Chemistry 30
Electrochemistry Notes
A. Redox Reactions
30-B1.1k  define oxidation and reduction operationally and theoretically
30-B1.2k  define oxidizing agent, reducing agent, oxidation number, half-reaction, disproportionation
30-B1.7k – write and balance equations for redox reactions in acidic and neutral solutions by…developing
simple half-reaction equations from information provided about redox changes

electrochemistry is the branch of chemistry that studies _____________________
__________________________________________________________________

_________________ is a _____________________________________________
eg)

__________________ is a ____________________________________________
eg)

oxidation and reduction reactions occur together, hence the term ______________

the reduction and oxidation reactions are called the ________________________

“adding” the half reactions together will give you the ______________________
____________________________ that takes place during the redox reaction

the e lost in the oxidation half reaction _______________________ the e gained
in the reduction half reaction

you may have to ________________________________________________ of
the half reactions to balance the e

_____________________________________ (ions not changing) are _________
included!

the substance that is ________________________ is called the
_____________________________________________…it causes the oxidation
by taking e

the substance that is _________________________ is called the
_____________________________________________…it causes the reduction
by giving up e
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Example 1
Given the following reaction, write the half reactions and the net ionic equation.
Na(s) + LiCl(aq)  Li(s) + NaCl(aq)
Example 2
Given the following reaction, write the half reactions and the net ionic equation.
3 Zn(s) + 2 Au(NO3)3(aq)  2 Au(s) + 3 Zn(NO3)2(aq)
Example 3
Given the following reaction, write the half reactions and the net ionic equation.
Br2(l) + 2 NaI(aq)  I2(s) + 2 NaBr(aq)
Your Assignment: pg 1 “Net Ionic Equations” #1-5
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B. Spontaneous Redox Reactions
30-B1.6k  predict the spontaneity of a redox reaction, based on standard reduction potentials, and
compare their predictions to experimental results

chemical reactions which occur on their own, without the input of
____________________________________________, are called
________________________________

not all reactions are spontaneous

in the table of redox half reactions (see pg 9 in Data Booklet), the
_____________________________________________ is at the top left and the
____________________________________________ is at the bottom right

the ________________________________________ rule states that a
spontaneous reaction occurs if the ____________________ agent is above the
_______________ agent in the table of redox half reactions
Try These:
For each of the following combinations of substances, state whether the reaction would
be spontaneous or non-spontaneous:
1. Cr3+(aq) with Ag(s)
2. I2(s) with K(s)
3. H2O2(l) with Au3+(aq)
4. Sn2+(aq) with Cu(s)
5. Fe2+(aq) with H2O(l)
C. Predicting Redox Reactions
30-B1.6k  predict the spontaneity of a redox reaction, based on standard reduction potentials, and
compare their predictions to experimental results
30-B1.7k – write and balance equations for redox reactions in acidic and neutral solutions by…using halfreaction equations from a standard reduction potential table
30-B1.2s  use a standard reduction potential table as a tool when considering the spontaneity of redox
reactions and their products
30-B2.7k – predict the spontaneity or nonspontaneity of redox reactions, based on the relative positions of
half-reaction equations on a standard reduction potential table

we will be predicting the strongest or most dominating reaction that occurs when
substances are mixed (other reactions do take place because of atomic collisions!)
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Steps:
1. List all species present as reactants
1. dissociate _____________________________________ and
_______________
2. do not dissociate ____________________________________
3. include ____________ ions if it is _________________
4. always include _________________
2. Identify each as ________ or _________ (***some can be both so memorize
them…_______________________________________)
3. Identify the _______ and __________ using the table (SOA is closest to top, SRA
is closest to bottom).
4. Write out the ______________________ for the SOA and SRA.
5. Determine the ________________________________________.
6. Determine ________________________ (SOA must be higher than SRA to be
spontaneous)
Example 1
Predict the most likely redox reaction when chromium is placed into aqueous zinc
sulphate.
Example 2
Predict the most likely redox reaction when silver is placed into aqueous cadmium
nitrate.
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Example 3
Predict the most likely redox reaction when potassium permanganate is slowly poured
into an acidic iron (II) sulphate solution.
Your Assignment: pgs 1-2 “Predicting Redox Reactions” #1-15
D. Generating Redox Tables



you can be given data for certain ions and elements then be asked to generate a
redox table like the one on pg 9 of you Data Booklet (a smaller version!)
you may have to generate a table from real or fictional elements and ions
the tables that we use are all written as _______________________ half reactions
Example 1
Generate a redox table given the following data:
 indicates no reaction
 indicates a reaction
Cu2+(aq)
Zn2+(aq)
Pb2+(aq)
Ag+(aq)
Cu(s)




