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Electrochemical Cells
Reference: Chapter 14 (pg. 610-669)
Assessment (Chapter 13)

Quizzes will be given in this chapter


You should use this to your advantage as they test smaller sections of the
curriculum and help prepare you for the Unit Exam
Electrochem Unit Exam at the end of this chapter
Electrochemical Cells
DAY 1 Objectives:
1.Define anode, cathode, anion, cation, salt bridge/porous cup, electrolyte,
and voltaic cell
2.Predict and write the half-reaction equation that occurs at each electrode
in an electrochemical cell
Electrochemical Cells
Today’s AGENDA:
1.Introduce Electrochemistry – Focus on Voltaic Cells
2.Voltaic Cell Worksheet – guided practice
Introduction to Electrochemistry


A VOLTAIC cell converts chemical energy into electrical energy

Alessandro Volta invented the first electric cell but got his inspiration from Luigi
Galvani. Galvani’s crucial observation was that two different metals could make the
muscles of a frog’s legs twitch. Unfortunately, Galvani thought this was due to some
mysterious “animal electricity”. It was Volta who recognized this experiment’s potential.

A votlaic cell produces very little electricity, so Volta came up with a better design:
A battery is defined as two or more voltaic cells connected in series to
produce a steady flow of current


Volta’s first battery consisted of several bowls of brine (NaCl(aq))
connected by metals that dipped from one bowl to another
His revised design, consisted of a sandwich of two metals
separated by paper soaked in salt water.
Introduction to Electrochemistry


Alessandro Volta’s invention was an immediate technological
success because it produced electric current more simply and
reliably than methods that depended on static electricity.
It also produced a steady electric current –something no
other device could do.
Introduction to Electrochemistry

Electrochemical cells are composed of two
electrodes – solid electrical conductors and at
least one electrolyte (aqueous electrical
conductor)

In many cells, the electrolyte is often a moist paste
(just enough water is added so that the ions can
move). Sometimes one electrode is the cell
container.

The positive electrode is defined as the cathode
and the negative electrode is defined as the anode

The electrons flow through the external circuit from
the anode to the cathode.

To test the voltage of a battery, the red(+) lead is
connected to the cathode (+ electrode), and the
black(-) lead is connected to the anode (- electrode)
Introduction to Electrochemistry


A voltmeter is a device that measures the energy difference, per unit charge, between
any two points in an electric circuit (called electric potential difference)

I.e. A 9V battery releases 6X as much energy compared with the electrons from a 1.5V battery.

The voltage of a cell depends mainly on the chemical composition of the reactants in the cell
An ammeter is a device that measures the rate of flow of charge past a point in an
electrical circuit (called electric current)

The larger the electric cell, the greater current that can be produced
Electric potential difference is like the potential
energy difference between 1kg of water at the top
of the dam and 1kg of water at the bottom of the
dam.
Electric current (or flow of electrons) is like the
flow of the water. A larger drain would be like a
larger electric cell, allowing more water (or
electrons) to flow.
Introduction to Electrochemistry

The power of a battery is the rate at which it produces electrical energy. Power is
measured in watts (W). Calculated using P = IV (Power = current x potential difference)

The energy density is a measure of the quantity of energy stored or supplied per unit
mass. Measured in J/kg.

Table 1 summarizes some important electrical quantities and their units
Secondary cells can be recharged
using electricity, but are expensive
Primary cells cannot be recharged,
but are relatively inexpensive
Fuel cells produce electricity
by the reaction of a fuel that
is continuously supplied.
More efficient, and used for
NASA vehicles, but still too
expensive for general or
commercial applications
Voltaic Cells (aka Galvanic or
Electric Cell)

A device that spontaneously produces electricity by redox

Uses chemical substances that will participate in a spontaneous redox reaction.


The reduction half-reaction (SOA) will be above the oxidation half-reaction (SRA) in the activity
series to ensure a spontaneous reaction.
Composed of two half-cells; which each consist of a metal rod or strip immersed in a
solution of its own ions or an inert electrolyte.





