Transcript Oxidation Numbers

Electrochemistry
Page 17
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Major Points
•
Electrochemistry or redox reactions involve a transfer of electrons.
•
Oxidizing agents gain electrons, and reducing agents lose
electrons.
•
If an element in a chemical equation changes oxidation number,
then it is a redox equation.
•
All redox reactions are balanced, so that the total number of
electrons gained is equal to the total number of electrons lost.
•
Carbon is reduced during photosynthesis and oxidized during
respiration.
•
A redox net ionic equation can be broken down into two halfreactions to illustrate which substances are either losing or gaining
electrons with the electrons. The electrons are included in each
half-reaction.
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Reduction is the gain of electrons.
Oxidation is the loss of electrons.
Oxidizing Agents and Reducing Agents
If the charge on an atom or ion increases, it has lost electrons and is a
reducing agent. If the charge on an atom or ion decreases, it has
gained electrons and is an oxidizing agent.
Tips
• Reduction half-reaction—electrons are on the reactants side.
• Oxidation half-reaction—electrons are on the products side.
• Metals tend to lose electrons and be oxidized (tend to be reducing
agents).
• Non-metals tend to gain electrons and be reduced (tend to be
oxidizing agents).
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Writing Reduction and Oxidation Half-Reactions
23. Write the two half-reactions and determine the oxidizing and reducing
agent for…
A. O2(g) + 4 H+(aq) + 2 Fe(s) → 2 Fe2+(aq) + 2 H2O(l).
oxidizing agent _______________ reducing agent _______________
Reduction: O2(g) + 4 H+(aq) + 4 e−  2 H2O(l)
Oxidation: 2 Fe(s)  2 Fe2+ (aq) + 2 e−
oxidizing agent: O2(g) + 4 H+(aq) reducing agent: Fe(s)
B. 2 Na(s) + 2 H2O(l) → 2 NaOH(aq) + H2(g).
oxidizing agent _______________ reducing agent _______________
Oxidation: 2 Na(s)  2 Na+(aq) + 2e−
Reduction: 2 H2O(l) + 2e−  2 OH−(aq) + H2(g)
oxidizing agent: 2 H2O(l) reducing agent: Na(s)
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Ranking Oxidizing Agents and Building Redox Tables—Redox tables
are built by comparing the relative strength of oxidizing agents to gain
electrons. If you follow sports, it is similar to the standings in the NHL or
CFL. The team with the most wins or points is placed on top.
Tips
•
If there is evidence of a chemical reaction, the oxidizing agent is
placed over the reducing agent.
•
If there is no evidence of a chemical reaction, the reducing agent is
placed over the oxidizing agent.
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24. A piece of metal is placed in a solution of nickel (II) nitrate and
then in a solution of iron (II) nitrate. Evidence of a spontaneous
reaction was observed with the nickel (II) nitrate, but not with the
iron (II) nitrate. The metal is most likely…
A.
B.
C.
D.
lead.
cadmium.
zinc.
nickel.
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Use the following information to answer the next question.
Br2(l) + 2 I−(aq) → 2 Br−(aq) + I2(s)
25. For the net ionic equations, the reduction half reaction is i and the
reducing agent is ii .
The row that completes the above statement is…
Row
i
ii
A.
Br2(l) + 2 e− → 2 Br−(aq)
Br2(l)
B.
2 I−(aq) + 2 e− → I2(s)
I−(aq)
C.
Br2(l) + 2 e− → 2 Br−(aq)
I−(aq)
D.
2 I−(aq) → I2(s) + 2 e−
Br2(l)
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Use the following information to answer the next question.
Indium, cerium, europium and tantalum are elements that are rapidly
increasing in demand due to their use in touch screens, batteries,
energy-efficient lighting and handheld electronics. Their reactions
relative to each other are:
In(s) + 3 Ce4+(aq) → In3+(aq) + 3 Ce3+(aq) (spontaneous)
3 Eu2+(aq) + 2 Ta(s) → 3 Eu(s) + 2 Ta3+(aq) (nonspontaneous)
In(s) + Ta3+(aq) → Ta(s) + In3+(aq) (nonspontaneous)
Species
1. In(s)
5. Eu(s)
2. Ce4+(aq)
6. Ta3+(aq)
3. Eu2+(aq)
7. In3+(aq)
4. Ta(s)
8. Ce3+(aq)
Numerical Response
26.
