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ELECTROCHEMISTRY
Background Material:
What is a NET IONIC EQUATION ?
A net ionic equation represents the active species in a chemical reaction while omitting the inactive species. It is a brief but accurate description of the chemical change.
Steps involved in writing a net ionic equation:
1. Write the chemical equation: (non-ionic equation)
Pb(NO3)2(aq) + 2KI(aq) Æ 2KNO3(aq) + PbI2(s)
2. Redo the equation and dissociate all electrolytes into their component ions.
Pb2+
(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) Æ 2K+(aq) + 2NO3-(aq) + PbI2(s)
3. Cancel all species that do not change. (these species are not really part of the chemical change)
What remains is the net ionic equation.
Pb2+
(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) Æ 2K+(aq) + 2NO3-(aq) + PbI2(s)
Pb2+
(aq) + 2I-(aq) Æ PbI2(s)
Write net ionic equations for the following reactions:
1. Zinc solid is added to a solution of lead II nitrate
2. A copper strip is placed in a solution of silver nitrate.
3. Zinc metal is added to a solution of copper II sulfate
AN INTRODUCTION TO REDOX
Procedure:
1. Four sets of the following solutions are prepared. Cu(NO3)2(aq) Zn(NO3)2(aq) Pb(NO3)2(aq) AgNO3(aq)
Cu(NO3)2 Zn(NO3)2 Pb(NO3)2 AgNO3 Cu(NO3)2 Zn(NO3)2 Pb(NO3)2 AgNO3
Strips of copper are placed in the 1st set of
solutions Strips of zinc are placed in the 2nd set of solutions
Strips of silver are placed in the 4th set of solutions
Strips of lead are placed in the 3rd set of solutions
Cu(NO3)2 Zn(NO3)2 Pb(NO3)2 AgNO3 Cu(NO3)2 Zn(NO3)2 Pb(NO3)2 AgNO3
Observations:
Cu2+
(aq) Zn2+(aq) Pb2+ (aq) Ag+aq)
Cu(s)
Zn(s)
Pb(s)
Ag(s)
QUESTIONS:
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2. What generalization can be made about a metal and its own aqueous ion ?
3. If the forward reaction is spontaneous will the reverse reaction also be
spontaneous?
4. List the metallic ions in order of their tendency to form metals, from greatest to
least.
5. List the metals in order of their tendency to form positive ions, from greatest to
least.
6. How do the lists in questions 4 and 5 compare ?
7. (a) Write the equations for the reactions that convert the metallic ions to metals.
(b) These reactions are called HALF-REACTIONS. What kind of half-reactions are these reactions ?
(b) What kind of half-reactions are these reactions ?
9. Define:
(a) OXIDATION
(b) REDUCTION
Notice that we can take the two tables of reactivity and combine them into one table:
Reactions of ions in order of reactivity: Reactions of metals in order of reactivity
Ag+
(aq) + e- Æ Ag(s) Zn(s) Æ Zn2+(aq) + 2e Cu2+
(aq) + 2e- Æ Cu(s) Pb(s) Æ Pb2+(aq) + 2e- Pb2+
(aq) + 2e- Æ Pb(s) Cu(s) Æ Cu2+(aq) + 2e Zn2+
(aq) + 2e- Æ Zn(s) Ag(s) Æ Ag+(aq) + e
-When combined, we get:
Ag+
(aq) + e- ÅÆ Ag(s)
Cu2+
(aq) + 2e- ÅÆ Cu(s) Pb2+(aq) + 2e- ÅÆ Pb(s)
Zn2+
(aq) + 2e- ÅÆ Zn(s)
Decreasing order of tendency tolose electrons
Decreasing order of
tendency togain electrons
This is the species most readily reduced
or the strongest oxidizing agent
Ag+
(aq) + e- <---> Ag(s)
Cu2+
aq) + 2e- <---> Cu(s)
Pb2+
(aq) + 2e- <---> Pb(s)
Zn2+
(aq) + 2e- <---> Zn(s)
This is the species most readily oxidized or the strongest reducing agent
1. State whether the following changes are oxidation or reduction:
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(b) 2O2- ---> O2 ___________________
(c) 2N3- ---> N2 ___________________
(d) X5- ---> X7- ___________________
(e) Sn4+ ---> Sn2+ ___________________
(f) Cu ---> Cu2+ ___________________
(g) Y3- ---> Y1- ___________________
(h) Br2 ---> 2Br- ___________________
2. State whether the following changes are oxidation or reduction and write the electrons in the equation:
(a) X2 ---> 2 X- __________________
(b) Mn4+ ---> Mn2+ __________________
(c) P4 ---> 4 P3- __________________
(d) Z5- ---> Z2- __________________
(e) 8 S2- ---> S8 __________________
(f) Cu1+ ---> Cu2+ __________________
(g) Al ---> Al3+ __________________
(h) 3 S8 ---> 24 S2- __________________
(i) Q2- ---> Q2+ __________________
(j) 5 P4 ---> 20 P- __________________
(k) J4+ ---> J2- __________________
(a) 2Al(s) + 6H+(aq) Æ 2Al3+(aq) + 3H2(g)
(b) Cl2(aq) + 2Br-(aq) Æ 2Cl-aq) + Br2(aq)
(c) 3Zn(s) + 2Fe3+aq) Æ 3Zn2+ (aq) + 2Fe(s)
(d) Ba(s) + 2H+(aq) Æ Ba2+(aq) + H2(g)
(e) 2Mn7+
(aq) + 10Cl-(aq) Æ 2Mn2+ (aq) + 5Cl2(aq)
(f) 2Al(s) + 3Zn2+(aq) Æ 2Al3+ (aq) + 3Zn(s)
(g) 2Z
4-(aq) + 3Q2+(aq) Æ 2Z1-(aq) + Q3(aq)
1. For each of the following reactions indicate:
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(b) substance oxidized
(c) oxidizing agent
(d) reducing agent
(a) Br2(aq) + Sn2+(aq) Æ 2Br-(aq) + Sn4+(aq)
(b) Al3+
(aq) + 3Fe2+(aq) Æ Al(s) + 3Fe3+(aq)
(c) 2Ag(s) + Cu2+(aq) Æ 2Ag+(aq) + Cu(s)
(d) 2Al(s) + 6H+(aq) Æ 2Al3+(aq) + 3H2(g)
(e) Br2(l) + 2I-(aq) Æ I2(aq) + 2Br-(aq)
(f) Pb2+
(aq) + 2K(s) Æ 2K+(aq) + Pb(s)
(g) Ni2+
(aq) + Ca(s) Æ Ni(s) + Ca2+(aq)
When several reducing agents and oxidizing agents are added together, a reaction will occur between the substance most readily oxidized and the substance most readily reduced. A redox half-reaction table must be used to determine the reacting species.
