Study Guides
Big Picture
Electrochemistry is an important field of chemistry that affects much of our everyday lives, from batteries to electric
motors. The type of reaction that occurs in electrochemistry is called a redox reaction, in which electrons are exchanged
between atoms. Many chemical reactions are redox reactions; all single-replacement reactions are redox, as well as
some double replacement reactions.
Key Terms
Oxidation: The process of losing electrons (becoming
more
positively
or
less
negatively
charged).
Oxygen is the most famous oxidizer, as it tends to
Half-Reaction: An equation showing just the oxidation
or just the reduction process of a redox reaction.
Electrochemical Cell: A device that converts chemical
take electrons from other elements, thus providing
energy into electrical energy or vice versa.
the origin of the term.
Anode:
Reduction: The opposite of oxidation, it is the process
Reaction:
A
Where
oxidation
takes
place
in
an
electrochemical cell.
Cathode: Where reduction takes place in an
of gaining electrons.
Redox
Chemistry
Electrochemistry
chemical
reaction
involves
electrons being transferred from one substance to
another. Involves both oxidation and a reduction
reactions.
electrochemical cell.
Salt Bridge: Links the half-cells and allows a salt to
travel between cells to maintain charge equilibrium.
Voltage: Strength of a redox reaction. Unit: volt
Oxidizing Agent: The substance that is reduced;
Standard Reduction Potential (E): The voltage
in other words, it aids the oxidation of the other
associated with a reduction half-reaction. Since
substance.
oxidation must occur for a reaction to happen, the
Reducing Agent: The substance that is oxidized.
oxidation of hydrogen (H++e-
Oxidation State: A measure of the degree to which an
standardize all measurements.
½ H2) is used to
atom in a substance is oxidized.
Redox Reactions
Oxidation and reduction always occur together, so reactions involving these processes are called redox reactions.
In a redox reaction, the oxidizing agent will oxidize another substance, the reducing agent.
To remember oxidation and reduction reactions, remember the phrase: “LEO the lion says GER”, which stands
for Losing Electrons is Oxidation, Gaining Electrons is Reduction.
Redox reactions result in a complete transfer of electron in ionic reactions. When the compounds are covalent, the
electrons are shifted away (oxidation) or toward (reduction) an atom in a covalent bond.
Oxidation States
Oxidation states help us determine which atom in a reaction
has oxidized and which has reduced. The oxidation state is a
hypothetical charge that an atom would have if every atom in
a substance were purely ionic.
• Oxidation
state of diatomics: Atoms in diatomic
substances, such as O2, H2, and N2 have an oxidation
state of 0.
state of -2, except in hydrogen peroxide (H2O2), where it
has an oxidation state of -1, and O2 , where it is 0.
• Oxidation
state of hydrogen: Hydrogen always has an
oxidation state of +1, except in alkali hydrides, such as
LiH, where it has a state of -1.
• All
other elements: Can have varying oxidation states,
both negative and positive. Often, it will be the charge
the atom would form if it were purely ionic. (ex: nitrogen
= -3).
• The
sum of the oxidation states of all the elements in a
compound must equal the total charge of the compound.
For neutral substances, it must equal 0.
Oxidation states can be used to balance redox
reactions.
1.Assign oxidation numbers to all the atoms in the
equation.
2.Identify which atoms are oxidized and which are
reduced.
3.Determine the oxidation number change for the
oxidized species and for the reduced species.
4.The total increase in oxidation number for the
oxidized species needs to equal the total decrease
in oxidation number for the reduced species. Use
appropriate coefficients so that this is true.
5.Make sure the equation is balanced for both
atoms and charge.
This guide was created by Steven Lai, Rory Runser, and Jin Yu. To learn more
about the student authors, visit http://www.ck12.org/about/about-us/team/
interns.
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v1.11.4.2011
Disclaimer: this study guide was not created to replace
your textbook and is for classroom or individual use only.
