Electrochemistry
Objectives:
1. Be able to balance net ionic equations for redox reactions:
(Half Reaction Method-Covered in Lab).
2. Understand what we mean by standard reduction potential (SRP).
3. Be able to determine cell voltages using SRP.
4. Understand Voltaic cells:
The Nernst equation and cell voltage under nonstandard conditions
The “good”-batteries and fuel cells
The “bad”-metal corrosion (rust) and methods used to inhibit corrosion
5. Understand electrolytic cells:
Electroplating
Quantitative calculations
6. Understand the relationship between cell voltage, thermodynamics and
equilibria:
∆G° and equilibrium constants.
7. Understand terms used for electrical power:
My PG&E bill is too high! What’s a watt anyway? What’s a kilowatt-hour?
Electrochemistry
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Review of Terms & Half Reactions
•
Oxidation: the loss or apparent loss of one or more electrons. An increase in oxidation number for
the atom.
Example: Na(s) —> Na+(aq) + e-
•
Reduction: the gain or apparent gain of one or more electrons. A decrease in oxidation number for
the atom.
Example: F2(g) + 2 e- —> 2 F-(aq)
•
Oxidizing Agent (oxidant): a species that will oxidize (take electrons from) another species. An
oxidizing agent is reduced as it oxidizes the other species.
Example: O2(g) + 4 H+(aq) + 4 e- —> 2 H2O(l)
•
Reducing Agent (reductant): a species that will reduce (give electrons to) another species. A
reducing agent is oxidized as it reduces the other species.
Example: Al(s) —> Al3+(aq) + 3 e-
•
Remember that the oxidation number of an atom does not generally equal the actual charge on the
atom, it is a convenient form of “electron bookkeeping” that enables us to decide if a redox reaction
has occurred. For covalently bonded atoms, oxidation numbers are assigned as if the electrons in
the bond “belong” to the more electronegative element.
•
Oxidation and reduction always occur together.
The number of electrons lost in oxidation must equal the number of electrons gained in reduction.
CHARGE BALANCE MUST BE MAINTAINED.
Electrochemistry
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Standard Reduction Potentials (SRP’s)
Electrochemistry
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Properties of E°red & SRP’s
1. All potentials are for the reduction
half-reaction.
2. The more (+) E°red the more
favorable is the reduction. Hence a
better oxidizing agent.
3. When the reaction is reversed to
indicate oxidization the sign of the
voltage is changed.
4. Oxidizing agents are the
REACTANTS on the LEFT side of
the table. STRONGEST at the top,
WEAKEST at the bottom.
5. Reducing agents are the
PRODUCTS on the RIGHT side of
the table. WEAKEST at the top,
STRONGEST at the bottom.
A spontaneous redox reaction will occur under
standard conditions when:
Any REACTANT (oxidizing agent) is combined
with any PRODUCT (reducing agent)
LOWER in the table.
The value of E°rxn will always be (+) under
these conditions:
E°rxn = E°red + E°ox
Spontaneous Rxns (positive E°rxn):
Cu(s) with Fe3+(aq)
Br2(l) with Ag(s)
Zn2+ with Al(s)
Non-spontaneous Redox Rxns (negative E°rxn):
Fe2+(aq) with Cu2+(aq)
Ni2+(aq) with Ag(s)
Al3+(aq) with Zn(s)
Electrochemistry
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Using the Table of SRP’s
1.
Choose the best oxidizing agent between
Fe2+(aq), Fe3+(aq) and Na(s).
2.
Choose the best reducing agent between
Fe2+(aq), Zn2+(aq) and Cu(s).
3.
Rank the halogens in order of their
strength as oxidizing agents.
4.
What halogen(s) can oxidize Ag?
5.
What halogens are stable in water?
6.
What metals in the table are unstable in
water?
7.
Based upon standard reduction
potentials, it appears as if MnO4– ions
should be spontaneously reduced by
water. However, we find that
permanganate is actually stable in water
solution. Why is this?
