Electrochemistry
“One day sir, you may tax it”
- Michael Faraday,
in response to a question posed about the
practical uses of electricity.
Chapter Overview and
Introduction
Chapter Learning Goals
Balancing redox reactions (these are fun ).
Finding voltages, gibbs free energy and equilibrium constants
for electrolytic and galvanic cells.
Relate standard and non-standard voltages
Determining what will and will not undergo redox reactions
And much much more!!!
Real Life Explanation Goals
Fuel cells
Corrosion
Electroplating
If we have time: battery alternatives:
bacteria, viruses and other cool
things
Electrochemistry: A Definition
Electro chemistry is the intersection of electrical and chemical
energy
One species is “reduced”, another is “oxidized”- a transfer of
electrons from one (or more) atoms to another.
Zn(s)+Cu2+(aq) Zn2+ (aq)+ Cu(s)
Electrons go from the Zn to the Cu2+
Electrochemistry
Definitions
Learning Outcomes
Define oxidation number, reduction, oxidation, reducing
agent and oxidation agent
Identify the oxidation number for species
Identify the species being oxidized
Identify the species being reduced
Identify the reducing agent
Identify the oxidation agent
Before moving on to harder things, make sure being able
to do these things is automatic!!!!
Definitions
Oxidation number: “imaginary charge if the compound
was broken down into ions”
Oxidation number of atom in its elemental state is zero
Oxidation number of a monatomic ion is equal to its charge
Oxidation numbers of individual atoms must add up to the
charge on the whole molecule.
Use this to find oxidation states of ions you don’t know, or that may have
multiple possibilities.
+0
+0
+1
-1
Oxidation state examples:
MnO4-
Na2CrO4
Na= +1 (2)= +2
O= -2 (4)= -8
-6
Neutral compound
so Cr=+6
O=
-2 (4)= -8
-8
Charge is -1, so
Mn=+7
Definitions
Oxidized: Losing e-, increased charge Reduced: Gaining e-, lowered charge
Oxidizing Agent: Species which oxidizes other compounds (and is thereby reduced)**
Reducing Agent: Species which reduces other compounds (and is thereby oxidized)**
Anode: Where oxidation takes place
+0
Oxidized
Reducing Agent
Anode
Cathode: where reduction takes place
+0
Memory Tricks:
LEO the lion goes GER
RED CAT / AN OX
+1
-1
Reduced
Oxidizing Agent
Cathode
**Note: The whole compound is the oxidizing/reducing agent.
Learning outcomes
Learn to balance redox reactions in acidic or basic
solutions
Using half reaction method
Leaving the reactions together
You should know how to do both methods:
Leaving reaction together is always possible, but
sometimes more difficult
Half reaction method is sometimes not possible.
Redox Reactions- Balancing
Some methods are easier for different situations, be
open to switching ways!
Two main methods
Half reaction method: couple of different algorithms, pick one
and stick with it. (I picked the one I think is easiest for students)
Leaving the reaction together: great for when the species can’t
be separated. Sometimes its just easier.
there will be one where you HAVE to do this on the exam.
Half method rules: Use this as a checklist
Step 1: write oxidation numbers for every species.
Step 2: identify which is being reduced and which is being oxidized
Step 3: separate into two separate half reactions, do steps 4-7 for each.
Step 4:First balance all elements EXCEPT
H and O
Step 5: Add H2O to balance oxygen
Step 6: Add H+ to balance H
If it’s a basic solution add as many OH- to each side as needed to cancel
H+ into H2O (can also be done after steps 7, 8, or 9, instead. Do wherever
is most convenient)
Step 7: add electrons to balance charges (oxidation states)
Step 8: Look at both half reactions and multiply each as needed to get
the number of electons on both sides to be equal.
Step 9: Combine reactions together.
Examples:
Balance the following, identify which species is oxidized, which is reduced,
which is the oxidizing agent and which is the reducing agent.
Acid solution
Basic
Balancing without separating
Step 1: Write oxidation numbers.
Step 2: Balance all species that ARE NOT H or O
Step 3: Identify which species is being oxidized and which is
reduced.
Step 4: Decide how many electrons are transferred in each.
Step 5: Multiply coefficients as needed to get the electrons
transferred in the Red/Ox to be equal
Step 6: Balance O by adding water
Step 7: Balance H by adding H
If it’s a basic solution add as many OH- to each side as needed to
cancel H+ into H2O
Examples:
Balance the following, identify which species is oxidized, which is reduced,
which is the oxidizing agent and which is the reducing agent.
