Chemistry: The Central Science Redox reaction … G11 Chemistry: The Central Science Chapter 20:...
Transcript of Chemistry: The Central Science Redox reaction … G11 Chemistry: The Central Science Chapter 20:...
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Chemistry: The Central Science
Chapter 20: Electrochemistry
Redox reaction power batteries
Electrochemistry is the study of the relationships between electricity and chemical
reactions
o It includes the study of both spontaneous and nonspontaneous processes
20.1: Oxidation States and Oxidation-Reduction Reactions
Oxidation numbers of all the elements involved in the reaction can be tracked to
determine whether the reaction is a redox reaction
In some reactions, the oxidation numbers change, but we cannot say that any
substance literally gains or loses electron
o E.g. Combustion of hydrogen to form water
In this reaction, hydrogen is oxidized from 0 to +1 oxidation state and
oxygen is reduced from the 0 to the -2 oxidation state
Water is not an ionic substance, however, and so there is not a
complete transfer of electrons from hydrogen to oxygen as water is
formed
o Using oxidation states is a convenient form of “bookkeeping,” but you should
not generally equate the oxidation state of an atom with its actual charge in a
chemical compound
The substance the oxidizes the other substance (thus becoming reduced) is called
the oxidizing agent or oxidant
The substance that reduces the other substance (thus becoming oxidized) is called
the reducing agent or reductant
20.2: Balancing Oxidation-Reduction Equations
When balancing a redox reaction, the gains and the losses of electrons must be
balanced
Half-Reactions
o Although oxidation and reduction must take place simultaneously, it is often
convenient to consider them as separate processes
o E.g.
Oxidation:
Reduction:
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Equations that show either oxidation or reduction alone are called half-
reactions
In the overall reaction, the number of electrons lost in the
oxidation half-reaction must equal to the number of electrons
gained in the reduction half-reaction
Balancing Equations by the Method of Half-Reactions
o The use of half-reactions to balance oxidation-reduction equations usually
begin with a “skeleton” ionic equation that shows only the substances
undergoing oxidation and reduction
o For balancing a redox reaction that occurs in acidic aqueous solution, the
procedure is as follows:
Divide the equation into two half-reactions, one for oxidation and the
other for reduction
Balance each half-reaction
First, balance the elements other than H and O
Next, balance the O atoms by adding H2O as needed
Then, Balance the H atoms by adding H+ as needed
Finally, balance the charge by adding e- as needed
At this point, you can check whether the number of electrons in
each half-reaction equals corresponds to the changes in
oxidation state
Multiply the half-reactions by integers, if necessary, so that the number
of electrons lost in one half-reaction equals the number of electrons
gained in the other
Add the two half-reactions and, if possible, simplify by canceling
species appearing on both sides of the combined equation
Check to make sure that atoms and charges are balanced
Balancing Equations for Reactions Occurring in Basic Solution
o One way to balance these reactions is to balance the half-reactions initially as
if they occurred in acidic solution
Then, count the H+ in each half-reaction, and add the same number of
OH- to each side of the half-reaction
The OH- will neutralize the protons on the side containing H+ and the
other side ends up with OH-
20.3: Voltaic Cells
The energy released in a spontaneous redox reaction can be used to perform
electrical work
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o This task is accomplished through a voltaic (or galvanic) cell, a device in which
the transfer of electrons takes place through an external pathway
E.g. spontaneous reaction occurs when a strip of zinc is placed in
contact with a solution containing Cu2+
Zn metal is in contact with Zn2+(aq) in one compartment of the
cell, and Cu metal is in contact with Cu2+(aq) in another
compartment
o Consequently, the reduction of the Cu2+ can only occur
by a flow of electrons through an external circuit
Two solid metals that are connected by the external circuit are
called electrodes
o Electrode at which oxidation occurs is called the anode
o Electrode at which reduction occurs is called the cathode
Each compartments of a voltaic cell is called a half-cell
o Anode:
o Cathode:
For a voltaic cell to work, the solutions in the two half-cells must
remain electrically neutral
As Zn is oxidized in the anode compartment, Zn2+ enter the
solution
As Cu2+ at the cathode reduces, the positive charge from the
solution is removed
A salt bridge serves this purpose
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o A salt bridge consists of a U-shaped tube that contains
an electrolyte, such as NaNO3(aq), whose ions will not
react with other ions in the cell or with the electrode
materials
Anions always migrate toward the anode and the cations
toward the cathode
In any voltaic cell the electron flow from the anode through the
external circuit to the cathode
A Molecular View of Electrode Processes
o Redox reaction between Zn(s) and Cu2+(aq) lead to an increase in Zn2+(aq) and
Cu, and a decrease in Zn(s) and Cu2+(aq)
o In the case of the voltaic cell, the Zn atom “loses” two electrons and becomes
a Zn2+(aq) in its compartment
The electron travels through the wire and attached to Cu2+(aq),
forming Cu(s) in its compartment
o The redox reaction between Zn and Cu2+ is spontaneous regardless of
whether they react directly or in the separate compartments of a voltaic cell
