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Chapter 19: Oxidation-Reduction Reactions

19.1 The Meaning of Oxidation and ReductionOxidation-reduction reactions- chemical changes that occur when electrons are transferred between reactants. also called redox reactions

Oxidation Originally meant the combination of an element with oxygen to

form oxides.4Fe + 3O2 2Fe2O3 (rusting of iron)C + O2 CO2 (burning of carbon)C2H5OH + 3O2 2CO2 + 3H2O (burning of ethanol)

Modern definition - loss of electrons or gain of oxygen

Reduction Originally meant the loss of oxygen from a compound. Metal

ores were heated to remove oxygen and their mass and volume was thus “reduced”.

2Fe2O3 + 3C 4Fe + CO2

Modern definition - gain of electrons or loss of oxygen

Oxidation and reduction always occur simultaneously. One process can not occur without the other.

Ex 1. 2Na + S Na2S Sodium goes from the neutral atom to the 1+ ion. Therefore, it has lost an electron (it was oxidized). Sulfur goes from the neutral atom to the 2- ion. Therefore, it has gained two electrons (it was reduced).

19.2: Assigning Oxidation NumbersAn easy way to determine what element is oxidized and which element is reduced in a reaction is to assign each element an oxidation number.

Rules for Assigning Oxidation NumbersEx:

1. The oxidation number of a monatomic ion is equal to its charge.

2. The oxidation number of hydrogen in a compound is always +1 unless it is a hydride, where it is -1.3. The oxidation number of oxygen in a compound is always -2 unless it is a peroxide, where it is -1.

4. The oxidation number of a free element is zero.

5. For any neutral compound, the sum of the oxidation numbers is zero.

6. For a polyatomic ion, the sum of the oxidation numbers must equal the charge of the ion.Ex 2.

To help remember these definitions, use one of these mnemonic devices:

OIL RIGOxidation Is Loss (of e-)Reduction Is Gain (of e-)

LEO the lion goes GERLose Electrons = OxidationGain Electrons = Reduction

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P2O5 H2O NH4+ Na2Cr2O7 Ca(OH)2

P:___ O:___ H:___ O:___ N:___ H:___Na:___ Cr:___ O:___

Ca:___ O:___ H:___

Oxidation is an increase in oxidation number. Reduction is a decrease in oxidation number.

Oxidizing agent- The substance in a redox rxn that accepts electrons (the substance that was reduced). Whereas only an element can be reduced, the entire compound is called the oxidizing agent.

Reducing agent- The substance in a redox rxn that donates electrons (the substance that was oxidized). Whereas only an element can be oxidized, the entire compound is called the reducing agent.

Ex 3. Identify what is being oxidized and reduced. Identify the oxidizing and reducing agent.

a.) 2Na + Cl2 2NaCl b.) 2HNO3 + 6HI 2NO + 3I2 + 4H2O

Reduced: Reducing Agent: Reduced: Reducing Agent:

Oxidized: Oxidizing Agent: Oxidized: Oxidizing Agent:

c.) 2PbSO4 + H2O Pb + PbO2 + 2H2SO4

Reduced: Reducing Agent:

Oxidized: Oxidizing Agent:

19.3: Balancing Redox ReactionsHalf-reaction Method

1. Write half reactions.2. Balance all elements other than oxygen and hydrogen.3. Balance oxygen by adding water.4. Balance hydrogen by adding hydrogen ions.5. Balance charge by adding electrons.6. Multiply half-reactions by factors to make the number of electrons lost equal to the number of electrons gained.7. Add half-reactions together.8. Cancel out anything that is the same on both sides.9. If the reaction occurs in basic solution, add hydroxide ions to both sides to cancel out the hydrogen ions. You will make water on the side with the hydrogen ions. Cancel water if necessary.10. Check to see that both charge and mass (atoms) are balanced.

Ex 4. Cr2O72- + CH3OH HCO2H + Cr3+ (acidic)

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Ex 5. NO3- + I2 IO3

- + NO2 (acidic)

Ex 6. MnO4

- + I- MnO2 + I2 (basic)

Galvanic and Voltaic Cells The interconversions of chemical and electrical energy are

called __________________ _________________, all of which involve the transfer of ____________. This process occurs in a device called an _________________ ___________.

