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Electrochemistry

Section 20.1 Voltaic Cells

Section 20.2 Batteries

Section 20.3 Electrolysis

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the corresponding slides.

Section 20-1


• Describe a way to obtain electrical energy from a redox reaction.

oxidation: the loss of electrons from the atoms of a substance; increases an atom’s oxidation number

• Identify the parts of a voltaic cell, and explain how each part operates.

• Calculate cell potentials, and determine the spontaneity of redox reactions.

reduction: the gain of electrons by the atoms of a substance; decreases the atom’s oxidation number

Section 20-1

Section 20.1 Voltaic Cells (cont.)

salt bridge

electrochemical cell

voltaic cell

half-cell

In voltaic cells, oxidation takes place at the anode, yielding electrons that flow to the cathode, where reduction occurs.

anode

cathode

reduction potential

standard hydrogen electrode

Section 20-1

Redox in Electrochemistry

• Electrochemistry is the study of the redox processes by which chemical energy is converted to electrical energy and vice versa.

• Redox reactions involve a transfer of electrons from the species that is oxidized to the species that is reduced.

Section 20-1

Redox in Electrochemistry (cont.)

Section 20-1


• A salt bridge is a pathway to allow the passage of ions from one side to another, so that ions do not build up around the electrodes.

• An electrochemical cell is an apparatus that uses a redox reaction to produce electrical energy or uses electrical energy to cause a chemical reaction.

Section 20-1


• A voltaic cell is a type of electrochemical cell that converts chemical energy to electrical energy by a spontaneous redox reaction.

CMC In Motion/CMC_P709.rm

Section 20-1

Chemistry of Voltaic Cells

• An electrochemical cell consists of two parts called half-cells, in which the separate oxidation and reduction reactions take place.

• The electrode where oxidation takes place is called the anode.

• The cathode is the electrode where reduction occurs.

Section 20-1

Chemistry of Voltaic Cells (cont.)

• Electric potential energy is a measure of the amount of current that can be generated from a voltaic cell to do work.

• Electric charge can flow between two points only when a difference in electric potential energy exists between the two points.

• A volt is a unit used to measure cell potential—the force from the difference in electric potential energy between two electrodes.

Section 20-1

Calculating Electrochemical Cell Potentials

• The tendency of a substance to gain electrons is its reduction potential.

• When two half-reactions are coupled, the voltage generated corresponds to the difference in potential between the half-reactions.

Section 20-1

Calculating Electrochemical Cell Potentials (cont.)

• The standard hydrogen electrode consists of a small sheet of platinum immersed in a hydrochloric acid solution that has a hydrogen ion concentration of 1 M. Hydrogen gas (H2), at a pressure of 1 atm, is bubbled in and the temperature is maintained at 25°C.

Section 20-1


• The standard hydrogen electrode is the standard against which all other reduction potentials are measured.

Section 20-1


Section 20-1


• Formula for cell potential

E0cell = E0

reduction – E0oxidation

Section 20-1

Use Standard Reduction Potentials

• Cell potentials can be used to determine if a proposed reaction under standard conditions will be spontaneous.

• If the calculated potential is positive, the reaction is spontaneous.

• If the calculated potential is negative, the reaction is not spontaneous.

A. A

B. B

C. C

D. D

Section 20-1

Section 20.1 Assessment

In electrochemistry, the site where oxidation occurs is called ____.

A. electrode

B. anode

C. cathode

D. ion

A B C D

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A. A

B. B

C. C

D. D

Section 20-1

A B C D

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The standard potential of a voltaic cell is the difference between the:

A. electrode voltage

B. standard reduction potential of the cell and hydrogen

C. standard reduction potentials of the half-cell reactions

D. half-cell reactions and the salt bridge

Section 20-2


• Describe the structure, composition, and operation of the typical carbon-zinc dry-cell battery.

reversible reaction: a reaction that can take place in both the forward and reverse directions

• Distinguish between primary and secondary batteries, and give two examples of each type.

• Explain the structure and operation of the hydrogen-oxygen fuel cell.

• Describe the process of corrosion of iron and methods to prevent corrosion.

Section 20-2

Section 20.2 Batteries (cont.)

battery

dry cell

primary battery

secondary battery

Batteries are voltaic cells that use spontaneous reactions to provide energy for a variety of purposes.

fuel cell

corrosion

galvanization

Section 20-2

Dry Cells

• A battery is one or more voltaic cells in a single package that generates electric current.