20.4: Cell EMF under Standard Conditions
The electrons flow spontaneously toward the electrode with the more positive
electrical potential
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The difference in potential energy per electrical charge (the potential difference)
between two electrodes is measured in units of volts
o
C is coulomb
V is volt
One electron has a charge of 1.60 × 10-19 C
Electromotive force (emf) – The potential difference between the two electrodes of
a voltaic cell providing the driving force that pushes the electron through the circuit
o The emf of a cell, denoted Ecell, is also called the cell potential
Ecell is measured in volts so it’s often referred to as cell voltage
o For any cell reaction that proceeds spontaneously such as that in a voltaic cell,
the cell potential will be positive
Under standard conditions (25°C and 1 M for aqueous or 1 atm for gases), the emf is
called the standard emf, or the standard cell potential, and is denoted E°cell
Standard Reduction (Half-Cell) Potentials
o Standard reduction potentials (E°red) – the standard electrode potentials
tabulated for reduction reactions
o
For all spontaneous reactions at standard conditions, E°cell > 0
o The reference half-reaction is the reduction of H+(aq) to H2(g) under standard
conditions, which is assigned a standard reduction potential of exactly 0 V
An electrode designed to produced this half-reaction is called a
standard hydrogen electron (SHE), or the normal hydrogen electrode
(NHE)
An SHE consists of a platinum wire connected to a piece of
platinum foil covered with finely divided platinum that serves as
an inert surface for the reaction
The electron is encased in a glass tube so that the hydrogen gas
under standard conditions (1 atm) can bubble over the platinum
The solution contains H+(aq) under standard (1 M) conditions
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o Whenever we assign an electrical potential to a half-reaction, we write the
reaction as a reduction
o Changing the stoichiometric coefficient in a half-reaction does not affect the
value of the standard reduction potential
o The more positive the value of E°red, the greater the driving force for reduction
under standard conditions
Strength of Oxidizing and Reducing Agents
o The more positive the E°red value for a half-reaction, the greater the tendency
for the reactant of the half-reaction to be reduced and oxidize another species
o The half-reaction with the smallest reduction potential is most easily reversed
as an oxidation
o Solutions of reducing agents are difficult to store for extended periods
because of the ubiquitous presence of O2, a good oxidizing agent
20.5: Free Energy and Redox Reactions
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E: A positive value of E indicates a spontaneous process; a negative value of E
indicates a nonspontaneous one
The activity series consists of the oxidation reactions of the metals, ordered from the
strongest reducing agent at the top to the weakest reducing agent at the bottom
o E.g.
Ni is oxidized and Ag+ is reduced
Positive value of E° indicates that the displacement of silver by
nickel is a spontaneous process
EMF and ΔG
o
ΔG is the change in Gibbs free energy
n is a positive number without units that represents the number of
electrons transferred in the reaction
F is called Faraday’s constant, which is the quantity of electrical charge
on one mole of electrons
o A positive value of E and a negative value of ΔG both indicate that a reaction is
spontaneous
o When the reactants and products are all in their standard states
20.6 Cell EMF under Nonstandard Conditions
As a voltaic cell is discharged, the reactants of the reaction are consumed, and the
products are generated, so the concentrations of these substances change
o The emf progressively drops until E = 0, at which point we say the cell is
“dead”
The Nernst Equation
o
This is the Nernst equation
At T = 298 K the quantity 2.303 RT/F equals 0.0592, with the
units of volts (V)
o
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o At E = 0, ΔG = 0
The system is at equilibrium
o In general, increasing the concentration of reactants of decreasing the
concentration of products increases the driving force for the reaction,
resulting in a higher emf and vice versa
Concentration Cells
o Cell emf depends on the concentration so a voltaic cell can be constructed
using the same species in both the anode and cathode compartments as long
as the concentration are different
A cell based solely on the emf generated because of a difference in a
concentration is called a concentration cell
o E.g. Nickel
Oxidation of Ni(s) occurs in the half-cell containing the more dilute
solution, thereby increasing the concentration of Ni2+(aq)
n (the number of electron being transferred) is equal to 2
20.7: Batteries and Fuel Cells
A battery is a portable, self-contained electrochemical power source that consists of
one or more voltaic cells
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o When cells are connected in series, the battery produces a voltage that is the
sum of the emfs of the individual cells
Higher emfs can also be achieved by using multiple batteries in series
o Some batteries are primary cells, meaning that they cannot be recharged
A secondary cell can be recharged from an external power source after
its emf has dropped
Lead-Acid Battery
o A 12-V lead-acid automotive battery consists of six voltaic cells in series, each
producing 2 V
The electrode reactions that occur during discharge are
Because the reactants are solids, there is no need to separate the cell
into anode and cathode compartments
Solids are excluded from the reaction quotient Q, the relative
amounts of Pb(s), PbO2(s), and the PbSO4(s) have no effect on