Remember the activity series of metals in Chapter 11? (below)

You have another version of it to the right. Electrochemical cells that convert chemical energy into

electrical energy are called ________ _____ or galvanic cells. The energy is produced by _____________ redox reactions.

Label the following voltaic cell:

Li+ + Fe Li + Fe3+

Li+ + e- Li(s) -3.05Cs+ + e- Cs(s) -2.92K+ + e- K(s) -2.92Rb+ + e- Rb(s) -2.92Ba2+ + 2e- Ba(s) -2.90Sr2+ + 2e- Sr(s) -2.89Ca2+ + 2e- Ca(s) -2.87Na+ + e- Na(s) -2.71Mg2+ + 2e- Mg(s) -2.37Be2+ + 2e- Be(s) -1.70Al3+ + 3e- Al(s) -1.66Mn2+ + 2e- Mn(s) -1.18Zn2+ + 2e- Zn(s) -0.76Cr3+ + 3e- Cr(s) -0.74Fe2+ + 2e- Fe(s) -0.44Cr3+ + e- Cr2+ -0.41Cd2+ + 2e- Cd(s) -0.40Co2+ + 2e- Co(s) -0.28Ni2+ + 2e- Ni(s) -0.25Sn2+ + 2e- Sn(s) -0.14Pb2+ + 2e- Pb(s) -0.132H+ + 2e- H2(g) 0.00Sn4+ + 2e- Sn2+ 0.15Cu2+ + e- Cu+ 0.15Cu2+ + 2e- Cu(s) 0.34Cu+ + e- Cu(s) 0.52Fe3+ + 3e- Fe2+ 0.77Ag+ + e- Ag(s) 0.80Hg2+ + 2e- Hg(l) 0.85

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Voltaic cells can be separated into two half cells, one in which oxidation occurs and one in which reduction occurs. A half cell consists of a metal rod or strip immersed in a solution of its ions. The two half-cells are connected by a porous partition or a salt bridge. A salt bridge is a tube containing a conducting solution. Ions pass through the salt bridge to keep the charges balanced. Electrons pass through an external wire. The metal rods in voltaic cells are called electrodes. Oxidation occurs at the anode and reduction occurs at the cathode. (An Ox and a Red Cat) Species undergoing reduction receive electrons from the cathode. Species undergoing oxidation donate electrons to the anode. The direction of electron flow is therefore from the anode to the cathode.

Line notation is a shorthand method for representing electrochemical cells. The general form is:

anode | anode ion || cathode ion | cathode

The double vertical line represents the salt bridge. For a voltaic cell involving Cu and Zn as shown in the previous diagram, the line notation would be represented as:

Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

When only ions (rather than solids) are involved in the redox process, "inert" electrodes such as graphite or platinum are used.

Ex7. Given the voltaic cell: Ag+ + Fe2+ Fe3+ + Ag(s); sketch the cell (label the electrodes, salt bride, direction of electron flow, anode and cathode) , write the line notation and the half reactions.

Line notation:

Oxidation half-reaction:

Reduction half-reaction:

______ ________ are voltaic cells in which the electrolyte is a paste. The flashlight battery is a common example. Alkaline Batteries:

o Anode: _______ powder; This metal is in a gel which is in contact with a concentrated solution of __________ .

o Cathode: mixture of _____________ and ____________, separated from the anode by a porous fabric.

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o The battery is sealed in a steel can to reduce the risk of leakage. o Zinc is __________________ and manganese is ____________________.

Batteryo one or more voltaic cells hooked together. o lead storage battery is commonly used in a ________.

o anode is _____________. o cathode is _________________. o electrolyte (battery acid) is _______________.

Modern Battery Chemistry (Do not try to memorize these)

Modern batteries use a variety of chemicals to power their reactions. Typical battery chemistries include:

Lithium-iodide battery - Lithium-iodide chemistry is used in pacemakers and hearing aides because of their long life.

Nickel-cadmium battery - The electrodes are nickel-hydroxide and cadmium, with potassium-hydroxide as the electrolyte (rechargeable).

Nickel-metal hydride battery - This battery is rapidly replacing nickel-cadmium because it does not suffer from the memory effect that nickel-cadmiums do (rechargeable).

