Chapter Eighteen Electrochemistry Today Turn in: Nothing Our
Plan: Test Results Notes Redox Equations
Worksheet #1 Homework (Write in Planner): Redox Equation WS due
Friday 6% Curve A 2 B 5 C 4 D 1 F Average 72.72% High Score 90%
(x2) OxidationReduction: The Transfer of Electrons
Silver metal is formed, and the solution turns blue from copper(II)
ions formed. Electrons from copper metal are transferred to silver
ions. Half-Reactions (Review)
In any oxidationreduction reaction, there are two half-reactions:
Oxidation: a species loses electrons to another species. (LEO)
Reduction: a species gains electrons from another species. (GER)
Both oxidation and reduction must occur simultaneously. A species
that loses electrons must lose them to something else (something
that gains them). A species that gains electrons must gain them
from something else (something that loses them). Oxidation Numbers
An oxidation number is the charge on an ion, or a hypothetical
charge assigned to an atom in a molecule or polyatomic ion.
Examples: in NaCl, the oxidation number of Na is +1, that of Cl is
1 (the actual charge). In CO2 (a molecular compound, no ions) the
oxidation number of oxygen is 2, because oxygen as an ion would be
expected to have a 2 charge. What is the oxidation number on carbon
in CO2? Rules for Assigning Oxidation Numbers
For the atoms in a neutral speciesan isolated atom, a molecule, or
a formula unitthe sum of all the oxidation numbers is 0. This
includes elements in their standard state (Cu (s), Cl2 (g), etc.)
For the atoms in an ion, the sum of the oxidation numbers is equal
to the charge on the ion. In compounds, the group 1A metals all
have an oxidation number of +1 and the group 2A metals all have an
oxidation number of +2. In compounds, the oxidation number of
fluorine is 1. In compounds, hydrogen has an oxidation number of
+1. In most compounds, oxygen has an oxidation number of 2. In
binary compounds with metals, group 7A elements have an oxidation
number of 1, group 6A elements have an oxidation number of 2, and
group 5A elements have an oxidation number of 3. Assign oxidation
numbers to the elements in each compound: NO3-1 SO2
Lets Try It Assign oxidation numbers to the elements in each
compound: NO3-1 SO2 Fe2O3 Cu (s) Identifying OxidationReduction
Reactions
In a redox reaction, the oxidation number of a species changes
during the reaction. Oxidation occurs when the oxidation number
increases (species loses electrons). LEO Reduction occurs when the
oxidation number decreases (species gains electrons). GER Another
mnemonic is OIL RIG! If any species is oxidized or reduced in a
reaction, that reaction is a redox reaction. Examples of redox
reactions: displacement of an element by another element;
combustion; incorporation of an element into a compound, etc.
Review Fe2O3 + S --> Fe + SO2
In the reactions below, write oxidation numbers for each substance
and identify which substance is oxidized and which is reduced? N2O
H > H2O NH3 K KNO >N K2O Fe2O S--> Fe+ SO2 The
Half-Reaction Method of Balancing Redox Equations
Separate a redox equation into two half-equations, one for
oxidation and one for reduction. Balance the number of atoms of
each element in each half-equation. Usually we balance O and H
atoms last. Balance each half-reaction for charge by adding
electrons to the left in the reduction half-equation and to the
right in the oxidation half-equation. Adjust the coefficients in
the half-equations so that the same number of electrons appears in
each half-equation. Add together the two adjusted half-equations to
obtain an overall redox equation. Simplify the overall redox
equation as necessary. Redox Reactions in Acidic and in Basic
Solution
Redox reactions in acidic solution and in basic solution may be
very different from one another. If acidic solution is specified,
we must add H2O and/or H+ as needed when we balance the number of
atoms. If basic solution is specified, the final equation may have
OH and/or water molecules in it. A simple way to balance an
equation in basic solution: Balance the equation as though it were
in acidic solution. Add as many OH ions to each side as there are
H+ ions in the equation. Combine the H+ and OH ions to give water
molecules on one side, and simplify the equation as necessary.