Section 20-2

Dry Cells (cont.)

• A dry-cell is an electrochemical cell in which the electrolyte is a moist paste. The paste in a zinc-carbon cell consists of zinc chloride, manganese(IV) oxide, ammonium chloride, and a small amount of water.

• The anode is the zinc shell.

• The cathode is a carbon rod, but reduction occurs in the paste.

Section 20-2

Dry Cells (cont.)

• In the alkaline cell, zinc is in a powdered form and mixed with potassium hydroxide contained in a steel case.

• Alkaline batteries are small and more useful in small devices.

Section 20-2

Dry Cells (cont.)

• Silver batteries are similar to alkaline but smaller.

Section 20-2

Dry Cells (cont.)

• Primary batteries produce electric energy by means of redox reaction that are not easily reversed.

• Secondary batteries depend on reversible redox reactions and are rechargeable.

Section 20-2

Lead-Acid Storage Battery

• Lead-acid storage batteries are common in automobiles.

• The electrolyte solution is sulfuric acid, hence the name.

• The anode consists of grids of porous lead.

• The cathode consists of lead grids filled with lead(IV) oxide.

Section 20-2

Lithium Batteries

• Lithium is the lightest known metal and has the lowest standard reduction potential of the metallic elements.

• Lithium batteries can be either primary or secondary.

Section 20-2

Fuel Cells

• A fuel cell is a voltaic cell in which the oxidation of a fuel is used to produce electric energy.

Section 20-2

Fuel Cells (cont.)

• How a fuel cell works

− Potassium hydroxide is often the electrolyte.

− The oxidation half reaction

2H2(g) + 4OH– → 4H2O + 4e–

− The reduction half reaction

O2(g) + 2H2O(l) + 4e– → 4OH–(aq)

− When combined, the equation is the same as burning hydrogen in oxygen to form water.

Section 20-2

Corrosion

• Corrosion is the loss of metal resulting from an oxidation-reduction reaction of the metal with substances in the environment.

• Rusting begins in a chip or pit in the iron surface, which become the anode.

Fe(s) → Fe2+(aq) + 2e–

• Iron(II) becomes part of the water solution.

Section 20-2

Corrosion (cont.)

• The cathode is usually at the edge of the water drop where water, iron, and air come into contact.

• The reduction reaction is O2(g) + 4H+(aq) + 4e– → 2H2O(l).

• Next the oxidation 4Fe2+(aq) + 2O2(g) + 2H2O(l) + 4e– → 2Fe2O3(s) + 4H+.

• Rusting is slow, but salts speed the process.

Section 20-2

Corrosion (cont.)

Section 20-2

Corrosion (cont.)

• Paint and other covers seal out moisture to prevent corrosion.

• Blocks of metal that are more easily oxidized than steel, such as magnesium, aluminum, or titanium, are often attached to the hulls of ships—they corrode while the iron in the hull is protected.

Section 20-2

Corrosion (cont.)

• Galvanization is the process of coating iron with a layer or zinc.

• Zinc oxidizes at the surface, creating a layer of metal-oxide that protects from further corrosion.

A. A

B. B

C. C

D. D

Section 20-2

A B C D

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Which type of battery has a reversible spontaneous reaction?

A. alkaline battery

B. secondary battery

C. primary battery

D. zinc-carbon battery

A. A

B. B

C. C

D. D

Section 20-2


A B C D

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Which is NOT a method of preventing corrosion?

A. painting

B. galvanization

C. coating with electrolytes

D. sacrificial anode

Section 20-3


• Describe how it is possible to reverse a spontaneous redox reaction in an electrochemical cell.

redox reaction: an oxidation-reduction reaction

electrolysis

electrolytic cell • Compare the reactions

involved in the electrolysis of molten sodium chloride with those in the electrolysis of brine.

• Discuss the importance of electrolysis in the smelting and purification of metals.

In electrolysis, a power source causes nonspontaneous reactions to occur in electrochemical cells.

Section 20-3

Reversing Redox Reactions

• The use of electrical energy to bring about a chemical reaction is called electrolysis.

• An electrochemical cell in which electrolysis occurs is called an electrolytic cell.

Section 20-3

Applications of Electrolysis

• Electrolysis of water is one method of obtaining hydrogen gas for commercial use.

Section 20-3

Applications of Electrolysis (cont.)

• Electrolysis can separate molten sodium chloride into sodium metal and chlorine gas in a chamber called a Down’s cell.