Lithium-ion battery - With a very good power-to-weight ratio, this is often found in high-end laptop computers and cell phones (rechargeable).

Zinc-air battery - This battery is lightweight and rechargeable. Zinc-mercury oxide battery - This is often used in hearing-aids. Silver-zinc battery - This is used in aeronautical applications because the power-to-weight ratio

is good.

A _________________ is an electrochemical device that, in most cases, combines hydrogen and oxygen to produce electricity, with water and heat as its by-product. As long as fuel is supplied, the fuel cell will continue to generate power.  Since the conversion of the fuel to energy takes place via an electrochemical process, not combustion, the process is clean, quiet and highly efficient – two to three times more efficient than fuel burning.

Ex8. What are some uses for fuel cells?

1. 2. 3.

The electrical potential, or cell potential (Eo), of a galvanic cell is a combination of the potentials of the two half-reactions. The standard reference that all half-cells are measured against is the standard hydrogen electrode, 2H+ + 2e- H2 Eo = 0.00 Volts

Table 23-2 in your book (and at the beginning of your notes) has a list of the reduction potentials of many reduction reactions. To change a reaction from reduction to oxidation, simply reverse the reaction and change the sign of the reduction potential to make it an oxidation potential.

Adding the potentials together for the oxidation and the reduction half-reactions will give you the overall cell potential.

Ex9. Using the Zn and Cu cell we used earlier, we can calculate the cell potential.

Galvanic cells require Eocell > 0 V

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One of the reduction potentials will have to be reversed (to form an oxidation half-reaction) in every Eo calculation.

To determine which reaction is to be reversed, the sum of the oxidation and reduction half-reactions must be > 0 V in a galvanic cell.

When you reverse a reaction, Eo gets the opposite sign. When you multiply a reaction by a coefficient (for purposes of balancing), the Eo is NOT changed. If the Eo

cell is negative, the reaction is not spontaneous.o The more positive a reduction potential is, the more easily that substance can be _______________. o The more negative a reduction potential is, the more easily that substance can be _______________.

Ex10. Write the half-reactions and calculate the cell potential for the following reaction:Co2+ + Fe Fe2+ + Co

Ex11. Given the following half-cells, decide which is the anode and the cathode, write the overall reaction and calculate Eo

cell.A. Ni2+ + 2e- Ni Eo = -0.23 V O2 + 4H+ + 4e- 2H2O Eo = +1.23 V

B. Ce4+ + e- Ce3+ Eo = +1.70 V Sn2+ + 2e- Sn Eo = -0.14 V

Electrolysis The process of forcing a current through a cell to produce a chemical change is ____________. In

order for electrolysis to occur, you must apply an external voltage that is greater than the potential of the galvanic cell if you want to force the reaction in the opposite (electrolytic) direction.

The direction of electron flow is always from _____ to ______. "FAT CAT" An electrolytic cell is used to change ______________ into _____________________.

When water is electrolyzed, an electrolyte such as sulfuric acid or sodium sulfate must be added to make the water conduct electricity. The two half reactions that occur in the electrolysis of water are:

anode:

cathode:

Overall reaction:

When brine (salt water) is electrolyzed, chlorine gas, hydrogen gas, and sodium hydroxide are produced. Solid sodium metal is not produced because water is more easily reduced than is the sodium ion. The two half-reactions that occur in the electrolysis of brine are:

anode:

cathode:

Overall reaction:

When molten sodium chloride is electrolyzed, chlorine gas and sodium metal are produced. The two half-reactions that occur in the electrolysis of molten sodium chloride are:

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anode:

cathode:

Overall reaction:

What is electroplating?

True or False (if false, correct the statement to be true)

____1. The half-cell that gains electrons contains the substance that is oxidized.

____2. The electrode at which oxidation occurs is the anode.

____3. In a galvanic cell, electrons flow from the half-cell where reduction occurs to the half-cell where oxidation occurs.