Zn(s)




Pb(s)




Ag(s)




Redox Table
Put the oxidizing agents in order from strongest to weakest.
Put the reducing agents in order from strongest to weakest.
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Example 2:
Generate a redox table given the following data:
Cu(s) + Ag+(aq)  Cu2+(aq) + Ag(s)
Zn2+(aq) + Ag(s)  no reaction
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
Hg(l) + Ag+(aq)  no reaction
Redox Table
Put the oxidizing agents in order from weakest to strongest.
Put the reducing agents in order from weakest to strongest.
Example 3
Generate a redox table given the following data:
2X(aq) + Y2(g)  spontaneous reaction
2Z(aq) + Y2(g)  no reaction
2Z(aq) + W2(g)  spontaneous reaction
Redox Table
Put the oxidizing agents in order from strongest to weakest.
Put the reducing agents in order from strongest to weakest.
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Your Assignment: pgs 2-3 “Generating Redox Tables” #1 - 5
E. Oxidation Numbers (States)
30-B1.2k  define oxidizing agent, reducing agent, oxidation number, half-reaction, disproportionation
30-B1.7k – write and balance equations for redox reactions in acidic and neutral solutions by…

an _________________________ is the charge an atom ___________ to have
when found in a __________________________________ or charged
_______________________________________

can be used when you have a ______________________________ where there
are no ________________________ to determine if oxidation or reduction is
occurring

how do you use a change in the number?
1. if the number __________________ then __________________ has occurred
2. if the number __________________ then __________________ has occurred
Rules for Assigning Oxidation Numbers:
1. In a pure element, the oxidation number is ____________.
2. In simple ions, the oxidation number is equal to the ________________.
3. In most compounds containing hydrogen, the oxidation number for hydrogen is
________. (Exception is the metal hydrides eg) LiH where the oxidation number of
hydrogen is _________)
4. In most compounds containing oxygen, the oxidation number for oxygen is
__________. (Exception is the peroxides eg) H2O2, Na2O2 where the oxidation
number of oxygen is __________)
5. The sum of oxidation numbers of all atoms in a substance must equal the
___________________ of the substance. (______________ for compounds and
__________________ of the polyatomic ion)
eg) sum of MgO = _______
sum of PO43 = ________
Example
What is the oxidation number (state) for the element identified in each of the following
substances?
a) N in N2O
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b) N in NO3
c) C in C2H5OH
d) C in C6H12O6
30-B1.3k  differentiate between redox reactions and other reactions, using half-reactions and/or
oxidation numbers
30-B1.4k  identify electron transfer, oxidizing agents and reducing agents in redox reactions that occur in
everyday life, in both living systems (eg. cellular respiration, photosynthesis) and nonliving
systems; i.e., corrosion

figuring out oxidation numbers can help to identify whether a reaction is a
________________________________________________ or not

for it to be a redox reaction, there has to be ___________ an ________________
in oxidation number and a ________________ in oxidation number seen in the
reaction
eg)
Ag(s) + NaNO3(aq)  Na(s) + AgNO3(aq)
PbSO4(aq) + 2 KI(aq)  PbI2(s) + K2SO4(aq)

electron transfer occurs in ___________________________________
eg)

also occurs in _____________________________________________
eg)
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F. Disproportionation
30-B1.2k  define oxidizing agent, reducing agent, oxidation number, half-reaction, disproportionation

disproportionation occurs when one element _______________________
__________________________________________________________
eg)
2 H2O2(aq)  2 H2O(l) + O2(g)
Cl2(g) + 2 OH(aq)  ClO(aq) + Cl(aq) + H2O(l)
Your Assignment: pg 4 “Oxidation Numbers” #1-5
G. Balancing Redox Reactions
30-B1.7k – write and balance equations for redox reactions in acidic and neutral solutions by…assigning
oxidation numbers, where appropriate, to the species undergoing chemical change