1. List all species mentioned
-all soluble ionic compounds must be dissociated into their component ions -strong acids dissociate into H+ ions and the negative partner ions -everything else must be left as is
2. Look at the redox table and determine the substance most readily reduced. (S.O.A.) Write the reduction half-reaction
3. Look again at the redox table and determine the substance most readily oxidized (S.R.A.) Write the oxidation half-reaction.
4. Multiply one or both the half-reactions by a number or numbers so that the number of electrons gained will equal the number of electrons lost.
5. Add the two half-reactions together.
Example problem:
Determine the net ionic equation for the reaction that occurs when a strip of chromium is placed in an aqueous solution of zinc nitrate.
1. List species: Cr(s) Zn2+(aq) NO3-(aq) H2O(l)
2. Check table for strongest oxidizing agent and write reduction half-reaction.
Zn2+
(aq) + 2e- Æ Zn(s)
3. Check the table for strongest reducing agent and write oxidation half-reaction:
Cr(s) Æ Cr2+(aq) + 2e
-4. Multiply one or both half-reactions by a number or numbers to make the number of electrons gained equal to the number of electrons lost (not necessary in this case)
5. Add the two half-reactions together to get the net ionic equation:
Cr(s) + Zn2+(aq) Æ Cr2+(aq) + Zn(s)
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1. For a reaction to be spontaneous, the reactants must be reactive.
Those species up high on the left hand side of the table are powerful oxidizing agents.
Those species low down on the right hand side of the table are powerful reducing agents.
--> --> --> --> --> --> --> --> When the oxidizing agent is above the reducing agent, both reactants are powerful enough for a spontaneous reaction --> --> --> --> --> --> --> --> When the reducing agent is above the oxidizing agent, the reactants are not powerful enough for a spontaneous reaction
Write the net ionic equation for each of the following chemical combinations and state whether the reaction is spontaneous or nonspontaneous.
1. A laboratory technician stores an aqueous solution of iron III chloride in a nickel plated container.
2. A chemistry teacher demonstrates the test for bromide ions by bubbling some chlorine gas through a sodium bromide solution.
3. Aqueous solutions of tin II bromide and iron III nitrate are mixed.
4. Sodium metal is added to some water in a typical demonstration of the reactivity of the alkali metals.
5. A student uses hydrobromic acid to acidify a solution of potassium dichromate for later use as an oxidizing agent.
11 7. Two students try to etch their initials on a copper plate using hydrochloric acid.
8. A student uses copper electrodes to test the conductivity of a nitric acid solution.
9. An iron bolt is exposed to air and water (a reaction which causes many millions of dollars of damage each year).
10. A solution of acidic potassium dichromate is mixed with an aqueous solution of hydrogen peroxide.
11. A solution of tin II chloride is added to a solution of iron II chloride.
Assorted redox questions:
1. The following metals and metallic ions react as follows:
A(s) + B+(aq) Æ B(s) + A+(aq)
B(s) + C+(aq) Æ no reaction
A(s) + C+(aq) Æ C(s) + A+(aq)
B(s) + D+(aq) Æ D(s) + B+(aq)
Using the above data: (a) The strongest reducing agent is ____________ (b) The strongest oxidizing agent is_____________
(c) The substance that undergoes oxidation most readily is____ (d) The substance that undergoes reduction most readily is____
2. In the redox reaction M + N Æ P + Q, which statement is false?
(a) If M is the reducing agent, then N is reduced. (b) If M is oxidized, then N is the oxidizing agent (c) If M is oxidized, then N is reduced.
(d) If M is the reducing agent, then N is oxidized.
3. A reducing agent is a substance which:
(a) loses electrons and becomes reduced (b) gains electrons and becomes reduced (c) loses electrons and becomes oxidized (d) gains electrons and becomes oxidized
4. The four elements W, X, Y and Z form diatomic molecules and also form ions with a single negative charge. The following observations were made in a series of experiments:
2X
-(aq) + Y2(l) Æ 2Y-(aq) + X2(l)
2W
-(aq) + Y2(l) Æ no reaction
2Z
-(aq) + X2(l) Æ 2X-(aq) + Z2(l)
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5. Metals K, L, M and N and their salts are selectively reacted and yield the following data:
L2+
(aq) + 2K(s) Æ L(s) + 2K+(aq)
K(s) + N+(aq) Æ no change
K+
(aq) + M(s) Æ K(s) + M+(aq)
N(s) + M+(aq) Æ N+(aq) + M(s)
Using the above data:
(a) The strongest oxidizing agent is_________ (b) The strongest reducing agent is _________ (c) The substance most readily oxidized is ________ (d) The substance most readily reduced is _________ (e) The weakest oxidizing agent is___________ (f) The weakest reducing agent is____________
6. Metals Ga(s) In(s) Mn(s) and Np(s)and their salts are selectively reacted and
yield the following data:
3Mn2+
(aq) + 2Np(s) Æ 3Mn(s) + 2Np3+(aq) In2+
(aq) + Ga(s) Æ In(s) + Ga3+(aq) Mn2+
(aq) + Ga(s) Æ no appreciable reaction
Using the above data: (a) The strongest oxidizing agent is__________
(b) The substance most readily oxidized is_________
7. Metals M, N, O, and P and their salts are reacted and yield the following results:
2M+
(aq) + N(s) Æ N2+(aq) + 2M(s)
O2+
(aq) + N(s) Æ no reaction
P3+
(aq) + 3M(s) Æ P(s) + 3M+(aq)
(a) The strongest reducing agent is________ (b) The strongest oxidizing agent is________ (c) The weakest reducing agent is _________ (d) The weakest oxidizing agent is_________
(e) The species that attracts electrons most readily is_______
Y2 + X2+ Æ 2Y- + X4+ 3Q2- + 2P Æ 3Q4- + 2P3+ 2P3+ + 6Y- Æ 2P + 3Y
2 3X4+ + 3V2- Æ V
3 + 3X2+
From the above data make a mini redox table. Place the strongest oxidizing agent in the top left hand corner and show the equations undergoing reduction from left to right. Place the electrons in the equations.
