electrochemical cells and half-reactions

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section 6electrochemical cells and half-reactionsSection overviewIn this section, students will use their red LED to build a conductivity tester. They test several solutions to determine which ones conduct electricity, and therefore contain electrolytes. They will construct a zinc-copper battery and use the LED to determine the direction of the electricity flow. To conclude the section, students will connect batteries in a series to create a circuit with more voltage. This should enable them to light LEDs that require greater voltage.

Background information

electricity Conduction through a Solution

In most electrochemical cells, part of the electrical circuit runs through a solution. Electrical conductivity has two conditions that must be satisfied simultaneously:

1) There must be charged particles in the solution – anions and cations.

2) The charged particles must be able to move around.

It follows that the more charged particles there are, the better the medium will conduct electricity. Similarly, the more easily the charged particles can move, the better the medium will conduct electricity. This places a limit on the increase in concentration in many circumstances. For example, it has been found that as the concentration of sulfuric acid increases, its

conductivity decreases when it reaches very high concentrations (about 97%) because at this point the solution is so viscous that the charged particles cannot move as easily.

There are many examples of media that do not conduct electricity well. These examples can confuse students, largely because applying two conditions simultaneously can be conceptually challenging. Pure distilled water is not a very good conductor of electricity, and its electrical conductivity is certainly not measurable with the low-tech conductivity meters students construct in this activity. Pure water is an example of a medium that satisfies the second criterion listed, but does not satisfy the first one. (Actually, there are charged particles, namely H3O+ and OH-, in pure water, but they are present in such low concentrations [about 10-7 M], that electrical conductivity is not measurable in this section.)

Molecular solids do not conduct electricity because they do not satisfy either of the criteria (molecules are neutral particles). Therefore, when molecular solids are dissolved in water, the resulting solutions do not conduct electricity either, because there are no charged particles.

Crystals of salts, which are ionic solids, can also confuse students if they have learned that ionic solids are composed of charged ions arranged in lattices. This is an example of a medium that satisfies the first but not the second criterion. However, when ionic solids are dissolved in water, the ions become solvated and can move through the liquid phase. Therefore, solutions containing dissolved ions do conduct electricity.

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SECTION 6 electrochemical cells and halF-reactionsch

apter 4

electrochemical Cells

An electrochemical cell can be thought of as a way of moving electrons through a wire for an oxidation-reduction (or redox) reaction to take place. The simplest version of an electrochemical cell is represented by the following diagram:

SaltBridge

ElectronFlow

Anode Cathode+–

Each beaker represents one-half of the redox reaction. The beaker on the left, with the anode, is where the oxidation half-reaction takes place (the anode cell is defined as the location of the oxidation half-reaction), while the beaker on the right, with the cathode, is where the reduction half-reaction takes place. The circuit is complete when electrically charged particles can make a complete circuit, though the same particles do not travel around the entire circuit. (See section titled The Path of Electricity in a Battery in the Chem Talk for this section in the Student Edition.) In particular, it is important to note that removing the salt bridge, or altering it so that electrically charged particles cannot move through it (such as filling it with distilled water), breaks the circuit and stops the flow of electricity.

reduction potentials and predicting Spontaneity

Two questions that apply to all of reaction chemistry are: Can the change occur? and, How fast does it occur? In the case of electrochemistry,

the answer to the first question lies in the thermodynamic quantity ΔG, as usual. It can be shown that:

ΔG = -nFE°cell

where n = least common multiple of electrons transferred (see Background Information in Section 4)

F = the Faraday constant = 96,500 Coulombs/mole of electrons, and

E°cell = the standard cell potential, measured in volts,

and the “naught” superscript (o) on ΔGo and E°cell indicates “standard” conditions (of 298 K temperature, 1 atm pressure, and 1.0 M concentrations of all solutions). The product nF is simply the quantity of electric charge transferred from anode to cathode. The main point to keep in mind regarding this equation is that spontaneity of a particular cell setup is determined by the sign of E°cell. Specifically, a cell setup will be spontaneous (ΔG<0) when the standard cell potential is positive (E°cell > 0).

