21 Electrochemistry Electrochem Planning Guide

726A Chapter 21

ElectrochemistryPlanning GuideElectrochemPlanning G21

Every chemical process uses or produces energy.

Lessons and Objectives Print Resources

For the Student For the Teacher

B-3, E-2 21.1 Electrochemical Cells p 728–73621.1.1 Identify the type of chemical reaction

that is involved in all electrochemical processes.

21.1.2 Describe how a voltaic cell produces electrical energy.

21.1.3 Identify the current applications that use electrochemical processes to produce electrical energy.

Reading and Study Workbook Lesson 21.1

Lesson Assessment 21.1 p 736

Teaching Resources, Lesson 21.1 Review

Teacher Demo, p 730: A Redox Reaction

Teacher Demo, p 732: Inside a Dry Cell

Teacher Demo, p 733: Making a Lead Cell

A-1, A-2, B-3 21.2 Half-Cells and Cell Potentials p 737–74321.2.1 Identify what causes the electrical

potential of an electrochemical cell.21.2.2 Determine the standard reduction potential

of a half-cell.21.2.3 Determine if a redox reaction is

spontaneous or nonspontaneous.




Teacher Demo, p 739: The Corrosion of Iron

Class Activity, p 741: Combining Half-Reactions

A-1, A-2, B-3, E-2

21.3 Electrolytic Cells p 745–75121.3.1 Distinguish between electrolytic and

voltaic cells.21.3.2 Describe some applications that use

electrolytic cells.



Quick Lab: Electrochemical Analysis of Metals, p 750

Small-Scale Lab: Electrolysis of Water, p 752


Teacher Demo, p 747: The Electrolysis of Water

Essential Questions1. How is energy produced in an electrochemical

process?2. How can energy be used to drive an

electrochemical process?

Study Guide p 753STP p 759Reading and Study

Workbook Self-Check and Vocabulary Review Chapter 21

Every chemical process uses or produces energy.E h i l d

Introducing the BIGIDEA: MATTER AND ENERGY

Essential Questions1. How is energy produced in an electrochemical

Study Guide p 753STP p 759

Assessing the BIGIDEA: MATTER AND ENERGY

NSES

Electrochemistry 726B

For the StudentQuick Lab, p 750

9-V battery, sodium • sulfate solutioncopper penny• nickel coin• iron nail• filter paper• aluminum foil• reaction surface•

Small-Scale Lab, p 752reaction surface• electrolysis device• droppers• 0.2• M Na2SO40.1• M KI0.2• M CuSO40.04% BTB• 1.0• M NaCl0.2• M KBrliquid starch•

For the TeacherTeacher Demo, p 730

200 mL of 0.1• M silver nitrate250-mL beaker• glass stirring rod• strip of polished copper • metal

Teacher Demo, p 7332 strips of lead• wooden rod• 2 connecting wires• 250-mL beaker• dilute sulfuric acid• 6-V DC power supply• doorbell•

Teacher Demo, p 732dry cell• voltmeter• 2 wire leads• zinc electrode• carbon electrode• manganese dioxide/• ammonium chloride paste

Teacher Demo, p 7392 pennies• metal wires from twist ties• table salt• paper towel•

Teacher Demo, p 747Hoffman apparatus or • apparatus constructed using iron nails as electrodes2 test tubes• 6• M NaOH250- or 500-mL beaker• 2 insulated wire leads•

Online Student Edition Online Teacher’s Edition20.2 Virtual Chem Lab 30: Redox Titrations:

Determination of Iron

Online Student Edition

Additional Digital Resources

Digital Resources

Editable Worksheets PearsonChem.com

Small-Scale Lab Manual Lab 36: Small-Scale Voltaic CellsLab 47: Corrosion

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An electrochemical process was used to produce the shiny chrome finish on this car.

726 Chapter 21

Focus on ELL

1 CONTENT AND LANGUAGE Begin the class by writing Electrochemistry on the board and pronouncing it slowly. Have students determine the meaning of the word from its prefix. Then, present a complete list of the vocabulary terms in this chapter. Have students examine how these terms are related to the chapter content.

BEGINNING: LOW/HIGH Create illustrations to represent the terms. If literate, create a native language definition.

INTERMEDIATE: LOW/HIGH Provide students with sentences in which key words have been replaced with blanks. Have students complete the blanks with words from a given word bank as they read through the chapter.

ADVANCED: LOW/HIGH Challenge students to create their own sentences in which key words have been replaced with blanks. Have students trade papers with a partner to help them review lesson terminology.

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virtual lab tour in which electrochemical reactions are studied in a simulated laboratory environment.”

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MATTER AND ENERGY

Essential Questions:1. How is energy produced in an

electrochemical process?2. How can energy be used to drive an

electrochemical process?

BIGIDEA

Trash or Treasure?Maria and her friend decided to spend the Saturday browsing a local flea market. At one vendor’s display, Maria spotted a beautiful, shiny gold ring. The vendor told Maria that the ring was an antique from the 1800s and made from solid gold. Enamored by the ring, Maria purchased it and placed it on her finger.

Several weeks later, while taking off the ring, Maria noticed that the ring was discol-ored in many places. It almost looked like the gold was peeling off the ring. Maria was disturbed because the ring was expensive, and she believed that it was a valuable antique. She decided to take the ring to a jeweler to see if it could be polished and restored to its origi-nal gold color. However, Maria became upset when the jeweler revealed the truth about her “gold” ring.

Connect to the BIGIDEA As you read about electrochemical processes, think about how a ring could be made to look like it was made out of pure gold when it was not.

CHEMYSTERY

NATIONAL SCIENCE EDUCATION STANDARDS

A-1, B-3, B-5, E-2, G-1, G-3

Electrochemistry 727Electrochemistry 727

Understanding by DesignStudents are building toward understanding electrochemistry using the relationship between matter and energy.

