Transcript Electrochemistry - Waterford Public Schools

Chapter 17
AP Chemistry
• As you watch the following video, answer the questions on the
sheet provided:
Decoding the Past - The Real Dr. Frankenstein
• Oxidation-reduction reaction (Redox)
• Involves a transfer of electrons from the reducing agent to the oxidizing agent
• Oxidation
• Loss of electrons
• LEO
• OIL
• Reduction
• Gain of electrons
• GER
• RIG
•
•
•
•
Oxidation number
Oxidizing agent
Reducing agent
Half-reactions
• Overall reaction is split into two half-reactions, one involving oxidation and one reduction
Work on the sample Review Problem
• Balance the following oxidation-reduction reaction IN BASIC
SOLUTION using the half-reaction method. Be sure to identify
the oxidizing agent and reducing agent.
HXeO4- (s)  XeO64- + Xe (g)
• Luigi Galvani
• Italian physician who observed a frog’s leg twitch
when it was touched with two different metals
• In attempting to explain what happened, Galvani thought
that the animal tissue in the frog’s leg was the source of
electricity
• Alessandro Volta
• Italian physicist who disputed Galvani’s hypothesis
• Resulting controversy resulted in discovery that electric
currents could be produced by chemical reactions
• Volta used this discovery to create the first chemical
battery
• In general, all chemistry is electrical in the sense
that it involves the behavior of electrons and
other charged particles
•The term electrochemistry is reserved
specifically for the study of the interchange of
chemical and electrical energy
• All of the following involve the principles of
electrochemistry:
• Remote controls for TVs, DVD players, CD players, stereos
• Itty bitty teeny tiny batteries
• Calculators
• Silverware
• Metal-plated jewelry
• Defined as a device in
which chemical energy is
changed to electrical
energy
• Examples – batteries and
fuel cells
• Name comes from the work
of Volta and Galvani
• Uses a spontaneous redox
reaction to produce a
current that can be used to
do electrical work
• Electrons flow from one terminal to the other when the terminals are
connected by an external circuit
• Terminals are called electrodes
• Anode
• Oxidation occurs here
• Electrons leave the cell here
• Cathode
• Reduction occurs here
• Electrons are accepted by the species being reduced and enter the
cell here
• Electrodes are submerged in an electrolyte
• A salt solution that contains ions
• Electrolyte may be involved in the reaction or the ions may be used
to carry the charge
• Can contain a salt bridge or a porous-disk connection in order to neutralize
charge buildups in electrode compartments
• Completes the circuit!
• Anode
• Oxidation occurs here
• Cathode
• Reduction Occurs here
Salt Bridge
Porous Disk
Galvanic Cells
• AN OX
• Oxidation occurs at the anode
• RED CAT
• Reduction occurs at the cathode
• FAT CAT
•The electrons in a galvanic (voltaic) cell always
flow From the Anode To the CATode
• Recall that a galvanic cell
consists of an oxidizing
agent in one compartment
that “pulls” electrons
through a wire from a
reducing agent in the other
compartment
• The “pull” or driving force on
the electrons is called the
CELL POTENTIAL or
electromotive force (emf)
• If pull occurs
spontaneously, cell is a
good battery!
• A VOLTMETER is used to measure cell potential
• The unit of electrical potential is the volt (V)
• Defined as 1 joule of work per coulomb of charge transferred
• The cell potential (always positive for a galvanic cell) and the
balanced cell reaction is written somewhere on the diagram
• The direction of electron flow is given
• Obtained by inspecting the half-reactions and using the direction that
gives a positive E0cell
• The anode and cathode are designated
• The nature of each electrode and the ions present in each
compartment are labeled
• A chemically inert conductor such as Pt is required if none of the
substances participating in the half-reaction is a conducting solid
• Example – Fe2+ and Fe3+
• Reaction in a galvanic cell is always an oxidationreduction reaction that can be broken down into two
half-reactions
• Each half-reaction has a cell potential
• We can obtain the overall cell potential by summing the halfcell potentials!
• A cell will always run spontaneously in the direction that
produces a POSITIVE cell potential
• Each potential is measured against a standard called
the STANDARD HYDROGEN ELECTRODE (SHE)
• SHE consists of a piece of inert Platinum that is bathed by
hydrogen gas at 1 atm
• SHE is assigned a potential of ZERO volts
Reduction Half-Reaction
Standard Hydrogen Electrode
• Half-reaction cell potentials are listed in a convenient table!
• The values in the table correspond to REDUCTION halfreactions with all solutions at 1 M, all gases at 1 atm, and 25°C
(298K) for all
Cu2+ + 2 e- → Cu
Symbol for
standard
E0 = -0.34 V versus SHE
conditions!
SO42- + 4 H+ + 2 e- → H2SO3 + H2O
E0 = 0.20 V versus SHE
• Elements that have the MOST POSITIVE reduction potentials are
easily REDUCED
• In general, non-metals
• Elements that have the LEAST POSITIVE reduction potentials are
easily OXIDIZED
• In general, metals
• Table can also be used to tell the strength of various oxidizing
and reducing agents
• Another form of the activity series
• Metals having LESS POSITIVE reduction potentials are MORE active
and will replace metals with more positive potentials
1. Decide which element is oxidized or reduced using the table
of reduction potentials
•
THE MORE POSITIVE REDUCTION POTENTIAL GETS TO BE REDUCED
2. Write both equations AS IS from the chart with their voltages
3. REVERSE the equation that will be OXIDIZED and change the
sign of the voltage!