sometimes most reactants and products are known but the complete reaction is not
given…called a ________________________ reaction
1. Half Reaction Method
1. Assign __________________________________________.
2. Balance the _________________ that changes in oxidation number.
3. Add _______ to balance the change in _________________ oxidation number
(______________________________________________).
4. Balance O using _________.
5. Balance H using _________.
6. Check that the half-reaction is balanced with respect to ___________.
Example 1
Balance the following half reaction:
CrO42(aq)
Chemistry 30 Electrochemistry
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CrO2(aq)
Example 2
Balance the following half reaction:
HClO2(aq) 
Cl2(g)
Your Assignment: pg 5 “Balancing Redox Half Reactions” #1-10

now we are going to combine coming up with our own half reactions with
figuring out the net redox reaction
Steps
1. Assign _________________________________________.
2. Separate the partial net equation into two _______________________ (omit any
_______ or _______).
3. Balance each half-reaction.
4. _________________________________________ of the equations so e lost = e
gained
5. Add the equations to produce a balanced _______________________.
6. __________________. Check to see if all elements and charges are balanced.
Example 1
Balance the following reaction using the half reaction method::
___H+(aq) + ___CrO42(aq) + ___SO32-(aq)  ___CrO2(aq) + ___SO42(aq) + ___H2O(l)
Your Assignment: pg 5 “Half Reaction Method” #1-6
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2. Oxidation Number Method
 we will use the unbalanced reaction and the change in oxidation numbers to work
out the balancing
Steps:
1. Assign oxidation numbers.
2. Balance the substances that change in oxidation number.
3. Use a _________ to join the reducing agent with its corresponding product (ignore
the H+(aq) and H2O(l)) and a ________ to join the oxidizing agent with its
corresponding product.
4. On each line, write the ___________ in oxidation number  # of atoms.
5. ___________ the RA and/or OA to balance the change in oxidation number.
6. ____________________ the H2O(l) and the H+(aq).
Example 1
Balance the following reaction using the oxidation number method.
__H+(aq) + __MnO4(aq) + __SO32(aq)  __MnO2(s) + __SO42(aq) + __ H2O(l)
Example 2
Balance the following reaction using the oxidation number method.
__ H2O(l) + __N2O4(g) + __ Br(aq)
 __NO2 (aq) + __ BrO3(aq) + __ H+(aq)
Your Assignment: pg 5 “Oxidation Number Method” #1-6
H. Redox Stoichiometry
30-B1.8k  perform calculations to determine quantities of substances involved in redox titrations
30-B1.4s  select and use appropriate numeric, symbolic, graphical and linguistic modes of representation
to communicate equations for redox reactions and answers to problems related to redox
titrations
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1. Calculations
 stoichiometry can be used to predict or analyze a quantity of a chemical involved
in a chemical reaction
 in the past we have used balanced chemical equations to do stoich calculations
 we can now apply these same calculations to balanced redox equations
Example 1
What is the mass of zinc is produced when 100 g of chromium is placed into aqueous
zinc sulphate.
Example 2
What volume of 1.50 mol/L potassium permanganate is needed to react with 500 mL of
2.25 mol/L acidic iron (II) sulphate solution?
Your Assignment: pgs 6 “Redox Stoichiometry” #1-6
Chemistry 30 Electrochemistry
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2. Titrations

a titration is a lab process used to determine the _________________ of a
substance needed to react completely with another substance

this volume can then be used to calculate an unknown
_______________________ using stoichiometry

one reagent (_________________ - ________) is slowly added to another
(____________________ - __________) until an abrupt change
(____________________) occurs, usually in colour

in redox titrations, two common oxidizing agents are used because of their
____________________and _____________________:
i.
ii.

as long as the sample (RA) in the flask is reacting with the
________________________________________________ the sample will be
______________________

when the reaction is complete, any unreacted permanganate ions will turn the
sample ____________________ (pink) (with dichromate, sample goes from
orange to green)

the volume of titrant (OA) needed to reach the endpoint is called the
________________________________

the __________________________________of the titrant must be accurately
known

the concentration of the permanganate solution must be calculated against a
______________________________________ (a solution of known
concentration) before it can be used in a titration itself

this is done just prior to the titration
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Example
Find the concentration of (standardize) the KMnO4(aq) solution by titrating 10.00 mL of
0.500 mol/L acidified tin (II) chloride with the KMnO4(aq).
Evidence
Trial
1
2
3
4
Final Volume (mL)
18.40
35.30
17.30
34.10
Initial Volume
1.00
18.40
0.60
17.30
pink
light pink
light pink
light pink
(mL)
Volume of
_________ (mL)
Endpoint Colour

endpoint average is calculated by using 3 volumes within 0.20 mL
Endpoint average =
Your assignment: pg 6 “Redox Stoichiometry” #7-9
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I. Electrochemical Cells
1. Voltaic Cells
30-B2.1k  define anode, cathode, anion, cation, salt bridge/porous cup, electrolyte, external circuit,
power supply, voltaic cell and electrolytic cell