PERFORMANCE GOALS At the end of Chapter 21, students will be able to answer the essential questions by applying their knowledge of electrochemistry. Students will also be able to calculate standard cell potentials.

ESSENTIAL QUESTIONS Read the essential questions aloud. Ask What two main types of energy do you think are involved in an electrochemical process? (electrical and chemical) Ask Why would electrical energy be involved in redox reactions? (It involves the flow of electrons.)

BIGIDEA Use the photo of the chrome on a car to help students connect to the

concepts they will learn in this chapter. Activate prior knowledge by asking students to name chrome items familiar to them. Ask Is chrome an element? (No. It is an alloy, or mixture of metals, containing chromium that is deposited onto a surface.) Ask How do you think they get the chrome onto the surface? (using a redox reaction)

CHEMYSTERY Have students read over the CHEMystery. Connect the

CHEMystery to the Big Idea of Matter and Energy by reminding students that matter can undergo both physical changes and chemical changes. Encourage students to think about whether a physical change or chemical change might have been used to make Maria’s ring look like solid gold. Then, have them predict what the change was. As a hint, suggest that they consider how electrical energy could have been used to change the ring.

Introduce the ChapterCONNECTIONS TO REDOX REACTIONS In Chapter 20, students learned how oxidation-reduction reactions can change materials. Use these activities to introduce students to the connection between redox reactions and electrochemistry.

Activity 1 You will need a simple circuit composed of wire, a dry cell battery, and a bulb. Have a student connect the wire to complete the circuit. Point out that an electrical current is needed to light the bulb. Explain that charge moves through the battery by a redox reaction. Tell students that they will learn about electrical applications of redox reactions in this chapter.

Activity 2 You will need rusted nails, a nut or bolt electroplated with zinc, and a shiny chrome material which has scratches with rust. Pass the materials around for students to observe. Point out that the coated metal did not rust. Ask How do you think the coating was applied to the metals? (Students may think it was painted on.) Explain that electrical energy was used to cause a redox reaction that coated the metal.

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Electrochemical Cells21.1

Key Questions What type of chemical

reaction is involved in all electrochemical processes?

How does a voltaic cell produce electrical energy?

What current applications use electrochemical processes to produce electrical energy?

Vocabulary

Q: Why do some kinds of jellyfish glow? On a summer evening, fireflies glow to attract their mates. In the ocean depths, angelfish emit light to attract prey. Luminous shrimp, squid, jellyfish, and even bacteria also exist. These organ-isms, and others, are able to give off energy in the form of light as a result of redox reactions.

Electrochemical Processes What type of chemical reaction is involved in all

electrochemical processes?Chemical processes can either release energy or absorb energy. The energy can sometimes be in the form of electricity. An electrochemical process is any conversion between chemical energy and electrical energy. All elec-trochemical processes involve redox reactions. Electrochemical processes have many applications in the home as well as in industry. Flashlight and automobile batteries are familiar examples of devices used to generate elec-tricity. The manufacture of sodium and aluminum metals and the silver-plating of tableware involve the use of electricity. Biological systems also use electrochemistry to carry nerve impulses.

Redox Reactions and the Activity Series When a strip of zinc metal is dipped into an aqueous solution of blue copper(II) sulfate, the zinc becomes copper-plated, as shown in Figure 21.1. The net ionic equation involves only zinc and copper.

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

Figure 21.1 Redox ReactionZinc metal oxidizes spontaneously in a copper-ion solution. a. A zinc strip is immersed in a solution of copper(II) sulfate. b. As the copper plates out onto the zinc, the blue copper(II) sulfate solution is replaced by a colorless solution of zinc sulfate. The copper appears black because it is in a finely divided state.

728 Chapter 21 • Lesson 1

Focus on ELL

1 CONTENT AND LANGUAGE Draw students’ attention to the words anode and cathode. Explain that the word cathode comes from the Greek kathodos meaning “way down” and that the word anode comes from the Greek anodos meaning “way up.” Thus the two words, cathode and anode, are opposites. Write the symbols – and + next to words.

2 FRONTLOAD THE LESSON Ask students to recall what they know about electricity, electrical current, and circuits. Then, have students examine Figure 21.3 and identify features that the voltaic cell shares with an electrical circuit. Explain that, in this lesson, they will learn how an electrical current can be generated using chemicals.

3 COMPREHENSIBLE INPUT Help students visualize a dry cell as a classic aqueous voltaic cell by drawing Figure 21.3 on the board. Have students use a think-aloud strategy to determine which of the dry cell components in Figure 21.4 are equivalent to the aqueous cell components as you write the labels on the drawing.

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Key Objectives21.1.1 IDENTIFY the type of chemical reaction

that is involved in all electrochemical processes.

21.1.2 DESCRIBE how a voltaic cell produces electrical energy.

21.1.3 IDENTIFY the current applications that use electrochemical processes to produce electrical energy.

Additional ResourcesReading and Study Workbook, Lesson 21.1

Available Online or on Digital Media:

• Teaching Resources, Lesson 21.1 Review• Laboratory Manual, Lab 47• Small-Scale Chemistry Laboratory Manual, Lab 36

EngageCHEMISTRY YOU YOYY U&& Have students read the

opening paragraph. Ask What have you learned that shows that energy can be released in the form of light? (When electrons drop from higher to lower energy levels in an atom they emit energy in the form of light.)

