GE-CHEM 1103 Module 4.3 Flashcards
the study of the
relationships between electricity
and chemical reactions. It includes
the study of both spontaneous and
nonspontaneous processes.
Electrochemistry
To keep track of what
loses electrons and
what gains them, we
assign
oxidation
numbers
If the oxidation number
increases for an
element
element
is oxidized
If the oxidation number
decreases for an
element
element
is reduced
Synopsis of Assigning Oxidation Numbers
- Elements = 0
- Monatomic ion = charge
- F: –1
- O: –2 (unless peroxide = –1)
- H: +1 (unless a metal hydride = –1)
- The sum of the oxidation numbers equals
the overall charge (0 in a compound).
LEO
Loses Electron Oxidized
GER
Gains Electron Reduced
causes something else to be
oxidized
oxidizing agent
causes something else
to be reduced
reducing agent
- The oxidation and reduction are written
and balanced separately. - We will use them to balance a redox
reaction.
Half-Reactions
Balancing Redox Equations: The
Half-Reactions Method (For acidic solution)
- Make two half-reactions (oxidation
and reduction). - Balance atoms other than O and H.
Then, balance O and H using H2O/H+. - Add electrons to balance charges.
- Multiply by common factor
to make electrons in half-reactions
equal. - Add the half-reactions.
- Simplify by dividing by common factor
or converting H+
to OH–
if basic. - Double-check atoms and charges
balance!
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.
- In spontaneous redox
reactions, electrons
are transferred and
energy is released. - That energy can do
work if the electrons
flow through an
external device.
Voltaic Cells
oxidation occurs
anode
reduction occurs
cathode
When electrons flow, charges aren’t balanced so a contraption is used to keep the charges balanced
a salt bridge, usually a U-shaped tube that contains a salt/agar solution
electrons leave
the anode and flow through
the wire to the cathode
cell
formed in the
anode compartment
Cations
electrons reach the
cathode
cations in solution
are attracted to the now
negative cathode
gain electrons
and are deposited as metal
on the cathode
cations
flow
spontaneously
one way in a
redox reaction,
from high to low
potential energy.
electrons
potential difference between the anode
and cathode in a cell
electromotive force (emf)
Other name of electromotive force (emf)
cell potential and is
designated Ecell
It is measured in volts (V). One volt is one
joule per coulomb (1 V = 1 J/C).
cell potential
Standard condition of a reduction potentials
1 M, 101.3 kPa,
25°C
reference for the reduction potentials is
called
standard
hydrogen electrode
(SHE)
the reduction potential
for hydrogen
0 V
The cell potential at standard conditions can be found through this equation:
E°cell = E°red(cathode) - E°red(anode)
cell potential is based on the potential
energy per unit of charge
intensive property
The more positive the
value of E° red
the greater the tendency
for reduction
most positive
reduction potentials
The strongest
oxidizers
most negative
reduction potentials
The strongest
reducers
produce a positive
cell potential, or emf.
Spontaneous redox reactions
Equation for spontaneous redox reactions
E° = E° red (reduction) – E° red (oxidation)
positive emf corresponds to
negative ΔG
Since Gibbs free energy is the measure of
spontaneity
How is Free Energy, Redox, and K related
ΔG° = –nFE° = –RT ln K
Nernst equation
E = E° – (RT/nF) ln Q
OR E = E° – (2.303 RT/nF) log Q
Nernst equation (Using standard thermodynamic temperature and the constants R and F)
E = E° – (0.0592/n) log Q
a cell could be created that has the same substance at both electrodes
concentration cell
as long as the concentrations are
different
E will not be 0
Some Applications of Cells
– Batteries
– Prevention of corrosion
– Electrolysis
Examples of batteries
- Lead–acid battery
- Alkaline battery
- Ni–Cd and Ni–metal hydride batteries
- Lithium-ion batteries
reactants and products are
solids, so Q is 1 and the potential is independent of concentrations; however, made with lead and sulfuric acid (hazards)
- Lead–acid battery
most common primary battery
- Alkaline battery
lightweight, rechargeable; Cd is toxic and heavy, so hydrides are replacing it
- Ni–Cd and Ni–metal hydride batteries
rechargeable, light; produce
more voltage than Ni-based batteries
- Lithium-ion batteries
the energy created can be
converted to electrical energy
fuel is burned ( this conversion is only 40% efficient, with the remainder lost as heat)
The direct conversion of chemical to electrical energy is expected to be more efficient and is the basis for it
fuel cells
are NOT batteries; the source of energy
must be continuously provided
fuel cells
- In this cell, hydrogen
and oxygen form water. - The cells are twice as
efficient as combustion. - The cells use hydrogen
gas as the fuel and
oxygen from the air.
Hydrogen Fuel Cells
it is an oxidation and its common name is rusting
Corrosion
occurs because zinc is more easily
oxidized, so that metal is sacrificed to keep the iron from rusting
Cathodic protection
A method to prevent corrosion is used for
underground pipes wherein the anode is oxidized before the pipe
sacrificial anode
Use of electrical energy to create chemical
reactions
electrolysis
1 coulomb
1 ampere × 1 second
equation for Q (relation to electrolysis and stoichiometry)
= It = nF
Q represents
charge (C)
I represents
current (A)
t represents
time (s)
n represents
moles of electrons that
travel through the wire in
the given time
NOTE: n is different than that
for the Nernst equation!
F represents
= Faraday’s constant (96,485 C/mol)