_______________ are devices that convert ____________ energy into
_____________ energy

in redox reactions, e are transferred from the ___________________ to the
___________________________

the transfer of e can occur through a __________________________ separating
the two substances in containers called ________________

a ___________________________________ is an arrangement where
_______________________________ are joined so that the ________ and
________________ can move between them

_________________________ are made of good conducting materials so e can
flow…can be the _________________of the solution or inert such as
_______________________

the ________________________________ is a solution that contains ions and
will transmit ions

the electrode where _______________ occurs is called the __________________

if the anode is a metal, it _________________ mass as the cell operates

the anode is labelled as _________________________ since it is the electrode
where electrons originate
the __________________ move to the _______________ since this electrode
________ electrons (leaving a net ___________________ charge in the electrode)

the electrode where _______________ occurs is called the __________________

if the cathode is a metal, it _________________ mass as the cell operates

the cathode is labelled as _________________________ since the anode is
labelled negative

the __________________ move to the _______________ since this electrode
_______________ electrons (leaving a net ______________ charge in the
electrode)
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
electrons flow from the ______________________________ to the
____________________________ through a connecting wire

ions must be able to ____________ to their attracting electrode (either through the
_______________________ or a _________________) otherwise a buildup of
charge will occur opposing the movement of e

the flow of ions through the solution and e through the wire maintains overall
__________________________________________
2. Standard Reduction Potentials (E)
30-B2.5k  explain that the values of standard reduction potential are all relative to 0 volts, as set for the
hydrogen electrode at standard conditions

________________________________ are the ability of a half cell to
______________________________________

these potentials are measured using a __________________________

each half reaction listed in the Data Booklet has an E value measured in
___________________assigned to it

all values in the table are arbitrarily assigned based on a standard

the _________________________________________ half reaction has been set
as the standard and has an E value of ______________
3. Predicting Voltage of a Voltaic Cell
30-B2.3k  predict and write the half-reaction equation that occurs at each electrode in an electrochemical
cell
30-B2.6k – calculate the standard cell potential for electrochemical cells

the standard cell potential ____________ is determined by _________ the
______________________ for the two half reactions

the ____________ on the E value for the ____________________ half reaction
must be ____________________________

if you multiply an equation to balance e, you ___________________ multiply
the E value (voltage is independent of number of e transferred)

a ____________________ E net is a _____________________ reaction

a ____________________ E net is a _____________________ reaction
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Example
Calculate the Enet for the reaction of Zn(s) with CuSO4(aq).
4. Shorthand Notation
or

line ___________ separates_______________

double line __________ represents the ____________________ or
_____________________ and separates the two _________________

comma _______separates ___________________________________
___________________________________
5. Drawing Cells
30-B2.3k  predict and write the half-reaction equation that occurs at each electrode in an electrochemical
cell
30-B2.6k – calculate the standard cell potential for electrochemical cells
30-B2.7k – predict the spontaneity or nonspontaneity of redox reactions, based on standard cell potential


when drawing a cell from the shorthand notation, you have to be able to label the
cathode, anode, positive terminal, negative terminal, electrolytes, direction of e
flow, and directions of cation and anion flow
you also have to show and label the reduction half reaction, oxidation half
reaction and net reaction including E values, E net and spontaneity
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Example
Draw and fully label the following electrochemical cell:
3+
2+
Al(s) /Al (aq)// Ni (aq) /Ni(s)
Your Assignment: pg 7 “Electrochemical Cells” #1-8
Electrochemical Cells Lab
J. Commercial Cells
30-B2.1sts  explain that scientific knowledge may lead to the development of new technologies, and new
technologies may lead to or facilitate scientific discovery
30-B2.3sts  explain that science and technology have influenced, and been influenced by, historical
development and societal needs

__________________________are made by connecting two or more voltaic cells
in ____________________________________________

the _______________________ of the battery is the ____________ of the
voltages of the _________________________________________

there are many types of batteries:
a) Dry Cell

common __________________________ batteries of clocks, remote controls,
noisy kids toys etc.
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Cathode (Red):

