ELECTROCHEMISTRY
Is the branch of chemistry that deals with the interconversion
of electrical energy and chemical energy.
It deals directly with electrons and their movement.
Whether the electrons move under the driving forces of
spontaneous chemical processes or respond to an applied
external electric potential.
KEY TERMS
 Oxygen almost always has an oxidation number of -2,
except in peroxides (H2O2) where it is -1 and in compounds
with fluorine (OF2) where it is +2.
 Hydrogen has an oxidation number of +1 when combined
with non-metals, but it has an oxidation number of -1 when
combined with metals.
 The algebraic sum of the oxidation numbers of elements in a
compound is zero.
 The algebraic sum of the oxidation states in an ion is equal
to the charge on the ion.
ELECTROCHEMICAL PROCESSES are (redox) oxidationreduction reactions in which the energy released by a
spontaneous reaction is converted to electricity
REDOX(oxidation-reduction) reactions-occurs when electrons
are transferred from an atom that is oxidized to an atom that is
reduced.
KEY TERMS
OXIDATION REACTION refers to the half-reaction that
involves the loss of electrons.
REDUCTION REACTION refers to the half reaction that
involves the gain of electrons.
OXIDIZING AGENT, the one that accepts electrons; the one
that is reduced
REDUCING AGENT, the one that donates electrons the one
that is oxidized
OILRIG Oxidation is loss (of electrons). Reduction is Gain (of
electrons).
OXIDATION NUMBERS
Signifies the number of charges the atom would have in a
molecule or in an atomic compound if electrons were
transferred completely.
LEOGER Loss of Electron – Oxidized. Gain electron –
Reduced.
SAMPLE PROBLEM
Defined as number of charges the atom would have in a
molecule (or an ionic compound) if electrons were transferred
completely.
RULES IN ASSIGNING OXIDATION STATE:
 The oxidation number of a free element is always 0
 The oxidation number of a monatomic ion equals the charge
of the ion.
 Fluorine in compounds is always assigned an oxidation
number of -1.
 The alkali metals (group 1 ) always have an oxidation
number of +2.
a. =+1Li2 -2O b.+1HxN-2O3
c. xCr22-O7
1+x+(-2x3)=0 2x+(-2x7)=-2
1+x-6=0
x=5
2x-14=-2
2x=+12
x=+6
BALANCING REDOX REACTIONS BY THE METHOF OF
HALF-REACTIONS
A. balancing of redox reaction that occurs in an acidic
aqueous solution.
1. Divide the equation into one oxidation half-reaction and one
reduction half-reaction.
2. Balance each half-reaction.
SAMPLE PROBLEM
Identify the oxidizing and reducing agents in the reaction:
3. Multiply half reactions by integers as needed to make the
number of electrons lost in the oxidation half-reaction equal the
number of electrons gained in the reduction half-reaction.
4. Add half-reactions and, if possible, simplify by cancelling
species appearing on both sides of the combined equation.
5. Check to make sure the atoms and charges are balanced.
a. Balance elements other than H and O.
b. Balance O atoms aby adding H2O as needed.
c. Balance H atoms by adding H+ as needed
d. Balance charge by adding e- as needed.
REDOX REACTIONS
Considered as electron transfer reactions. Oxidation –
reduction reactions are very much a part of the world around
us. They range from the burning of fossil fuels to the action of
household bleach.
Half Reactions – equations that show either oxidation or
reduction alone.
B. balancing of redox reaction that occurs in a basic
aqueous solution.
1. Divide the equation into one oxidation half-reaction and one
reduction half-reaction.
2. Balance each half reaction.
a. Balance elements other than H and O.
b. Balance O atoms by adding H2O as needed.
c. Balance H atoms by adding H+ as needed.
d. Balance charge by adding e- as needed.