Determining the value of the standard cell potential is a matter of knowing how to use tables of standard reduction potentials and knowing the oxidation and reduction half-reactions in the system. Tables of standard reduction potentials provide voltage values for half-cells. The standard cell potential is the sum of the voltages for each half-cell. When looking up a reduction half-reaction in the tables, simply use the voltage provided in the table. When looking up an oxidation half-reaction in the tables, reverse the sign of the voltage provided, since the oxidation is the reverse of the reduction. The logic is no different from the logic applied when using the Metal Activity Series to judge whether a given redox reaction will be spontaneous. In fact, it is the order of the metals in the standard reduction potentials that gives us the Metal Activity Series.

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For example, consider the following problem: What is the spontaneous direction in which electricity will flow in the following electrochemical cell, and what voltage would register if a voltmeter is connected across the wire?

To determine the spontaneous direction, look at the half-reactions and determine which cell will be anode and which will be cathode to provide a positive value for the standard cell potential. The two half-reactions are either:

case 1: reduction Al e Al3+ + →−3 oxidation Ni Ni e2+→ + −2

or

case 2: oxidation Al Al e3+→ + −3 reduction Ni e Ni2+ + →−2

(This should look similar to the approach in Section 4.) According to tables of standard reduction potentials, the standard reduction potential for the aluminum reduction is –1.66 volts, and the standard reduction potential for the nickel reduction is –0.25 volts. Therefore, the two possibilities translate to the following potentials (sign switches when it’s an oxidation):

case 1: reduction Al e Al V3+ o+ → =−−3 1 66E . oxidation Ni Ni e V2+ o→ + =+−2 0 25E .

or

case 2: oxidation Al Al e V3+ o→ + =+−3 1 66E . reduction Ni e Ni V2+ o+ → =−−2 0 25E .

Case 1 leads to a negative sum of the half-cell potentials (E°cell = -1.66 + 0.25 = -1.41 V), and case 2 leads to a positive sum for E°cell = +1.41 V. Therefore, case 2 is the spontaneous direction, and the voltage that should register on a voltmeter is 1.41 V. The cell in spontaneous operation should look like this:

non-Standard Conditions

There are two ways to deviate from standard conditions: change the temperature, and change the concentrations of the solutions (or pressure, if it’s a gas). Laboratory conditions are very close to standard temperature (and pressure), so only changing solution concentrations will be considered.

The value of n = 6 is the least common multiple of electrons required to arrive at the balanced redox reaction. The expression for the reaction quotient is:

Q = [Al3+]2 ______

[Ni2+]3

Under standard conditions, Q = 1, because all solution concentrations are 1.0 M. Continuing with the same aluminum-nickel cell example, the value of Q will decrease (be a value less than 1.0) from standard when either the aluminum ion concentration is less than 1.0 M or when the nickel ion concentration is greater than 1.0 M.

Ni

voltage = 1.41 V

v

1.0 Msolution ofAl2(SO4)3

1.0 Msolution ofNiSO4

Al

ElectronFlow

Anode Cathode+–Salt

Bridge

Ni

voltage = ?

v

1.0 Msolution ofAl2(SO4)3

1.0 Msolution ofNiSO4

AlSalt

Bridge


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SECTION 6 electrochemical cells and halF-reactions

learninG outComeS

learninG outComeS loCation in SeCtion eViDenCe of unDerStanDinG

Determine if a substance will conduct electricity when dissolved in water.

Investigate Part A Chem Talk, Checking Up Questions 1-2 Chem to Go Questions 1-4

Students successfully complete Part A, and provide answers that match those provided in this Teacher’s Edition.

Construct a galvanic cell and explain the function of its components.

Investigate Part B Checking Up Questions 3-5 Chem Essential Questions

Students successfully complete Steps 1-8 of Part B, and can itemize the battery’s components and functions.