•
This is now E0oxidation
4. Balance the two half-reactions using integers
• Number of electrons lost must equal number gained
• DO NOT MULTIPLY VOLTAGE VALUES
5. Add the two half reactions and the voltages together
0
0
0
𝐸𝑐𝑒𝑙𝑙
= 𝐸𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛
+ 𝐸𝑟𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛
At 25ºC (298 K)
• Consider a galvanic cell based on the reaction:
𝐴𝑙3+ 𝑎𝑞 + Mg s → 𝐴𝑙 𝑠 + 𝑀𝑔2+ (aq)
• Give the balanced cell reaction and calculate E0 for the cell
• Line notation can be thought of as an “Ion Sandwich” in alphabetical
order
Anode metal | Anode ion || Cathode ion | Cathode metal
• “|” indicates phase boundary (solid → solution or gas or solution or gas →
solid)
• “||” indicates salt bridge
𝑀𝑔 𝑠
𝑀𝑔2+ 𝑎𝑞
𝐴𝑙 3+ 𝑎𝑞
𝐴𝑙 (𝑠)
• Phase (AND concentration if not 1M) specified in parentheses
• ZnSO4 (aq, 0.5M)
• A comma should be used to separate 2 components in the same
phase
𝑃𝑡 𝑠 𝐹𝑒 2+ 𝑎𝑞 , 𝐹𝑒 3+ (𝑎𝑞) 𝐴𝑔+ 𝑎𝑞 𝐴𝑔 (𝑠)
• Calculate the cell voltage for the following reaction. Draw a
diagram of the galvanic cell for the reaction and label
completely
• See previous slide for requirements of a complete diagram
Fe3+ (aq) + Cu (s) → Cu2+ (aq) + Fe2+ (aq)
Combining thermodynamics, electrochemistry, and not to mention a bit of physics!
• The work that can be accomplished when electrons are transferred
through a wire depends on the “push” or emf (electromotive force)
• emf is defined in terms of potential difference (in volts) between two points in
the circuit
• Recall that a volt represents a joule of work per coulomb of charge
transfered
𝑒𝑚𝑓 =
𝐸0
𝑤𝑜𝑟𝑘 (𝐽)
= 𝑝𝑜𝑡𝑒𝑛𝑡𝑖𝑎𝑙 𝑑𝑖𝑓𝑓𝑒𝑟𝑒𝑛𝑐𝑒 𝑉 =
𝑐ℎ𝑎𝑟𝑔𝑒 (𝐶)
• Thus, one joule of work is produced (or required) when one coulomb
of charge is transferred between two points in the circuit that differ
by potential of one volt
• Work is viewed from the point of view of the system
• Therefore, if work flows OUT of the system, it is assigned a MINUS sign
• When a cell produces a current (aka a battery), the cell
potential is positive and the current can be used to do work (like
running a motor)
• Therefore, emf and work have opposite signs!
𝑒𝑚𝑓 =
𝐸0
−𝑤𝑜𝑟𝑘 (𝐽) −𝑤
=
=
𝑐ℎ𝑎𝑟𝑔𝑒 (𝐶)
𝑞
𝑤=
0
−𝑞𝐸
• q is the quantity of charge in coulombs transferred
• The charge of 1 mole of electrons is a constant called the faraday (F)
• Has the value of 96,485 coulombs of charge per mole of electrons
𝐶
𝐹 = 96,485
𝑚𝑜𝑙 𝑒 −
• So,
−
𝑞 (𝐶) = # 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑒 × 𝐹 = 𝑛𝐹
• In chemistry, we often refer to processes as either spontaneous
or nonspontaneous
• A spontaneous process is said to occur if it occurs without outside
intervention
• Has nothing to do with the speed of the reaction
• To explore the idea of spontaneity, consider the following
physical and chemical processes:
• A ball rolls down a hill but never spontaneously rolls back up the hill
• If exposed to air and moisture, steel rusts spontaneously. However, the
iron oxide in rust does not spontaneously change back to iron metal and
oxygen gas
• Heat flow always occurs from a hot object to a cooler one. The reverse
process never occurs spontaneously
• The driving force for a spontaneous process is an increase in
the entropy of the universe
• Entropy can be viewed as a measure of molecular randomness or disorder
• Natural progression of things is from order to disorder (from lower
entropy to higher entropy)
• Entropy is related to another thermodynamic quantity called
Gibb’s Free Energy (symbolized by G) – more on this later!