9. Four metal strips and four metallic ion solutions Be(s) Cd(s) Ra(s) V(s) and Be2+(aq)
Cd2+
(aq) Ra2+(aq) and V2+(aq) are reacted to give the following results:
Be reacts spontaneously with Cd2+
(aq) and V2+(aq) but not with Ra2+(aq).
Cd(s) does not react with any solution.
Ra(s) reacts with the Be2+(aq), Cd2+(aq) and V2+(aq) solutions.
V(s) reacts with only the Cd2+(aq) solution.
List the metals with the strongest reducing agent first and the weakest reducing agent last.
10. The elements A, B, C, and D, form diatomic molecules and negative ions. The following observations were made:
A2 + 2B- Æ 2A- + B2 D2 + 2C- Æ no reaction B2 + 2C- Æ 2B- + C2
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11. Using the data from the following table, rank the oxidizing agents from the strongest to the weakest.
W+(aq) X+(aq) Y+(aq) Z+(aq)
W(s)
NR
9 9 NR
X(s) NR NR NR NR Y(s) NR
9 NR NR
Z(s)
9 9 9 NR
12. In an oxidation-reduction reaction, electrons are
A gained by the oxidizing agent B gained by the reducing agent C lost by the oxidizing agent
D transferred from the oxidizing agent to the reducing agent
13. A reducing agent
A is the product of an oxidation B becomes reduced
C loses electrons D gains electrons
14. In a redox reaction, the oxidizing agent
A gains electrons and is oxidized B loses electrons and is oxidized C loses electrons and is reduced D gains electrons and is reduced
15. In the reaction 2Fe3+
(aq) + Sn2+(aq) Æ Sn4+(aq) + 2Fe2+(aq)
Reduction and Oxidation Potentials:
Reduction potentials measure the ability of a species to gain an electron or electrons.
Hydrogen ions were chosen as the standard and arbitrarily assigned the value of zero. Species that gain electrons more readily than hydrogen ions have a positive reduction
potential. Species that gain electrons less readily than hydrogen ions are assigned a negative reduction potential.
The oxidation potential of a species is the reverse of the reduction potential
The sum of the reduction potential plus the oxidation potential equals the net potential or the voltage.
A positive net voltage indicates a spontaneous reaction. A negative net voltage indicates a nonspontaneous reaction.
Electric Potentials in Redox Reactions
1. Calculate the electric potential difference (voltage) and write the net equation for the following:
(a) A cell is composed of: zinc, copper, zinc sulfate, copper (II) sulfate (components of the Daniel cell used in the 1800’s to operate doorbells and telegraphs.)
(b) A cell is composed of: lead (IV) oxide, sulfuric acid (assume sulfate ions present) and lead (components of a car battery).
(c) A cell is composed of: oxygen with hydrogen ions and hydrogen gas (components of a hydrogen fuel cell used on the space shuttle).
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(a) When lithium is oxidized to form lithium ions and an electric potential difference of 3.84V, metal ion Y undergoes reduction.
i What is the reduction potential of metal ion Y?
ii What is metal ion Y?
(b) When fluorine gas is reduced to produce fluoride ions and an electric potential difference of 2.73V, metal X undergoes oxidation.
i What is the oxidation potential of metal X ?
ii What would be the reduction potential of X?
iii What metal is X?
3. Calculate the electric potential difference when the following pairs of half-reactions are combined together:
a) A+ + e- Æ A
(s) E = 0.20V B+ + e- Æ B
(s) E = 0.50V
b) Z(s) Æ Z+ + e- E= 0.30 V X(s) Æ X+ + e- E= 0.40V
c) P(s) Æ P2+ + 2e- E= -0.60V Q(s) Æ Q3+ + 3e E= -0.40V
(a) If Cl2(g) + 2e- Æ 2Cl- is assigned a reduction potential of 0.00V. What would be the reduction potential for Cu2+ + 2e- Æ Cu
(s) ?
(b) If Al3+ + 3e- Æ Al
(s) is assigned a reduction potential of 0.00V, what would the reduction potential of I2(s) + 2e- Æ 2I- ?
(c) If Ni2+ + 2e- Æ Ni
(s) is assigned a reduction potential of 0.00V, what is the oxidation potential of Ba(s) Æ Ba2+ + 2e- ?
(a) If, instead of choosing the hydrogen half-reaction as the standard, chemists had chosen the half-reaction F2(g) + 2e- Æ 2F-(aq) to be the standard, what would be the net potential for the reaction:
2H+
(aq) + Pb(s) Æ Pb2+(aq) + H2(g)
5. Build a redox table from the following evidence:
PbO2(s) + SO42-(aq) + 4H+(aq) + X(s) Æ 2X+(aq) + PbSO4(s) + 2H2O(l) E= 2.00V 2X(s) + Z2+(aq) Æ 2X+(aq) + Z(s) E= 0.19V
P2(l) + Z(s) Æ 2P-(aq) + Z2+(aq) E= 0.88V
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1. Calculate the net potential expected for the following combinations of half-reactions:
Note: These are just skeleton equations to give you an idea of the two half-reactions involved. You do not need to complete the equations to figure out the voltage. Use your data booklet to look up the reduction and oxidation
potentials by noting the reactants and products.
(a) Co(s) + Fe3+(aq) Æ Co2+(aq) + Fe2+(aq)
(b) I
-(aq) + MnO4-(aq) Æ I2(aq) + Mn2+(aq) (acidic sol'n)
(c) Pb2+
(aq) + Zn(s) Æ Pb(s) + Zn2+(aq)
2. Write balanced equations for the following skeleton equations and calculate the voltage.
(a) Cr2O72-(aq) + Hg(l) Æ Cr3+(aq) + Hg2+(aq) (acidic solution)
(b) MnO4-(aq) + H2O2(l) Æ Mn2+(aq) + O2(g) (acidic solution)
(c) Br2(aq) + Zn(s) Æ
1. From the statement that X(s) reacts with H+(aq) to give H2(g) and the additional
information that zinc metal reacts readily with X2+
(aq), decide where to place the X2+(aq) + 2e- Æ X
(s) equation in our list.
Ag+
(aq) + e- Æ Ag(s) Cu2+
(aq) + 2e- Æ Cu(s) 2H+
(aq) + 2e- Æ H2(g) Zn2+
(aq) + 2e- Æ Zn(s)
2. Can acid be stored in a cobalt container ?
3. (a) Use the redox table to decide which substances in the following list tend to oxidize
the bromide ion Br-(aq).