Describe the effect of adding cells in series.

Investigate Part B Chem to Go Questions 5-7 Inquiring Further

Students successfully complete Steps 9-12 of Part B, and can provide answers that match those in this Teacher’s Edition.

When Q is less than 1, the logarithm of it takes a negative value, and therefore the non-standard cell potential is greater than E°cell.

Generalizing from this, to increase the voltage, make the solution in the anode cell (Al3+ ions in our example) more dilute, or else make the solution in the cathode cell (Ni2+ ions in our example) more concentrated. This makes sense because the potential to react is greater when there is more drive to arrive at products (Al3+ ions) or more drive to use up reactants (Ni2+ ions). This is an illustration of a general principle called Le Châtelier’s principle, which is useful in explaining much of equilibrium chemistry.

A battery begins with concentrations in half-cells in a non-equilibrium state. The reactions run spontaneously in the direction that approaches equilibrium. As the battery runs, the voltage continually decreases. Once the concentrations in the half-cells reach equilibrium, there is no more drive for the battery to run spontaneously in only one direction. This means, in effect, that the voltage drops to zero. Batteries eventually cease to produce voltage because they “run out.” Rechargeable batteries can be connected to an electrical source and forced to run backwards (with the input of electrical energy), building up concentrations again in the half-cells to a non-equilibrium state where the half-reactions will run spontaneously again as if the battery were new.

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Section 6 materials, Chemicals, preparation, and Safety(“per Group” quantity is based on group size of 4 students)

materials and equipment Chemicals

materials (and equipment)

Quantity per Group

(4 students)

Tongue depressor 1

Resistor, 100-ohm 1

D-cell battery 1

Battery holder, D-cell 1

Red LED 1

Copper strip, 5 mm x 5 cm x 1 mm 1

Zinc strip, 5 mm x 5 cm x 1 mm 1

Steel wool or sandpaper 1

U-tube salt bridge 1

Cotton balls 4

Voltmeter 1

Graduated cylinder, 10 mL 1

Graduated cylinder, 100 mL 1

Beakers, 50 mL 2

Beakers, 100 mL 5

Wires with alligator clips on each end.

4

materials (and equipment)Quantity per Class

Scissors 1

Electrical tape, roll ½” 1

Beakers, 250 mL 7

Chemicals Quantity per Class

(24 students)

Sodium chloride solution, NaCl, 0.1 M

200 mL

Potassium nitrate solution, KNO3, 0.1 M

200 mL

Sucrose solution, C12H22O11, 0.1 M 200 mL

Fructose solution, C6H12O6 ,0.1 M 200 mL

Zinc nitrate solution, Zn(NO3)2 , 0.1 M

200 mL

Copper (II) nitrate solution, Cu(NO3)2 ,0.1 M

200 mL

Sodium nitrate solution, NaNO3, 0.1 M

200 mL


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teacher preparation

0.1 M NaCl—Dissolve 1.2 g of sodium chloride (NaCl) in 180 mL deionized water and then adjust volume to 200 mL. Store in a labeled 250-mL beaker for class use.

0.1 M KNO3—Dissolve 2.0 g of potassium nitrate (KNO3) in 180 mL of deionized water and then adjust volume to 200 mL. Store in a labeled 250-mL beaker for class use.

0.1 M Sucrose—Dissolve 6.8 g of sucrose (C12H22O11) in 180 mL deionized water and then adjust volume to 200 mL. Store in a labeled 250-mL beaker for class use.

0.1 M Fructose—Dissolve 3.6 g of fructose (C6H12O6) in 180 mL of deionized water and then adjust volume to 200 mL. Store in a labeled 250-mL beaker for class use.

1.0 M Zn(NO3)2—Dissolve 59.5 g of zinc nitrate hexahydrate (Zn(NO3)2•6H2O) in 180 mL of deionized water and then adjust volume to 200 mL. Store in a labeled 250-mL beaker for class use.