____4. A substance with a high positive reduction potential is a good oxidizing agent.

____5. In electrolysis, oxidation always occurs at the anode.

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Chapter 19 Worksheet #1

Balance the following equations using the half-reaction method of balancing:1. Cr(s) + NO3

-(aq) Cr3+(aq) + NO(g) (acidic solution)

2. MnO4

-(aq) + I-(aq) MnO2(s) + I2(s) (basic solution)

3. NO3

-(aq) + I2(s) IO3-(aq) + NO2(g) (acidic solution)

4. As(s) + NO3

-(aq) H3AsO3(aq) + NO(g) (acidic solution)

5. H2O2(aq) + ClO2(aq) ClO2

-(aq) + O2(g) (basic solution)

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Chapter 19: Worksheet #2

1. Determine whether these redox reactions will occur spontaneously. Calculate the standard cell potential in each case.

a. Cu(s) + 2H+(aq) Cu2+ + H2(g)

b. 2Ag(s) + Fe2+(aq) 2Ag+ + Fe(s)

2. Use the information in Table 23.2 at the end of your notes to calculate standard cell potentials for these voltaic cells.

a. Ni|Ni2+||Cl2|Cl-

b. Sn|Sn2+||Ag+|Ag

3. Describe briefly how you would electroplate a teaspoon with silver.________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________

4. Calculate E0cell and write the overall cell reaction for these cells.

a. Sn|Sn2+||Pb2+|Pb

b. H2|H+||Br2|Br-

5. Write the overall cell reactions and calculate E0cell for voltaic cells composed of the following

sets of half-reactions. a. AgCl(s) + e- Ag(s) + Cl-(aq)

b. Al3+(aq) + 3e- Al(s)

6. This spontaneous redox reaction occurs in the voltaic cell illustrated below.Ni2+(aq) + Fe(s) Ni(s) + Fe2+(aq)

a. Identify the anode and the cathode.

b. Assign charges to the electrode.

c. Write the half reactions.

d. Calculate the cell potentials.

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Chapter 19: Worksheet #3

1. Using the table of standard reduction potentials, arrange the following substances in decreasing order of ability as reducing agents (Hint: A good reducing agent is easily oxidized): Al, Co, Ni, Ag, H2, Na.

2. Using the table of standard reduction potentials, arrange the following substances in order of decreasing ability as oxidizing agents (Hint: A good oxidizing agent is easily reduced): Fe3+, F2, Pb2+, I2, Sn2+, O2

3. Calculate Eocell for the following cells using this data:

A. Cadmium and hydrogen

B. Silver and hydrogen

C. Cadmium and silver

4. For the following redox reaction, write out the two half-reactions and balance the equation, calculate Eo, and determine whether the reaction will occur spontaneously as written.

Fe2+ + MnO4- Fe3+ + Mn2+ (acidic solution)

5. For the following voltaic cell, write the half-reactions, designating which is oxidation and which is reduction. Write the overall cell reaction and calculate the voltage of the cell made from standard electrodes. The cell has electrodes of solid cobalt and nickel, each immersed in a 1 M solution of their ions.

6. Determine the cell reaction and the standard cell potential, Eocell, for the voltaic cells composed of the

following half-cells.Keep Going!

Reaction E o volts CdCd2+ + 2e- 0.403H2 2H+ + 2e- 0.000AgAg+ + e- -0.799

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a. Mg2+(aq) + 2e- Mg(s) Cl2(g) + 2e- 2Cl-(aq)

b. Ni2+(aq) + 2e- Ni(s) Ag+(aq) + e- Ag(s)

c. MnO4

-(aq) + 8H+(aq) + 5e- Mn2+(aq) + 4H2O(l) Cd2+(aq) + 2e- Cd(s)

d. Br2(l) + 2e- 2Br-(aq) Na+(aq) + e- Na(s)

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Chapter 19 Practice TestMULTIPLE CHOICE: (3 PTS EACH)1. Which of these chemical expressions illustrates reduction?