Electrodes: solid conductors connecting the cell to an external circuit
Anode: electrode where oxidation occurs (-)
Cathode: electrode where reduction occurs (+)
The electrons flow from the anode to the cathode (“a before c”) through an electrical
circuit rather than passing directly from one substance to another
A porous boundary separates the two electrolytes while still allowing ions to flow to
maintain cell neutrality

Often the porous boundary is a salt bridge,
containing an inert aqueous electrolyte
(such as Na2SO4(aq) or KNO3(aq)),

Or you can use a porous cup containing
one electrolyte which sits in a container of a
second electrolyte.
Voltaic Cells

Voltaic cells can be represented using cell notation:
The single line represents a phase
boundary (electrode to electrolyte)
and the double line represents a
physical boundary (porous boundary)


The SOA present in the cell always undergoes reduction at the cathode
The SRA present in the cell always undergoes oxidation at the anode
Match the cell notation to the descriptions
a)
Sn(s) Sn4+(aq)
b)
Mg(s) MgCl2(aq)
c)
Sn(s) SnCl2(aq)
d)
Mg(s) Mg2+(aq)
e)
Mg(s)
Mg2+(aq)
f)
Sn(s) SnCl2(aq)
Cu2+(aq) Cu(s)
SnCl4(aq) Sn(s)
1.
Copper placed in a solution of
copper(II) chloride and tin metal placed
in a solution of tin(II) ions
2.
A copper-magnesium cell
3.
Magnesium in a solution of magnesium
chloride and tin in a solution of tin(II)
chloride
4.
A tin(II) ion solution containing a tin
electrode and a solution of magnesium
ions containing a magnesium electrode.
5.
Two tin electrodes in solution of tin(II)
chloride and tin (IV) chloride
respectively
6.
Copper placed in a copper(II) solution
and tin placed in a tin(IV) solution
CuCl2(aq) Cu(s)
Cu2+(aq) Cu(s)
Sn2+(aq)
Sn(s)
SnCl4(aq) Sn(s)
Voltaic Cells – What is going on?

Example: Silver-Copper Cell
Cu(s) Cu2+(aq) Ag+(aq) Ag(s)
1.
Determine which of the entities is the SOA.
The SOA present in the cell always undergoes a reduction at the cathode.
Write the reduction half reaction
2.
Determine which of the entities is the SRA.
The SRA present in the cell always undergoes an oxidation as the anode.
Write the oxidation half-reaction
3.
The cathode is the
electrode where the
strongest oxidizing agent
present in the cell reacts.
Balance the half-reactions and add together to create the net equation.
Memory device:
“An ox ate a red cat”
Anode oxidation; reduction
cathode
The anode is the
electrode where the
strongest reducing agent
present in the cell reacts.
Voltaic Cells – What is going on?

Example: Silver-Copper Cell

Silver ions are the strongest oxidizing agents in the cell, so they undergo a reduction halfreaction at the cathode, creating more Ag(s)

Copper atoms are the strongest reducing agents in the cell, so they give up electrons in an
oxidation half-reaction and enter the solution (Cu2+ = blue ions) at the anode.

Electrons released by the oxidation of copper atoms at the anode travel through the
connecting wire to the silver cathode. (Ag+(aq) win in the tug of war for e-’s over Cu2+(aq))

Since the positive silver ions are being removed from solution, you would assume that the
solution would become negatively charged. This does not happen. Why?

Cations (positively charged ions) move from the salt bridge into the solution in the cathode
compartment to maintain an electrically neutral solution.

Anions (negatively charged ions) move from the salt bridge into the solution in the anode
compartment to maintain an electrically neutral solution.
Voltaic Cell
Animation
Voltaic Cells with Inert Electrodes

Inert electrodes are needed when the SOA or SRA involved in the reaction is not solid.
If this is the case, usually a graphite (C(s)) rod or platinum strip is used as the electrode.
 Inert (unreactive) electrodes provide a location to connect a wire and a surface on
which a half-reaction can occur.

Example: a) Write equations for the half-reactions and the overall reaction that occur in
the following cell:
C(s)
cathode:
anode:
Cr2O72-(aq) H+(aq)
Cu2+(aq) Cu(s)
Cr2O72-(aq) + 14 H+(aq)+ 6 e-  2Cr3+(aq) + 7H2O(l)
3 [Cu(s) Cu2+(aq) + 2e- ]
3Cu (s) + Cr2O72-(aq) + 14 H+(aq)  3Cu2+(aq + 2Cr3+(aq) + 7H2O(l)
b) Draw a diagram of the cell labeling
electrodes, electrolytes, the direction of
electron flow and the direction of ion
movement.
The copper electrode will decrease in mass and the
blue colour of the electrolyte increases (Cu2+), which
indicates oxidation at the anode.
The carbon electrode remains unchanged, but the
orange colour of the dichromate solution becomes
less intense and changes to greenish-yellow (Cr3+),
evidence that reduction is occurring in this half cell
Voltaic Cell Summary