Activate Prior Knowledge Review the concepts of reversible reactions and spontaneous and nonspontaneous reactions that were introduced in Chapter 18. Ask If a reaction releases free energy, is the reaction spontaneous or nonspontaneous? (spontaneous)

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Electrons are transferred from zinc atoms to copper ions. This is a spon-taneous redox reaction. Zinc atoms lose electrons as they are oxidized to zinc ions, while copper ions in solution gain the electrons lost by the zinc. The copper ions are reduced to copper atoms and are deposited as metallic copper. As the copper ions in solution are gradually replaced by zinc ions, the blue color of the solution fades. The balanced half-reactions for this redox reaction can be written as follows:

Oxidation: Zn(s) Zn2 (aq) 2e

Reduction: Cu2 (aq) 2e Cu(s)

In the activity series of metals in Table 21.1, zinc is above copper on the list. For any two metals in an activity series, the more active metal is the more readily oxidized. Zinc is more readily oxidized than copper. When zinc is dipped into a copper(II) sulfate solution, zinc becomes plated with copper. In contrast, when a copper strip is dipped into a solution of zinc sulfate, the cop-per does not spontaneously become plated with zinc. This is because copper metal is not oxidized by zinc ions.

Remember: The more active metal will be oxidized; the less active metal will be reduced.

ized;

Table 21.1 The half-reaction for the oxidation of each metal is shown.

Read Tables a. What is the half-reaction for the oxidation of nickel?

Compare b. Which metal is more readily oxidized, lead or magnesium?

Relate Cause and c. Effect What will happen if a strip of copper is dipped in a solution of silver nitrate? If a reaction occurs, write the half-reactions.

Activity Series of Metals

Element Oxidation half-reaction

s) aq) e

s) K aq) e

s) Ba2 aq) 2e

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s) Al3 aq) 3e

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s) Fe2 aq) 2e

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H2 g) 2H aq) 2e

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s) aq) e

s) 2 aq) 2e

s) 3 aq) 3e*Hydrogen is included for reference purposes.

Electrochemistry 729

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Foundations for ReadingBUILD VOCABULARY Have students create a concept map from the vocabulary terms for the lesson. Encourage students to include drawings or symbols to enhance the descriptions used to explain the connections made in the map.

READING STRATEGY Show students how to make an outline as they read this section. Have them use the red headings as their main entries and the blue headings as secondary entries. Ask them to include, under each heading, a sentence summarizing the important idea.

Explain

Electrochemical ProcessesUSE VISUALS Direct students to Figure 21.1. Ask What is the color of the copper (II) sulfate solution in Figure 21.1a? (blue) Is the color different in Figure 21.1b? (It’s still blue but lighter in color.) What is the significance of the lighter color? (Blue copper (II) ions have been removed from the solution and changed into the copper atoms on the zinc.) Next, direct students’ attention to Table 21.1. Ask In regard to what you have learned previously, what does most active mean? (The ionization energy needed for an atom to lose an electron is less, therefore it is more likely to occur.) Have students locate Zn and Cu in the table. Ask Look back at the net ionic equation for the copper-plating reaction. Which species was oxidized in that reaction? (Zn) Ask If a gold bar was placed in the CuSO4 solution, would you expect it to become copper-plated? Why? (No, it is lower on the activity series than copper, so it would be more likely to be reduced than oxidized.)

APPLY CONCEPTS Explain that redox reactions make it possible for energy interconversion between electrical energy and chemical energy. Review oxidation and reduction half-cell reactions and stress that the electrons must be balanced. Show how these reactions are combined to form the net ionic equation for an electrochemical process. Point out that the net equation summarizes the transfer of electrons from the species being oxidized to the species being reduced.

AnswersINTERPRET DATAa. Ni(s) → Ni2+(aq) + 2e−

b. magnesiumc. The copper becomes plated with silver.

Cu(s) → Cu2+(aq) + 2e−

Ag+(aq) + e− → Ag(s)

Differentiated InstructionL3 ADVANCED STUDENTS Challenge students to research the electrochemical nature of the nervous system. Students may want to focus on different aspects of the system and combine their findings into a visual display and/or oral report. Possible areas to be addressed include the role of sodium and potassium ions, how signals are transmitted across synapses, what an EEG measures, and the value of squid for nervous system research.

L1 STRUGGLING STUDENTS Have students think about the half-reactions in Table 21.1 in terms of where the element is located in the periodic table. Remind them of what they learned in Chapter 6 about periodic trends in chemical properties. Point out that the elements on the left side of the periodic table are more easily oxidized (lose electrons) than the elements on the right side of the periodic table. Explain that Group 1A and 2A metals, listed at the top of the activity series, are some of the most reactive elements in nature.

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Figure 21.2 Voltaic CellA voltaic cell powers this mp3 player.Predict What other items contain voltaic cells?

Electrochemical Cells When a zinc strip is dipped into a copper(II) sulfate solution, electrons are transferred from zinc atoms to copper ions. This flow of electrons is an electric current. If a redox reaction is to be used as a source of electrical energy, the two half-reactions must be physically separated. In the case of the zinc-metal–copper-ion reaction, the electrons released by the zinc atoms must pass through an external circuit to reach the copper ions if useful electrical energy is to be produced. In that situation, the system serves as an electrochemical cell. Alternatively, an electric current can be used to produce a chemical change. That system also serves as an electrochemical cell. Any device that converts chemical energy into electrical energy or elec-trical energy into chemical energy is an electrochemical cell. Redox reactions occur in all electrochemical cells.

Voltaic Cells How does a voltaic cell produce electrical energy?

In 1800, the Italian physicist Alessandro Volta built the first electrochemical cell that could be used to generate a direct electric current (DC). Named after its inventor, a voltaic cell is an electrochemical cell used to convert chemical energy into electrical energy. Electrical energy is produced in a voltaic cell by a spontaneous redox reaction within the cell. You can find voltaic cells everywhere. They power your flashlight and your mp3 player, as shown in Figure 21.2.