2 MnO2(s) + H2O(l) + 2e  Mn2O3(aq) + 2 OH (aq) E= +0.79 V
Anode (Oxid):
E =

2+
Net: 2 MnO2(s) + H2O(l) + Zn(s)  Mn2O3(aq) + 2 OH (aq) + Zn (aq)

Enet = +1.55 V
the __________ produced causes irreversible side reactions to occur making
recharging impossible
b) Nickel-Cadmium

one type of ___________________________ battery

Cathode (Red): 2 NiO(OH)(s) + 2 H2O(l) + 2e  2 Ni(OH)2(s) + 2 OH (aq) E= +0.49 V
Anode (Oxid):
Net:

Cd(s) + 2 OH (aq) 
2 Cd(OH)2(s) + 2e
2 NiO(OH)(s) + 2 H2O(l) + Cd(s)  2 Ni(OH)2(s) + Cd(OH)2(s)
E = +0.76 V
Enet = +1.25 V
c) Lead Storage Battery

__________________________________ where ___________ serves as the
anode, and ____________________________________ serves as the cathode

both electrodes dip into an electrolyte solution of ________________
______________________

_______________________________ are connected in series

Cathode (Red): PbO2(s) + HSO4 (aq) + 3 H+(aq) + 2e  PbSO4(s) + 2 H2O(l) E= +1.68 V
Anode (Oxid):
Net:

Pb(s) + HSO4 (aq)  PbSO4(s) + H+(aq) + 2e
+

Pb(s) + PbO2(s) + 2 H (aq) + 2 HSO4 (aq)  2 PbSO4(s) + 2 H2O(l)
E = +0.36 V
Enet = +2.04 V
d) Fuel Cells

cells where reactants are _____________________________________________
Chemistry 30 Electrochemistry
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
the energy from this reaction can be used to ______________________________

one type is the _____________________________________________________

_________________________ is pumped in at the _______________ while
______________________ is pumped in at the _____________ (which both have
a lot of surface area)




pressure is used to push the H2 through a platinum catalyst which splits the H2
into 2H+ and 2e
the 2e move through an external circuit towards the cathode generating electrical
energy
the O2 is also pushed through the platinum catalyst forming two oxygen atoms
the H+ ions and oxygen atoms combine to form water
Cathode (Red):
O2(g) + 4 H+(aq) + 4e-  2 H2O(l)
Anode (Oxid):
2 H2(g) 
4 H+(aq) + 4e
E=
E =
V
V
Net:

need a source of hydrogen…reformers are used to convert CH4 or CH3OH into
___________ and ______________

unfortunately, __________ is a _______________________________________

about 24-32% efficient where gas-powered car is about 20% efficient
K. Electrolytic Cells
1. The Basics
30-B2.1k  define anode, cathode, anion, cation, salt bridge/porous cup, electrolyte, external circuit,
power supply, voltaic cell and electrolytic cell

in an electrolytic cell, ____________________ energy is used to force a
____________________________________ chemical reaction to occur (opposite
of a voltaic cell)

these reactions have a __________________________________Enet
Chemistry 30 Electrochemistry
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
commonly used to _________________________________________________
(eg. gold, silver, bronze, chromium etc), _________________________________
and split compounds into ____________________________ (eg. H2, O2, Cl2 etc)

the electrolytic cell is hooked up to a ____________________________________
_________________________________(instead of load or external circuit) so
the flow of e is __________________________________

the __________________________ of the electrolytic cell is connected to the
____________________ of the battery and therefore is _____________________

the __________________________ of the electrolytic cell is connected to the
________________________ of the battery and therefore is ________________
30-B2.2k  identify the similarities and differences between the operation of a voltaic cell and that of an
electrolytic cell