A voltaic cell consists of two-half cells
separated by a porous boundary with solid
electrodes connected by an external circuit

SOA undergoes reduction at the cathode
(+ electrode) – cathode increases in mass

SRA undergoes oxidation at the anode
(- electrode) – anode decreases in mass

Electrons always travel in the external circuit
from anode to cathode

Internally, cations move toward the cathode,
anions move toward the anode, keeping the
solution neutral
Homework
•
Pg 619 Q 13,14
•
Pg 621 Q 1-3 & 5
•
Pg 626 1-7
•
Voltaic Cells WS
•
What is coming up tomorrow?
•
Standard Cells and Cell Potentials
Electrochemical Cells
Day 2 Objectives:
1.Explain that the values of standard reduction potential are all relative to the E r
= 0.00 V set for the hydrogen electrode at standard conditions
2.Calculate the standard cell potential for electrochemical cells
3.Predict the spontaneity of redox reactions based on standard cell potentials
Electrochemical Cells
Today’s AGENDA:
1.Review Homework
2.Standard Cells and Cell Potentials – guided practice
Standard Cells and Cell Potentials

A standard cell is a voltaic cell where each ½ cell contains all entities necessary at SATP
conditions and all aqueous solutions have a concentration of 1.0 mol/L
 Standardizing makes comparisons and scientific study easier

Standard Cell Potential, E0 cell = the electric potential difference of the cell (voltage)
E0 cell = E0r cathode – E0r anode
•
Where E0r is the standard reduction potential, and is a measure of a standard ½ cell’s
ability to attract electrons.
•
The higher the E0r , the stronger the OA
•
All standard reduction potentials are based on the standard hydrogen ½ cell being 0.00V.
This means that all standard reduction potentials that are positive are stronger OA’s than
hydrogen ions and all standard reduction potentials that are negative are weaker.
•
If the E0 cell is positive, the reaction occurring is spontaneous.
•
If the E0 cell is negative, the reaction occurring is non-spontaneous
Rules for Analyzing Standard Cells
1.
Determine which electrode is the cathode. The cathode is the electrode where the
strongest oxidizing agent present in the cell reacts.
I.e. The OA that is closest to the top on the left side of the redox table = SOA
If required, copy the reduction half-reaction for the strongest oxidizing agent and its reduction potential
2.
Determine which electrode is the anode. The anode is the electrode where the
strongest reducing agent present in the cell reacts.
I.e. The RA that is closest to the bottom on the right side of the redox table = SRA
If required, copy the oxidation half-reaction (reverse the half-reaction)
3.
Determine the overall cell reaction. Balance the electrons for the two half reactions
(but DO NOT change the E0r) and add the half-reaction equations.
4.
Determine the standard cell potential, E0cell using the equation:
E0 cell = E0r cathode – E0r anode
Standard Cells and Cell Potentials #1

Example: What is the standard potential of the cell represented below:
Pt(s) │H+(aq) H2(g)║Cu2+(aq) │Cu(s)
1.
Determine the cathode and anode
2.
Determine the overall cell reaction
3.
Determine the standard cell potential
Standard Cells and Cell Potentials #2

Example: What is the standard potential of an electrochemical cell made of a cadmium
electrode in a 1.0 mol/L cadmium nitrate solution and chromium electrode in a 1.0 mol/L
chromium(II) nitrate solution?
SOA
SRA
Cd2+(aq)
Cd(s)
Cr2+(aq)
cathode
Cr(s)
H2O(l)
anode
E0 cell = E0r cathode – E0r anode
= (-0.40V) - (-0.91V)
= + 0.51V
The E0 cell is positive, therefore the reaction is spontaneous.
Standard Cells and Cell Potentials #3

Example: A standard lead-dichromate cell is constructed. Write the cell notation, label
the electrodes, and calculate the standard cell potential.
SRA
Pb(s) Pb2+(aq)
SOA
Cr2O72-(aq) H+(aq) Cr3+(aq)
anode
C(s)
cathode
E0 cell = E0r cathode – E0r anode
= (+1.23V) - (-0.13V)
= + 1.36V
The E0 cell is positive, therefore the reaction is spontaneous.
Cell Potential Animation
Standard Cells and Cell Potentials #4