Constructing a Voltaic Cell A voltaic cell consists of two half-cells. A half-cell is one part of a voltaic cell in which either oxidation or reduction occurs. A typical half-cell consists of a piece of metal immersed in a solu-tion of its ions. Figure 21.3 on the following page shows a voltaic cell that makes use of the zinc–copper reaction. In this cell, one half-cell is a zinc strip immersed in a solution of zinc sulfate. The other half-cell is a copper strip immersed in a solution of copper(II) sulfate.

The half-cells are connected by a salt bridge, which is a tube containing a strong electrolyte, often potassium sulfate (K2SO4). Salt bridges also con-tain agar, a gelatinous substance. A porous plate may be used instead of a salt bridge. The salt bridge or porous plate allows ions to pass from one half-cell to the other but prevents the solutions from mixing completely. A wire car-ries the electrons in the external circuit from the zinc strip to the copper strip. A voltmeter or light bulb can be connected in the circuit. The driving force of such a voltaic cell is the spontaneous redox reaction between zinc metal and copper ions in solution.

The zinc and copper strips in this voltaic cell serve as the electrodes. An electrode is a conductor in a circuit that carries electrons to or from a substance other than a metal. The reaction at the electrode determines whether the electrode is labeled as an anode or a cathode. The electrode at which oxidation occurs is called the anode. Electrons are produced at the anode. Therefore, the anode is labeled the negative electrode in a voltaic cell. The electrode at which reduction occurs is called the cathode. Electrons are consumed at the cathode in a voltaic cell. As a result, the cathode is labeled the positive electrode. Neither electrode is actually charged, however. All parts of the voltaic cell remain balanced in terms of charge at all times. The moving electrons balance any charge that might build up as oxidation and reduction occur.

Q: Jellyfish and other creatures that glow contain compounds that undergo redox reactions. What do these reactions have in common with redox reactions that occur in electrochemical cells?

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Teacher DemoPURPOSE Students will observe a redox reaction.

MATERIALS 200 mL 0.1M silver nitrate (AgNO3) in a 250 mL beaker, glass stirring rod, strip of polished copper metal

SAFETY Wear gloves, apron, and safety goggles.

PROCEDURE Put a glass stirring rod across the top of the beaker and suspend a polished copper strip from the rod so that the strip dips into the AgNO3

solution. The metal surface will darken and appear “fuzzy” as silver metal is deposited. Over time, a layer of silver will form on the copper. Ask What is the significance of the blue color of the solution? (Blue copper(II) ions are being produced.) Write on the board the net reaction for the electrochemical process:

2Ag+(aq) + Cu(s) → 2Ag(s) + Cu2+(aq) Point out that the reaction involves a transfer of electrons from copper atoms to silver ions. Ask What was oxidized? What was reduced? (Copper was oxidized; silver ions were reduced.) Combine liquid wastes and add a 50% molar excess of NaCl. Filter or decant and dry the AgCl residue. Put in a plastic container and bury in an approved landfill. Flush the filtrate down the drain with excess water.

EXPECTED OUTCOME Silver crystals form on the copper wire. The solution turns blue because of dissolved copper(II) ions.

Check for UnderstandingWhat type of chemical reaction is involved in all electrochemical

processes?

Assess students’ understanding that electrochemical processes are redox reactions by asking the following questions. Ask What is transferred between ions in a redox reaction? (electrons) Ask What is transferred between ions in an electrochemical reaction? (electrons) Ask What happens to the ion that gives up electrons in an electrochemical reaction? (It is oxidized.) Ask What happens to the ion that receives electrons in an electrochemical reaction? (It is reduced.)

ADJUST INSTRUCTION If students are confusing oxidation and reduction, have them reread Lesson 20.1 on the Meaning of Oxidation-Reduction as well as the text following the section on Electrochemical Processes in Lesson 21.1. Then, write the chemical reaction shown at the bottom of the previous page on the board and diagram the movement of electrons between the reactants and products to reinforce the concept that this electrochemical reaction is an oxidation-reduction reaction.

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Figure 21.3 Zinc–Copper Voltaic CellIn this voltaic cell, the electrons generated from the oxidation of Zn to Zn2 flow through the external circuit (the wire) into the copper strip. These electrons reduce the surrounding Cu2 to Cu. To maintain neutrality in the electrolytes, anions flow through the salt bridge.Explain What is the purpose of the salt bridge?

How a Voltaic Cell Works The electrochemical process that occurs in a zinc–copper voltaic cell can best be described in a number of steps. These steps actually occur at the same time.

Step 1 Electrons are produced at the zinc strip according to the oxidation half-reaction:

Zn(s) Zn2 (aq) 2e

Zinc is oxidized at the zinc strip, so the zinc strip is the anode, or negative electrode, in the voltaic cell.

Step 2 The electrons leave the zinc anode and pass through the external cir-cuit to the copper strip. (If a bulb is in the circuit, the electron flow will cause it to light. If a voltmeter is present, it will indicate a voltage.)

Step 3 Electrons enter the copper strip and interact with copper ions in solu-tion. There, the following reduction half-reaction occurs:

Cu2 (aq) 2e Cu(s)

Copper ions are reduced at the copper strip, so the copper strip is the cathode, or positive electrode, in the voltaic cell.

Step 4 To complete the circuit, both positive and negative ions move through the aqueous solutions via the salt bridge. The two half-reactions can be summed to show the overall cell reaction. Note that the electrons must cancel.

Zn(s) Zn2 (aq) 2e

Cu2 (aq) 2e Cu(s)

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

See voltaic cells animated online.

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AnswersFIGURE 21.2 Most handheld electronic devices are

powered by voltaic cells, including calculators and flashlights.

