SCH4U1
UNIT 5  Electrochemistry
1. Relating Science to Technology,
STSE Focus:
F1.1 assess, on the basis of research, the viability of using
electrochemical technologies as alternative sources of energy (e.g., fuel
cells for emergency power generation or as power sources in remote
locations), and explain their potential impact on society and the
environment
Sample issue: Hydrogen fuel cells use hydrogen as the fuel and oxygen as the
oxidant, and produce water, rather than environmentally harmful greenhouse
gases, as waste. Although some cars run on such cells, practical problems
must be resolved before this source of energy is commonly used in the
transportation sector.
Sample questions: What is the capacity of a standard rechargeable battery before
it has to be recharged? What methods should be used to dispose of depleted
batteries? What impact has the use of rechargeable batteries in portable
electronic devices had on society?
F1.2 analyse health and safety issues involving electrochemistry (e.g.,
corrosion of metal pipes in drinking water systems)
Sample issue: Corrosion is a leading cause of structural degradation of bridges
and roadways. Not only does rust weaken metal structures, but as it builds up it
forces apart connecting parts of the structure, causing the structure to fail and
risking public safety. Yet, methods used to prevent corrosion may also
have negative effects on human health.
Sample questions: What health and safety hazards are associated with waste
generated by electroplating companies? Why do metal orthodontic braces not
corrode? What are some of the toxic substances that can escape from electronic
waste into the environment? What are the potential effects of these
poisons on our health?
SCH4U1
UNIT 5 – Electrochemistry (Chapter 9)
9.1 Oxidation and Reduction
Historical definitions:
*Oxidation  to combine with ___________ [e.g. 4Fe(s) + 3O2(g) 
]
*Reduction the opposite of oxidation (i.e. the formation of a _________ from its compounds / the
process of obtaining a metal from an ore) [e.g. 2Fe2O3(s) 
]
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*Ore  a naturally occurring solid compound or mixture of compounds from which a _________ can be
extracted
Modern Definitions: (*based on the idea of ___________ transfers)
*Oxidation  the ________of electrons (LEO =
)
*Reduction  the ________ of electrons (GER=
)
*Redox Reaction (i.e Oxidation – Reduction rxn)  a rxn involving both an _________ and a _________
Example: Redox Reactions (e.g. Single Displacement)
Ex.1: Overall Equation: Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
 Total ionic equ’n:
 Net ionic equ’n:
 Oxidation half – reaction:
[Oxidation(LEO) = e- on _____ side]
 Reduction half – reaction:
[Reduction(GER) = e- on _____side]
 Sum of half –reactions =
More Terms:
*The oxidizing agent (OA) is the entity that causes _____________; it is the entity that gets __________.
In the above example, the OA =________.
*The reducing agent (RA) is the entity that causes _____________; it is the entity that gets __________.
In the above example, the RA = _______.
*A half -reaction is a balanced equation that shows the number of _____________ involved in either
oxidation or reduction. *Two half – reactions are needed to represent a ________ reaction. The two half –
reaction “add up to” the ____________equation.
Practice: a) Write balanced half - reactions for the following. b) Identify the OA and RA in each.
a) Al(s) + Fe3+(aq)  Al3+(aq) + Fe(s)
b) Cu(s) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)
c) 3Zn(s) + Fe2(SO4)3(aq)  3ZnSO4(aq) +2Fe(s)
Try: p. 653 # 2, 3; pp.656-657 #8-11
9.1 con’d - Oxidation States / Numbers
 Oxidation numbers are used to keep track of the electron transfers taking place in redox reactions.
 For reactions involving covalent reactants and products, you cannot use ionic charges to describe the
transfer of electrons. However, to assign oxidation numbers to the elements in a covalent molecule or
polyatomic ion, you can pretend the bonds are ionic.
 Oxidation numbers, therefore, are hypothetical charges, assigned using a set of arbitrary rules. In a
covalently bonded molecule or polyatomic ion, the more electronegative atoms are considered to be
negative and the less electronegative atoms are considered to be positive.
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 The oxidation state of an atom in an entity is defined as the apparent net electric charge that an atom
would have if electron pairs in covalent bonds belonged entirely to the more electronegative atom.
e.g Consider H2O
*Common Oxidation Numbers (Table 1 p. 658)
Atom or ion
Oxidation #
1. All atoms in elements
Examples
Na in Na(s)
Br in Br2(l)
P in P4(s)
2. a) Hydrogen in all compounds
H in HCl
H in H2S
H in CH4
b)Except hydrogen in metal
H in NaH
hydrides
H in CaH2
H AlH3
3. a) Oxygen in all compounds
O in H2O
O in Li2O
O in KNO3
O in MnO4b) Except oxygen in peroxides
O in H2O2
4. All monatomic ions
Na+ is +1
S2- is –2
Ca2+ is +2
*Remember: oxidation numbers are not units of electric charge. For this reason, the sign __________
the number.