Example: A standard scandium-copper cell is constructed and the cell potential is
measured. The voltmeter indicates that the copper electrode is positive.
Sc(s)
Sc3+(aq)
Cu2+(aq)
anode
Cu(s)
E0 cell = +2.36V
cathode
Write and label the half-reaction and net equations, and calculate the standard reduction
potential of the scandium ion.
E0 cell
= E0r cathode - E0r anode
2.36V = (+0.34V) - (x)
E0r anode = -2.02V
Standard Reference Half Cell




The standard hydrogen half cell
Eor values were determined by using a reference half cell
which was arbitrarily assigned a value of 0 volts. This is
called the standard hydrogen half-cell it’s reaction is
shown below.
2H+(aq) + 2e- → H2(g
Eor = 0.00V
A positive Eor means that the the half cell connected to
the hydrogen half cell is a stronger OA than H+ ( is
stronger at attracting e-). A negative Eor means the half
cell is a weaker OA than the H+.
Standard Reference Half Cell
Varying Reference Cells
You must be able to calculate relative Eor using any half
cell.
 Note: the individual reduction potentials will be different,
but their relative values or ∆Eo for the net reaction will stay
the same.
 a)Revise the half cell reduction potentials for the half
reactions if copper was the reference half cell and
assigned a reduction potential of 0.00V.

Varying Reference Cells
b) What if zinc was assigned the reference half cell value
of 0.00V.
What would be the revised have cell reduction potentials
for hydrogen and copper?


c) Calculate the ∆Eo for a copper-zinc cell using their
original Eor values and those determined in parts a and b.
Homework
•
Pg 631 Q 11
Pg 633 Q 12,13
Pg 637 Q 17,18,24 (On Corrosion, be sure to read this section in your text)
Cell Potentials Extra Exercises WS
•
What is coming up tomorrow?
•
•
•
•
•
Start Electrochemical Cells
Voltaic Cell and Cell Potential Quiz (in two days)
Electrochemical Cells
Day 3 Objectives:
1.Identify the similarities and differences between a voltaic cell and an
electrolytic cell
2.Predict the spontaneity of redox reactions based on standard cell potentials.
3.Recognize that predicted reactions do not always occur.
Electrochemical Cells
Today’s AGENDA:
1.Review Homework
2.Electrolytic Cells
Electrolytic Cells

The term “electrochemical cell” is a umbrella terms used for

Voltaic Cells – one with a spontaneous reaction
SOA over SRA on the activity series
Eocell greater than zero = spontaneous

Electrolytic cells – one with a nonspontaneous reaction
SOA below SRA – i.e. zinc sulfate and lead solid cell
Eocell less than zero= nonspontaneous

Why would anyone be interested in a cell that is not spontaneous?

This would certainly not a good battery choice, but by supplying electrical energy to
a nonspontaneous cell, we can force this reaction to occur.

This is especially useful for producing substances, particularly elements. I.e. the zinc
sulfate cell discussed above is similar to the cell used in the
industrial production of zinc metal.
Electrolytic Cells

Electrolytic Cell – a cell in which a nonspontaneous redox
reaction is forced to occur; a combination of two electrodes, an
electrolyte and an external power source.


Electrolysis – the process of supplying electrical energy to force a
nonspontaneous redox reaction to occur
The external power source acts as an “electron pump”; the electric energy
is used to do work on the electrons to cause an electron transfer
Electrons are pulled from the anode
and pushed to the cathode by the
battery or power supply
Comparing Electrochemical Cells:
Voltaic and Electrolytic
It is best to think of “positive” and “negative” for electrodes as labels, not charges.
Procedure for Analyzing Electrolytic Cells

Use the redox table to identify the SOA and SRA

Don’t forget to consider water for aqueous electrolytes.

Write equations for the reduction (cathode) and oxidation (anode) halfreactions. Include the reduction potentials if required.

Balance the electrons and write the net cell reaction including the cell
potential. E0 cell = E0r cathode - E0r anode

If required, state the minimum electric potential (voltage)
to force the reaction to occur. (The minimum voltage is
the absolute value of E0 cell)

If a diagram is requested, use the general outline in
Figure 6, and add specific labels for chemical entities.
Analyzing Electrolytic Cells #1

Example: What are the cell reactions and the cell potential of the aqueous
potassium iodide electrolytic cell?

