Electrochemistry

Electrochemistry is a branch of chemistry that deals with
electrically related applications of redox reactions.
•
•

Reduction-oxidation reactions involve the transfer of
electrons.
Oxidation means the losing of electrons and reduction
means the gaining of electrons. The 2 occur together, they
are opposite sides of the same coin.
For example: when zinc is in contact with a copper II
sulfate solution, the zinc strip loses electrons to the
copper ions in solution. The copper ions accept the
electrons and fall out of solution. As electrons are
transferred between zinc atoms and copper ions energy
is released as heat – when the reactions are separated
we can set them up so that instead of heat energy we can
get electrical energy.
Example 1

Net ionic equation:
Zn (s) + Cu+2 (aq)
Oxidation:
Reduction:

Cu (s)
+
Zn+2 (aq)
Voltaic (Galvanic) Cells


An electrochemical cell, such as a voltaic cell,
consists of 2 electrodes. Each electrode is in
contact with an electrolyte. The 2 electrodes
are connected by a conducting wire or a circuit.
And a porous barrier separates the 2 half
reactions (or half cells).
A voltaic cell specifically deals with a
spontaneous redox reaction as the source of
energy. It converts chemical energy into
electrical energy.
How it works……

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In the “wet” voltaic cell represented in the previous slide
an electric current can run through an external connecting
wire so that the electric current moves in closed loop path
(closed circuit).
The electrode where oxidation occurs is called the anode.
The electrode where reduction occurs is called the
cathode.
The 2 half-reactions occur at the same time but in
different places at the cell (the porous barrier separates
them). A salt bridge is necessary to keep the half cells
electrically balanced so that a charge does not build up in
the cell and stop the electrochemical reaction prematurely
– this salt bridge allows for the passage of ions in the cell.


ANode, OXidation; REDuction, CAThode
AN OX and a RED CAT:
Zinc & Copper
Cell Notation

When these cells are represented they
are written as follows:
Anode electrode anode solution

cathode solution
cathode electrode
Example 2: Write the cell notation for
the following reaction:
Zn (s) + Cu+2 (aq)

Cu (s)
+
Zn+2 (aq)
Zn(s)|Zn2+ (aq) || Cu2+ (aq) | Cu(s)
Inactive Electrodes
graphite| I−(aq) | I2(s) || H+(aq), MnO4−(aq), Mn2+(aq) |graphite
Practice Problems

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•
•
•
•
Write the half reaction in which I- (aq)
changes to I2 (s). Identify if this occurs at the
anode our cathode.
Nickel solid is oxidized in to Ni+2 ions in a
voltaic cell while the Cu+2 ion are being
reduced in to copper solid atoms.
Write the half reactions
Write the net ionic equation
Identify the anode and cathode
Write the cell notation
Cell Voltage

The cell voltage from a redox reaction is
referred to as the standard voltage, Eo (unit
volts, V).
• Standard conditions are 1 atm and 1 M solutions
• Example 3: Zinc metal is placed in hydrochloric
acid. Zinc is the anode and hydrogen gas forms
at the cathode. The reaction gives off a standard
voltage of 0.762 V. Write the net ionic equation
for this reaction:
Standard Voltage

When calculating the standard voltage
the standard voltage from the reduction
and oxidation reactions must be
considered. So when calculating the
standard voltage you must use the
following formula:
Eo = Eored + Eoox
Standard Potentials


The standard half cell (half redox
reaction) voltages are referred to as
standard potentials and are used to
calculate the standard voltage.
Standard Reduction Potentials
(one of your equation sheets)
• This table gives the standard reduction
potentials, the standard oxidation potentials
are the same magnitude but the reverse sign.
When Calculating Cell Voltage

When calculating cell voltages there are
2 main points to remember:
• 1. the calculation of E , is always a positive
•
quantity for a voltaic cell (spontaneous
reaction).
2. The standard cell voltage is independent of
how the equation for the cell reaction is
written. This means you must never multiply
the voltage by the coefficients used to
balance the chemical equation.
Example
Example 4: Use standard reduction potentials
to calculate the standard voltage for the ZnH+ cell from example 3.

The Pain of a Dental
Voltaic Cell
•
Have you ever felt a jolt of
pain when biting down with
a filled tooth on a scrap of
foil left on a piece of food?
Here’s the reason. The
aluminum foil acts as an
active anode (E° of Al = −
1.66 V), saliva as the
electrolyte, and the filling
(usually a
silver/tin/mercury alloy) as
an inactive cathode. O2 is
reduced to water, and the
short circuit between the
foil in contact with the
filling creates a current
that is sensed by the nerve
of the tooth.
Finding Eored or Eoox

If the standard voltage and the cell
voltage from either the reduction or
oxidation reaction the other maybe found
by a simple rearrangement.
• Example 4: If the standard voltage gathered
from the standard Zn-Cu+2 cell is 1.101 V and
the Eoox = 0.762 V, then find the Eored.
Reducing & Oxidizing Agents


If a species undergoes reduction (gains
electrons) then it is the oxidizing agent. If it
undergoes oxidation (loses electrons) then it is
the reducing agent.
The stronger the attraction for electrons the
stronger the oxidizing agent. Or if using the
standard reduction potentials, the more
positive the Eored the stronger the oxidizing
agent (oxidizing strength would be the opposite
if using the reduction potential table).
Cell Voltage Gibbs Free Energy & Equilibrium