*Determining Oxidation Numbers
Step 1: Assign _______________ oxidation numbers.
Step 2: The total of the oxidation numbers of atoms in a molecule or ion ________ the value of the net
electric charge on the molecule or ion.
a) The sum of the oxidation numbers for a compound is ________.
b) The sum of the oxidation numbers for a polyatomic ion equal the __________ on the ion.
Step 3: Any unknown oxidation number is determined algebraically from the sum of the
known oxidation numbers and the net charge on the entity.
Try: p. 659 #12 – 16
Oxidation Numbers and Redox Reactions:
*An oxidation is an ______________ in the oxidation number.
*A reduction is a _____________ in the oxidation number.
*Oxidation numbers __________ when redox reactions take place.
Try: p.662 #18 - 20
9.2 Balancing Redox Reactions
Method #1  Using Oxidation Numbers
Procedure:
Step 1: Assign oxidation numbers and identify the atoms/ions whose oxidation numbers change.
Step 2: Using the change in oxidation numbers, write the number of electrons transferred per atom.
Step 3: Using the chemical formulas, determine the number of electrons transferred per reactant.
Step 4: Calculate the simplest whole number coefficients for the reactants that will balance the total
number of electrons transferred. Balance the reactants and products.
Step 5: Balance the O atoms using H2O(l), and then balance the H atoms using H +(aq)
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*For basic solutions only:
Step 6: Add OH-(aq) to both sides equal in number to the number of H +(aq) present.
Step 7: Combine H+(aq) and OH-(aq) on the same side to form H2O(l), and cancel the same number of H2O(l)
on both sides.
Examples: Balance the following using the oxidation number method.
1. H2S(g) + O2(g)  SO2(g) + H2O(g)
2. ClO3-(aq) + I2(aq)  Cl-(aq) + IO3-(aq)
3. CH3OH(aq) + MnO4-(aq)  CO32-(aq) + MnO42-(aq)
[In acidic solution.]
[In basic solution.]
Try: p. 668 #2 & 3
Method #2  Using Half-Reactions
“Getting Ready” Procedure for Writing Half-Reactions:
Step 1: Write the chemical formulas for the reactants and products.
Step 2: Balance all atoms, other than O and H.
Step 3: Balance O by adding H2O(l).
Step 4: Balance H by adding H+(aq).
Step 5: Balance the charge on each side by adding e- and cancel anything that is the same on both sides.
*For basic solutions only:
Step 6: Add OH-(aq) to both sides equal in number to the number of H +(aq) present.
Step 7: Combine H+(aq) and OH-(aq) on the same side to form H2O(l), and cancel the same number of H2O(l)
on both sides.
*Practice: Writing Half- Reactions. (Try #5 on p. 671)
Method #2 (Balancing Redox Reactions Using Half-Reactions)
Procedure:
Step 1: Separate the skeleton equation into the start of two half-reaction equations.
Step 2: Balance each half-reaction equation.
Step 3: Multiply each half-reaction equation by simple whole numbers to balance the electrons lost and
gained.
Step 4: Add the half-reaction equations, cancelling the electrons and anything else that is exactly the same
on both sides of the equation.
*For basic solutions only:
Step 5: Add OH-(aq) to both sides equal in number to the number of H +(aq) present.
Step 6: Combine H+(aq) and OH-(aq) on the same side to form H2O(l), and cancel the same number of H2O(l)
on both sides.
Examples: Balance the following using the half-reaction method:
1. Fe2+(aq) + CrO72-(aq)  Fe3+(aq) + Cr3+(aq)
2. MnO4-(aq) + C2O42-(aq)  CO2(g) + MnO2(s)
[In acidic solution.]
[In basic solution.
Try: p. 673 #6-8; More Practice p. 673(bottom) # 3 & 4
9.3 Predicting Redox Reactions
*If one particle is able to pull electrons away from another particle, a _________________redox reaction
will occur. (i.e. one that does not require a continuous input of energy in order to continue.)
Ex. 1: Overall Equation
Net equation:
Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
RA = _______; OA =________
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Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)
Ex. 2: Overall Equation:
Net Equation:
RA = _______; OA =________
*Both examples 1 & 2 reactions occur spontaneously.
*However…
Ex. 3: Overall Equation:
Cu(s) + Ni(NO3)2(aq)  No Reaction
*Reaction does not occur spontaneously…why?