Cl2(aq)
H+(aq)
Ni2+
(aq)
Ni(s)
O2(g)
(b) Which of the above would tend to reduce bromine, Br2(aq).
4. When a shiny iron nail is placed in a mixture of Zn(NO3)2(aq)and AgNO3(aq) solutions, the
products formed would be:
(a) Fe2+
(aq) + Ag(s)
(b) Fe2+
(aq) + Zn(s)
(c) Fe3+
(aq) + H2(g)
(d) OH-(aq) + NO2(g)
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I2(aq) + 2e- Æ 2I-(aq) Eo = 0.00 V
(a) Determine what Eo would be for:
Na+(aq) + e- Æ Na(s) Eo = ______
(b) Determine how much the net potential would have been changed for:
2Na(s) + I2(aq) Æ 2Na+(aq) + 2I-(aq) ___________________
2. (a) What will happen if an aluminum spoon is used to stir a Fe(NO3)2(aq)solution?
(b) What will happen if an iron spoon is used to stir an AlCl3solution ?
3. Can a 1.00 M Fe2(SO4)3 solution be stored in a container made of nickel metal?
4. A student placed single drops of Cd(NO3)2(aq), Hg(NO3)2(aq) and Pt(NO3)2(aq)separately on
a Cr metal strip. a spontaneous reaction was observed for each drop. In a similar test, the student observed that Pt(NO3)2(aq) reacted spontaneously with Hg metal, but
Cd(NO3)2(aq)and Cr(NO3)2(aq) did not. In decreasing order of strength, the student's list of
oxidizing agents should be:
5. The reaction 2X- + R Æ R2- + 2X takes place spontaneously generating a voltage
of 1.20 V. If the Eo value for the half-reaction
2X- Æ 2X + 2e- is -0.80 V the potential for the half-reaction
R + 2e- Æ R2- would be____________
1. What mass of bromine will be produced if a potassium bromide solution reacts with 50.0 mL of a 1.20 mol/L solution of acidified NaClO4? (34.4 g)
2. The silver in a 5.90 g chunk of material was removed and then reacted with an acidified solution of K2Cr2O7. 35.9 mL of a 0.200 mol/L solution of K2Cr2O7 was required for the
reaction. Calculate:
(a) The number of moles of Cr2O72-(aq) used in the reaction. (0.00718 mol)
(b) The number of grams of silver reacted. (4.65 g)
(c) The percentage of silver in the sample. (78.8%)
3. A 0.400 mol/L solution of Sn(NO3)2(aq)was titrated with an acidified KMnO4solution.
The following data was obtained:
Volume of Sn(NO3)2(aq) = 10.0 mL
Trial 1 Trial 2 Trial 3
Initial buret reading 8.70 mL 17.30 mL 26.00mL
Final buret reading 17.30 mL 26.00 mL 34.80 mL
Using the above data, calculate the molar concentration of the KMnO4solution.
(0.184 mol/L)
4. In a titration experiment all the Br-(aq)ions in an acidic solution were oxidized to Br2(aq)
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and the volume of the KMnO4(aq) solution required was 15.0 mL. Calculate the
concentration of the Br-ions in the solution.
(0.0600 mol/L)
5. In a redox titration 12.50 mL of 0.0800 mol/L K2Cr2O7(aq)was used in an acidic solution
to oxidize Sn2+
(aq)ions to Sn4+(aq)ions. The volume of K2Cr2O7(aq)used was just sufficient to
oxidize all the Sn2+
(aq)in 10.0 mL of solution. Calculate the concentration of the Sn2+(aq)ions
in the solution. What color change indicated the endpoint of this reaction ? (0.300 mol/L orange Æ green)
6. What volume of 0.0500 mol/L KMnO4(aq) is needed to oxidize all the Br-(aq)ions in 25.0
mL of an acidic 0.200 mol/L NaBr solution ? (20.0 mL)
Tricky Redox Stoichiometry Problems:
1.The mass of iodine produced by the reaction of 400 mL of 0.200 mol/L magnesium iodide solution with excess chlorine gas is ? (20.3 g)
2. 200 mL of ferric sulfate solution is required to react with 50.0 g of zinc. What is the concentration of the ferric sulfate solution ? (3.82 mol/L)
3. 20.0 mL of a solution of calcium iodide are required to reduce 100 g of bromine. Calculate the concentration of the calcium iodide solution.
(31.3 mol/L)
4. 100 mL of a 0.500 mol/L solution of MgBr2is used to reduce 50.0 mL of a solution of
acidic potassium permanganate. Calculate the concentration of the permanganate solution.
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5. 20.0 g of calcium metal is required to reduce 250 mL of a solution of Cr2(SO4)3(aq).
Calculate the concentration of the chromium III sulfate solution. (2.00 mol/L)
6. 50.0 mL of a 0.600 mol/L aluminum bromide solution is required to reduce 30.0 mL of a gold III nitrate solution. Calculate the concentration of the gold ions in solution. (1.00 mol/L)
Titration Questions
Volume of sulfurous acid samples = 100.0 mL Concentration of KMnO4(aq) = 0.310 mol/L
Volume of KMnO4(aq)
Trial 1 2 3 4
Final buret reading (mL) 9.50 18.15 26.75 34.75
Initial buret reading 1.00 9.50 18.15 26.75
Final colour of mixture pink pink pink pink
(a) To determine the concentration of the sulfurous acid, the average volume of potassium permanganate used is _____________ mL. (8.58 mL)
(b) Calculate the concentration of sulfurous acid in the sample. (6.65 x 10-2 mol/L)
2. In a laboratory experiment 0.0500 mol/L K2Cr2O7(aq) was used in an acidic solution to
oxidize Fe2+
(aq) ions to Fe3+(aq) . The following data were obtained.
Volume of Fe2+(aq) solution 25.0 mL
Final buret reading (K2Cr2O7) 48.7 mL
Initial buret reading (K2Cr2O7) 3.7 mL
The Fe2+
(aq) solution was then used to titrate a solution of acidic potassium permanganate.
50.0 mL of the Fe2+
(aq) solution was required to completely react with 20.0 mL of the
potassium permanganate solution. Calculate the concentration of the potassium permanganate solution. (0.270 mol/L)
OXIDATION NUMBERS
27 this assumption is to be able to determine if oxidation or reduction has taken place in
situations where it is not obvious.