Identify major entities and identify the SOA and SRA.

Write the half-reaction equations and calculate the cell potential.

State the minimum electric potential (voltage) to force the reaction to occur.
Electrons must by supplied with a
minimum of +1.37 V from an external
battery or other power supply to
force the cell reactions.
Potassium-Iodide Electrolytic Cell

In the potassium iodide electrolytic cell, litmus
paper does not change colour in the initial solution
and turns blue only near the electrode from which
gas bubbles. Why?

At the other electrode, a yellow-brown colour and
a dark precipitate forms. The yellow brown
substance produces a purplish-red colour in the
halogen test (pg. 805). Why?
Analyzing Electrolytic Cells #2

Example: An electrolytic cell containing cobalt(II) chloride solution and lead
electrodes is assembled. The notation for the cell is as follows:
a)
Predict the reactions at the cathode and anode, and in the overall cell.
b)
Draw and label a cell diagram for this electrolytic cell, including the power supply.
c)
What minimum voltage must be applied to make this cell work?
Analyzing Electrolytic Cells #3

Example: An electrolytic cell is set up with a power supply connected to two nickel
electrodes immersed in an aqueous solution containing cadmium nitrate and zinc nitrate.

Predict the equations for the initial reaction at each electrode and the net cell reaction.
Calculate the minimum voltage that must be applied to make the reaction occur.
The Chloride Anomaly (******Diploma)

Some redox reactions predicted using the SOA and
SRA from a redox table do not always occur in an
electrolytic cell.

The actual reduction potential required for a
particular half-reaction and the reported halfreaction reduction potential may be quite different
(depending on the conditions or half-reactions)


This difference is known as the half-cell overvoltage.
“As an empirical rule, you should recognize that
chlorine gas is produced instead of oxygen gas in
situations where chloride and water are the only
reducing agents present.”
Applications of Electrolytic Cells

Read pg. 646-650

Summary:



In molten-salt electrolysis, metal cations are reduced to metal
atoms at the cathode and nonmetal anions are oxidized at the
anode.
Electrorefining is a process used to obtain high grade metals at
the cathode from an impure metal at the anode.
Electroplating is a process in which a metal is deposited on the
surface of an object placed at the cathode of an electrolytic cell.
Electrolytic Cells

Summary:

An electrolytic cell is based upon a reaction that is nonspontaneous; the
Eocell for the reaction is negative.
An applied voltage of at least the absolute value of Eocell is required to
force the reactions to occur.

The SOA undergoes reduction at the cathode (- electrode)

The SRA undergoes oxidation at the anode (+ electrode)

Electrons are forced by a power supply to travel from the anode to the
cathode through the external circuit.

Internally, anions move toward the anode and cations move toward the
cathode
Homework
•
•
Pg 644 Q 5-6
Pg. 645 #7-14
•
•
•
Remember molten means liquid state – so there is no water present
Don’t forget about the chloride anomaly for one
What is coming up tomorrow?
•
•
Quiz Tomorrow everything up to now 
Cell Stoich
Electrochemical Cells
Today’s AGENDA:
1.Review Homework
2.Cell Worksheet #1-6
3.Start STS Connections
Homework
•
Cell Worksheet #1-6
Study for Cell Quiz tomorrow
•
What is coming up tomorrow?
•
•
•
Cell Quiz
Finish STS Connections
Electrochemical Cells
Today’s Objectives:
1.Calculate mass, amounts, current, and time in single voltaic and
electrolytic cells by applying Faraday’s law and stoichiometry
2.Predict and write the half-reaction equation that occurs at each electrode
in an electrochemical cell
Electrochemical Cells
Today’s Agenda:
1.Cell Quiz
2.Finish STS Connections
Cell Quiz
Electrochemical Cells
Today’s Agenda:
1.Introduce Cell Stoichiometry
Cell Stoichiometry Summary
1.
Write the balanced equation for the half-cell reaction of the substance
produced or consumed. List the measurements and conversion factors
for the given and required entities.
2.
Convert the given measurements to an amount in moles by using the
appropriate conversion factor. (M, c, F)
3.
Calculate the amount of the required substance by using the mole ratio
from the half-reaction equation.
4.
Convert the calculated amount to the final quantity
by using the appropriate conversion factor. (M, c, F)
Background