1.0 M Cu(NO3)2—Dissolve 48.4 g of copper (II) nitrate trihydrate (Cu(NO3)2•3H2O) in 180 mL of deionized water and then adjust volume to 200 mL. Store in a labeled 250-mL beaker for class use.

1.0 M NaNO3—Dissolve 17.0 g of sodium nitrate in 180 mL of deionized water and then adjust volume to 200 mL. Store in a labeled 250-mL beaker for class use.

Look at the packaging from the red LEDs. If the package indicates a forward voltage maximum that is higher than 3 volts, you will not need to use resistors. If no maximum is indicated, you can still avoid resistors if you have extra LEDs and can replace those that burn out (life spans of LEDs will vary when operated at 3 volts). If you need or prefer to use resistors, the calculation necessary to determine which resistor(s) to use is below. Look for the forward voltage value and the forward current value on the package. Assuming a 3-volt battery, calculate the resistance needed by the following equation:

Rneeded = 3-VLED _______

ILED =

(battery voltage)-(LED voltage) ____________________________

(LED current)

Resistors are only available in certain increments (10-ohm, 47-ohm, etc.). Use the resistors that you can, in series, to sum up to a resistance value near, but not more than, the resistance calculated. To save time, test this in advance.

Safety requirements

Goggles and aprons are required in the •laboratory area.

Solutions can be disposed of down the drain.•

Wash hands and arms before leaving the •laboratory area.

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meeting the needs of All studentsDifferentiated instruction

learninG iSSue referenCe auGmentation anD aCCommoDationS

Constructing a conductivity tester, fine motor skills, following a diagram

Investigate Part A, 1.

accommodations• Some students may have the ability to construct something real with their

hands from given parts and a diagram, but for others it may be daunting. One way to divide the groups for this activity is to identify those students who enjoy putting things together and place one of them in each group.

• A simple way to help students with this task would be to provide a model they can refer to at their lab tables. Using multiple models that illustrate each step would provide an even better scaffold.

• The simplest scaffold would be to construct a conductivity tester step-by-step with the class.

Organizing observations in a table

Investigate Part A, 5.

augmentation• Check students’ tables before they begin their observations to make sure they

have the space for all the information they need. Help them correct their table format before they proceed.

accommodations• List the components of the data table for students. They need to chart the

name of the solution, the chemical formula, the appearance at the beginning and end, and the results of the conductivity test. Be sure to include a line for distilled and tap water.

• Give the students a table already formatted that they can complete.

Draw pictures to represent particles in distilled water and in tap water

Investigate Part A, 6.a) ChemTalk

augmentation• Show students sample pictures of particles in solution. Ask them to discuss

which picture represents particles in distilled water, which could be particles in tap water, and which might be a third category altogether. Discuss the reasons the solutions may be different.

• Have students compare the drawings of solutions and the conditions necessary for a solution to conduct electricity with their observations in Investigate, Part A. Point out what they should have been observing according to the text. Allow them to modify their observations to improve the accuracy of their descriptions and the vocabulary used, but do not allow them to copy the observation unless their actual reaction was similar.

Multi-tasking, attention to task, sequencing

Investigate Part B

augmentation • In Step 2, students are asked to measure copper nitrate, place it into a specific-

sized beaker, label the beaker, place their copper electrode, observe what happens and record it. Students who have difficulty following directions, sequencing a series of steps and paying attention could bring chaos to this activity. The requirements for simultaneously processing all those tasks at once is significantly diminished when the activity is structured so that students can complete one task at a time and work from a model. A quiet, orderly classroom will help students maintain their focus. The activity remains quite complex in the next steps. Monitor closely to see that all groups stay on task and follow correct procedures. Built-in checks will prevent the procedure from failing.

augmentation and accommodations


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Applying previously learned concepts and skills in a new problem

Investigate Part B Chem Talk

augmentation • Many students may not remember how to write the half-reactions and

oxidation-reduction equations and may not be able to find the information in the previous section. Review and model balancing these equations with other chemicals before students get started. Make sure all students have the essential concept of how ions can be expected to react.