RULES FOR DETERMINING OXIDATION NUMBERS:
1. All pure elements have an oxidationnumber of zero.
Mg(s) = 0 Ca(s) = 0 S8(s) = 0 P4(s) = 0 He(g) = 0 Cl2(g) = 0
2. All simple ions have an oxidation number of the charge on the ion.
Mg2+ = 2+ O2- = 2- Al3+ = 3+ Br- = 1- P3- = 3- S2- = 2- N3- = 3- Ca2+ = 2+ F- = 1-
3. The oxidation number of the alkalimetals in any compound is 1+, as in K+, and Li+.
4. The oxidation number of alkaline-earthmetals in compounds is 2+ as in Ca2+, Mg2+, and Ba2+.
5. The oxidation number for oxygen in compounds is almost always 2- (In peroxides O is 1-)
6. The oxidation number for hydrogen in compounds is almost always 1+.
(In metallic hydrides like NaH, CaH2, or BaH2, hydrogen is 1-)
7. All other oxidation numbers are assigned so that the sum of the oxidation numbers equals the net charge on the molecule or complex ion.
EXAMPLE: Give the oxidation number of each of the underlined species:
H2SO4 NO3- P2O5 CaI2 PO43- CO32-
6+ 5+ 5+ 1- 5+ 4+
EXERCISE: Give the oxidation number of the following underlined species:
(a) H3BO3 CH4 OBr3- SO32- HCO3- MnO H2O2
(b) VO2+ HNO2 C2O42- XeO46- NaH S2O62- HXeO
_____ _____ _____ _____ _____ _____ ____
(c) NO+ C
2H5OH NH3OH+ Cr2O72- CaH2 Na2O2 ClNO
_____ _____ _____ _____ _____ _____ ____
2. State whether the following changes involve oxidation or reduction:
(a) AsO3- ----> AsO2 ________
(b) MnO4- ----> MnO2 ________
(c) H2O2 ----> O2 ________
(d) C2H5OH ----> CO2 ________
(e) NH4+ ----> N2O5 ________
(f) C2O42- ----> CH4 ________
(g) CrO3+ ----> Cr
2O72- ________
(h) H4IO6- ----> I2 ________
(i) Au ----> Au(CN)4- ________
(j) H2SO3 ----> H2S ________
(k) NH3OH+ ----> NO2 ________
3. For each of the following reactions identify the oxidizingagent and the reducing agent:
(a) Sn + NO3- ----> SnO2 + NO
29
(c) Fe3+ + NH
3OH+ ----> Fe2+ + N2O
(d) H3PO2 + Cr2O72- ----> H3PO4 + Cr3+
(e) VO2+ + Sn2+ ----> VO2+ + Sn4+
(f) H2 + N2O4 ----> H2O + N2
(g) CrO2- + S2O8 ----> CrO42- + SO42-
(h) XeO3 + I- ----> Xe + I2
4. State whether the following changes are oxidation or reduction and indicate how many electrons must be lost or gained in the process:
(a) C2O42- ----> 2CO2 ______________
(b) NiO2 ----> Ni(OH)2 ______________
(c) ClO3- ----> Cl- ______________
(d) C2H5OH ----> 2CO2 ______________
(e) Na2HPO4 ----> PO2 ______________
(f) N2O3 ----> 2NH3 ______________
(g) XO54- ----> H3X ______________
(h) KMnO4 ----> MnO2 ______________
(i) 2P2O5 ----> P4 ______________
BALANCING REDOX REACTION EQUATIONS (Without the use of a table)
THE HALF- REACTION METHOD:
(a) ACIDIC SOLUTIONS:
EXAMPLE: Balance the following equation by the half-reaction method:
HClO2 + I- Æ Cl2 + HIO (acidic solution)
STEP 1 Break the unbalanced equation into 2 half-reactions:
HClO2 ----> Cl2
I- ----> HIO
STEP 2 Balance the main actors first!( everything but hydrogen and oxygen)
2HClO2 ----> Cl2
I- ----> HIO
STEP 3 Balance the oxygen in each half-reaction by adding H2O to the side low in oxygen: 2HClO2 ----> Cl2 + 4 H2O
I- + H2O ----> HIO
STEP 4 Balance the hydrogen atoms by adding H+ to the side low in hydrogen 2HClO2 + 6H+ ----> Cl2 + 4H2O
I- + H2O ----> HIO + H+
STEP 5 Balance the charge by adding electrons to the side low in + charge 2HClO2 + 6H+ + 6e- ----> Cl2 + 4H2O
I- + H2O ----> HIO + H+ + 2e-
STEP 6 Add the two half-reactions together to cancel the electrons: 2HClO2 + 6H+ + 6e- ----> Cl2 + 4H2O
(M3) 3I- + 3H2O ----> 3HIO + 3H+ + 6e-
________________________________________________________
31
EXERCISE:
Balance the following equations:
1. Zn + NO3- Æ Zn2+ + N2O (acidic solution)
2. MnO4- + Fe2+ Æ Mn2+ + Fe3+ (acidic solution)
3. MnO4- + SO32- Æ Mn2+ + SO42- (acidic solution)
4. AsO33- + BrO3- Æ AsO43- + Br- (acidic solution)
BASIC SOLUTIONS
RULES:
STEP 1 To balance a reaction in a basic solution, first go through all the steps pretending you have an acidic solution:
Referring the example on page 29, the balanced equation for an acidic solution is: 2HClO2 + 3H+ + 3I- ----> Cl2 + H2O + 3HIO
STEP 2 An equation must be added to the acidic equation in order to cancel out the H+ ions.
The 3H+ ions on the left hand side must be cancelled with 3H+ ions on the right of our new equation. These 3H+ ions come with 3OH- ions from the dissociation of 3 molecules of water.