Voltaic Cells
Voltaic vs. Electrolytic Cells
Electrolytic Cells














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30-B1.2sts  explain that technological problems often require multiple solutions that involve different
designs, materials and processes and that have both intended and unintended consequences

some processes are used in industry to produce gases, for example:
1.
the ___________________________________ for producing
_________________________…aluminum oxide is electrolyzed using
carbon electrodes…liquid aluminum is collected
2.
a _______________________________________________ for producing
_____________________________…a saturated sodium chloride solution
is electrolyzed…chlorine gas is formed and collected at the anodes
Example 1
An electric current is passed through a solution of nickel (II) nitrate using inert
electrodes. Predict the anode and cathode reactions, overall reaction, and minimum
voltage required.
Example 2
An electric current is passed through a solution of potassium iodide using inert
electrodes. Predict the anode and cathode reactions, overall reaction, and minimum
voltage required.
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Example 3
An electric current is passed through a solution of copper(II) sulphate using a carbon
electrode and a metal electrode. Predict the anode and cathode reactions, overall
reaction, and minimum voltage required.
30-B2.4k  recognize that predicted reactions do not always occur; eg) the production of chlorine gas
from the electrolysis of brine
 _________________ is an exception to our rules…when water and chlorine are
competing as reducing agents, water is the stronger RA but _________________ is
chosen because the transfer of e from H2O to O2 is more difficult…called
____________________________
Example 4
An electric current is passed through a solution of sodium chloride. Predict the anode
and cathode reactions, overall reaction, and minimum voltage required.
Your Assignment: pg 7 “Electrochemical Cells” #1-8
Electrolysis Lab
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2. Quantitative Study of Electrolysis
30-B2.8k – calculate mass, amounts, current and time in single voltaic and electrolytic cells by applying
Faraday’s law and stoichiometry

quantitative analysis (stoich) provides information on necessary quantities,
current and/or time for electrolytic reactions

the unit for charge _________ is the _______________________

one e carries __________________________ of charge

this means that one _________of e carry _____________________ of charge

____________________________________ is called the
_________________________ (see Data Booklet pg 3)
and
where:

the above equations can be combined into one equation:

we can use these equations in stoichiometric calculations for current, time, mass,
moles of a substance and moles of e-
Example 1
An electrochemical cell caused a 0.0720 mol of e to flow through a wire. Calculate the
charge.
Chemistry 30 Electrochemistry
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Example 2
Determine the number of moles of electrons supplied by a dry cell supplying a current of
0.100 A to a radio for 50.0 minutes.
Example 3
If a 20.0 A current flows through an electrolytic cell containing molten aluminum oxide
for 1.00 hours, what mass of Al() will be deposited at the cathode?
Your Assignment: pgs 7-8“Quantitative Study of Electrolysis” #1-17
L. Rust and Corrosion

corrosion can be viewed as the process of ______________________
________________________________________________________

the metal is oxidized causing the loss of ________________________
_____________________________________

most metals develop a thin oxide coating which then _____________
________________________________________________________

commonly, the oxide coating will scale off leaving new metal exposed an
susceptible to corrosion

salt will _____________________________________________ by acting as a
________________________________
Chemistry 30 Electrochemistry
25
Cathode: (Reduction) O2(g) + 2H2O(l) + 4e  4 OH(aq)
Fe(s)  Fe2+(aq) + 2e
Anode: (Oxidation)
Net:
M. Prevention of Corrosion
30-B1.1sts  describe the methods and devices used to prevent corrosion; i.e., physical coatings and
cathodic protection
30-B2.2sts  describe science and technology applications that have developed in response to human and
environmental needs

apply a coating of _____________to protect metal from oxidation

other metals (eg. Zn, Cr, Sn) can be _________________ onto metals that you
don’t want to corrode (eg. steel (Fe))

this coating is of a metal that is a ____________________________ than the
metal that is to be protected…the coating metal will react instead and is called the
____________________________________
Fe
Zn coating

this method is also called ___________________________________

has been in use before the science of electrochemistry was developed

Sir Humphrey Davy first used cathodic protection on British naval ships in 1824!

can be used to protect any metal but steel (iron) is most commonly protected

we use steel (iron) for many structures such as buried fuel tanks, septic tanks,
pipelines, hulls of ships, bridge supports etc

to protect these structures by cathodic protection, an ______________
_________________________________ is connected by a ________ to the
structure

because the attached metal is a _______________________________ than the
iron in the steel, the more active metal supplies the
______________________________________________ and therefore the steel
(iron) becomes the cathode and is protected
Chemistry 30 Electrochemistry
26

another protection method is alloying pure metals, which changes their
________________________________________

stainless steel contains chromium and nickel, changing steels reduction potential
to one characteristic of ______________________
_____________________________________ (basically unreactive)

_____________________________ is the process of _____________
______________________________________________________ by
_____________________metal ions in solution

an object can be plated by making it the ____________________ in an
_____________________________________ containing ions of the plating metal
Electrolytic Cell
Chemistry 30 Electrochemistry
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