*Electron transfers can be thought of as a “competition” for electrons. (“Tug-of-War”) *In example 1, the
oxidizing agent, ______, is able to pull electrons away from _____ atoms. (i.e. Cu2+ ions have a stronger
attraction for Zn’s valence electrons than Zn does.)or … Cu2+ ions are able to __________ Zn atoms.
Therefore:
Zn(s) + Cu2+(aq)  spontaneous
Cu(s) + Zn2+(aq)  nonspontaneous (Zn2+ ions are not able to oxidize Cu atoms)
*In ex.2, Ag+ ions are able to oxidize Cu atoms. (i.e. reaction is _________________)
*In ex.3, Ni2+ ions are unable to oxidize Cu atoms. (i.e. reaction is __________________)
*Is there a way to determine whether a redox reaction will occur spontaneously or not?...without doing
an experiment.
*We can rank metal ions (_____________ agents) according to their ability to react with solid metals
(_____________ agents).
Example: Consider the reactivity of the following metals: Ag, Cu, Pb, and Zn. In an experiment, each of
these metals was put into four different solutions containing various ions.
Summary of results:
Solution
Containing
Reacted with
Ag+
Cu2+
Pb2+
Zn2+
# of Reactions
Reactivity Order
Most ------------------------------------------------------------------- least
*Which of the four oxidizing agents in this example is the strongest?  _____ ; weakest? _____
*Recall: Strong oxidizing agents are “good at” attracting electrons away from other metals; that is, they
are readily reduced.
*What classes of substances (e.g. metals, nonmetals, acidic, basic) usually behave as
a) reducing agents?
b) oxidizing agents?
Constructing a table Relative Strengths of Oxidizing Agents and Reducing Agents
*The strongest oxidizing agent (SOA) is always located in the upper ______ corner of this type of table.
The strongest reducing agent (SRA) (i.e entities that give up electrons readily) is located in the lower
_________part of the table.
*By convention, the half- reactions represented in this type of table are always written as ___________.
*Double arrows are used to indicate that the half-reaction can be read from left to right (reduction) or from
right to left (oxidation). Therefore, the double arrows do not represent chemical ______________ in this
case. For the example above the table would be:
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The Spontaneity Rule: A spontaneous reaction only occurs if the OA is ______________ the RA in a
table of relative strengths of oxidizing agents and reducing agents. (See p. ______ of text.)
1. Practice: From the following evidence, construct a table of relative strengths of OAs and RAs. Label the
SOA and SRA.
a)
 Co2+(aq) + Pd(s)
Co(s) + Pd2+(aq)   
Spontan eous
 Pd2+(aq) + Pt(s)
Pd(s) + Pt2+(aq)   
Spontan eous
 Mg2+(aq) + Co(s)
Mg(s) + Co2+(aq)   
Spontan eous
b)
 Cd2+(aq) + H2(g)
Cd(s) + 2H+(aq)   
Spontan eous
Hg(s) + 2H+(aq)     Hg2+(aq) + H2(g)
Nonspontan eous
 Be2+(aq) + Cd(s)
Be(s) + Cd2+(aq)   
Spontan eous
Be(s) + Ca2+(aq)     Be2+(aq) + Ca(s)
Nonspontan eous
c)
Ag(s) + Br2(l)  AgBr(s)
Ag(s) + I2(s)  no evidence of reaction
Cu2+(aq) + I-(aq)  no evidence of reaction
Br2(l) + Cl-(aq)  no evidence of reaction
2. A student is required to store an aqueous solution of iron(III) nitrate. She has a choice of a copper, tin
iron, or silver container. Use the table of relative strengths of oxidizing and reducing agents in your Data
Booklet to predict which container would be most suitable for storing the solution.
Nelson Try: p. 676 #1-3; p. 678 #10, 11, 13
9.5 Galvanic/Voltaic Electrochemical Cells
*An arrangement of _____ half–cells that can produce electricity ___________________.
*A complete _______can be split into two parts connected by a porous boundary or ______ bridge. Each
part is called a half-cell. (*a group of two or more cells connected in series = a ___________)
*Half-cell  consists of an ____________ (___________ or ___________) and an __________________
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*Ex. of Galvanic Cell:
*Galvanic Cells Connected in Series (i.e. a battery)
Ex. Shorthand Galvanic cell notation:
*Voltaic Pile
Zn(s)  ZnSO4(aq)  CuSO4(aq)  Cu(s)
*Half-cell #1 = Electrode = _________ ; Electrolyte = ___________
*Half-cell #2 = Electrode = _________ ; Electrolyte = ___________
*In a Galvanic cell the anode is labeled ____________ and the cathode is the labeled ____________.