Charge (Q) is determined by multiplying the electric current (I),
(measured in C/s) by the time (t) measured is seconds.
Q = It
(C) = (Ampere)(second)
(Coulomb) = (Coulombs per second) x (second)
One coulomb is the quantity of charge transferred by a current of 1
Ampere during 1second.
Example: Calculate the charge that passes through one 300kA cell in a 24 hour period.
Q = It
= (300kA x 1000A/kA)(24 h x 3600s/h)
= (300000C/s )(86400s) = 2.6 x 1010C
Practice: Calculating Charge

Pg. 653 #1-4
Faraday’s Law

“The mass of an element produced or consumed at an electrode is directly
proportional to the time the cell operated, as long as the current was constant.”

He also found that “9.65 x 104C of charge is transferred for every mole of
electrons that flows in the cell. This value is the molar charge of electrons, also
called Faraday’s constant”
F = 9.65 x 104 C _.
mol e-

This constant can by used as a conversion factor in converting electric charge to
moles, very similarly to the way that molar mass is used to convert mass to a
chemical amount.

I.e.
2.5 mol x 2.02 g = 5.05 g
mol
15000 C x 1 mol e- = 0.16 mol e9.65 x 104 C
Practice: Faraday’s Law #1

Convert a current of 1.74 A for 10.0 min into an amount
of electrons
We do not have a direct way of measuring amount of electrons, so we
need to calculate charge first (Q = It), and then use Faraday’s constant to
calculate the amount of electrons.
Q = It = (1.74 C ) (10.0min x 60 s ) = 1044 C x 1 mol = 0.0108 mol
s
min
9.65 x 104 C
You can also write this as a single equation (using cancellation of units):
Practice: Faraday’s Law #2

How long, in minutes, will it take a current of 3.50 A to
transfer 0.100 mol of electrons?
Remember: Amperes = Coulombs/second
Practice: Calculating Charge

Pg. 654 #5-7
Half Cell Calculations

Since the mass of an element produced at an electrode depends
on the amount of transferred electrons (Faraday’s Law), a halfreaction equation showing the number of electrons involved is
necessary to do stoichiometric calculations.


This applies to all electrochemical cells, whether voltaic or electrolytic.
Separate calculations are carried out for each electrode, although
the same charge and therefore the same amount of electrons
passes through each electrode in a cell.

Remember: The only new part of the stoichiometry is the calculation
involving the amount of electrons based on Faraday’s constant
Practice: Half-Cell Calculations #1


What is the mass of copper deposited at the cathode of a copper
electrorefining cell operated at 12.0 A for 40.0 min?
First identify and write the appropriate half-cell.



Copper is being deposited at the cathode, copper(II) ions must be gaining
electrons to form copper metal.
Write the equation for this reduction and list all the information given.
Do we have enough information to solve for the amount of electrons
(moles of electrons)?
Practice: Half-Cell Calculations #1

What is the mass of copper deposited at the cathode of a copper
electrorefining cell operated at 12.0 A for 40.0 min?

Yes, we can solve for the number of moles, and then use the mole ratio
to convert from a chemical amount of one substance to another.

The last step is to convert to the quantity requested in the question, in
this case the mass of the copper metal
Could we do this as one equation instead?

Practice: Half-Cell Calculations #1

What is the mass of copper deposited at the cathode of a copper
electrorefining cell operated at 12.0 A for 40.0 min?
Practice: Half-Cell Calculations #2



Silver is deposited on objects in a silver electroplating cell. If 0.175 g of silver is
to be deposited from a silver cyanide solution in a time of 10.0 min, predict the
current required.
Write the balanced equation for the half-cell reaction, list the measurements
and conversion factors.
Convert to moles, use the mole ratio, convert to the current (C/s)
Stoichiometry Calculations
(Measured quantity)
solids/liquids
m
gases
V,T, P
solutions
c,V
electrochemical cells
Q
solutions
gases
solids/liquids
(Required quantity)
c,V
V,T,P
m

n
mole
ratio

n
Homework
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Cell Stoich WS
Pg. 657 #1-8
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What is coming up tomorrow?
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Chapter 14 Review
Electrochemistry Unit Exam coming soon!