FIGURE 21.3 to prevent significant build-up of charges at the electrodes

Explain

Voltaic CellsUSE VISUALS Direct students’ attention to Figure 21.3. Explain that voltaic cells can be used as sources of electrical energy because the two half-reactions are physically separated. The reaction in the illustration could take place in a single beaker, but it would not be possible to produce a stream of electrons. Ask At which electrode does oxidation (loss of electrons) take place? (Oxidation occurs at the anode (negative electrode).) Ask Where does reduction (gain of electrons) take place? (Reduction occurs at the cathode (positive electrode).) Ask What path do the electrons given up by zinc follow? (They go through the wire and the electric light to the copper electrode.) Ask What happens to the electrons at the copper electrode? (They reduce copper ions to copper.)

APPLY CONCEPTS In discussing the need for a salt bridge in the Zn/Cu voltaic cell, explain that as zinc is oxidized at the anode, Zn2+ ions enter the solution. Explain that they have no negative ions to balance their charges, so a positive charge tends to build up around the anode. Similarly, point out that at the cathode, Cu2+ ions are reduced to Cu and taken out of the solution, leaving behind unbalanced negative ions. Thus, a negative charge tends to develop around the cathode. Convey that the salt bridge allows negative ions, such as SO4

2−, to be drawn to the anode compartment to balance the growing positive charge. Likewise, positive ions, such as K+, are drawn from the salt bridge to balance the growing negative charge at the cathode.

Differentiated Instruction L1 LESS PROFICIENT READERS Students may have difficulty remembering that oxidation occurs at the anode of a voltaic cell and reduction occurs at the cathode. Tell them that consonants go together and vowels go together. Both cathode and reduction begin with consonants and anode and oxidation start with vowels.

ELL ENGLISH LANGUAGE LEARNERS Clarify students’ understanding of the word cell as it relates to this lesson. Make sure they are not looking for a living cell as they might have learned from biology class. Stress that an electrochemical cell is a chemical system that generates electrical energy.

L1 SPECIAL NEEDS STUDENTS For developmentally delayed students, provide simple DC devices, such as flashlights, cell phones, walkie-talkies, etc. Have them determine how the battery must be inserted in order for the object to power up. Tell students to pay particular attention to the markings on the batteries and the battery compartments.

Positive button (+) Positive button (+)

Dry Cell Alkaline Battery

Steel case

MnO2 in KOH paste

Graphite rod (cathode)

Absorbent separator

Zinc (anode)

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Graphite rod(cathode)

Moist paste of MnO2,ZnCl2, NH4Cl, H2O,and graphite powder

Zinc (anode)

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Figure 21.4 Dry CellsBoth dry cells and alkaline batteries are single electrochemical cells that produce about 1.5 V. a. The dry cell is inexpensive, has a short shelf life, and suffers from voltage drop when in use. b. The alkaline battery has a longer shelf life and does not suffer from voltage drop.Apply Concepts What is oxidized in these cells and what is reduced?

Representing Electrochemical Cells You can represent the zinc–cop-per voltaic cell by using the following shorthand form.

Zn(s) | ZnSO4(aq) || CuSO4(aq) | Cu(s)

The single vertical lines indicate boundaries of phases that are in con-tact. The zinc strip, Zn(s), and the zinc sulfate solution, ZnSO4(aq), for example, are separate phases in physical contact. The double vertical lines represent the salt bridge or porous partition that separates the anode compartment from the cathode compartment. The half-cell that undergoes oxidation (the anode) is written first, to the left of the double vertical lines.

Using Voltaic Cells as Energy Sources What current applications use electrochemical processes to

produce electrical energy?Although the zinc–copper voltaic cell is of historical importance, it is no longer used commercially. Current applications that use electro-chemical processes to produce electrical energy include dry cells, lead storage batteries, and fuel cells.

Dry Cells When a compact, portable electrical energy source is required, a dry cell is usually chosen. A dry cell is a voltaic cell in which the electrolyte is a paste. In one type of dry cell, a zinc container is filled with a thick, moist electrolyte paste of manganese(IV) oxide (MnO2), zinc chloride (ZnCl2), ammonium chloride (NH4Cl), and water (H2O). As shown in Figure 21.4a, a graphite rod is embedded in the paste. The zinc container is the anode, and the graphite rod is the cathode. The thick paste and its surrounding paper liner prevent the contents of the cell from freely mixing, so a salt bridge is not needed. The half-reactions for this cell are shown below.

Oxidation: Zn(s) Zn2 (aq) 2e (at anode)

Reduction: 2MnO2(s) 2NH4 (aq) 2e Mn2O3(s) 2NH3(aq) H2O(l) (at cathode)

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Using Voltaic Cells as Energy Sources

Teacher DemoPURPOSE Students will observe the construction of a dry cell.

MATERIALS dry cell, voltmeter, 2 wire leads, zinc electrode, carbon electrode, manganese dioxide/ammonium chloride paste (Prepare the paste by adding saturated NH4Cl(aq) to powdered MnO2

until thick.)

PROCEDURE Before class, cut through a dry cell with a hacksaw from the top down on the right side of the central carbon electrode. In class, have students identify the three parts of the cell. (The graphite rod is the cathode, the zinc container is the anode, and the manganese dioxide/ammonium chloride paste is the electrolyte.) Construct a dry cell by placing a zinc electrode and a carbon electrode into a manganese dioxide/ammonium chloride paste. Connect the cell to a voltmeter and measure the potential.

EXPECTED OUTCOME Students should be able to identify the anode, cathode, and electrolyte in each cell.

ExplainAPPLY CONCEPTS Explain to students that, by itself, a dry cell does not provide a complete circuit, that is, electrons cannot flow from the anode to the cathode. When devices using dry cells are turned on, an external circuit is completed, allowing the flow of electrons from the anode to the cathode.

USE VISUALS Direct students to Figure 21.4. Ask students to describe the difference between the common dry cell and the alkaline battery. (In the alkaline dry cell, the electrolyte is a basic KOH paste, thus the distinction of this dry cell as alkaline.)