Standard cell voltage and standard free energy are
related by the following equation:
Go = -nFEo
•
When Go < 0 and Eo > 0 the reaction is spontaneous.
Standard cell voltage and equilibrium are related by the
following equation:
Eo = RTlnK
nF
or at standard conditions (25o C)
Eo = 0.0257 V lnK
n

•
When K > 1 the reaction is spontaneous
Effect of Concentration


Voltage will increase for a reaction if the
concentration of the reactants is
increased or that of the products is
decreased. This makes the reaction
more spontaneous.
Voltage will then decrease if the
concentration of the reactants is
decreased or that of the products is
increased. This makes the reaction less
spontaneous.
The Nernst Equation


Offers a quantitative relationship
between cell voltage & concentration:
E = Eo - RT lnQ
nF
or at standard conditions (25oC)
E = Eo - 0.0257 V lnQ
n
Interpreting Q
• If Q > 1
- concentration of the products
are high so E < Eo

(meaning lnQ is positive)
• If Q < 1
- concentration of the reactants
are high so E > Eo

(meaning lnQ is negative)
• If Q = 1
- reaction at standard conditions
for cell voltage so E = Eo

(meaning lnQ = 0)
Electrolytic Cells

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An electrolytic cell is a non-spontaneous
redox reaction that made to occur by
pumping electrical energy into the
system.
When carried out in an electrochemical
cell this is referred to as electrolysis.
This is the procedure used when
electroplating. Electrons are pushed into
the cathode and removing them from the
anode.
Quantitative Relationships


Quantitative relationships between the amount
of electricity passed through an
electrochemical cell:
For the reaction Cu+2 (aq) + 2e-  Cu (s)
2 mol e- = 1 mol Cu (s) = 63.55 g Cu
•
Coloumb (C) – the quantity of electrical charge (or
electrical current).
• 1 mol e- = 9.648 x 104 C
•
Determining Current flow:
•
Joule (J) – the amount of electrical energy, 1 J = 1 C*V
I = q/t
• Ampere (A) – unit for the rate of current flow, (1 A = 1
C/s)
• Current (C) over time (s)
Electrochemistry Stoichiometry
From left to right, Walther Nernst, Albert Einstein, Max Planck, Robert Millikan, and
Max von Laue.
FRQ #1
It is observed that when silver metal is placed in aqueous thallium(I) fluoride, TlF,
no reaction occurs. When the switch is closed in the cell, the voltage reading is
+1.14 V.
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(a) Write the reduction half-reaction that occurs in the cell.
(b) Write the equation for the overall reaction that occurs in the cell.
(c) Identify the anode in the cell. Justify your answer.
(d) On the diagram above, use an arrow to clearly indicate the direction of
electron flow as the cell operates.
(e) Calculate the value of the standard reduction potential for the Tl+/Tl halfreaction.
The standard reduction potential, E°, of the reaction Pt2+ + 2 e− → Pt is 1.20 V.
 (f) Assume that electrodes of pure Pt, Ag, and Ni are available as well as 1.00 M
solutions of their salts.
Three different electrochemical cells can be constructed using these materials.
Identify the two metals that when used to make an electrochemical cell would
produce the cell with the largest voltage. Explain how you arrived at your
answer.
 (g) Predict whether Pt metal will react when it is placed in 1.00 M AgNO3(aq).
Justify your answer.
FRQ #2
2 H2(g) + O2(g)  2 H2O(l)


In a hydrogen-oxygen fuel cell, energy is produced by the
overall reaction represented above.
(a) When the fuel cell operates at 25˚C and 1.00 atm for
78.0 minutes, 0.0746 mol of O2(g) is consumed. Calculate
the volume of H2(g) consumed during the same time period.
Express your answer in liters measured at 25˚C and 1.00
atm.
(b) Given that the fuel cell reaction takes place in an acidic
medium,
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•

(i) write the two half reactions that occur as the cell operates,
(ii) identify the half reaction that takes place at the cathode, and
(iii) determine the value of the standard potential, E˚, of the cell.
(c) Calculate the charge, in coulombs, that passes through
the cell during the 78.0 minutes of operation as described in
part (a).
FRQ #3
An external direct-current power supply is
connected to two platinum electrodes immersed
in a beaker containing 1.0 M CuSO4(aq) at
25˚C, as shown in the diagram above. As the
cell operates, copper metal is deposited onto
one electrode and O2(g) is produced at the other
electrode. The two reduction half-reactions for
the overall reaction that occurs in the cell are
shown in the table below.
(a)
Half-Reaction
E0(V)
O2(g) + 4 H+(aq) + 4 e-  2 H2O(l)
+1.23
Cu2+(aq) + 2 e-  Cu(s)
+0.34
On the diagram, indicate the direction of electron flow in the wire.
(b) Write a balanced net ionic equation for the electrolysis reaction that occurs in the cell.
(c) Predict the algebraic sign of ∆G˚ for the reaction. Justify your prediction.
(d) Calculate the value of ∆G˚ for the reaction.
An electric current of 1.50 amps passes through the cell for 40.0 minutes.
(e) Calculate the mass, in grams, of the Cu(s) that is deposited on the electrode.
(f) Calculate the dry volume, in liters measured at 25˚C and 1.16 atm, of the O2(g) that is
produced.