the emf, helping the battery maintain a relatively constant emf
o Lead-acid batter can be recharged
During recharging, an external source of energy is used to reverse the
direction of the overall cell reaction
Alkaline Battery
o Alkaline batteries are nonrechargeable (primary battery)
o The anode of this battery consists of powdered zinc metal immobilized in a gel
in contact with a concentrated solution of KOH
The cathode is a mixture of MnO2(s) and graphite
Nickel-Cadmium, Nickel-Metal-Hydride, and Lithium-Ion Batteries
o Nickel-cadmium (nicad) battery
During discharge, cadmium metal is oxidized at the anode of the
battery while nickel oxyhydroxide is reduced at the cathode
Cadmium is a toxic heavy metal
Its use increases the weight of batteries and provides an
environmental hazard
o Nickel-metal-hydride (NiMH) batteries
Cathode reaction of NiMH is the same as that for the nickel-cadmium
batteries
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The anode consists of a metal alloy that has the ability to absorb
hydrogen ions
During the oxidation at the anode, the hydrogen atoms lose
electrons, and the resulting H+ ions react with OH- ions to form
H2O
Due to the robustness of the batteries toward discharge and recharge,
the batteries can last up to 8 years
o Lithium-ion (Li-ion) battery
Lithium is a very light element and therefore achieve a greater energy
density—the amount of energy stored per unit mass—than nickel-
based batteries
It is based on the ability of Li+ ions to be inserted into and removed
from certain layered solids
Hydrogen Fuel Cells
o The direct production of electricity from fuels by a voltaic cell could, in
principle, yield a higher rate of conversion of the chemical energy of the
reaction
Voltaic cells that perform this conversion using conventional fuels,
such as H2 and CH4 are called fuel cells
Strictly speaking, fuel cells are not batteries
o In the fuel cell for the reaction of hydrogen and oxygen, the anode and
cathode are separated by a thin polymer
Protons are able to pass through these polymers but electrons cannot
20.8: Corrosion
Corrosion reactions are spontaneous redox reactions in which a metal is attacked by
some substance in its environment and converted to an unwanted compound
For nearly all metals, oxidation is a thermodynamically favorable process in air at
room temperature
o When oxidation process is not inhibited in some way, it can be very
destructive to whatever object is made from the metal
o Oxidation can form an insulating protective oxide layer that prevents further
reaction of the underlying metal
Corrosion of Iron
o Rusting of iron requires both oxygen and water
Other factors—such as the pH of the solution, the presence of salts,
contact with metal more difficult to oxidize than iron, and stress on the
iron—can accelerate rusting
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o Corrosion of iron is electrochemical in nature
Electrons can move through the metal from a region where oxidation
occurs to another region where reduction occurs
Preventing the Corrosion of Iron
o Iron is often covered with a coat of paint or another metal such as tin or zinc
to protect its surface against corrosion
If the coating is broken and the iron is exposed to oxygen and water,
corrosion will begin
o Galvanized iron, which is iron coated with a thin layer of zinc, uses the
principles of electrochemistry to protect the iron from corrosion even after
the surface coat is broken
The Zn(s) is easier to oxidize than Fe(s)
Thus, even if the zinc coat is broken, the zinc will serves as the anode
and is corroded instead of iron
o Protecting a metal from corrosion by making it the cathode in an
electrochemical cell is known as cathodic protection
The metal that oxidized while protecting the cathode is called the
sacrificial anode
20.9: Electrolysis
Electrical energy can be used to cause nonspontaneous redox reactions to occur
o Such processes, which are driven by an outside source of electrical energy, are
called electrolysis reactions and take place in electrolytic cells
An electrolytic cell consists of two electrodes in a molten salt or a
solution
A battery or some other source of direct electrical current acts as an
electron pump, pushing electrons into one electrode and pulling them
from the other
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The electrode of the electrolytic cell that is connected to the negative terminal of the
voltage source is the cathode of the cell
Several practical applications of electrochemistry are based on active electrodes—
electrodes that participate in the electrolysis process
o Electroplating, for example, uses electrolysis to deposit a thin layer of one
metal on another metal to improve beauty or resistance to corrosion
Quantitative Aspects of Electrolysis
o For any half-reaction, the amount of a substance that is reduced or oxidized in
an electrolytic cell is directly proportional to the number of electrons passed
into the cell
o A coulomb is the quantity of charge passing a point in a circuit in 1 s when the
current is 1 ampere (A)
Coulombs = amperes × seconds
o Electrons can be thought of as reagents in electrolysis reactions
Electrical Work
o
-wmax means that a voltaic cell does work on its surrounding
o Eext > Ecell is needed to bring about a nonspontaneous electrochemical process
o When an external potential Eext is applied to a cell, the surroundings are doing
work on the system
n is the number of moles of electrons forced into the system by
the external potential
o n × F is the total electrical charge supplied to the system
by the external source of electricity
o watt (W) is a unit of electrical power
Watt-second is a joule
Kilowatt-hour (kWh) is equal to 3.6 × 106 J