When the reducing agents are ranked in order from strongest to
weakest, their corresponding numbers are ____ , ____ , ____ ,
and ____ .
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Writing Redox Equations
A.
Using a Reduction Potentials Table (page 7 of data booklet)
i.
List all species present. Strong acids and soluble ionic compounds
are written in dissociated form.
ii. Choose the strongest oxidizing agent closest to the top left of the
table and write the reduction half-reaction. Read from left to right.
iii. Choose the strongest reducing agent closest to the bottom right of
the table and write the oxidation half-reaction. Read the equation
from right to left.
iv. Balance for electrons so that electrons gained equals electrons lost
and write the net ionic equation. The reaction is spontaneous if the
oxidizing agent is above the reducing agent.
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Tips
•
When looking for the strongest oxidizing agent look only on the left
side of the table on page 7, and on the right side of the table for the
strongest reducing agent.
•
Be sure to multiply all the species in the half-reaction by the
coefficient used to balance for electrons.
•
Be sure to cancel species that are the same on both sides.
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Use the following information to answer the next question.
A solution of potassium permangate is acidified and reacted
with a sulfurous acid solution.
Numerical Response
27. When the net ionic equation is determined for the above statement
and balanced using lowest whole number coefficients, the coefficient
for…
MnO4−(aq) is ________ .
H2SO3(aq) is ________ .
H2O(l) is ________ .
H+(aq) is ________ .
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28. Which of the following aqueous ions can either gain or lose electrons in
a redox reaction?
A.
B.
C.
D.
Cr2+(aq)
Cl−(aq)
Ca2+(aq)
S2−(aq)
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B. Balancing Equations Using Half-Reaction Method
(ion electron method)
Divide the overall reaction into two half-reactions. The two half-reactions
are balanced separately and then added together to get a balanced net
ionic equation.
Tip
Every half-reaction and redox reaction must be balanced for charge,
so that the total charge on the reactants side equals the total charge
on the products side.
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29. Platinum is a precious metal and highly resistant to corrosion and
used in electrical contacts and catalytic converters. Platinum
reacts with aqua regia (a mixture of nitric and hydrochloric acid),
according to the following unbalanced equation.
Pt(s) + HNO3(aq) + HCl(aq) → H2PtCl6(aq) + NO2(g) + H2O(l)
Determine the oxidation and reduction half-reactions, and the
oxidizing and reducing agents.
Oxidation: Pt(s) + 6 HCl(aq)  H2PtCl6(aq) + 4 H+(aq) + 4 e−
Reduction: HNO3(aq) + H+(aq) + 1 e−  NO2(g) + H2O(l)
The reducing agent is Pt(s) + 6 HCl(aq) and the oxidizing agent
is HNO3(aq) + H+(aq).
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30. Of the following unbalanced skeleton half-reactions, which one
represents an oxidation?
A.
B.
C.
D.
CH3COOH(aq) → C2H5OH(aq)
MnO4−(aq) → MnO2(s)
HCOOH(aq) → CO2(g)
IO2−(aq) → I2(s)
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Oxidation Numbers (Oxidation States)
Major Points
Oxidation states represent the charge that an atom would have if all the
atoms making up a molecule were 100% ionic. Oxidation numbers are
an accounting method to keep track of electrons during redox reactions
and are used to balance redox equations.
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Rules
i.
The oxidation number of an atom or simple ion is the charge on
the atom or ion.
ii.
In a molecule containing only two elements, the most
electronegative atom will have a negative oxidation number and
the least electronegative atom will have a positive oxidation
number.
iii. The sum of the positive charges or oxidation numbers, and the
negative charges or oxidation numbers, must equal the total
charge on the ion or molecule.
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Tips
•
Hydrogen is usually +1, except when bonded to metals (then
it is −1).
•
Group 1 elements are almost always +1.
•
Group 17 elements are almost always −1, except when in
elemental form.
•
Oxygen is almost always −2, except in peroxides (where it
is −1), and when bonded to fluorine.