Rebalance using this method in an acidic solution
Balance keeping it together in an acidic solution
Extra examples
(podcasted)
Cinnabar/vermillion
Both Acidic and Basic
Review
Balancing redox reactions require that you account for
the exchange of electrons as well.
You must have the same number of electrons being
transferred from one reaction to the other.
Learn at minimum the method for leaving reactions
together. The half reaction method is often very useful
though too!
Introduction to
electrochemical cells
Learning Outcomes
Identify the components of a galvanic cell
Identify the components of a cell.
Determining which reaction happens in which part
will be covered in the next section.
Define the difference between a galvanic and
electrolytic cell.
Electrochemical cells
Galvanic cell: an electrochemical cell where a spontaneous chemical reaction is used
to generate an electric current
Electrolytic cell: an electrochemical cell where a non-spontaneous chemical reaction
is occurring due to a supplied electrical source.
Oxidation occurs at the Anode
Reduction occurs at the Cathode
Memory trick:
Red Cat
An Ox
𝐶𝑢2+ 𝑎𝑞 + 2𝑒 − → 𝐶𝑢(𝑠)
Cathode
𝑍𝑛 𝑠 → 𝑍𝑛2+ 𝑎𝑞 + 2𝑒 −
Anode
Galvanic cells: Questions
Where is metal deposited?
Copper
Which direction do the electrons
flow?
Anode to cathode
Why must we separate the
reactions?
In order to have the electrons do
work
𝐶𝑢2+ 𝑎𝑞 + 2𝑒 − → 𝐶𝑢(𝑠)
Cathode
𝑍𝑛 𝑠 → 𝑍𝑛2+ 𝑎𝑞 + 2𝑒 −
Anode
Why is the salt bridge necessary?
Allow flow of anions and cations
Review
Galvanic cells separate reactions.
This allows the electrons to flow from one area to the
other doing work.
You should be able to identify the anode and cathode
given the reactions
But how would we know which is which if not
given the reactions occurring?
How do we know how much voltage we can get?
Cell Voltage.
Learning Outcomes
Define EMF (electromotive “force”) or voltage
Identify the standard hydrogen electrode and how it is used as
a standard.
Use table of values and half reactions to calculate the EMF of a
cell.
Use the cell voltage to decide if a cell is spontaneous or not.
Electromotive (EM) “force”
Not really a force it is a voltage
Synonymous with cell voltage
Listed values are Eored aka
reduction voltages
Eored=-Eooxid
These are standard values,
1atm or 1M
Electromotive (EM) “force”
Listed values are Eored aka
reduction voltages
Eored=-Eooxid
EMF of a cell is the difference in
potential energy between the
anode and cathode.
Standard Hydrogen Electrode
Standard potentials are calculated against
the standard hydrogen electrode!
We need a voltage to compare
everything too.
The H2 half reaction cell was
chosen.
Why do we have a Pt electrode?
Need an inert metal to conduct
electrons
Standard Hydrogen Electrode
Standard potentials are calculated against
the standard hydrogen electrode!
Why do we have a Pt electrode?
Need an inert metal to conduct electrons
Calculating Cell Voltage
Two equivalent ways. Pick your favorite and stick with it.
Be very careful not to interchange them!!!
Method 1:
Method 2:
Why are these equivalent?
Eored=-Eooxid
positive Eocell is a spontaneous reaction (galvanic cell)
negative Eocell is a non-spontaneous reaction (electrolytic cell)
Calculating Cell Voltage
Calculate the cell voltage for a spontaneous Zn/Zn2+, Cu/Cu2+ cell.
Which is Anode/Cathode
How do we decide?
spontaneous: Eocell = positive
non-spontaneous: Eocell =negative
When reversing the sign of the
anode (oxidation), they must add
to be positive.
Copper=cathode
Zinc= anode
Eocu2+/Cu= +0.34V
EoZn2+/Zn=
-0.44V
Calculating Cell Voltage: Calculate the cell voltage for a
spontaneous Zn/Zn2+, Cu/Cu2+ cell
Logic 1:
Eocell=Eored_Cu2+/Cu - Eored_Zn
Eocell=(+0.34V) - (-0.44V)= +.078 V
Logic 2:
Eocell=Eored_Cu2+/Cu + Eooxidation_Zn
Eocell=(+0.34V)+(+0.44)= +0.78V
Example: Calculating Cell Voltage
A typical alkaline battery has the following half reactions, identify the cathode and anode
and write the complete reaction and find E0.