• Oxidation state of oxygen: Oxygen always has an oxidation
An increase in oxidation number indicates oxidation,
while a decrease indicates reduction. A redox
reaction can be identified because the oxidation
number of the reacting species changes.
Chemistry
Electrochemistry
cont .
Half-Reactions
For some redox reactions, it is not easy to balance by oxidation states. In these cases, use half-reactions to balance
the reaction.
1.Write the unbalanced equation in ionic form.
2.Write half-reactions for oxidation and reduction.
3.Balance the atoms in each half-reaction.
4.Add enough electrons to one side of each half-reaction to balance charges.
5.Multiply each half-reaction so that the number of electrons in each reaction are equal.
6.Add the half-reactions together.
7.Add the spectator ions and balance the equation.
Electrochemical Cell
Redox reactions occur in electrochemical cells. The
oxidation and reduction reactions needs to be physically
separated if a redox reaction is used to provide electrical
energy.
Voltaic cells are electrochemical cells where electrical
energy is produced by spontaneous redox reactions.
• A
half-cell is one part of the voltaic cell where either
the oxidation or the reduction reaction occurs. The
anode is a metal where oxidation reaction occurs,
and the cathode is a metal where reduction reaction
occurs.
Remember the phrase “An Ox Red Cat.” Anode =
Oxidation and Reduction = Cathode.
• The
two half-cells are then connected with a salt
bridge that allows electrons to flow to maintain
charge equilibrium. The electrons flow from the
anode, where electrons are produced, to the cathode,
where electrons are consumed.
Below is a standard setup of a voltaic cell. Two metals,
zinc and copper, are used. Copper is reduced, so it is
at the cathode, while zinc is oxidized at the anode. The
two half reactions are shown beneath each cell, and the
voltmeter in the middle measures the total voltage. The
nitrate ion (NO3-) is used to balance charge, and moves
through the salt bridge.
Figure: A zinc copper battery
Image Credit: Rory Runser, CC-BY-SA 3.0
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A battery is a group of voltaic cells connected together.
A fuel cell is similar to a battery, except it consumes fuel
from an environment. Thus, it is an open system, unlike
a battery, which is closed. The most common fuel cell is
a hydrogen fuel cell, which uses hydrogen and oxygen to
produce water and energy.
An electrochemical cell can also be used to convert
electrical energy into chemical energy. This type of
electrochemical cell is an electrolytic cell. Unlike a voltaic
cell, electrons flow as a result of an outside power source.
Electroplating is a process that uses an electrolytic cell to
deposit a thin layer of metal on an object.
Electrical Potential
The voltage is a measure of the electrical potential,
or the cell’s ability to produce an electric current. The
standard reduction potential is the tendency for a
half-reaction to be the reduction process. The difference
between the reduction potential of the half-cell where
reduction occurs and the half-cell where oxidation occurs
is the cell potential, which can be measured.
Since temperature and concentrations of the ion
solutions affect the reaction rate and the voltage of the
cell, the standard reduction potential is chosen so that
the temperature is 25° C, the ion concentrations are 1 M,
and gases are at 1 atm.
Reduction half-reactions can be organized into a redox
table, where the half-reactions are listed in order of
increasing standard reduction potentials (decreasing
oxidation potentials). If the difference of the potentials
of one reduction half-reaction and one oxidation halfreaction (flip the sign on the reduction potential) is
positive, a reaction will occur. If not, no redox reaction
will occur.
Basic Redox Reactions
Example: Balance the basic redox reaction
copper, while aluminum oxidizes to form aluminum(III) ions.
, where copper(II) ions reduce to form solid
1.Write the half-reactions, including electrons:
• • 2.In order for the electrons to cancel out, multiply both half-reactions by the smallest common multiple of the number
of electrons:
• • (multiply by 3)
(multiply by 2)
3.Add the reactions together and cancel the electrons:
• • Determining Oxidation States
Determining an atom’s oxidation state is important for determining which atoms become oxidized and reduced.
