Electrochemistry Unit Summary:
Introduction and the Basics
Electrochemistry is the study of electron transfers in chemical reactions (Redox Chemistry)
REDUCTION  gain of electrons
OXIDATION  loss of electrons
Redox reactions are involved in many processes such as: cellular respiration, photosynthesis,
combustion, corrosion and electroplating metals.
Reducing agents promote oxidation and oxidizing agents promote reduction.
LEO the lion says GER ... RAO/OAR
LEO = lose electrons oxidation
| GER = gain electrons reduction
RAO = reducing agent undergo oxidation | OAR = oxidizing agent undergo reduction
Balancing Redox Reactions using Complex Half-Reactions:
 Polyatomic ions and molecular compounds often undergo more complex oxidation and reduction.
They will most often occur in an aqueous environment that is acidic [H+]
 Steps to balance:
1) Write out the Reactant  Products equation of what is undergoing change.
2) Balance all species other than oxygen [O] and hydrogen [H].
3) Balance oxygen’s by adding H20 to appropriate side.
4) Balance hydrogen’s by adding H+ to appropriate side.
5) Balance charge by adding electrons to appropriate side.
6) Once you have made both half-reactions, balance the electrons and get net reaction.
Remember that the electrons must cancel out from the 2 half-reactions.
Building Redox Tables:
 By conducting a series of experiments testing the relative oxidizing and reducing strengths of
various substances, a redox table can be constructed.
 A redox table lists half-reactions by placing the strongest oxidizing agent (SOA) at the top of the
table and the weakest oxidizing agent (WOA) at the bottom of the table.
 In a redox reaction, electrons are transferred from RA to OA. This transfer is a competition for
electrons; if a transfer is successful a spontaneous reaction occurs.
 The OA or RA with MOST SPONTANEOUS reaction is strongest.
 Spontaneous Rule: On a table of half-reactions if OA is above the RA, the reaction will be
spontaneous. If OA is under RA, the reaction is non-spontaneous.
Predicting Redox Reactions:
 A redox reaction involves a competition for electrons and when multiple OA’s and RA’s exist, the
strongest OA will react with strongest RA.
 Steps for predicting spontaneity:
1) List all entities present and classify them as a possible OA, RA or both.
Note: Aqueous solutions contain water.
If acidic solution, hydrogen ions are present.
If basic solution, hydroxide ions are present.
Ionic compounds dissociate
Soluble molecular compounds DO NOT dissociate.
2) Find the strongest oxidizing agent and write out the reduction half-reaction.
3) Find the strongest reducing agent and write out the oxidation half-reaction.
4) Balance electrons and determine net reaction.
5) Predict spontaneity.
 Disproportionation (auto oxidation): A redox reaction in which one species is BOTH oxidized and
reduced.
Oxidation States:

The apparent net electrical charge an atom would have if electron pairs in covalent bonds
belonged entirely to the more electronegative atom.
Common Rules:
Atom/ Ion:
Oxidation number:
Examples:
Elements
0
0
Mg(S)
0
Ns(G)
+1 -1
Hydrogen
+1
Oxygen
―2 / +2
Monoatomic ions
charge
HCl
*Exception is when hydrogen is bonded to a
less EN atom
―2
+2
CO
OF2
*Exception is in peroxides; oxygen would
then have oxidation number of ―1; and when
oxygen is bonded to group 17 atoms then
oxygen has oxidation number +2.*
+1
Na+
Special Notes:
 After assigning common oxidation numbers, you must find the remaining
oxidation numbers for entities.
 A compound / ion’s total sum of oxidation numbers should equal its
charge.



If oxidation number of an atom or ion changes during a chemical reaction, then an electron
transfer has occurred.
An increase in oxidation # == OXIDATION
A decrease in oxidation # == REDUCTION
Balance Redox Reactions using oxidation #s:
1)
2)
3)
4)
5)
Determine all oxidation numbers and identify the atoms that are changing.
Determine the # of electrons transferred per atom / ion.
Record the number of electrons per molecule.
Balance electrons gained or lost by adding coefficients.
Balance remaining including addition of [H2O] and [H+] if required.
Redox Stoichiometry:
o Recall the titration process whereby a titrant is slowly added to a sample until an abrupt change in
property occurs.
o In redox titration the titrant is often a strong OA or RA.
o Indicators are not required as color of reactants and products indicate the change.
o The volume of titrant required to reach the endpoint is called the equivalence point.
Electrochemical and Voltaic Cells:
Primary Cells: cells that cannot be recharged. (dry cell, mercury cell- such as Energizer/ Duracell
batteries found in the store)
Secondary Cells: cells that can be recharged by using electricity to reverse the chemical reaction.
(Ni-Cd / Ni-MH / Lithium ion)
Fuel Cells: Produce electricity by reaction of a fuel that is continuously supplied to keep the cell
operating. (Hydrogen – Oxygen / Aluminum – Air cells)
A voltaic cell is an electric cell that spontaneously produces energy. It has two electrodes, an
anode and a cathode. The anode is place where oxidation occurs and reduction occurs at the
cathode.
Sometimes we need an inert electrode for voltaic cell for the desired reaction to take place.
An inert electrode is an unreactive solid conductor in a cell that provides a location to connect a
wire and a surface on which a half-reaction can occur. Commonly used inert electrodes are a
carbon (graphite) rod and a platinum metal foil.
General Cell Notation:
Beaker 1
Beaker 2
Anode (―) | electrolyte || electrolyte | cathode (+)
←anions
cations→
Example: copper and zinc cell
Cu(S) | Cu(SO4)(aq) || Zn(NO3)2 (aq) | Zn(S)
 If we were to keep the copper-zinc cell running for many hours, we will find 4 thing:
 Copper solution will appear to have a lighter blue color
 Copper solid will become larger
 Zinc solid will be smaller
 Voltmeter would show a voltage being produced
Therefore, we know that the redox reaction will have the strongest oxidizing agent (SOA) will react
with the strongest reducing agent (SRA).
SOA = Reduction occurs at the cathode (+)
SRA = Oxidation occurs at the anode (―)
Electrons always flow from anode (―) to cathode (+).
The U-Tube or porous cup allows ions to flow between the half-cells to maintain electrical
neutrality.
Standard Cell Potentials:
 Standard Cell: a voltaic cell in standard conditions (SATP and 1.0 mol/L)
 Standard Reduction Potential: the ability of a standard half-cell to attract electrons, thus
undergoing a reduction.
E°cell
net




