• Using nonstandard conditions will give you an Ecell that is
different from the E°cell. Also, the Ecell changes as the reaction
progresses.
• Ex. Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
• Think of it as an LeChatelier’s problem
• As the reaction proceeds, the conc. of products ↑, and the
conc. of reactants ↓, so the reaction shifts to the left, and Ecell
decreases. Ecell reaches 0 at equilibrium (when the cell is
“dead”)
• In general, increasing the concentration of the reactants or
decreasing the concentration of the products drives the
reaction towards the products, increasing Ecell.
• Conversely, decreasing the concentration of the reactants or
increasing the concentration of the products results in a
smaller Ecell.
• Note: changing the size of the electrodes has no effect on the
Ecell since the concentration of solids doesn’t change (as long
as some electrode remains in contact with the solution).
20.9 -- Electrolysis
• Electrolytic Cell: Electricity from an external source causes a
nonspontaneous reaction to occur (electrolysis). The external
source of electricity acts like an “electron pump” and pulls
electrons away from the anode, where the oxidation occurs,
and pushes them into the cathode, where the reduction
occurs.
• Example: in this example, Na+
and Cl- are forced to become
neutral Na and Cl2. The
electrodes are inert.
• Electroplating: uses electrolysis to deposit a thin layer of one
metal onto another to improve the appearance or resist
corrosion.
• Quantitative Aspects of
Electrolysis:
• The quantity of reactant
consumed or product formed
during electrolysis depends
on:
• 1. the molar mass of the
substance (reactant or
product)
• 2. the quantity of electric
charge used (measured in
Coulombs)
• 3. the # of electrons
transferred
• The equation to use is: I = q/t, where I is current
(amps), q = the charge (Coulombs), and t = time
(seconds)
• To calculate the quantitative outcomes of electrolysis,
the strategy is as follows:
• 1. Solve for q
• 2. convert q to moles of e- (1 mole e- = 96,485 C)
• 3. change moles of e- to moles of reactant or product
• 4. convert moles of reactant or product to the unit
needed
• see sample exercise 20.14