E0red= -1.28V
E0red= 0.15V
One needs to be oxidized, when added together Ered and Eoxid need to add to be positive
Reverse so ZnO (aka Zn2+) needs to be oxidized and is therefore it is the anode
Logic 1:
Logic 2:
E0red cathode= 0.15V
E0red anode= -1.28V
E0cell= 0.15V-(-1.28)= 1.43V
E0red= 0.15V
E0ox= +1.28V
E0cell= 0.15V+(+1.28)= 1.43V
Review
Cell Voltages are calculated as the difference in potential
energy between the anode and cathode. (Pick your favorite
way of thinking about it, and stick with that)
A spontaneous cell will have a positive Eocell, while a nonspontaneous cell will have a negative Eocell.
Electrochemistry:What will
react?
By looking at the table, which species will react with each
other.
Learning Outcomes
Use what we already know to determine how you can
tell what species will readily react spontaneously by
looking at the table of Eored values.
Diagonal Rule
Diagonal rule: Species on left reacts with any species on the
right that is lower than it.
Easily reduced
Why?
Easily oxidized
Example: Lets take 5 reactions
Draw arrows next to the table excerpt identifying the most likely substances to be reduced or
oxidized, and the best oxidizing and reducing agents. Give one spontaneous combinations, and
one non-spontaneous reactions, using the above reactions (lots of available options).
Easily
reduced
Easily
oxidized
Review
On a standard reduction table with the highest
reduction values listed on top, species on the left will
react with species on the right of a reaction that is
lower than it.
This is because species that are easily reduced make
good oxidizing agents.
Cell Notation
Learning Outcomes
Write the cell notation for a given cell.
Given the cell notation, write the half reactions.
Cell Diagram/Notation
Single line denote phase change
Double line denote salt bridge
Start with Anode on far left, work forward in order you’d
encounter it
Cell Diagram Examples
Write the cell diagram for a cell consisting of Zn2+/Zn and Pt/H+/H2
Anode
Cathode
Review Example
Combining everything we’ve learned.
Zinc/Tin Cell Example
The standard reduction potential of a zinc electrode is -0.76 V.
Given that the standard potential of the cell where Zn is oxidized
and Sn4+is reduced to Sn2+ is +0.91V find the standard potential
of the Sn4+/Sn2+ half reaction and write the cell diagram.
Electrochemistry and
Thermodynamics
Learning Outcomes
Using thermodynamic data (gibbs free energy or enthalpy
and entropy) calculate the Eocell.
Calculate the Gibbs free energy of a cell using Eocell.
Calculate the equilibrium constant from Eocell.
Calculate Eocell from the equilibrium constant.
Thermodynamics: Gibbs Free Energy
n= number of moles
F=faraday’s constant= 96,500 J/(V*mol)
We can use this to find the Gibb’s free energy of a reaction if we
can measure/find Ecell (or of course vice versa)
Example
Calculate the standard free-energy change for the following reaction at 25oC
using standard reduction potentials.
2𝐴𝑢3+(𝑎𝑞) + 3𝑀𝑔 (𝑠) 2𝐴𝑢(𝑠) + 3𝑀𝑔2+ (𝑎𝑞)
Thermodynamics: Gibbs Free Energy
What else does DG equal?
How can we use this to relate K and G?
Set them equal to each other.
Cleaning the equation up a bit:
Thermodynamics G and K
Now put all the constants on one side
Filling in constants and solving
Example- Relating E and K;
Calculate the pressure of H2 in atm, required to maintain equilibrium
with respect to the following reaction at 25oC, Given that
[Pb2+]=0.035 and the solution is buffered at pH 1.6
Nernst Equation
Think way back to thermodynamics in Chem1B, how did we relate
DG and DG0?
But using the relation between G and E…..
Rearranging
OR
NOTE: The difference between E and Eo or G and Go is that the o symbolizes standard state
Example
Calculate the reaction quotient Q for the cell reaction, given the
measured values of the cell potential.
Review
We can relate between E, G and
K by using one of the three
equations we introduced in this
section.
For non standard conditions,
use the Nernst equation.
Applications of electrochemical
cells.
Learning Outcomes
Introduction to two types of batteries that are
commonly used.
Specifics of these will not be tested. I.e., don’t memorize
the reactions.
Identify anode and cathode of batteries given the
reactions that are occurring.
Introduction to fuel cells.