Section 20-3


• In the decomposition of brine, an aqueous solution of NaCl, electrolysis is used to produce hydrogen gas, chlorine gas, and sodium hydroxide.

Section 20-3


• The Hall-Héroult process requires large amounts of energy, which is the primary reason for recycling aluminum.

Section 20-3


• Objects can be electroplated with a metal such as silver.

• The cathode is the object to be electroplated, where reduction occurs.

• The anode is a bar of silver, where silver is oxidized and silver ions transferred to the cathode.

A. A

B. B

C. C

D. D

Section 20-3

A B C D

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Which of the following is NOT a product of electrolysis of brine?

A. Chlorine gas

B. Hydrogen gas

C. Sodium hydroxide solution

D. Sodium metal

A. A

B. B

C. C

D. D

Section 20-3


A B C D

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What is required to drive a nonspontaneous reaction in an electrolytic cell?

A. electrodes

B. additional ions

C. an energy source

D. an electrolyte

Study Guide 1


Key Concepts

• In a voltaic cell, oxidation and reduction take place at electrodes separated from each other.

• The standard potential of a half-cell reaction is its voltage when paired with a standard hydrogen electrode under standard conditions.

• The reduction potential of a half-cell is negative if it undergoes oxidation when connected to a standard hydrogen electrode. The reduction potential of a half-cell is positive if it undergoes reduction when connected to a standard hydrogen electrode.

Study Guide 1

Section 20.1 Voltaic Cells (cont.)

Key Concepts

• The standard potential of a voltaic cell is the difference between the standard reduction potentials of the half-cell reactions.

Study Guide 2


Key Concepts

• Primary batteries can be used only once; secondary batteries can be recharged.

• When a battery is recharged, electric energy supplied to the battery reverses the direction of the battery’s spontaneous reaction.

• Fuel cells are batteries in which the substance oxidized is a fuel from an external source.

• Methods of preventing corrosion are painting, coating with another metal, or using a sacrificial anode.

Study Guide 3


Key Concepts

• In an electrolytic cell, an outside source of power causes a nonspontaneous redox reaction to occur.

• The electrolysis of molten sodium chloride yields sodium metal and chlorine gas. The electrolysis of brine yields hydrogen gas, sodium hydroxide, and chlorine gas.

• Metals such as copper are purified in an electrolytic cell.

• Electrolysis is used to electroplate objects and to produce pure aluminum from its ore.

A. A

B. B

C. C

D. D

Chapter Assessment 1

A B C D

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A salt bridge is used for what?

A. electrolysis

B. allowing ions to move from one half-cell to the other

C. increasing the reduction potential

D. linking multiple voltaic cells together

A. A

B. B

C. C

D. D


A B C D

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Batteries that can be used more than once are also called ____.

A. primary batteries

B. secondary batteries

C. tertiary batteries

D. fuel cells

A. A

B. B

C. C

D. D


A B C D

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Corrosion is loss of metal by what mechanism?

A. electrolysis

B. galvanization

C. oxidation-reduction reaction

D. salt bridge

A. A

B. B

C. C

D. D


A B C D

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Electrolysis of H2O results in what products?

A. H+ and OH–

B. H2O2 and O2

C. NaOH and H2

D. H2 and O2

A. A

B. B

C. C

D. D


A B C D

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Which of the following is NOT a dry-cell?

A. a car battery

B. a silver battery

C. a zinc-carbon battery

D. an alkaline battery

A. A

B. B

C. C

D. D

STP 1

A B C D

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Which is NOT a characteristic of an acid?

A. sour taste

B. conduct electricity in solution

C. reactivity with some metals

D. slippery feel

A. A

B. B

C. C

D. D

STP 2

A B C D

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Which metal has the lowest standard reduction potential?

A. Sodium

B. Magnesium

C. Lithium

D. Silver

A. A

B. B

C. C

D. D

STP 3

A B C D

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Which statement is NOT true about batteries?

A. Batteries are composed of two half-cells.

B. Secondary batteries are also known as storage batteries.

C. Redox reactions in zinc-carbon batteries are easily reversible.

D. Batteries are compact forms of voltaic cells.

A. A

B. B

C. C

D. D

STP 4

A B C D

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Which of the following is the equation used to calculate a cell potential?

A. E0cell = E0

oxidation – E0reduction

B. E0cell = E0

reduction – E0oxidation

C. E0cell = E0

reduction + E0oxidation

D. E0cell = E0

reduction x E0oxidation

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