E°RC ― E°RA
cathode anode
E° stands for Electric Potential
A positive E°cell means cell reaction is spontaneous
A negative E°cell means cell reaction is non- spontaneous
If a question asks what would happen is hydrogen is not the reference half-cell, two things apply:
 If looking for cell potential of voltaic cell, it will not change if reference half-cell
changes.
 If looking for a reduction potential for a half-reaction using a different reference,
you will need to find it as it will be different.
Corrosion n’ rust:
 Corrosion is an electrochemical process in which a metal reacts with substances in the
environment returning the metal to an ore-like state.
 An example is a freshly cleaned surface of aluminum rapidly oxidizing in air to form aluminum
oxide.
 Corrosion is such an important problem in our society that many technologies have been
developed and continue to be developed to minimize the problems.
 PROTECTIVE COATING: paint and other similar coatings are a simple method of
corrosion prevention. This method works well as long as the surface is completely
covered and the coating remains intact.
 CATHODIC PROTECTION: the iron is forced to become the cathode by supplying the
iron with electrons using either an impressed current or sacrificial anode.
Electrolytic Cells:
 Electrolytic cells are a combination of 2 electrodes and only 1 electrolyte and are nonspontaneous.
 These cells use a process called electrolysis, which is a process of supplying electrical energy to
force a non- spontaneous redox reaction to occur.
 The external power supply (such as a battery) acts as an “electron pump” and the electric energy
is used to do work on the electrons to cause an electron transfer inside the electrolytic cell.
Comparing Voltaic Cells and Electrolytic Cells:
Voltaic Cells:
Electrolytic Cells:
Spontaneous Reaction
Non- spontaneous Reaction
Positive (+) E°cell
Negative (―) E°cell
Exothermic  produces electricity
Endothermic  absorbs electrical
energy from a power supply
Cathode  Strongest oxidizing agent
Cathode
Anode
 Strongest oxidizing agent
present undergoes reduction
 Positive electrode
 Strongest reducing agent
present undergoes oxidation
 Negative electrode
Electron move: ANODE  CATHODE
Anode
present undergoes reduction
 Negative electrode
 Strongest reducing agent
present undergoes oxidation
 Positive electrode
Electron move: ANODE  CATHODE
Ion Movement: ANIONS  ANODE
CATIONS  CATHODE
Ion Movement: ANIONS  ANODE
CATIONS  CATHODE
The
Chloride Anomaly:
When your only reducing agents are chloride and water, even though water is a stronger reducing agent
(SRA), chloride actually REACTS FIRST!!! This is called over-voltage or The over Potential Effect.
Electrolysis Application:
1) METAL PRODUCTION
o Most elements exist in nature combined with other elements in compounds. So to produce
different metals, electrolysis of metal ion is used.
o Many ionic compounds have low solubility in water and are often weaker OA’s than water
itself.
o To get around the problem, you can take insoluble ionic compounds and melt them into
molten rock, but for some compounds, the temperature is very high!
2) CHLOR-ALKALI PROCESS
o The electrolysis of brine (sodium chloride) to produce chlorine gas and sodium hydroxide.
3) REFINING METALS
o When producing metals, the initial product is impure as there are other metals present in
the ore. Electrolysis is used to purify the metal and is called electro-refining.
4) ELECTROPLATING
o This is the process of plating one metal on top of the other metal in order to prevent
corrosion. This can also make the products look nicer and end up selling for a lower cost.
Cell Stoichiometry:
In electrolysis, the quantity of electricity that passes through a cell determines the mass of
substances that react or are produced at the electrodes.
Faraday’s constant relates charge transferred to mass of e-: F = 9.65 x 104 C/ mol eFormulas: q = It where q is charge; I is current (A for amperes); and t is time (s)
ne
where ne is # of electrons; and F is Faraday’s constant
― END OF ELECTROCHEMISTRY UNIT SUMMARY ―
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