Check for UnderstandingHow does a voltaic cell produce electrical energy?

Assess students’ understanding of how a voltaic cell produces electrical energy by listing the following vocabulary terms on the board: voltaic cell, half-cell, salt bridge, electrode, anode, and cathode. Then, have students write a brief paragraph using the vocabulary terms that answers the key question.

ADJUST INSTRUCTION If students are having difficulty with this concept, have students reread the Voltaic Cells section. Then, draw or project Figure 21.3 on the board. Read each numbered step in the section How A Voltaic Cell Works, and have students indicate where on the figure that step takes place. Provide students with an opportunity to revise their paragraphs.

Lead grid filledlead(IV) oxide (cathode)

Sulfuric acid(H2SO4 (aq))electrolyte

Lead grid filled withspongy lead (Pb) (anode)

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Figure 21.5 Lead Storage BatteryOne cell of a 12-V lead storage battery is illustrated here. Current is produced when lead at the anode and lead(IV) oxide at the cathode are both converted to lead(II) sulfate. These processes decrease the sulfuric acid concentration in the battery. Reversing the reaction recharges the battery.

In an ordinary dry cell, the graphite rod serves only as a conductor and does not undergo reduction, even though it is the cathode. The manganese in MnO2 is the species that is actually reduced. The electrical potential of this cell starts out at 1.5 V but decreases steadily during use to about 0.8 V. Dry cells of this type are not rechargeable because the cathode reaction is not reversible.

The alkaline battery, shown in Figure 21.4b on the previous page, is an improved dry cell. In the alkaline battery, the reactions are similar to those in the common dry cell, but the electrolyte is a basic KOH paste. This change in design eliminates the buildup of ammonia gas and maintains the zinc elec-trode, which corrodes more slowly under basic, or alkaline, conditions.

Lead Storage Batteries People depend on lead storage batteries to start their cars. A battery is a group of voltaic cells connected together. A 12-V car battery consists of six voltaic cells connected together. Each cell produces about 2 V and consists of lead grids, as shown in Figure 21.5. One set of grids, the anode, is packed with spongy lead. The other set, the cathode, is packed with lead(IV) oxide (PbO2). The electrolyte for both half-cells in a lead stor-age battery is sulfuric acid. Using the same electrolyte for both half-cells allows the cell to operate without a salt bridge or porous separator. The half-reactions are as follows:

When a lead storage battery discharges, it produces the electrical energy needed to start a car. The overall spontaneous redox reaction that occurs is the sum of the oxidation and reduction half-reactions.

This equation shows that lead(II) sulfate forms during discharge. The sulfate slowly builds up on the plates, and the concentration of the sulfuric acid electrolyte decreases.

Oxidation: Pb(s) SO42 (aq) PbSO4(s) 2e

Reduction: PbO2(s) 4H (aq) SO42 (aq) 2e

PbSO4(s) 2H2O(l)

Pb(s) PbO2(s) 2H2SO4(aq) 2PbSO4(s) 2H2O(l)

READING SUPPORTBuild Reading Skills: Compare and Contrast

How are alkaline batteries and lead storage batteries similar? How are they different?


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AnswersFIGURE 21.4 In both the alkaline battery and the dry cell, zinc is oxidized and MnO2 is reduced.

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Teacher DemoPURPOSE Students will observe a simple version of the lead storage battery.

MATERIALS 2 strips of lead, wooden rod, 2 connecting wires, 250-mL beaker, dilute sulfuric acid (H2SO4), 6-V DC power supply, doorbell

SAFETY Sulfuric acid is corrosive. Wear safety goggles, gloves, and a lab apron.

PROCEDURE Attach two lead strips to a wooden rod so that the strips hang vertically in a 250-mL beaker. Place the strips 4 cm apart. Pour sufficient dilute sulfuric acid into the beaker to cover two-thirds of each strip. Connect a 6-V DC power supply to the strips with wires, and charge the battery for a few minutes. Then, connect the cell to a doorbell.

EXPECTED OUTCOME Students observe that after charging the battery, the bell rings. Discuss the half-cell reactions that occur. Ask students to write the shorthand notation for the electrochemical cell.

[Pb(s) | PbSO4(s) || PbO2(s) | PbSO4(s)]

ExplainSTART A CONVERSATION Direct students’ attention to Figure 21.5. Call attention to the grids, two containing Pb and two containing PbO2. Ask What are the oxidation numbers of the metal lead (Pb) and the lead in lead(IV) oxide (PbO2)? (oxidation number of elemental Pb is 0; oxidation number of Pb in PbO2 is +4.) Ask Which is the anode and which is the cathode? (Pb is the anode and PbO2 is the cathode.) How do you know? (Pb can only lose electrons and be oxidized, and Pb4+ can gain electrons and be reduced.) Write on the board the overall reaction for the battery when it is discharging. Call attention to the products (PbSO4 and H2O) and note that Pb is oxidized to form PbSO4 and Pb4+ (in PbO2) is reduced to also form PbSO4. Ask What is the oxidation number of the lead in PbSO4? (+2; Pb is oxidized to Pb2+ and Pb4+ is reduced to Pb2+.)

Car Battery

The forerunner of today’s car battery was built by the French physicist Gaston Plant in 1859. Plant’s first battery consisted of a single cell containing two sheets of lead separated by rubber strips and immersed in a 10% sulfuric acid solution. A year later, he presented a battery consisting of nine cells enclosed in box with outside terminals to the Académie des sciences (French Academy of Sciences). Plant’s invention was significant because it provided the first means of storing electricity.