•
In polyatomic ions, or in dealing with compounds that have
three different elements, it is easier to set up an algebraic
expression.
•
When determining oxidation numbers in compounds that have
three different elements, assume that the most electronegative
atom has a negative oxidation number, and the least
electronegative atom has a positive oxidation number, and solve
for the element in the middle.
•
You only need to be able to balance equations in acidic
conditions.
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Use the following information to answer the next question.
When a person suspected of intoxication blows into an alcohol
breathalyzer device, the alcohol in the person’s breath reacts with the
acidic dichromate in the breathalyzer, converting the alcohol into acetic
acid, according to the following skeleton equation.
C2H5OH(l) + Cr2O72−(aq)  CH3COOH + Cr3+(aq)
31. In the above reaction equation, the oxidation number of the…
A.
B.
C.
D.
oxidizing agent decreases by three.
oxidizing agent decreases by six.
reducing agent increases by four.
reducing agent decreases by two.
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C. Balancing Equations Using Oxidation Numbers
In any redox equation, the overall change in oxidation number must equal 0.
Reduction—an increase in oxidation number.
Oxidation—a decrease in oxidation number.
Steps
i.
Initial balancing—balance everything, except H and O.
ii. Determine the oxidation number of all elements and identify the ones
that have changed oxidation number.
iii. Multiply the half-reactions, so that the overall change in oxidation
number is 0.
iv. Balance oxygen by adding water to one side.
v. In acidic conditions, balance hydrogen by adding H+(aq) to one side.
vi. Check to see whether the equation is balanced, by checking for
charge.
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32.
Use oxidation numbers to balance the following equation.
BrO3−(aq) + HSO3−(aq) → SO42−(aq) + Br2(l) (acidic)
2 BrO3–(aq) + 5 HSO3–(aq)  5 SO42–(aq) + Br2(l) + H2O(l)
+ 3 H+(aq)
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Tip
For disproportionation reactions (the same element reacts with itself
and is both oxidized and reduced), it is best to balance for electrons on
the side where the elements appear twice. Once the electrons are
balanced, make sure that all elements are balanced.
33.
Balance the following.
Br2(l)
→
BrO3–(aq)
+
Br–(aq)
3 Br2(l) + 3 H2O(l)  BrO3–(aq) + 5 Br(aq) + 6 H+(aq)
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Use the following equations to answer the next question.
1. 2 Ca(s) + O2(g) → 2 CaO(s)
2. H2SO4(aq) + 2 KOH(aq) → K2SO4(aq) + 2 H2O(l)
3. C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g)
4. 2 NaI(aq) + Cl2(g) → 2 NaCl(aq) + I2(s)
5. 2 KBrO3(s) → 2 K(s) + Br2(l) + 3O2(g)
6. Cr(NO3)3(aq) + Na3PO4(aq) → CrPO4(s) + 3 NaNO3(aq)
34. The equations that are redox equations are _____ , _____ ,
_____ , _____ .
1
3
4
5
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Use the following equations to answer the next question.
A solution of potassium permanganate was standardized by titration with
an acidic solution of oxalic acid (H2C2O4). A 0.200 g sample of oxalic acid
was dissolved in excess sulfuric acid and titrated with potassium
permanganate according to the unbalanced equation…
MnO4−(aq) + H2C2O4(aq) → Mn2+(aq) + CO2(g)
Trial 1
Trial 2
Trial 3
Trial 4
Final burette reading (mL)
47.6
35.4
23.6
31.9
Initial burette reading (mL)
35.4
23.6
11.9
20.2
35. The concentration of the potassium permanganate solution is…
A.
B.
C.
D.
0.0757 mol/L.
0.189 mol/L.
0.473 mol/L.
0.947 mol/L.
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Electrochemical Cells
Electrochemical cells convert chemical energy into electrical energy, or
electrical energy into chemical energy.
Major Points
• Electrodes are solid conductors of electricity.
• Electrolytes are solutes that conduct electricity when dissolved in water.
• Voltaic cells are electrochemical cells in which chemical energy is converted
into electrical energy.
• In redox reactions, electrons are transferred from the substance oxidized to
the substance reduced. The reactants can be arranged so that the electrons
are transferred through a conducting wire.