• A process at constant temperature and pressure is spontaneous in the
direction in which the free energy decreases
−∆𝐺 = +∆𝑆𝑢𝑛𝑖𝑣
• Gibb’s free energy is qualitatively useful by telling us whether a
process is spontaneous or not
• It is quantitatively useful because it can tell us how much work
can be done with a given process
• Thus, G is defined as the energy available in a system that is available to
do useful work
• Maximum possible work obtainable from a process at constant
temperature and pressure is equal to the change in free energy:
𝑤𝑚𝑎𝑥 = ∆𝐺
• ∆G for a spontaneous process represents the energy that is free
to do useful work
• ∆G for a nonspontaneous process represents the minimum
amount of work that must be expended to make the process
occur
• Recall that for a galvanic cell:
𝑤 = −𝑞𝐸
0
• And:
𝑞 = 𝑛𝐹
• Since:
𝑤𝑚𝑎𝑥 = ∆𝐺
• We can make some substitutions to come up with a relationship between
Gibb’s Free Energy, work, and cell potential at constant temperature and
pressure:
∆𝐺 0 = 𝑤𝑚𝑎𝑥 = −𝑞𝐸 0 = −𝑛𝐹𝐸 0
0
∆𝐺
=
0
−𝑛𝐹𝐸
G = Gibb’s Free Energy
n = number of moles of electrons
F = Faraday constant = 96, 485 coulombs per mole of electrons
• This relationship is important because it confirms that a galvanic
cell will run in the direction that gives a positive value for E0
• +E0 corresponds to a negative ∆G value (spontaneous)
• -E0 corresponds to a positive ∆G value (nonspontaneous)
Using the Table of Standard Reduction Potentials, calculate ∆G0
for the reaction:
𝐶𝑢2+ 𝑎𝑞 + 𝐹𝑒 𝑠 → 𝐶𝑢 𝑠 + 𝐹𝑒 2+
Is this reaction spontaneous?
• So far, we have described galvanic cells under standard
conditions
• All solutions at 1 M, all gases at 1 atm, and 25°C (298K) for all
• What would happen to the cell potential if the solutions were
not at 1M?
• Can be answered qualitatively in terms of Le Chȃtelier’s Principle
• Virtually everything you encounter, including your own
bodily processes and senses, is the result of one or
more equilibrium reactions
• Chemical equilibrium can be defined as the condition where
reactant and product concentrations remain constant
• Occurs when a forward reaction and its reverse reaction
proceed at the same rate
• While concentrations do not change, products and reactants
continue to interconvert at equal rates
• If a system at equilibrium is disturbed by a
change in temperature, pressure, or the
concentration of one of the components,
the system will shift its equilibrium position
so as to counteract the effect of the
disturbance
• An increase in reactant concentration will favor the forward
reaction and thus, increase the driving force on the electrons
• E will increase
• An increase in product concentration will oppose the forward
reaction and thus, decrease the driving force on the electrons
• E will decrease
• For the cell reaction:
2 𝐴𝑙 𝑠 + 3 𝑀𝑛2+ 𝑎𝑞 → 2 𝐴𝑙3+ 𝑎𝑞 + 3 𝑀𝑛 (𝑠)
Predict whether E is larger or smaller than E0 for the following
cases:
• [Al3+] = 2.0 M, [Mn2+] = 1.0 M
• [Al3+] = 1.0 M, [Mn2+] = 3.0 M
• Because cell potentials
depend on
concentration, we can
construct galvanic cells
where both
compartments contain
the same components but
at different
concentrations
• An increase in
concentration of reactant
will increase cell potential
• Increase in product
concentration will decrease
cell potential
• This equation is used to calculate the potential of a cell in which some
or all of the components are not in their standard states
• Remember, standard states are 1M and gases at 1 atm
𝑅𝑇
𝐸=𝐸 −
ln 𝑄 @ 25℃ (298 𝐾)
𝑛𝐹
0
𝑜𝑟
0.0591
𝐸=𝐸 −
log 𝑄 @ 25℃ (298 𝐾)
𝑛
0
𝑅𝑇
𝐸=𝐸 −
ln 𝑄
𝑛𝐹
0
0.0591
𝐸=𝐸 −
log 𝑄
𝑛
0
• R = gas constant 8.315 J/K·mol
• F = Faraday constant
• 96, 485 colombs per mole of electrons
•
•
•
•
E = Energy produced by reaction
T = Temperature in Kelvin
n = number of electrons exchanged in BALANCED redox equation
Q (Reaction Quotient)
• A reaction quotient is an expression that is obtained by applying the
law of mass action
• The rate of a chemical reaction is directly proportional to the products of the
reactants
• Expression uses using initial concentrations of substances rather than
equilibrium concentrations
• Always written like so:
𝐶𝑜𝑒𝑓𝑓𝑖𝑐𝑖𝑒𝑛𝑡
[𝑃𝑟𝑜𝑑𝑢𝑐𝑡𝑠]
[𝑅𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠]𝐶𝑜𝑒𝑓𝑓𝑖𝑐𝑖𝑒𝑛𝑡
• Potential calculated from the Nernst equation is the maximum
potential before any current flow has occurred
• As the cell discharges and current flows from anode to cathode,
the concentration will change
• As reactants are being converted to products, Ecell will decrease
• Eventually, the cell potential reaches zero
• Zero potential means reaction is at equilibrium (a dead battery)
Also,
Q = K (equilibrium constant)
And
∆G = 0 as well