MnO4-1 + S2O3-2 --> Mn+2 + SO4-2
Example 18.1 Balance the equation in acidic solution: MnO4-1 +
S2O3-2 --> Mn+2 + SO4-2 P4 (s) + H+1 (aq) + NO3-1 (aq) -->
H2PO4-1 (aq) + NO (g)
Example 18.1 B Write a balanced equation for the oxidation of
phosphorus by nitric acid, which is described by P4 (s) + H+1 (aq)
+ NO3-1 (aq) --> H2PO4-1 (aq) + NO (g) Br2 (l) --> Br-1 (aq)
+ BrO3-1 (aq)
Example 18.2 Balance the following in basic solution: Br2 (l)
--> Br-1 (aq) + BrO3-1 (aq) CrO4-2 (aq) + AsH3 (g) -->
Cr(OH)3 (s) + As (s)
Example 30a Balance the following in basic solution: CrO4-2 (aq) +
AsH3 (g) --> Cr(OH)3 (s) + As (s) HCl + K2Cr2O7 --> KCl +
CrCl3 + H2O + Cl2
WS Example 1 HCl + K2Cr2O7 --> KCl + CrCl3 + H2O + Cl2 FeCl2 +
KMnO4 + HCl --> FeCl3 + KCl + MnCl2 + H2O
WS Example 2 FeCl2 + KMnO4 + HCl --> FeCl3 + KCl + MnCl2 + H2O
S-2 + MnO4-1 --> S + MnO2 (basic solution)
WS Example 3 S-2 + MnO4-1 --> S + MnO2 (basic solution) CuS +
NO3-1 --> Cu+2 + S + NO (acidic)
WS Example 4 CuS + NO3-1 --> Cu+2 + S + NO (acidic) Stop! Do the
Redox Equations Worksheet. Today Turn in: Get out Redox WS Our
Plan: Questions on Redox WS Quiz
Notes Galvanic Cells WS Homework (Write in Planner): Worksheet Due
Wednesday A Qualitative Description of Voltaic Cells
A voltaic cell uses a spontaneous redox reaction to produce
electricity. A half-cell consists of an electrode (strip of metal
or other conductor) immersed in a solution of ions. This Zn2+
becomes a Zn atom. Both oxidation and reduction occur at the
electrode surface, and equilibrium is reached. This Zn atom leaves
the surface to become a Zn2+ ion. Important Electrochemical
Terms
An electrochemical cell consists of two half-cells with the
appropriate connections between electrodes and solutions. Two
half-cells may be joined by a salt bridge (U shaped tube suspended
in gel) that permits migration of ions, without completely mixing
the solutions. The anode is the electrode at which oxidation
occurs. The cathode is the electrode at which reduction occurs.
Helpful Mnemonic AUTO (anode = oxidation) CAR (cathode =
reduction)
OR: To remember the charge: Ca+ions are attracted to the Ca+hode
(the t is a plus sign) To remember which reaction occurs at which
terminal: An Ox and Red Cat - Anode Oxidation, Reduction Cathode
Important Electrochemical Terms
In a voltaic cell, a spontaneous redox reaction occurs and current
(electricity) is generated. Cell potential (Ecell) is the potential
difference in volts between anode and cathode. Ecell is the driving
force that moves electrons and ions. A ZincCopper Voltaic
Cell
Positive and negative ions move through the salt bridge to equalize
the charge. the electrons produced move through the wire to the
Cu(s) electrode, where they are accepted by Cu2+ ions to form more
Cu(s). Zn(s) is oxidized to Zn2+ ions, and Reaction:Zn(s) +
Cu2+(aq)--> Cu(s)+ Zn2+(aq) Cell Diagrams A cell diagram is
shorthand for an electrochemical cell. The anode is placed on the
left side of the diagram. The cathode is placed on the right side.