• The oxidation and reduction half-reactions occur in separate containers
called half-cells.
• A battery is a group of voltaic cells connected in series.
• Electrode potential refers to the voltage generated by a voltaic cell.
• Voltage is the difference in electrical potential between two electrodes in a
circuit.
• Cell notation:
(–) anode electrode/anode electrolyte // cathode electrolyte/cathode electrode (+)
{one half-cell}
{other half-cell}
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Tips
• Electrochemical cells can be either voltaic or electrolytic.
• Voltaic cells are spontaneous reactions, in which the E°net is positive,
and usually have either a porous cup, or a salt bridge. Electrolytic cells
are nonspontaneous reactions, in which the E°net is negative, and
requires a power supply.
• Reduction occurs in the cathode and oxidation occurs in the anode.
• Electrons always move through a wire from the anode to the cathode,
through the external circuit.
• E°net = E°reduction – E°oxidation or E°cell = E°reduction − E°oxidation.
• Cations migrate to the cathode and anions migrate to the anode through
the porous cup or salt bridge (must contain an electrolyte).
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Use the following information to answer the next question.
A student constructed the following cell and recorded observations.
Observations: I.
II.
III.
IV.
The
The
The
The
voltmeter registers 0.56 V.
mass of electrode M decreases.
concentration of X−(aq) increases.
colour of M+(aq) ion becomes more intense.
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36.
The reduction half-reaction for the Voltaic cell shown would be i and
the anode half-reaction would be ii .
The row that completes the above statement is…
Row
i
ii
A.
2 X−(aq) → X2(g) + 2e−
M(s) → M+(aq) + e−
B.
X2(g) + 2e− → 2 X−(aq)
M+(aq) + e− → M(s)
C.
X2(g) + 2e− → 2 X−(aq)
M(s) → M+(aq) + e−
D.
M+(aq) + e− → M(s)
2 X−(aq) → X2(g) + 2e−
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Page 25
Use the following information to answer the next question.
Tellurium is one of a number of rare earth elements that scientists
suggest are critical for green technologies, such as solar cells. In
the notation below, the silver is the cathode.
Te(s) / Te2−(aq) // Ag+(aq) / Ag(s) E°cell = +1.94 V
37. The standard reduction potential for the Te(s)/Te2−(aq) half-cell is…
A.
B.
C.
D.
−2.74 V.
−1.14 V.
+1.14 V.
+2.74 V.
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Use the following information to answer the next question.
The net ionic equation for a redox reaction is…
OCl−(aq) + 2 NO(g) + 2 OH−(aq) → Cl−(aq) + H2O(l) + 2 NO2−(aq).
Numerical Response
38.
1.30
The E°net for this reaction is __________
V.
Tip
If the electrode is NOT one of the reactants, then a non-reactive
electrode such as carbon, C(s), or platinum, Pt(s), can be used. An inert
electrode provides a location to connect a wire, and a surface upon which
the half-reaction can occur.
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Changing the Standard
Oxidizing agents each have a different potential for the attraction of
electrons. Reduction potentials are obtained by comparing each halfreaction to a standard. Chemists have selected the hydrogen half-cell
as the standard and use it to compare all other half-cells.
Tips
• If the new standard is changed to something above the hydrogen
half-cell, then subtract the voltage of the new standard.
• If the new standard is changed to something below hydrogen, then
add the absolute value of the voltage of the new standard to obtain
the new reduction potentials.
• The E°net for a cell does not change if the standard is changed.
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39. If the lithium half-reaction, Li+(aq) + e− → Li(s), had been assigned
an E° value of 0.00 V, then the predicted E°net value for the reaction
Cu(s) + Zn2+(aq) → Cu2+(aq) + Zn(s) would be…
A.
B.
C.
D.
+1.10 V.
−2.28 V.
−0.42 V.
−1.10 V.
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Tips
• Fuel cells are cells that operate with a continual supply of a fuel such
as hydrogen or methane.
• Two types of hydrogen fuel cells when hydrogen and oxygen are
bubbled through either an acidic or basic solution. The voltage
produced is the same for both and is 1.23 V.