Identify the anode and the cathode of a given fuel cell
Dry Cell and Alkaline Batteries
Dry Cell
“Dry Cell”: Originally developed
Problems with corrosion and unstable
current and voltage lead to the development
of the “alkaline dry cell”
Used in a variety of applications, but are not
reachargable, limiting their utility in many
others
Alkaline dry cells replace NH4Cl with NaOH
or KOH
Anode
Cathode
Lead Storage Batteries:
i.e car battery
Grids provide large surface area, low specific
energy- allows high current for short periods
Anode
Cathode
Overall
Charged
Discharged
Batteries
Primary cells: non-rechargable
Secondary cells: rechargeable
Specific energy: energy that can be generated
divided by mass
Dry cell and
alkaline batteries
Lead acid cell: car battery
Grids provide large surface area, low specific energyallows high current for short periods
Fuel Cells
Runs similarly to a battery
However, reactants are supplied
continuously and only products
are H2O
Leads to less waste and less
weight
Originally used space
applications, now provides
backup power to many industrial
applications, as well as fuel cell
cars.
Review, Recap and Note
There are many many types of batteries. I’ve shown you
two variations.
Batteries can be tweaked to provide more voltage, more
current or various other desired outcomes.
Suggestion, look up some emerging research on
batteries that interest you and see if you can pick out
how the reactions work, and are similar or different to
the ones we talked about.
Electrolytic Cells
Review and Applications
Learning Outcomes
Review the difference between an electrolytic and
galvanic cell.
Determine the products at the anode and cathode of
an electrolytic cell.
Electrolytic cells: Reminder
A cell that requires an outside source of power.
i.e. a non spontaneous cell.
These can be separated, like in previous cells, but can
also be located in one container.
Power must be supplied!
More about electrolytic cells
Anode and Cathode are often in
the same container and often only
contain one electrolyte.
Reduction still occurs at cathode
and oxidation still occurs at the
anode.
Many uses, electrolysis,
electroplating, ore purification,
ect…
More about electrolytic cells
Reminder:
E, DG, K are all calculated the
same.
Non-spontaneous so E is negative.
Applications are the new:
interesting part!
Electroplating
Cathode
Anode
Electroplating Example
Use dimensional analysis.
Important things to know:
1 C= 1 amp*sec
F=9.6485x104 C/mol
How many grams of aluminum can be deposited by the passage of
105 C through an electrolytic cell?
How long does it take to deposit 0.63 g Ni on a decorative drawer
handle when 8.7 A are passed through a Ni(NO3)2 solution
Corrosion
Fe3+ precipitates out as Fe2O3 H2O
How can we stop this?
Give the O2 something else to react
with!!!!
Stopping Corrosion: Galvanizing
EoZn2+/Zn= -0.76
EoFe2+/Fe= -0.44
EoFe3+/Fe2+= +0.77
Galvanizing: Put a Zinc coating on it.
Zinc is more reactive, it becomes the anode,
Aka gets oxidized instead of the Iron
Zinc is “sacrificial” metal
Boats, wire, roofing, anything where you need to stop
corrosion.
Review
Electrolytic cells are much like galvanic cells only because
they are non-spontaneous, they must have a power
source.
The Eo is negative.
Electrolytic cells are used for applications such as
galvanizing or electroplating metals.
Long electrochem problem
(done on podcast)
Many important biological reactions involve electron transfer. Because the pH of bodily fluids is close to 7,the
“biological standard potential” of an electrode E, E*, is measured at pH=7. a) Calculate the biological standard
potential for the reducion of hydrogen ions to hydrogen gas, and the reduction of nitrate ions to NO gas.
Calculate the biological standard potential E* for the reduction of the biomolecule NAD+ to NADH in aqueous
solution. The reduction half reaction under thermodynamic standard conditions is NAD+(aq)+H+(aq)+2e-NADH
(aq), with Eo= -0.099V.
The pyruvate ion, CH3C(=O)CO2-, is formed during the metabolism of glucose in the body. The ion has a chain of
three carbon atoms. The central carbon atom has a double bond to a terminal oxygen atom and one of the end
carbon atoms is bonded to two oxygen atoms in a carboxylate group. Draw the Lewis structure of the pyruvate
ion and assign a hybridization scheme to each carbon atom.
The lactate ion has a similar structure to the pyruvate ion, except that the central carbon is no attached to an –
OH group: CH3CH(OH)CO2-. Draw the Lewis structure of the lactate ion and assign a hybrization scheme to the
central carbon atom.
During exercise the pyruvate ion is converted to lactate ion in the body by coupling to the half reaction for NADH
given above. For the half reaction pyruvate+2H++2e-lactate, E*=-0.190 V. Write the cell reaction for the
spontaneous reaction that occurs between these two biological couples and calculate E* and Eo for the overall
reaction.
Calculate the standard Gibbs free energy of reaction for the overall reactions in the above reaction.
Calculate the equilibrium constant at 25oC for the reaction.