Principles of Reactivity:
Electron Transfer Reactions
Chapter 20
Electron Transfer Reactions
• In order to produce electricity, the
electron transfer between
substances must be carried out in an
apparatus that allows the electrons
to be transferred through an
electrical circuit.
• Devices that use chemical reactions
to produce an electric current are
called voltaic cells or galvanic cells.
Electrochemistry
• refers to the interchange of electrical
and chemical energy
– Voltaic cells use product favored
reactions to convert chemical energy to
electrical energy.
– Electrolysis is when electrical energy is
used to effect a chemical change.
• Example: splitting water into its component
elements
20.1 Oxidation-Reduction Reactions
• One reactant is oxidized and one is
reduced.
• Oxidation and reduction must
balance.
• Oxidizing agent is reduced.
• Reducing agent is oxidized.
• Oxidation numbers can be used to
figure out what is oxidized (#
increases) and what is reduced (#
decreases).
Balancing Redox Equations
ALL REDOX REACTIONS MUST BE
BALANCED FOR BOTH MASS AND
CHARGE!
– The method most often used to balance
redox equation is by writing halfreactions.
– One half-reaction describes oxidation,
one describes reduction.
• Balance electrons and add together!
– The net ionic equation will therefore be
balanced for both mass and charge.
Practice Problem
• Aluminum reacts with nonoxidizing
acids to give Al3+(aq) and H2(g). The
(unbalanced) equation is
Al(s) + H+(aq)  Al3+(aq) + H2(g)
Write balanced half-reactions and
the balanced net ionic equation.
Identify the oxidizing agent, the
reducing agent, the substance
oxidized, and the substance
reduced.
Balancing Redox
• Sometimes water must be added and
either hydrogen or hydroxide ions if
the reaction occurs in acidic or basic
solution.
– equations with sulfate, nitrate,
chromate, permanganate, etc…
Practice Problem
• The permanganate ion, MnO4-, is an
oxidizing agent. A common
laboratory analysis for iron is to
titrate aqueous iron (II) ion with a
solution of potassium permanganate
of precisely known concentration.
Use the half-reaction method to write
the balanced net ionic equation for
the reaction in acidic solution.
MnO4-(aq) + Fe2+(aq)  Mn2+(aq) + Fe3+(aq)
Practice Problem
• Voltaic cells based on the oxidation
of sulfur are under development.
One such cell involves the reaction
of sulfur with aluminum under basic
conditions.
• Al(s) + S(s)  Al(OH)3(s) + HS-(aq)
• Balance this equation showing each
balanced half-reaction.
• Identify the oxidizing and reducing
agents, the substance oxidized, and
the substance reduced.
20.2 Simple Voltaic Cells
• In a voltaic cell, the two halfreactions are separated so that
electrons cannot be directly
transferred between reactants.
• The two cells are connected with a
salt bridge that allows cations and
anions to move between them.
Simple Voltaic Cells
• The anode is the electrode at which
oxidation occurs. The cathode is the
electrode at which reduction occurs.
• In the salt bridge, cations move from
anode to cathode, and anions move
from cathode to anode in order to
maintain electrical neutrality.
• Electrons flow spontaneously from
anode to cathode.
Practice Problem
• Describe how to set up a voltaic cell
using the following half-reactions:
• Reduction: Ag+(aq) + e-  Ag(s)
• Oxidation: Ni(s)  Ni2+(aq) + 2e• Which is the anode and which is the
cathode? What is the overall cell
reaction? What is the direction of
electron flow in an external wire
connecting the two electrodes?
Describe the ion flow in a salt bridge
connecting the cell compartments.
Inert Electrodes
• Sometimes there has to be an inert
electrode because there aren’t two
conductive metals as reactants.
• Shorthand is often used to
symbolize electrochemical cells.
– The anode is written on the left. A
single vertical line indicates a phase
boundary, and double vertical lines
indicate a salt bridge.
Cu(s)ICu2+(aq, 1.0M)IIAg+(aq, 1.0M)IAg(s)
20.4 Standard Electrochemical Potentials
• Electrons move from anode toward
the cathode due to the difference in
potential energy of electrons at the
two electrodes.