A single vertical line ( | ) represents a boundary betweenphases,
such as between an electrode and a solution. A double vertical line
( || ) represents a salt bridge or porous barrier separating two
half-cells. Reaction: Zn(s) + Cu2+(aq) --> Cu(s) + Zn2+(aq)
Another look at Cell Diagrams Example 18.3 Describe the
half-reactions and the overall reaction that occur in the voltaic
cell represented by the cell diagram: Pt(s) | Fe2+(aq), Fe3+(aq) ||
Cl(aq) | Cl2(g) | Pt(s) Standard Electrode Potentials
Since an electrode represents only a half-reaction, it is not
possible to measure the absolute potential of an electrode. The
standard hydrogen electrode (SHE) provides a reference for
measurement of other electrode potentials. The SHE is arbitrarily
assigned a potential of V. Standard Electrode Potentials
The standard electrode potential, E, is based on the tendency for
reduction to occur at an electrode. (Sometimes it is referred to as
standard reduction potential) All standard reduction potentials are
for 25 C and 1 M solutions. Efor the standard hydrogen electrode is
arbitrarily assigned a value of V. All other values of E are
determined relative to the standard hydrogen electrode. Standard
Electrode Potentials
The standard cell potential (Ecell) is the difference between E of
the cathode and E of the anode. Ecell= E(cathode)E(anode) Remember
AUTO CAR! Measuring the Standard Potential of the Cu2+/Cu
Electrode
The voltmeter reading and the direction of electron flow tell us
that Cu2+ is more easily reduced than H+, by volts. Standard
hydrogen electrode Cu2+ + 2e --> Cu E = V Measuring the Standard
Potential of the Zn2+/Zn Electrode
The voltmeter reading and the direction of electron flow tell us
that Zn2+ is harder to reduce than H+, by volts. Standard hydrogen
electrode Zn2+ + 2e --> Zn E = V F2 is the strongest oxidizing
agent F is the weakest reducing agent
Li is the strongest reducing agent Important Note The formula
Ecell= E(cathode)E(anode) allows us to simply use the numbers in
the table, but if you were asked for the oxidation value of a
reaction, you would have to reverse the sign since the values are
all reduction potentials. Important Points about Electrode and Cell
Potentials
Standard electrode potentials and cell voltages are intensive
properties; they do not depend on the total amounts of the species
present. A flashlight battery (D-cell) and a penlight battery (AA
cell) produce the same potential1.5 volts. E does depend on the
particular species in the reaction (or half-reaction). As we shall
learn later, cell and electrode potentials can depend on
concentration of the species present. Example 18.4 Determine E for
the reduction half-reaction
Ce4+(aq) + e --> Ce3+(aq), given that the cell voltage for the
voltaic cell Co(s) | Co2+(1 M) || Ce4+(1 M), Ce3+(1 M) | Pt(s) is
Ecell = V. Example 18.5 Balance the following oxidationreduction
equation, and determine Ecell for the reaction. O2(g) + H+(aq) +
I(aq)-->H2O(l) + I2(s) Electrode Potentials, Spontaneous Change,
and Equilibrium
An electrochemical cell does work. welec = nFEcell n = number of
electrons in the balanced equation F = 96,485 coulombs per mole.
The amount of electrical work is also equal to DG: DG = nFEcell
Under standard conditions: DG = nFEcell Criteria for Spontaneous
Change in Redox Reactions
If Ecell is positive, the forward reaction is spontaneous. If Ecell
is negative, the forward reaction is nonspontaneous (the reverse
reaction is spontaneous). If Ecell = 0, the system is at
equilibrium. When a cell reaction is reversed, Ecell and DG change
signs. Example 18.6 Will copper metal displace silver ion from
aqueous solution? That is, does the reaction Cu(s) + 2 Ag+(1
M)-->Cu2+(1 M) + 2 Ag(s) occur spontaneously from left to right?
The Activity Series Revisited
In the activity series of metals (Section 4.4), any metal in the
series will displace a metal below it from a solution of that
metals ions. Theoretical basis: The activity series lists metals in
order of their standard potentials. Displacement of a metal from a
solution of its ions by a metal higher in the series corresponds to
a positive value of Ecell and a spontaneous reaction. Visual
Example Equilibrium Constants in Redox Reactions
Whereas potential and free energy are related, and free energy and
equilibrium are related, equilibrium and potential must be related
to one another. DG = nFEcell and DG = RT ln Keq thereforeRT ln Keq
= nFEocell RT ln Keq RT Ecell == ln Keq nF nF R and F are constant,
therefore at 298 K: V Ecell = ln Keq n Example 18.8 Calculate the
values of G and Keq at 25 C for the reaction Cu(s) + 2 Ag+(1
M)-->Cu2+(1 M) + 2 Ag(s) Thermodynamics, Equilibrium, and
Electrochemistry: A Summary
From any one of the three quantities Keq, G, Ecell, we can
determine the others. Stop! Begin working on Worksheet #2! Today
Turn in: Nothing Our Plan: Quick Review
Notes Concentration & Electrochemical Cells Finish
Electrochemical Cells WS Homework (Write in Planner): Worksheet Due
Wednesday Effect of Concentrations on Cell Voltage
A nonstandard cell differs in potential from a standard cell (1 M
concentrations, 1 atm partial pressures). Effect of Concentrations
on Cell Voltage
From the previous relationships we can show that: RT Ecell = Ecell
ln Q nF At 25 C, and converting to common logarithms: V Ecell =
Ecelllog Q n This Nernst equation relates a cell voltage for
nonstandard conditions, Ecell, to a standard cell voltage, Ecell,
and to the concentrations of reactants and products expressed
through the reaction quotient, Q. We can use the Nernst equation to
find cell potential from concentrations, or we can measure Ecell
and determine the concentration of a species in the cell. Example
18.9 Calculate the expected voltmeter reading for the voltaic cell
pictured in Figure Another Nernst Equation Example
Use the Nernst equation to determine Ecell at 25C for the following
voltaic cells. Zn(s) | Zn2+(2.0 M) || Cu2+(0.050 M) | Cu (s) Zn(s)
| Zn2+(0.050 M) || Cu2+(2.0 M) | Cu (s) Batteries: Using Chemical
Reactions to Make Electricity
We often call any device that stores chemical energy for later
release as electricity a battery. Technically, a D, C, or AA
battery is actually a single electrochemical cell. A battery
consists of several cells joined together to produce higher current
or higher voltage. A 9-volt transistor battery, an automobile
battery, and a rechargable battery pack are all true batteries. The
Dry Cell A primary cell employs an irreversible chemical
reaction.