Copper dissolves in a nitric acid solution
with the evolution of NO(g). In contrast,
copper does not dissolve in a hydrochloric
acid solution. Explain these observations.
Write the balanced net ionic equation for
the reaction that occurs when copper is
placed in a nitric acid solution.
8.
The oxidation of Zn(s) with Cu2+(aq) ions.
Electrochemistry
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Predicting Spontaneous Redox Rxn’s, E°rxn
Determine E°rxn for the following reactions. Is the
reaction spontaneous as written?
E°rxn = E°red + E°ox
1. 2 H+(aq) + Cu(s) —> Cu2+ (aq) + H2(g)
2. 2 Fe3+(aq) + 2 Br–(aq) —> 2 Fe2+(aq) + Br2(l)
Use SRB’s to find the potential for the following redox reactions:
1.
Zn(s) + 2 H+(aq) <—> Zn2+(aq) + H2(g)
2.
2 Ag+(aq) + Cu(s) <—> Cu2+(aq) + 2 Ag(s)
3.
Zn(s) + Cu2+(aq) <—> Zn2+(aq) + Cu(s)
Electrochemistry
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How do we Measure E°rxn?
We use a VOLTAIC CELL to measure the emf
(voltage) between two half-reactions. A voltaic cell
has the following construction:
1. Anode 1/2 cell
2.
Cathode 1/2 cell
3.
4.
Salt bridge
Voltmeter or load
Play Movie
Electrochemistry
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Understanding How Voltmeters Work
Important for the Voltaic Cell Experiment!
How do we know which half-cell is the cathode and which is the anode? We use the sign
of the voltage and the placement of the voltmeter probes.
Remember: the voltmeter measures a voltage difference between the two probes The
sign of the voltage is interpreted as follows:
Observed (+) voltage:
The 1/2 cell at the red probe has a higher
reduction potential (greater tendency to be
reduced) than the 1/2 cell at the black probe.
Red probe = cathode (where reduction occurs)
Black probe = anode (where oxidation occurs)
red
Anode
Cathode
1/2-cell
1/2-cell
Observed (–) voltage:
The 1/2 cell at the red probe has a lower
reduction potential (less tendency to be
reduced) than the 1/2 cell at the black probe.
Red probe = anode (where oxidation occurs)
Black probe = cathode (where reduction
occurs)
+ voltage
blk
blk
– voltage
red
Cathode
Anode
1/2-cell
1/2-cell
Electrochemistry
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Simple Voltaic Cells and Voltmeters, Ecell = Erxn
Two half-cells connected with a
salt bridge. The salt bridge
maintains charge neutrality in
each half-cell.
Reference
•
Anode (-):
•
Cathode (+):
1/2 Cell
Black probe
Red probe
This system can do work! One
Coulomb of electrons loses 0.46
J of potential energy from
anode to cathode.
The cathode is always at a
lower potential energy than
the anode in a voltaic cell.
The cathode always has a
higher reduction potential
than the anode in a voltaic
cell.
1M
Solutions
Electrochemistry
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Standard Reduction (Half-Cell) Potentials
To determine the emf of a half-reaction we use the standard hydrogen electrode (SHE) as
the reference 1/2 cell. Thus, all TEXTBOOK measurements of reduction 1/2 reactions are
taken relative to the SHE. By convention the SHE 1/2 cell emf is set to 0 V.
Measurements are made under standard conditions for each half-cell:
Solid and liquid reactants and products are pure.
Solutes are 1 M.
Gaseous species are 1 atm (1 bar).
SHE: 2 H+(aq, 1M) + 2 e- —> H2(g, 1 atm) E°red= 0 V
The standard hydrogen
electrode uses an inert
electrode (one that just
serves to carry the
electrons). The most
common inert electrodes
used are Pt metal and
carbon (graphite).