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Electrochemical cells
Voltaic cells
Electrolytic cells
converts chemical energy
into electrical energy
converts electrical energy into
chemical energy
positive E°net
negative E°net
spontaneous
nonspontaneous
usually needs a porous cup, salt
bridge
needs a power supply
reduction occurs at the cathode and oxidation occurs at the anode
electrons move from the anode to the cathode through an external wire
cations migrate to the cathode and anions to the anode
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In order to obtain metals such as sodium, electrolysis must occur in
the absence of water. If water is present, water is the strongest
oxidizing agent. Therefore, the electrolysis occurs in molten salts.
The ionic compound is heated until it melts, and then the
electrodes are placed in the hot liquid and connected to a power
supply.
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Use the following information to answer the next question.
A current is passed through molten magnesium chloride. The diagram is a
schematic of the cell.
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40. In the above cell, the…
cathode is _______ .
direction of anion movement is _______ .
strongest oxidizing agent is _______ .
direction of electron movement is _______ .
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Tip
The Chloride Anomaly: Using the Table of Standard Electrode
Potentials for electrolytic cells does not always predict the correct
reaction. In the electrolysis of a concentrated sodium chloride
solution (brine), the reaction occurring at the anode is the oxidation
of chloride ions, not the oxidation of water. This is the only
exception you must remember for Chemistry 30.
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Cell Stoichiometry
Major Points
• Electric Charge is measured in Coulombs (C).
• One Ampere equals one Coulomb per second. 1 A = 1 C/s.
• Faraday Constant: One mole of electrons has a charge of 9.65 × 104 C.
• The number of moles of electrons can be found from the current…
Where:
ne– =
It
I is current in Amperes
F
t is time is seconds
ne− is number of moles of electrons
F is 9.65 × 104 C/mol (in data booklet)
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Where:
ne– =
It
I is current in Amperes
F
t is time is seconds
ne− is number of moles of electrons
F is 9.65 × 104 C/mol (in data booklet)
Tips
• Time must be in seconds. 1 h = 3600 s.
• Use the mole ratio to convert between moles of electrons and
moles of substance.
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41. The mass of copper plated onto a teapot from a solution of
copper(II) sulfate that operates at 6.2 A for 2.0 h is…
A.
B.
C.
D.
29.4 g.
14.7 g.
7.35 g.
2.04 g.
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Numerical Response
42. During the electrolysis of concentrated sodium chloride, the time
required to produce 25 g of product at the anode with a current of
9.15
124 A is ________
min.
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Corrosion—is an electrochemical process in which the metal is
oxidized back to the metal ion or ore-like state. For example, acid rain
will corrode iron:
O2(g) , H2O(l), H+(aq) , Fe(s)
cathode: O2(g) + 4 H+(aq) + 4e− → 2 H2O(l)
anode: 2 (Fe(s) → Fe2+(aq) + 2e−)
net ionic: O2(g) + 4 H+(aq) + 2 Fe(s) → 2 Fe2+(aq) + 2 H2O(l)
Even if the rain is not acidic (in Alberta rain is slightly basic), a
spontaneous reaction still occurs. In order for corrosion to occur all that
is needed is O2(g) and H2O(l).
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Corrosion Prevention
Major Points
• In protective coatings such as waxing or painting, the wax or paint acts
as a physical barrier that prevents contact of the oxygen and water
with the iron.
• Cathodic Protection protects iron from reacting by supplying an
alternate source of electrons in one of two ways.
i. Sacrificial anode—Metal, such as zinc or magnesium, is bolted to the
steel. The zinc or magnesium protects the iron by reacting in place
of the iron, because it is a stronger reducing agent, and therefore,
oxidizes more readily than the iron. The zinc or magnesium would
need to be replaced periodically.
ii. Impressed current—Use a source (such as a car battery) to provide
a small electric current through a wire to the metal object to prevent
corrosion by supplying another source of electrons.
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43. The iron hull of a ship can be protected against corrosion by
attaching i to the iron. This protects the iron, because it is
a ii than iron.
The row that best completes the above statements is…
Row
i
ii
A.
B.
zinc metal
zinc ions
stronger reducing agent
weaker oxidizing agent
C.
silver metal
weaker reducing agent
D.
silver ions
stronger oxidizing agent
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