2HClO2 + 3H+ + 3I- ----> Cl2 + H2O + 3HIO
3HOH ----> 3H+ + 3OH
-STEP 3 Now add the two equations together:
2HClO2 + 3H+ + 3I- ----> Cl2 + H2O + 3HIO
3HOH ----> 3H+ + 3OH-
________________________________________________________ 2HClO2 + 3H+ + 3I- + 3HOH ----> Cl2 + H2O + 3HIO + 3H+ + 3OH-
or 2HClO2 + 3I- + 2H2O ----> Cl2 + 3HIO + 3OH-
EXERCISE:
Balance the following equations:
1. CuO + NH3 ----> Cu + N2 (basic solution)
33
Balance the following equations:
1. N2O4 + Br- Æ NO2- + BrO3-(acidic solution)
2. H2O2 + CrO2- Æ H2O + CrO42-(acidic solution)
3. CrO42- + SO32- Æ CrO2- + SO42- (basic solution)
4. BrO3- + I- Æ Br- + I2 (basic solution)
5. Zn + NO3- Æ NH4+ + Zn2+ (basic solution)
Balance the following half-reactions:
1. P4 Æ H2PO4- (acidic)
2. IO3- Æ I- (acidic)
3. TeO32- Æ Te (basic)
4. Cr2O72- Æ Cr3+ (acidic)
5. NO Æ N2O (basic)
35
For the following unbalanced partial equations, determine the half-reactions that would balance each redox reaction:
1. Cr2O72- + Sn2+ Æ Cr3+ + Sn4+ (acidic)
2. NO3- + Zn Æ Zn2+ + NH4+(basic)
3. MnO4- + SO32- Æ MnO2 + SO42- (acidic)
4. H3AsO4 + Sn2+ Æ AsH3 + Sn4+(basic)
5. Bi(OH)3 + SnO22- Æ SnO32- + Bi (acidic)
MORE PRACTISE QUESTIONS: Balance the following equations:
1. IO3- + SO2 + H2O Æ I2 + SO42- + H+
2. Cr + ClO4- + OH- Æ CrO2- + ClO3- + H2O
3. Cr2O72- + Cl- + H+ Æ Cr3+ + Cl2 + H2O
4. MnO4- + Sn2+ + H+ Æ Mn2+ + Sn4+ + H2O
5. S8 + NO3- + OH- Æ SO42- + NO2- + H2O
6. H2C2O4 + MnO4- + 6H+ Æ CO2 + Mn2+ + H2O
7. H2SO3 + HIO3 Æ H2SO4 + HI
8. Fe2+ + MnO
4- + H2O Æ Mn2+ + Fe3+ + OH-
9. NO + H5IO6 Æ HNO3 + HIO3 + H2O
OH-37
Photosynthesis
is
a
redox
reaction
that
occurs
in
plants.
energy + 6CO
2(g)+ 6H
2O
(l)Æ
C
6H
12O
6(s)+ 6O
2(g)C6H12O6(s) (glucose) is manufactured and stored
1. Show that the above reaction is a redox reaction by assigning oxidation numbers to all species.
2. Show the half reactions:
3. One redox reaction important to many people is the combustion of natural gas (CH4(g))
(a) Write the equation for the combustion of methane.
(b) Justify that the methane reaction is a redox reaction by assigning oxidation numbers to all species.
Disproportionation Reactions
A disproportionation reaction is a reaction in which some atoms of a single element in a reactant are oxidized and others are reduced. For such reactant behavior to be possible, the reactant must contain an element that is capable of having at least three oxidation numbers - its original number plus one higher and one lower oxidation number. Note that any given atom is not both oxidized and reduced. Some atoms are oxidized and other atoms of the same element reduced.
Example:
3Br2 + 3H2O Æ HBrO3 + 5HBr 0 +5 -1
Some bromine atoms have oxidized and some have reduced.
In balancing these equations, we simply write the formula for the substance undergoing disproportionation twice. ( We pretend it is two different substances and proceed as before. )
Questions: Balance the following disproportionation equations:
1. MnO42- Æ MnO2 + MnO4- (acidic solution)
2. NO2 Æ NO3- + NO (acidic solution)
39
4. Cl2 Æ Cl- + ClO- (acidic solution)
5. S Æ S2- + SO32- (basic solution)
6. NO + NO3- Æ N2O4 (acidic solution)
7. H5IO6- + I- Æ I2 (acidic solution)
Voltaic Cells (Batteries)
In a spontaneous redox reaction electrons are transferred from the reducing agent
to the oxidizing agent.
Cu2+
(aq) + Zn(s) Æ Zn2+(aq) + Cu(s)
e-
If the reactants are arranged in a special way, these moving electrons can be made to flow through a wire. The reactants must be separated and yet in contact with each other so that the reaction will happen slowly.
1. The "main actors" (Cu2+
(aq) and Zn(s)) must be separated into two half-cells.
Cu2+
(aq) (CuSO4(aq)) Zn(s)
Each half-cell must have a solution and a conducting solid strip.
Cu2+
(aq) Zn(s)
A copper strip is added so that it will not react with the
coppersolution.
A zinc solution (Zn(NO3)2(aq) is
added so it will not react with the zinc strip
The solutions must be connected and the metallic strips must be connected.
41 The solutions are connected
by a salt bridge. (This allows the ions to move between the two half-cells)
A voltmeter is inserted in the wire
We now have a voltaic cell or a battery!
The half-cell where oxidation takes place is called the anode area. The metal bar is
called the
anode
itself.
The half-cell where reduction takes place is called the cathode area. The metal bar is the
cathode
itselfThe ions in both solutions and the salt bridge are called the electrolytes.
The positively charged ions migrate toward the cathode so they are called cations.
(Cu2+
(aq) and Zn2+(aq))
Cu2+
Zn2+
The negatively charged ions migrate toward the anode so they are called anions
(NO3-(aq))
Note:
The ions can only move through the solution - not through the wire!
The electrons can move only through the electrodes and wire! (anode to cathode)
Cell Notation: In order to indicate the ingredients of a cell the following notation is used:
Cu(s)/Cu(NO3)2(aq) // Ag(s)/AgNO3(aq)
The single slash
separates the electrode from the electrolytes
Example question:
H+
(aq), K2Cr2O7(aq)/C(s)// FeCl2(aq)/ Fe(s)
For the above cell indicate the following:
(a) Cathode Reaction (l)
(b) Anode Reaction:
(c) Net Reaction:
(d) Anode: (the solid bar where oxidation takes place) (e) Cathode: (the solid bar where reduction takes place)
Step 3:
The SRA is Fe(s)
This oxidation reaction will occur at the anode (Fe(s)) Step 2:
The SOA is H+ with Cr
2O72-(aq)
This reduction reaction will occur at the cathode (C(s))
Step 1:
List your species and get out your table!!