Galvanic Cell Activity: Draw and label a Silver-Copper Galvanic Cell after completing the following:
Shorthand Notation:
Net Notation:
Memorize the Following Procedure/Theory:
*Using Appendix C11 (p.805) determine the strongest oxidizing agent present in the cell. *SOA = ______
*The SOA always undergoes ________________at the ____________. (SOARC)
*Remember: “REDuction CAThode”; Cathode = electron “_________”
*Using Appendix C11 determine the strongest reducing agent present in the cell. *SRA = ______
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*The SRA always undergoes _______________ at the ______________.
*Remember: “ANode OXidation” ; Anode = electron “ ___________”
*Electrons released by ________ travel through a wire to the ___________ where they are accepted by the
_________. The energy produced by the electron flow can be harnessed and used to do work.
*Electrons always move from the __________ towards the ___________. (a c)
*Function of the Salt bridge (“U-tube”) …or porous boundary in a Daniell Cell
 enables electrolyte solution to maintain neutrality (i.e. prevents ___________ ) *Without the salt bridge,
current quickly ceases.
 anions migrate to the __________ via the salt bridge (to offset the buildup of __________)
 cations migrate to the ___________(to offset the buildup of __________as more and more of the OA is
reduced)
*This process gives the impression that negative charges are attracted to the negatively labeled ________,
and that positive charges are attracted to the positively labeled _____________.
Standard Cells and Cell Potentials
*Standard Cell  a galvanic cell in which each half-cell contains all entities shown in the half reaction at
_____________ conditions, with a concentration of _____ mol/L for aqueous solutions
*Standard Reduction Potentials (Er˚)  the Er˚ represents the ability of a standard half-cell to attract
electrons in a reduction half reaction
*Standard Cell Potentials (ΔE˚)  the ____________ (i.e. the maximum potential difference) of a cell
operating under standard conditions
ΔE˚cell =
(i.e. Standard cell potential = the difference between the reduction potentials of the two standard halfcells…the voltage)
*Standard Hydrogen Half-Cell (the reference half-cell)  The Er˚ for all other half-cells are measured
relative to the standard hydrogen half-cell. (2H+(aq) + 2e- ↔ H2(g)
Er˚ = 0.00) *See p. 805
Example: Zn(s)  Zn2+(aq)  H2(g), H+(aq)  Pt(s)
Try: p. 700 #3, 4, 5, 6(a), 7, 8; p. 708 # 10(a), 11, *16
Complete the following:
1. What is the overall reaction for this cell ?
2. Using the appropriate formula, show how the voltage produced by this standard cell is equal to 0.76 V.
*A positive ΔE˚cell represents a spontaneous process. (a negative ΔE˚cell = nonspontaneous)
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3. Predict the voltage produced by the Silver-Copper standard cell drawn previously.
10.1 Electrolysis
*Electrolysis  the process of supplying _____________ energy to force a ____-spontaneous reaction to
occur
*Applications:
i) Production of Elements (e.g. Na, Li, Al, Cl2)
ii) Refining of metals
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iii) Electroplating
*Electrolytic Cell:
- a cell that consists of a combination of two ______________, an _____________, and an external
___________or power source
- power source acts as an “electron pump”, (i.e power source _________ an electron transfer)
Examples:
Electrochemical Cells: Galvanic vs. Electrolytic
Galvanic Cell
Spontaneity
Standard Cell potential (ΔE˚)
Cathode
Electrolytic Cell
Anode
Direction of e- movement
Direction of ion movement
*Try: p. 735 # 1a), 2a)
10.3 Faraday’s Law
 the amount of a substance produced or consumed in an electrolytic reaction is __________ proportional
to the quantity of _____________ that flows through the circuit
*Determining the quantity of electricity charge:
q=
Where: q = quantity of ___________ in ____________;
I = electric ___________ in _____________ (1 A = 1 ______)
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t = time in __________
Examples: Try #1-4 on p.748
*Faraday’s Constant (F)
 the charge of one _______ of __________ (F = ____________________)
*Determining the number of moles of electrons:
ne-=
q
F
; ne-=
It
F
Examples: Try #5-7 on p.749
Half-Cell Stoichiometry
Examples:
1. What is the mass of Aluminum deposited at the cathode of a aluminum electrorefining cell operating at
12.0 A for 40.0 min?
Steps:
i) Write balanced half-reaction.
ii) Convert given measurements to moles
iii) Calculate the amount of the required substance using the mole ratio
iv) Convert the calculated quantity to the quantity required (i.e. answer the question)
2. In a copper electroplating cell, 0.275 g of copper is to be deposited from a copper(II) sulfate solution in a
time of 20.0 min. Predict the current required.
More Practice: p. 751 # 8; p. 753 #3, 4, 5, 7
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