When the reactants inside the cell are largely used up, the cell is
dead. The LeClanch cell or dry cell (right) is the ordinary type of
flashlight battery. Alkaline cells cost more than the LeClanch cell
but they have a longer shelf life and longer service life. The
LeadAcid Storage Battery
A leadacid storage battery used in an automobile uses secondary
cells; they are rechargeable. By connecting the cell to an external
electric energy source, the discharge reaction is reversed. Cell
reaction:Pb(s) + PbO2(s) + 2 H+(aq) + 2 HSO4(aq) --> 2 PbSO4(s)
+ 2 H2O(l) Charging reaction: 2 PbSO4(s) + 2 H2O(l) --> Pb(s) +
PbO2(s) + 2 H+(aq) + 2 HSO4(aq) Other Secondary Cells The
nickelcadmium (NiCd) cell uses a cadmium anode and a cathode
containing Ni(OH)2. A NiCd cell can be recharged hundreds of times.
It produces 1.2 V (a Leclanch cell produces 1.5 V). Nickelmetal
hydride cells (NiMH) use a metal alloy anode that contains
hydrogen. In use, the anode releases the hydrogen, forming water.
Like the NiCd cell, a NiMH cell produces 1.2 V. Lithium-ion cells
use a lithiumcobalt oxide or lithiummanganese oxide material as the
anode. The electrolyte is an organic solvent containing a dissolved
lithium salt. Many modern laptop computers and cellular phones use
lithium-ion cells. Fuel Cells In a fuel cell, the cell reaction is
equivalent to a combustion reaction. The reactants are supplied
externally; the cell does not go dead as long as the oxidizing and
reducing agents are provided. Fuel cells are generally operated
under nonstandard conditions and at temperatures considerably
higher than 25 C. H2O2 fuel cells are seeing use in some
automobiles. Corrosion: Metal Loss Through Voltaic Cells
In moist air, iron is oxidized to Fe2+, particularly at scratches,
nicks, or dents. These areas are referred to as anodic areas. Other
regions of the iron serve as cathodic areas, where the electrons
from the anodic areas reduce O2 to OH. Iron(II) ions migrate from
the anodic areas to the cathodic areas where they combine with the
hydroxide ions. The iron(II) is then further oxidized to iron(III)
by atmospheric oxygen. Common rust is Fe2O3 x H2O. Corrosion of an
Iron Piling
One way to minimize rusting is to provide a different anode
reaction. Protecting Iron from Corrosion
The simplest defense against corrosion of iron is to coat it (with
paint or metal) to exclude oxygen from the surface. An entirely
different approach is to protect iron with a more active metal.
Galvanized iron has been coated with zinc. The zinc provides an
alternative anode reaction; the zinc corrodes, protecting the iron.