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Power supplyRecharge

Cathode( )

Anode( )

Cathode( )

Anode( )

AOxidation half-reaction at BReduction half-reaction at

AReduction half-reaction at BOxidation half-reaction at

Pb(s) SO42 (aq) PbSO4(s) 2e

PbSO4(s) 2e Pb(s) SO42 (aq)

PbO2(s) 4H (aq) SO42 (aq) 2e

PbO2(s) 4H (aq) SO42 (aq) 2e

PbSO4(s) 2H2O(l )

PbSO4(s) 2H2O(l )

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Figure 21.6 Discharge and Recharge of a Lead-Acid BatteryThe lead-acid battery in an automobile acts as a voltaic cell (top) when it supplies current to start the engine. Some of the power from the running engine is used to recharge the battery, which then acts as an electrolytic cell (bottom). You will learn more about electrolytic cells in Lesson 21.3.

The reverse reaction occurs when a lead storage battery is recharged. This reaction occurs whenever the car’s generator is working properly.

This is not a spontaneous reaction. To make the reaction proceed as written, a direct current must pass through the cell in a direction opposite that of the current flow during discharge. The processes that occur during the dis-charge and recharge of a lead-acid battery are summarized in Figure 21.6. In theory, a lead storage battery can be discharged and recharged indefinite-ly, but in practice its lifespan is limited. Small amounts of lead(II) sulfate fall from the electrodes and collect on the bottom of the cell. Eventually, the electrodes lose so much lead(II) sulfate that the recharging process is inef-fective or the cell is shorted out. The battery must then be replaced.

Fuel Cells To overcome the disadvantages associated with lead storage bat-teries, cells with renewable electrodes have been developed. Such cells, called fuel cells, are voltaic cells in which a fuel substance undergoes oxida-tion and from which electrical energy is continuously obtained. Fuel cells do not have to be recharged. They can be designed to emit no air pollutants and to operate more quietly and more cost-effectively than a conventional electri-cal generator.

Perhaps the simplest fuel cell involves the reaction of hydrogen gas and oxygen gas. The only product of the reaction is liquid water. In the hydrogen–oxygen fuel cell shown in Figure 21.7a, there are three compartments sepa-rated from one another by two electrodes. The electrodes are usually made of carbon. Oxygen (the oxidizing agent) from the air flows into the cathode compartment. Hydrogen (the fuel) flows into the anode compartment. The anode and cathode are separated by a thin membrane that allows hydrogen ions to pass through but not electrons. The membrane therefore acts as a salt bridge. Electrons from the oxidation half-reaction at the anode pass through an external circuit to enter the reduction half-reaction at the cathode.

2PbSO4(s) 2H2O(l) Pb(s) PbO2(s) 2H2SO4(aq)


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Using Voltaic Cells as Energy SourcesMAKING CONNECTIONS Point out that fuel cells were developed for space travel where lightweight, reliable power systems were needed. Explain that fuel cells differ from lead storage batteries in that they are not self-contained. Their operation depends on a steady flow of fuel and oxygen into the cell—where combustion takes place—and the flow of the combustion product out of the cell.

Explain that in the case of the hydrogen fuel cell, the product is pure water. Convey that both the electricity generated and the water produced is consumed in space flights. Tell students that fuel cells convert 75% of the available energy into electricity, in contrast with a conventional electric power plant that converts about 35 to 40% of the energy of coal to electricity.

Point out fuels other than hydrogen can also be used, for example, ammonia (NH3), hydrazine (N2H4), and methane (CH4). Students may be interested in the equations for the reactions that take place in these cells.

4NH3(g) + 3O2(g) → 2N2(g) + 6H2O(g)

2N2H4(g) + 2O2(g) → 2N2(g) + 4H2O(g)

CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

Explain that in each case, the products are gases and water vapor which are normally found in Earth’s atmosphere.

ExtendConnect to TECHNOLOGY

Divide the class into groups of three or four students. Have them research the use of fuel cells in the space shuttle and by utility companies. Ask them to prepare posters showing different types of fuel cells and their applications. Where possible, have students include the half-cell and overall reactions for each fuel cell. Have them describe the advantages gained by using a fuel cells rather than more conventional sources of electrical energy.

Check for UnderstandingThe Essential Question How is energy produced in an electrochemical process?

To assess students understanding of voltaic cells, give students one to two minutes to explain, (orally, written, or visually) how a voltaic cell produces energy.

ADJUST INSTRUCTION If students are struggling in their response, direct them to Figure 21.3. Project the figure on the board or a screen, and use various colors of highlighters to visually show each step of how a voltaic-cell works.

Anode ( ) (+) Cathode

H2(g) fuel O2(g) (from air)

Unused H2(g) Water vapor (H2O(g))

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Figure 21.7 Hydrogen–Oxygen Fuel CellThe hydrogen–oxygen fuel cell is a clean source of power. a. The membrane allows H (aq) ions produced by the oxidation of H2(g) at the anode to migrate to the cathode, where H2O(g) is formed. b. Such cells can be used to fuel vehicles.

The half-reactions in this type of hydrogen–oxygen fuel cell are as follows:

The overall reaction is the oxidation of hydrogen to form water.

Other fuels, such as methane (CH4) and ammonia (NH3), can be used in place of hydrogen. Other oxidizing agents, such as chlorine (Cl2) and ozone (O3), can be used in place of oxygen.

Since the 1960s, astronauts have used fuel cells as an energy source aboard spacecraft. Hydrogen–oxygen fuel cells with a mass of approximately 100 kg each were used in the Apollo spacecraft missions. Fuel cells are well suited for extended space missions because they offer a continuous energy source that releases no pollutants. On space shuttle missions, for example, astronauts drink the water produced by onboard hydrogen–oxygen fuel cells.