Simple Problem
Example: Given the molecule CaO, find the oxidation state of each atom.
1.Oxygen has an oxidation state of 2-.
2.The overall charge of the atom is zero, so calcium has a state of 2+.
More Difficult Problems
Example: Given the molecule KClO4, find the oxidation state of each atom.
1.Since this is not a peroxide, we know the oxidation state of each oxygen is 2-. This gives a total oxidation so far of
-8, since there are 4 oxygens.
2.Since this is an ionic substance, we know we could break this down into K+ and ClO4-. Thus, the oxidation state of
the potassium must equal its ionic charge of 1+.
3.Since the overall molecule is neutral, the oxidation state of chlorine must equal 7+, since 1 + 7 - 8 = 0.
Example: Given the molecule MgBr2, find the oxidation state of each atom.
1. Here, we have no hydrogen or oxygen to help us get started. In this case, it is reasonable to assume that the
oxidation states of atoms will be equal to the charge of their most common ionic form. Bromine tends to form
ions with a 1- charge, and magnesium forms ions with a 2+ charge.
2.Assign magnesium a 2+ charge and each bromine a 1- charge for a total charge of 2- on bromine and 2+ on
magnesium. Adding the charges, 2 - 2 = 0, proves that this is a neutral compound.
Calculating Voltage
Knowing the voltage of a redox reaction is important in making batteries.
Example: Calculate the voltage of an electrochemical cell reaction between aluminum and copper given these
standard reduction potentials: E = 0.34 V for Cu2+ + 2e- = Cu and E = -1.66 V for Al3+ +3e- = Al. Determine which
reactions occur at the cathode and anode.
1.A redox reaction will only occur if the overall voltage is positive. One of these reactions needs to be reversed.
Reversing the reaction of copper will yield an oxidation potential of -0.34 V, and an overall potential of -1.66 V- 0.34
V = -2.0 V. This fails.
2.Try reversing aluminum and having it oxidize instead. Here, voltage = 0.34 V + 1.66 V = 2.0 V. Success!
3.Since copper is reducing, it occurs at the cathode, and aluminum oxidizes at the anode.
Even though the reaction involves 2 aluminum and 3 copper, do NOT multiply the standard reduction potentials
by 2 and 3, respectively. They remain as is.
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Chemistry
Electrochemistry Problem Guide
Chemistry
Electrochemistry Problem Guide
cont .
Redox Reactions in Acid/Base Solutions
Many times, a redox reaction will occur in a solution, where water and either H+ or OH- ions are present in abundance.
These ions, and water, can be used to balance out a redox reaction.
Example: Balance the reaction
that occurs in an acidic solution.
1.Write the oxidation states for all the elements, including those attached to other elements. Oxidation states are
written as exponents for the elements.
2.Write the half-reaction for the molecule that oxidizes.
3.Write the half-reaction for the molecule that reduces.
Note: Here, 5 electrons are added to the left because the magnesium atom gains five electrons during the reaction
(from an oxidation of +7 to +2, 5 electrons are gained)
4.Multiply each half-reaction to balance the number of electrons.
5.Add the half-reactions together, and cancel out the electrons.
6.Balance the charge by adding hydrogen or hydroxide ions, depending on the problem. In this case we add H
because it is in an acidic solution. The left side has a total charge of 1-, and the right side has a total charge of 7+.
Thus, you add 8H to the left side, making the charge 7+ on each side. If this were in a basic solution, you would
add 8OH- to the right side instead.
7.Now you need to finish balancing the chemical reaction. Add to either side to finish balancing the equation. In this
case, adding 4s to the right side balances both the hydrogens and the oxygens in the permanganate ion (if you did
it correctly, this should always happen).
8.Check one more time to make sure all the atoms and charges are balanced.
9.You’re done!
Notes
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