READOX REACTIONS
e. Count the number of H+ in each half-reaction and add the
same number of OH-to each side of the reaction.
f. Neutralize H+ and OH-by forming H2O when they are on the
same side of either half-reaction.
g. Cancel water molecules that both appear in the product and
reactant sides.
SAMPLE PROBLEM
Balance the following equation in an acidic medium:
SAMPLE PROBLEM
Balance the following equation in a basic medium:
BALANCING REDOX REACTIONS
1. The amount of each element must be the same on both
sides of the equation.
2. The gains and losses of electrons must be balanced.
ELECTROCHEMICAL CELLS
Systems that incorporate redox reaction to produce or utilize
electrical energy.
TYPES OF ELECTROCHEMICAL CELLS:
1. VOLTAIC (GALVANIC CELL): cells that uses spontaneous
reactions resulting from the difference in chemical potential
energy between higher-and lower-energy products to
generate electrical energy
2. ELECTROLYTIC CELL: uses electrical energy to drive
nonspontaneous reaction. The electrical energy from the
external source converts lower energy reactants to higher
energy products.
PARTS OF ELECTROCHEMICA CELLS
1.ELECTRODES: conduct electricity between the cell and its
surroundings
a) ANODE: the cell at which the oxidation half-reaction occurs
b) CATHODE: e cell at which the reduction half-reaction
occurs
2. ELECTROLYTES: a mixture of ions usually in aqueous
solution that are involved in the reaction or that carry the
charges. This is where the electrodes are dipped.
GALVANIC CELLS
•The experimental apparatus for generating electricity through
the use of a spontaneous reaction.
•The anode in a galvanic cell is the electrode at which
oxidation occurs. The cathode is the electrode at which
reduction occurs.
SAMPLE PROBLEM
•The two half-reaction in a voltaic cell are:
(a) Indicate which reaction occurs at the anode and which at
the cathode.
an: Zn -> Zn2- + 2ecat: ClO3- + 6H+ + 6e- -> Cl- + 3H2O
(b) Which electrode is consumed in the cell reaction?
The electrode that is consumed is the anode
CELL DIAGRAM
•are the conventional notation used for representing galvanic
cells.
Salt v bridge
CELL POTENTIALS UNDER STANDARD CONDITIONS
•Potential energy is higher at the anode than at the cathode;
thus, electrons flow spontaneously toward the electrode with
the more positive electrical potential.
•Anions always migrate toward the anode and cations toward
the cathode through an external circuit.
•One volt (V) is the potential difference required to impart 1
Joule of energy to a charge of 1 coulomb, C
•Because of this directional flow, the anode in a voltaic cell is
labeled with a negative sign and the cathode is labeled with a
positive sign.
•Cell potential (Ecell) is the potential difference between the
two electrodes of a voltaic cell. Also called the electromotive
force, emf and voltage of the cell.
•The magnitude of the cell potential depends on the reactions
that occur at the cathode and anode, the concentrations of
reactants and products and the temperature.
STANDARD CELL POTENTIAL
•E°cell, is the cell potential under standard conditions (1M
concentrations of products and reactants in solution, 1 atm
pressure and 25°C.)
Note:
•Standard reduction potentials are intensive properties.
•The more positive the value of E°red, the greater tendency for
reduction under standard conditions.
DETERMINING THE STANDARD REDUCTION POTENTIAL
THE DANIELL CELL
SAMPLE PROBLEM
•determine which of the following reactions are spontaneous
under standard conditions:
A.
B.
C.
EMF AND FREE ENERGY
∆𝐺=−𝑛𝐹𝐸
Where: n, is a positive number without units that represents the
number of moles of electrons transferred according to the
balanced equation for the reaction F, is the Faraday’s constant,
96,485 C/mol = 96,485 J/V mol.
EMF, FREE ENERGY WITH THE EQUILIBRIUM CONSTANT
(K)
•At standard conditions,
where K is:
where K = equilibrium constant
a, b, c and d = stoichiometric
coefficients
•At 𝑇=298𝐾,
FREE ENERGY AND REDOX REACTIONS
•A positive value of Eo indicates a spontaneous process; a
negative value of Eo indicates a non-spontaneous process.