H+
(aq) K+(aq) Cr2O72-(aq) Fe2+(aq)
Fe(s) H2O(l) C(s) Cl-(aq)
(f) Net Eo
(g) Anions: (h) Cations:
Answers:
(a) Cathode Reaction: Cr2O72- + 14H+(aq) + 6e-Æ 2Cr2+(aq) + 7H2O(l)
(b) Anode Reaction: Fe(s) Æ Fe2+(aq) + 2e- (M3)
(c) Net Reaction: Cr2O72- + 14H+(aq)+ 3Fe(s)Æ 2Cr2+(aq)+ 7H2O(l) + 3Fe2+(aq)
(d) Anode: (the solid bar where oxidation takes place) Fe(s)
(e) Cathode: (the solid bar where reduction takes place) C(s)
(f) Net Eo 1.78 V
(g) Anions: Cr2O72-(aq), Cl-(aq)
(h) Cations: Fe2+
(aq), K+(aq), H+(aq)
NOTE:
The anode area consists of a solid bar and a solution around it. Sometimes the anode itself reacts and sometimes the solution around the anode reacts!
The cathode area consists of a solid bar and a solution around it. Sometimes the cathode itself reacts and sometimes the solution around the cathode reacts!
When we want the solution to react and not one of the electrodes, we choose an
43
Exercise:
7
8
1 Fe(s) 6 Ag(s)
9
2 Fe2+(aq) 10 4 Ag+(aq)
3 SO42-(aq) 5 NO3-(aq)
1. Referring to the above cell, find the correct number for the following:
(a) anode ________ (b) cathode ________ (c) anions ________ (d) cations ________
(e) direction of electron flow ________ (f) direction of anion movement ________ (g) direction of cation movement ________
(h) What would we see happening at the cathode ?
(i) What would we see happening at the anode ?
(j) which electrode could be replaced without changing the voltage
7
1 C(s) 8 6 Zn(s)
9
2 MnO4-(aq) 4 Zn2+(aq) 3 K+
(aq) 10 5 NO3-(aq) 11 H+
(aq)
2. Referring to the above cell, find the correct number for the following:
(a) anode ________ (b) cathode ________ (c) anions ________ (d) cations ________
(e) direction of electron flow _________ (f) direction of anion movement ________ (g) direction of cation movement _______
(h) What would we see happening at the cathode ?
(i) What would we see happening at the anode ?
(j) which electrode could be replaced without changing the voltage
45
Label the parts for the following cells:
1. Ag(s) / Ag+(aq) // Cd(s) / Cd2+(aq)
V
2. K2Cr2O7(aq), H+(aq) / C(s) // FeSO4(aq) / Fe(s)
3. HNO3 (aq),H2(g) / Pt(s) // Zn(NO3)2(aq) / Zn(s)
V
4. MnO4-(aq),H+(aq) / C(s) // PbSO4(aq) / Pb(s)
47
Exercise:
1. What is a salt bridge ?
2. Consider the following anode reactions:
(a) Cu(s) Æ Cu2+(aq) + 2e
(b) Zn(s) Æ Zn2+(aq) + 2e
(c) H2(g) Æ 2H+(aq) + 2e (d) Cl-(aq)Æ Cl2 + 2e (e) Pb(s) Æ Pb2+(aq) + 2e (f) Ag(s) Æ Ag+(aq) + e-
(i) Which of the above anode reactions involve the "dissolving" of a metal ?
(ii) Anodes often ______________ in size while anodes often _____________ in size
3. Consider the following cathode reactions:
(a) Cu2+(aq)
+ 2e- Æ Cu(s) (b) Br2(l) + 2e- Æ 2Br-(aq) (c) Al3+(aq) + 3e- Æ Al(s)
(d) Ag+(aq) + e- Æ Ag(s)
(e) 2H2O(l) + 2e- Æ H2(g) + 2OH-(aq)
(i) Which of the above reactions involves the production of a solid from a solution ?
4. Complete the prediction section of the following investigation report.
Problem: What is the complete cell description (half-reactions, net reaction, cell potential, electrodes, electrolytes, cathode and anode,
and the electron and ion flow) of the maximum-voltage cell assembled from the materials available ?
Materials: metal strips of iron, copper, chromium, lead and tin.
5. Consider the diagram below and answer the following questions
V
Ag(s) Mg(s)
AgNO3(aq) Mg(NO3)2(aq)
(a) When the cell is operating the anode will be______________
(b) As the cell operates the _______________electrode will gain in mass. (c) The Eo net for this cell is_____________
(d) The anions would be_________________ (e) The cations would be______________
6. If you wish to replate a silver spoon would you make it the anode or the cathode ? ___
7. A half-cell consisting of a palladium rod dipping into a 1.00 mol/L solution Pd(NO3)2(aq)
solution is connected with a standard hydrogen half-cell. The cell voltage is 0.99 volt and the platinum electrode in the hydrogen half-cell is the anode. Determine the Eo for the reaction: Pd2+
(aq) + 2e- Æ Pd(s)
8. Consider the cell: Ni2+
(aq) / Ni(s) // Cl2(g),Cl-(aq) / C(s)
(a) The anode is_____________
(b) The cathode reaction is:__________________________
(c) The net cell voltage is_____________
49
Cell Questions:
1. The reduction half reactions and potentials for a silver-cadmium cell are given below. Write down and label the anode and cathode half reactions, the net equation, and determine the Eo for the cell.
2AgO(s) + H2O(l) + 2e- Æ Ag2O(s) + 2OH-(aq) Eo = +0.57 V
Cd(OH)2(s) + 2e- Æ Cd(s) + 2OH-(aq) Eo = -0.76 V
2. The reduction half reactions and potentials for a methane-oxygen fuel cell are given below. Write down and label the anode and cathode half reactions, the net equation, and determine the Eo net for the cell.
CO3
2-(l) + 7H2O(g) + 8e- Æ CH4(g) + 10OH- Eo = +0.17 V
O2(g) + 2H2O(g) + 4e- Æ 4OH-(l) Eo = +0.40 V
3. The equations involved in the operation of a silver-zinc battery are given below along with their oxidation potentials. Write and label the anode and cathode half reactions and write the net ionic equation. Include the Eo values and the Eo net.
2Ag(s) + 2OH-(aq) Æ Ag2O(s) + H2O(l) + 2e- Eo = -0.342 V
4. Scientists isolated a previously undiscovered metal of which they managed to produce a few thin strips. They predicted that the metal would fit between tin and lead on the table of standard electrode potentials. Assume they give you the metal strips and you have access to a lab with various common metals and solutions of their ions. Use M(s) to represent the new metal.