The use of fuel cells is no longer limited to space travel. Scientists and engineers have developed fuel-cell cars. These vehicles, such as the one shown in Figure 21.7b, are propelled by electric motors, which are powered by fuel cells. Fuel-cell vehicles can be fueled with pure hydrogen gas, which is stored in high-pressure tanks. However, more research and development is needed before fuel-cell vehicles will predominate the roadways. Currently, fuel cells are expensive to make, and it is difficult to store hydrogen. Nevertheless, you may soon be seeing fuel-cell cars, buses, and bicycles. You may even one day own a cellphone or laptop that is powered by a miniature fuel cell.

Oxidation: 2H2(g) 4H (aq) 4e (at anode)

Reduction: O2(g) 4H (aq) 4e 2H2O(g) (at cathode)

2H2(g) O2(g) 2H2O(g)

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ExplainConnect to ENGINEERING

Tell students that electric cars are quiet and nonpolluting—no noisy engines spew noxious gases into the atmosphere. So why have car makers not built, promoted, and sold many electric vehicles? Have interested students research the current status of electric cars in the marketplace. Have them find out what obstacles stand in the way of mass usage of these cars. Students could choose to write a report or prepare an oral presentation. (One obstacle that students may discover is the need for inexpensive, lightweight batteries for efficient storing of electrical energy to lengthen the distance of travel without recharging.)

Evaluate

Informal AssessmentHave students make a sketch of a tin/lead voltaic cell (Sn | SnSO4 || PbSO4 | Pb). Students should label the cathode and anode, and indicate the direction of electron flow.

(Tin is the anode; lead is the cathode. The electrons flow from tin to lead.) Ask students to then write the equations for the half-reactions.

[Sn(s) → Sn2+(aq) + 2e−

Pb2+(aq) + 2e− → Pb(s)]

Then, have students complete the 21.1 Lesson Check.

Reteach Emphasize that a chemical reaction can produce a flow of electrons or a flow of electrons can cause a chemical reaction to occur. Note that reduction always occurs at the cathode, and oxidation always occurs at the anode. Discuss the half-reactions for the charging process in a lead cell and compare them to the half-reactions when the cell is producing electric current.

Electroplating

Electroplating is the use of electrolysis to deposit a thin coating of metal on an object, often to enhance its appearance or value. Materials often used for electroplating include gold, silver, copper, nickel, and chromium. In an electrolytic cell, used for plating, the object to be plated is the cathode and the anode is the plating metal. The cathode in an electroplating apparatus must be able to conduct electricity. Thus, one might think that nonmetallic substances such as wood, plastic, and leather could not be electroplated. However, chromium-plated plastic automobile parts enhance the appearance of most cars. Interested students may want to research the pretreatment of plastics that makes electroplating possible, or how a base metal can be electroplated with an alloy, such as brass or bronze.

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O E 21.11. Identify What type of reaction occurs during

an electrochemical process?

2. Describe What is the source of electrical energy produced in a voltaic cell?

3. List What are three examples of technolo-gies that use electrochemical processes to supply electrical energy?

Compare 4. Which metal is more easily oxidized, lead or calcium?

Apply Concepts 5. What is the electrolyte in a lead storage battery? Write the half-reactions for such a battery.

Describe 6. Write the overall reaction that takes place in a hydrogen–oxygen fuel cell. What product(s) are formed? Describe the half-reactions in this cell.

Predict 7. What happens when a strip of copper is dipped into a solution of iron(II) sulfate?

Take It FurtherExplain 1. What was the function of the

pasteboard soaked in saltwater in Volta’s battery? Infer 2. What is one possible reason why Volta’s

“crown of cups” would generate more current than a “voltaic pile”?

Alessandro VoltaIn the late 1770s, Italian physicist Alessandro Volta (1745–1827) discovered that contact between two different metals could produce electricity. Using this knowledge, Volta began experimenting with ways to produce a steady electric current. In 1799, he built a stack of alternating zinc and copper discs, separated by pasteboard soaked in saltwater. When he connected a wire to both ends of the pile, a steady current flowed. This device, called the “voltaic pile,” was the first battery.

Volta found that different types of metals could change the amount of current produced and that he could increase the current by adding disks to the stack. Later, he improved on the pile by creating a “crown of cups”—separate cups of salt solution linked by metal straps. In 1810, Napoleon Bonaparte gave Volta the title of Count in honor of his work, and in 1881 the unit of electrical potential was named the “volt” in Volta’s honor.

THE BATTERY REINVENTED


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CHEMISTRY YOU YOYY U&&Tell students that Italian physicist Alesandro Volta had already been experimenting with static electricity when, in 1780, his friend Luigi Galvani demonstrated that a current of electricity could be generated if two different metals, copper and zinc, were placed in contact with the muscle of a frog. At that time, scientists believed that nerves acted as water pipes. Galvani was convinced that he’d discovered a new form of electricity—one that was generated by animal tissues and transmitted through the nerves as an electrical fluid. Volta later showed that animal tissue was not necessary—current flowed with only the two dissimilar metals. Galvani’s research led to the eventual discovery that nerves transmit electrical impulses. Pose the following challenge to students. Compare and contrast the electrical current generated by a nerve cell and the current generated by a voltaic cell. You may need to assist students in the following ways:

• Nerve cells contain sodium and potassium ions.• Nerve impulses are generated when sodium ions

flow into the cell and potassium ions flow out of the cell in response to a stimulus, creating an imbalance in charge across the membrane.

Lesson Check Answers 1. a redox reaction2. spontaneous redox reactions within

the cell3. fuel cells, lead storage batteries, and

dry cells4. calcium5. concentrated sulfuric acid;

Anode: Pb(s) + SO4

2–(aq) → PbSO4(s) + 2e–

Cathode:PbO2(s) + 4H+(aq) + SO4

2–(aq) + 2e– →PbSO4(s) + 2H2O(l)

6. 2H2(g) + O2(g) → 2H2O(l); water is the product. H2 is oxidized at the anode; O2 is reduced at the cathode.

7. no reaction

21 Electrochemistry Electrochem Planning Guide

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