– electromotive force (emf) = difference
in potential energy
– units of volts (V)
• 1 Joule = 1 volt x 1 coulomb
• One coulomb is the quantity of charge that
passes a point in an electric circuit when a
current of one ampere flows for one second
(1 coulomb = 1 amp x 1 sec)
Standard Conditions
• Half-cell potentials assume the
following:
– Reactants and products are present as
pure liquids or solids.
– Solutes in aqueous solution have a
concentration of 1.0M.
– Gaseous reactants or products have a
pressure of 1.0 atm or 1.0 bar.
Standard Potentials are measured under
these conditions; Eocell
Deviations from Standard Conditions
• The farther the reaction is from
equilibrium, the greater the
magnitude of the cell potential
– As the system approaches equilibrium,
the magnitude of cell potential
decreases, reaching zero at equilibrium
– Deviations that take the cell further
from equilibrium increase cell potential
– Deviations that take the cell closer to
equilibrium decreases the cell potential
Standard Cell Potentials
• Predict that the reaction occurring is
the one in which the reactants are
stronger reducing and oxidizing
agents than the products.
• Electrons move from the electrode
of higher potential energy to the one
of lower potential energy.
• Cell potential can be calculated to
determine the relative oxidizing or
reducing ability.
Standard Reduction Potentials
• Eocell = Eocathode – Eoanode
• When the Eocell has a positive value,
the reaction is predicted to be
product-favored as written.
Tables of Standard Reduction Potentials
• Table on page 920
• All potentials are for reduction
reactions.
• Best (strongest) oxidizing agent is
written at the top. (most positive
Eocell)
• Best (strongest) reducing agents are
written at the bottom. (most negative
Eocell)
• Reversing a half-reaction reverses
the sign of Eo.
• Northwest-southeast Rule.
Practice Problems
• The net reaction that occurs in a
voltaic cell is
Zn(s) + 2Ag+(aq)  Zn2+(aq) + 2Ag(s)
Assuming standard conditions,
identify the half-reactions that occur
at the anode and the cathode and
calculate a potential for the cell.
Practice Problem
• Rank the halogens in order of their
strength as oxidizing agents.
• Decide if hydrogen peroxide in
acidic solution is a stronger
oxidizing agent than Cl2.
• Decide which of the halogens is
capable of oxidizing gold metal to
Au3+(aq).
Practice Problem
• Determine which of the following
redox equations are productfavored. Assume standard
conditions.
• Ni2+(aq) + H2(g)  Ni(s) + 2H+(aq)
• 2Fe3+(aq) + 2I-(aq)  2Fe2+(aq) + I2(s)
• Br2(l) + 2Cl-(aq)  2Br-(aq) + Cl2(g)
• Cr2O72-(aq) + 6Fe2+(aq) + 14H+(aq) 
2Cr3+(aq) + 6Fe3+(aq) + 7H2O(l)
Recognizing Redox Reactions
• One reactant is a metal and the other
is an aqueous metal ion. Metal will
be oxidized to form an aqueous ion,
usually by not always having a
charge of 2+ and the aqueous ion
will be reduced to the corresponding
metal.
– A piece of solid zinc is placed in an
aqueous solution of copper (II) sulfate.
Recognizing Redox Reactions
• One reactant is a polatomic anion
with a metallic element displaying its
highest oxidation number and the
other is an anion displaying an
oxidation number lower than its max.
The polyatomic ion reduces to an ion
displaying the metal in a lower
oxidation state. Anion is oxidized to
a higher oxidation state.
– Acidic aqueous sodium dichromate is
mixed with a solution of potassium
bromide.
Recognizing Redox Reactions
• An organic compound burned in air
(or oxygen) produces carbon dioxide
and water.
– Ethanol is burned in air.
Recognizing Redox Reactions
• A metal reacts with a non-metal to
produce a binary salt.
– Solid sodium is mixed with chlorine
gas.
Recognizing Redox Reactions
• An active metal reacts with water to
produce hydrogen gas and an
hydroxide base.
• Solid lithium is placed in water.
Faraday’s Laws
• Used to determine stoichiometry of
the redox reactions in cells with
respect to:
– Number of electrons transferred
– Mass of material deposited or removed
from an electrode
– Current
– Time elapsed
– Charge of ionic species
Homework
• After reading Chapter 20, you should
be able to do the following
problems…
• P. 944 (50-59)