Cathodic Protection In cathodic protection, an iron object to be
protected is connected to a chunk of an active metal. The iron
serves as the reduction electrode and remains metallic. The active
metal is oxidized. Water heaters often employ a magnesium anode for
cathodic protection. Stop! Finish Worksheet #2 Electrochemical
Cells Worksheet Today Turn in: Get out Electrochemical Cells WS Our
Plan:
Questions on WS Electrochemical Cell Review Notes Electrolysis WS
Homework (Write in Planner): Worksheet Due Friday (LAST HOMEWORK)
Electrolytic Cells A voltaic cell corresponds to a spontaneous cell
reaction. An electrolytic cell corresponds to a nonspontaneous cell
reaction. The reaction is called electrolysis. An electrolytic cell
is the opposite of a galvanic cell. The external source of
electricity acts like an electron pump. It pulls electrons away
from the anode, where oxidation takes place, and pushes them toward
the cathode, where reduction takes place. The polarities of the
electrodes are reversed from those in the voltaic cell, because now
the external source controls the flow of electrons. Electrolysis of
Molten Sodium Chloride
2 NaCl(l) --> 2 Na(l)+Cl2(g) The nonspontaneous reaction is
driven by external potential. Molten NaCl (around 1000 C)
Predicting Electrolysis Reactions
In an electrolytic cell, all combinations of cathode and anode
half-reactions give negative values of Ecell. The reaction most
likely to occur is the one with the least negative value of Ecell
(requires the lowest applied voltage from the external electricity
source).HOWEVER In many half-reactions, particularly those
involving gases, various interactions at electrode surfaces make
the required voltage for electrolysis higher than the voltage
calculated from E data. Overvoltage is the excess voltage above the
voltage calculated from E values that is required in electrolysis.
Quantitative Electrolysis
The unit of electric charge is the coulomb (C), and the charge
carried by one electron is x 1019 C. The conversion factor well use
is 96,485 C/1 mole electrons Electric current, expressed in amperes
(A), is the rate of flow of electric charge (C/s). Quantitative
Electrolysis
To calculate the quantitative outcome of an electrolysis reaction:
Determine the amount of charge (C)the product of current and time.
Convert the amount of charge to moles of electrons. Use a
half-equation to relate moles of electrons to moles of a reactant
or a product. Convert from moles of reactant or product to the
final quantity desired. Example 18.13 We can use electrolysis to
determine the gold content of a sample. The sample is dissolved,
and all the gold is converted to Au3+(aq), which is then reduced
back to Au(s) on an electrode of known mass. The reduction
half-reaction is Au3+(aq) + 3e Au(s).What mass of gold will be
deposited at the cathode in 1.00 hour by a current of 1.50 A? Try
It Out! 18.13 A:For how many minutes must the electrolysis of a
solution of CuSO4 (aq) be carried out, at a current of 2.25 A, to
deposit 1.00 g of Cu (s) at the cathode? Example 18.14 An
Estimation Example
Without doing detailed calculations, determine which of the
following solutions will yield the greatest mass of metal at a
platinum cathode during electrolysis by a 1.50-A electric current
for 30.2 min: CuSO4(aq), AgNO3(aq), or AuCl3(aq). Producing
Chemicals by Electrolysis
Electrolysis plays an important role in the manufacture and
purification of many substances, including chlorine, copper,
silver, magnesium, aluminum, lead, zinc, sodium, fluorine,
titanium, sodium hydroxide, hydrogen Electrolysis of NaCl(aq) is
used to produce H2, Cl2, and NaOH, all of which have important
industrial uses. Electroplating Electrolysis can be used to coat
one metal onto another, a process called electroplating. Usually,
the object to be electroplated, such as a spoon, is cast of an
inexpensive metal. It is then coated with a thin layer of a more
attractive, corrosion-resistant, and expensive metal, such as
silver or gold. Complete the Electrolysis Worksheet
Stop Complete the Electrolysis Worksheet Today Turn in: Get out
Electrolysis WS Our Plan: Questions on WS
Electrochemical Cells Online Lab Homework (Write in Planner): Take
Home Test Due Monday OR prepare for the AP Exam! A Little Redox
Humor Today Turn in: Take Home Test Our Plan:
Begin Qualitative Analysis Lab (read the pre-lab and procedure
FIRST) Homework (Write in Planner): Work on Lab Report Lab
Preparation Read the Packet, including the entire procedure.
Answer the Pre-Lab Questions Begin working on your lab report I
recommend recording data in the handout that I gave you and then
transferring it to your final copy. You could begin working on the
title, purpose, pre-lab, and procedure, though. Reminders Safety
Goggles, gloves, apron at ALL TIMES!
Clean all glassware and lab supplies before and after use. Label
things well (include your initials)! Seal containers to store until
next class period.Place them on the back bookshelf. All solids go
in the trash and liquids down the drain EXCEPT barium (anion step 2
&3).Put them in the waste container in the hood! Run tests
until you get clear and reproducible results (you can repeat steps
if you need to). Report is due 5/14 at the beginning of class!
Today Homework (Write in Planner): Turn in: Nothing Our Plan:
Qualitative Analysis Lab Homework (Write in Planner): Report Due on
8/14