(a) List two steps you could use as an experimental procedure that would test the scientists prediction.
(b) Assume that the scientists prediction was verified and state the observations that you would expect for each step.
(c) Assume that the scientists prediction was verified and write a possible balanced equation for each reaction that occurred involving the unknown metal.
(d) Give a probable Eo value for the reduction half-reaction of the metal.
5. Explain the function of the salt bridge in the voltaic cell.
51
Fuel Cells
A fuel cell is a voltaic cell for which the reactants are continuously supplied.
Example:
Methane is usually combined with oxygen to produce carbon dioxide and water vapor and heat energy.
CH4(g) + 2O2(g) Æ CO2(g) + H2O(g) + 803 kJ
Usually the energy from this reaction is used to warm houses or run machines. In a fuel cell designed to use this reaction the energy is used to produce an electric current. The electrons flow from the reducing agent (CH4(g)) to the oxidizing agent (O2(g)).
The U.S. space program uses a fuel cell based on the reaction of hydrogen and oxygen to produce water. A diagram of such a cell is
e- wet, acidified plastic membrane
H2(g) Æ Å O2(g)
H+
Carbon and platinum
Steam
Anode reaction:
Cathode reaction:
Net Reaction:
Disadvantages:
ELECTROLYSIS
In Voltaic cells chemical energy is converted into electrical energy.
In electrolytic cells electricity is used to force a nonspontaneous chemical reaction to occur.
Write equations for the following reactions:
(a) Electrolysis of aqueous aluminum bromide.
(b) An aqueous solution of potassium sulfate is electrolyzed.
(c) An aqueous solution of lead II nitrate is electrolyzed.
(d) A solution of aqueous sodium bromide and aqueous zinc chloride are mixed in an electrolytic cell using inert electrodes.
53
ELECTROLYSIS OF MOLTEN IONIC COMPOUNDS
Solid ionic compounds cannot conduct electricity.
When an ionic compound is melted, the ions are free to move around and therefore molten ionic compounds are conductors of electricity.
In a molten ionic compound there are only two species, the positive ion and the negative partner ion.
EXAMPLE:
Write the cathode, anode and net reactions for the electrolysis of molten aluminum oxide:
Species: Al3+(l) O2-(l)
Cathode reaction:
Al3+
(l) + 3e- Æ Al(l)
Anode reaction:
2O2
-(l) Æ O2(g) + 4e-
Net Reaction:
4Al3+
(l) + 6O2-(l) Æ 4Al(l) + 3O2(g)
EXERCISE:
Write equations for:
1. Electrolysis of molten sodium iodide
2. Electrolysis of molten calcium phosphide
Electrolytic Cell
Electrochemical Cell
Reaction at Cathode
Reaction at Anode
Anion Movement
Cation Movement
Electron Movement
Use of Salt Bridge
Use of Power Source
Voltage
Position of Reducing Agent
Position of Oxidizing Agent
Spontaneous or Nonsponeous
Anode Appearance
55
QUANTATIVE ASPECTS OF ELECTROLYSIS
EQUATIONS:
m = I.t.M
substance reacted 96,500 C/mol.chg
ne- = I.t
96,500 C/mol
PROBLEMS:
1. Determine the number of moles of electrons supplied by a dry cell supplying a current of 0.100 A for 50.0 min to a portable tape deck.
2. An electrochemical cell caused 0.0720 mol of electrons to flow through a wire during a 3.00 h period. Calculate the average current.
3. If 20.0 A of current flows through an electrolytic cell containing molten aluminum oxide for 1.00 h, what mass of aluminum will be deposited at the cathode.
4. A student wishes to plate 10.0 g of nickel onto a hub cap from a 2.00 mol/L solution of NiCl2. How long must the student run a 0.500 a current in order to produce the desired
5. Determine the mass of magnesium deposited at the cathode of a molten MgCl2
electrolytic cell if 10.0 A flow through the cell for 2.00 h ?
6. An electroplating firm wishes to plate 10.0 g of copper onto a pair of baby shoes from a Cu(NO3)2solution. If a 2.00 A current is used, calculate the time required. At which
electrode would the shoes be attached ?
7. If 80 g of fluorine gas are required, what current would have to flow for 10.0 hours to produce the fluorine from molten NaF ? At which electrode would this reaction occur ?
8. Calculate the current required to plate out 35.0 g of nickel from an aqueous solution of nickel II sulfate if the current runs for 10.0 h ?
57
10. Which ions will need the largest number of Coulombs to plate out 200 g of metal ?
Na+(aq) Cu2+(aq) Ag+(aq) Ca2+(aq)
11. A Cl2(aq) / Cl-(aq) // Cd2+(aq) / Cd(s)electrochemical cell delivers 1.93 A for 2.00 hours.
During this time the greatest mass change would occur where ?
12. The 500 g lead head of a gavel is attached with a wire to a negative terminal of a power supply and lowered into a nickel II nitrate solution. If the circuit is completed with an anode made of nickel metal attached to the positive terminal and also lowered into the solution and if a current of 5.00 A flows for 10.0 h, the final mass of the gavel head will be ?
13. Which of the following ions will need the largest quantity of electricity to electroplate out (deposit on the cathode) 200 g of the metal ?
Al3+
(aq) Mg2+(aq) K+(aq) Fe2+(aq)
14. Four electrolytic cells contain CsCl(aq), AgCl(aq), PbCl2(aq), and AuCl3(aq).
15. If 0.200 mol of electrons flow through an electrolytic cell containing molten aluminum oxide, what mass of product is collected at the anode ?
16. In the electrolysis of molten aluminum chloride, if 20.0 g of chlorine are collected at the anode, what mass of aluminum is produced at the cathode ?
17. Three electrolytic cells are connected in series. Cell 1 contains molten magnesium nitride, cell 2 contains molten sodium chloride, and cell 3 contains molten aluminum sulfide.
(a) If cell 1 produces 22.2 g of magnesium metal at the cathode, what masses of metals are produced at the cathodes in cells 2 and 3 ?
(b) What masses of what product are produced at each anode ?
18. 50.0 g of copper (from Cu2+(aq))are produced at the cathode of an electrolytic cell. How
59
1. (a) What is a sacrificial anode ? (cathodic protection)
(b) Give some examples of cathodic protection.
2. What are some industrial applications of electrolysis ?
3. What is the original meaning of the terms oxidation and reduction
4. Could you store a solution of acidified potassium permanganate ? Explain.