Electrochemistry
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Voltaic Cells Using a SHE, E°cell
Play Movie
From the textbook:
E°cell = E°red (cathode) + E°ox (anode)
or
E°cell = E°red (cathode) – E°red (anode)
E°ox = +0.76 V
E°red = 0.00 V
Spontaneous redox:
2 H+(aq) + Zn(s) —> H2(g) + Zn2+(aq)
E°cell = (0.00 V) + (+0.76 V) = +0.76 V
Nonspontaneous redox:
H2(g) + Zn2+(aq) —> 2 H+(aq) + Zn(s)
E°cell = (-0.76 V) + (0.00 V) = –0.76 V
Electrochemistry
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Understanding Electromotive Force, emf, and the Volt
There is an electromotive force (“causing electron motion”) that pulls/pushes the electrons in a
redox reaction.
Reduction 1/2 rxn voltage: The “pull” of free electrons relative to SHE.
Oxidation 1/2 rxn voltage: The “push” of free electrons relative to SHE.
The emf of a “cell”, E°cell, is the difference in emf between the two half-reactions in the cell. There
are no “absolute” emf measurements.
We measure this emf as a voltage, the difference in e– potential energy per unit charge (Joule/
Coulomb).
1V =
1J(energy)
; 1C = the (+) charge of 6.241x1018 e −
1C(charge)
Coulombs are used to count total charge or electrons in the system. What is the total
charge in coulombs of a mole of electrons?
The emf of a “cell”, E°cell, is also called the cell potential or cell voltage.
NOTE: Cell voltage( E°cell) is an INTENSIVE property (J/C), therefore does not depend upon the extent of the reaction. Changing
the balancing coefficients in a 1/2 reaction does not change the value of E°red or E°ox. Likewise, changing the balancing
coefficients in a balanced redox reaction does not change E°cell.
Electrochemistry
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Shorthand Cell Notation for Voltaic Cells
Examples:
Balanced redox: Short-hand Notation:
Cu(s) + 2 Ag+(aq) —> Cu2+(aq) + 2 Ag(s)
Cu(s)|Cu2+(aq, 1M) || Ag+(aq, 1M)|Ag(s)
Anode
Cathode
1.
The anode half-cell is usually written on the left side, cathode on the right.
2.
3.
4.
A single vertical line represents phase boundaries within a 1/2 cell.
A comma is used to separate half-cell components within the same phase that appear in
the balanced half reactions.
A double vertical line separates the half-cells.
5.
H2O is generally omitted, unless it is being oxidized or reduced.
6.
Inert electrodes are designated when present.
•
Example redox: MnO4–(aq) + H+(aq) + Fe(s) —> Mn2+(aq) + Fe3+(aq) + H2O(l)
Short-hand Notation: Fe(s)| Fe3+(aq, 1 M) || MnO4– (aq, 1M), H+(aq,1 M), Mn2+(aq, 1M) |Pt
Write the shorthand notation for the standard Zn/Zn2+, SHE voltaic cell.
Electrochemistry
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Voltaic Cell Diagrams
Consider the following voltaic cell:
Fe
1 M Fe2+
a)
Where does oxidation occur?
b)
Where are electrons consumed?
c)
Which electrode is labeled (–)?
d)
In which direction do electrons flow?
Ni
1 M Ni 2+
e)
Suggest a solution for the anode electrolyte.
f)
Suggest a pair of ions for the salt bridge.
g)
Which electrode will decrease in mass?
h)
What is E°cell?
i)
Write the balanced net-ionic chemical equation for the overall reaction for this cell.
Electrochemistry
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Constructing a Voltaic Cell
A voltaic cell consists of a strip of lead in a solution of 1M Pb(NO3)2 in one
beaker, and in the other beaker a platinum electrode is immersed in a
1M NaCl solution, with 1 atm Cl2 gas bubbled around the electrode.
The two beakers are connected with a salt bridge. Assume standard
conditions.
1.
Which electrode serves as the anode, and which as the cathode?
2.
Does the Pb electrode gain or lose mass as the cell reaction occurs?
3.
Write the net-ionic chemical equation for the overall cell reaction.
4.
What is the emf (voltage) generated by the cell under standard
conditions?
Electrochemistry
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