Base-10 logarithm of 𝐾,
•*Eo is used to represent the emf under non-standard
conditions.
SAMPLE PROBLEM
For the reaction:
a) What is the value of “n”?
b) Calculate ΔG°
c) Calculate K at T=298K
SAMPLE PROBLEM
In a certain experiment, the 𝒆𝒎𝒇(𝑬)of the cell is found to be
0.54𝑉at 25°C. Suppose that [𝑍𝑛2+] =1.0𝑀 and 𝑃H2=1.0𝑎𝑡𝑚.
Calculate the molar concentration of 𝐻".
THE NERNST EQUATION
BATTERY
•is a galvanic cell, or a series of combined galvanic cells, that
can be used as a source of direct electric current at a constant
voltage. A battery has the advantage of being completely selfcontained and requiring no auxiliary components such as
salt bridges.
DRY CELL BATTERY
A cell without a fluid component, is
the Leclanché cell used in flashlights
and transistor radios.
SAMPLE PROBLEM
The electrolyte consists of
ammonium chloride and zinc chloride
in water, to which starch is added to
thicken the solution to a paste like
consistency so that it is less likely to leak.
The voltage produced by a dry cell is 1.5 V.
MERCURY BATTERY
•The Mercury battery is used
extensively in medicine and
electronic industries and is more
expensive than the common dry cell.
Contained in a stainless steel cylinder,
the mercury battery consists of a zinc
anode (amalgamated with mercury) in
contact with a strongly alkaline
electrolyte.
•The advantage of the battery is that lithium has the most
negative standard reduction potential.
•the mercury battery provides a more constant voltage (1.35 V)
than the Leclanché́ cell.
•It also has a considerably higher capacity and longer life
which makes it ideal for use on pacemakers, hearing aids,
electric watches, and light meters.
LEAD ACID BATTERY
A 12-V lead-acid automotive battery consists of six voltaic cells
in series, each producing 2V. The cathode of each cell is lead
dioxide (PbO2) packed on a lead grid. The anode of each cell is
lead. Both electrodes are immersed in sulfuric acid.
•A lithium-ion battery can be recharged literally hundreds of
times without deterioration. These desirable characteristics
make it suitable for use in cellular telephones, digital cameras,
and laptop computers.
HYDROGEN OXYGEN FUEL CELLS
•A fuel cell, a galvanic cell that requires a continuous supply of
reactants to keep functioning.
•A hydrogen-oxygen fuel cell consists of an electrolyte solution,
such as potassium hydroxide solution, and two inert
electrodes.
•A hydrogen-oxygen fuel cell is usually used in space
exploration programs. The pure water produced by the cell is
consumed by the astronauts.
•Unlike the Leclanché cell and the mercury battery, the lead
storage battery is rechargeable.
•Recharging the battery means reversing the normal
electrochemical reaction by applying an external voltage
at the cathode and the anode.
LITHIUM ION BATTERY
The anode is made of a conducting carbonaceous material,
usually graphite, which has tiny spaces in its structure that can
hold both Li atoms and 𝐿𝑖+ ions. The cathode is made of a
transition metal oxide such as CoO2, which can also hold 𝐿𝑖+
Ions.
CORROSION
 it is the degradation of metals by chemical reactions with the
environment. Corrosion generally involves a slow combination
of oxygen with metals to form oxides.
 Uniform corrosion -one of the most visible forms of corrosion.
 Galvanic corrosion-which occurs only when two different
metals contact each other in the presence of an appropriate
electrolyte.
 Crevice corrosion-When two pieces of metal touch each
other, they tend to leave a small gap (unless the joint is
covered with a coating, such as paint). At that gap, or crevice,
the metals are more likely to corrode.
MECHANISM OF CORROSION