A. Fe2+ Fe3+ + e- C. Mg Mg2+ + 2e-

B. Ca2+ + 2e- Ca D. 2Cl- Cl2 + 2e-

2. When an element is reduced, its atoms:A. gain protons D. lose electronsB. lose protons E. increase in atomic numberC. gain electrons

3. During an oxidation-reduction reaction, the reducing agent:A. is itself reduced. C. is itself oxidizedB. accepts electrons D. undergoes a decrease in oxidation number

4. The oxidation state of nitrogen in the compound HNO3 is:A. +1 B. +2 C. +3 D. +4 E. +5

5. What is the oxidation number of chlorine in the compound sodium perchlorate, NaClO4?A. +1 B. -1 C. +7 D.-7

6. What is the oxidation number of platinum in PtCl62-?

A. -2 B. +2 C. -4 D. +4 E. +12

7. What is the oxidation number of sulfur in sodium sulfite?A. +5 B. -3 C. +3 D. +4

8. In the reaction represented by the equation Mg + CuSO4 MgSO4 + Cu:A. copper is oxidized D. the sulfate ion is the reducing agentB. magnesium is reduced E. the oxidation # of the Cu stays the sameC. magnesium is the reducing agent

9. Which element acts as the oxidizing agent in the reaction? 2HBr + H2SO4 2H2O + SO2 + Br2

A. S in H2SO4 B. H in HBr C. H in H2SO4

D. Br in HBr E. O in H2SO4

10. What is the oxidation number of selenium in SeO42-

A. +6 B. +2 C. +8 D. +4

11. What happens to bromine in the reaction shown by the equation? 2BrO-(aq) 2Br-(aq) + O2(g)A. It is reduced D. It is both oxidized and reducedB. It is oxidized E. It is neither oxidized nor reducedC. It loses electrons

12. What element is reduced in the reaction represented by this equation? Cr2O3 + H2 2CrO + H2OA. Cr B. O C. H D. all three elements

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13. Given the net ionic equation: Zn(s) + Cu2+(aq) Zn2+ + Cu(s)for an electrochemical cell, which is the equation for the reaction at the anode of this cell?A. Zn Zn2+ + 2e- C. Zn2+ + 2e- ZnB. Cu Cu2+ + 2e- D. Cu2+ + 2e- Cu

14. In the electrolysis of water, what happens at the cathode?A. H2 gas bubbles form C. Ions are dischargedB. O2 gas bubbles form D. H2O2 liquid is formed

15. When this half reaction is balanced, Cl- ClO2- the coefficient for the number of electrons is:

A. 1 B. 2 C. 3 D. 4

16. The process by which an electric current produces a chemical change is :A. electrolysis B. electrophoresis C. electrostatics D. electrokinetics

17. Which of the following situations will result in a reaction?A. zinc metal added to a solution of copper(II) sulfateB. copper metal added to a solution of lead(II) sulfateC. lead metal added to a solution of aluminum nitrateD. silver metal added to a solution of copper(II) nitrate

18. A clean iron nail is dipped into a solution of silver nitrate. Predict what you will observe.A. No reaction occurs.B. Te iron nail will become covered with silver.C. Bubbles of oxygen gas will form on the iron nail.D. The iron will be reduced.

19. Balance the following reaction by the half-reaction method: (8 pts) H2SeO3 + Br- Se + Br2 (acidic solution)

20. Balance the following reaction by the half-reaction method: (8 pts) N2O4 + Br- NO2

- + BrO3- (basic solution)

21. In the following unbalanced reactions, determine the reducing agent, the element being reduced, the oxidizing agent, and the element being oxidized. (8 pts)A. 2Na + 2H2O 2NaOH + H2

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oxidized reduced

oxidizing agent reducing agent

B. Sb + H2SO4 Sb2(SO4)3 + H2O + SO2

oxidized reduced

oxidizing agent reducing agent

22. Write and label the half-reactions for oxidation and reduction, write the complete balanced equation and calculate the potential in volts for each of the following reactions. (6 pts each)

Zn + Fe3+ Zn2+ + Fe

Fe2+ + MnO4

- Fe3+ + Mn2+

23. What is the standard cell potential of a cell made of Na/Na+ and Cl-/Cl2? (5 pts)

24. What is the standard cell potential of a voltaic cell made of K/K+ and Zn/Zn2+? (5 pts)

Extra Credit:Titration of the iron(II) ion from a 0.5811g sample of iron ore required 22.49 mL of 0.05347 M K2Cr2O7 to reach the equivalence point. What was the mass percentage of iron in the ore? In acidic solution, dichromate ions are reduced to Cr3+.