• Refer students to The Path of Electricity in a Battery in the Chem Talk section for a model of the equations and an explanation of the reaction. Have them compare their observations to the ones in the text.

Following directions Investigate Part B, 5.

augmentation • Model this step for students.

Defining vocabulary and connecting it to key concepts

Chem Talk, Checking Up

augmentation • How do a metal’s reactivity and its number of electrons affect its charge and

the direction in which electricity flows? How is this related to the anode and cathode of a battery? Students will need to synthesize the information in the Chem Talk section and refer back to Investigate, Part B, Step 8 where anode and cathode are first mentioned. They may need help putting this information together to understand the concepts.

Strategies for Students with limited english language proficiency


Background knowledge

What Do You Think?

Ensure that students understand the meaning of “eventually” as an adverb. Share derivatives.

Vocabulary Following directions

Investigate Check understanding and pronunciation of new words such as “distilled.” Have students work in small groups to explain what the questions are asking for. Clarify the direction in Part B, Step 1., when students are asked to “construct.” Also, clarify how students are to “record” answers.

Background knowledge Vocabulary Comprehending text

Chem Talk Have students add words that are in block print to their personal dictionaries. Refer to the word “conduct” as a derivation of “conductivity.” Students might benefit from multiple meanings of the word “conduct.” Check for understanding of signal words such as “since,” and “while.” Also, check for understanding of the word “spontaneously,” and share derivatives. The teacher may want to separate sections of the text and either share the reading aloud or have small groups read through the sections. Asking literal questions will allow the teacher to assess understanding.

Supporting detailsResearch skills

What Do You Think Now?

Refer to another use of the word “eventually.”

Comprehension Vocabulary

Chem Essential Questions, Chem to Go

Share derivatives of the word “phenomenon.” Ensure that students understand the word “resemble.” Make sure that students know the distinction between “compare,” and “contrast.” Allow students to discuss the possible answers to the questions. Have students work in groups to achieve a consensus on the possible answers.

Following directions Inquiring Further Check for understanding of the word “affect” and explain its difference from “effect.” Students can answer questions in small groups or with a partner. Clarify what is expected when they are asked to “determine” something.

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What Do you See?

The primary purpose here is to arouse students’ interest and encourage them to share their perceptions and ideas based on the illustration. You can expect all kinds of responses, ranging from “the cat and the mouse are dancing” to more interpretive observations. For example, it appears that the phonograph is being powered by something coming from two beakers. Both responses are fine and it is your task to elicit responses without judgment.

Some of the features in the illustration that students will later find significant are: a galvanic battery, a zinc strip, a copper strip, and a diagram explaining How to Make Batteries.

Some stumbling blocks to learning can include the following misconceptions:

1. electrons can flow through solutions.

Most people think of electricity as involving the flow of electrons, and for metals, this is correct. This probably comes from the root of the word electricity being electron, and both teachers and students capitalize on this when first learning about electricity. However, in solutions, it is not the electrons that are in motion. Instead, charge balance is maintained in the solution by the movement of cations and anions toward the electrodes where charge transfer takes place at the interface between an electrode and the solution. Without ions, there is little or no electrical current.

2. water is a good conductor of electricity.

Water is, in fact, a very poor conductor of electricity. The reason it is dangerous to insert a light bulb while standing in a puddle of water, or to use a hair dryer while in the bathtub, is that water is a great solvent for ionic compounds. Tap water and fresh water typically contain dissolved ions in sufficiently high concentrations to enable the solution to be electrically conductive. It is the ions in the solution that carry the charge, however, not the water itself. The water simply provides a medium through which the charged ions can move.

section 6 teaching suggestions and sample Answers

StuDentS’ prior ConCeptionS


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What Do you think?

These are broad questions that will have varied answers. They should lead to a good discussion. The purpose of these questions is to elicit students’ prior understanding and to uncover misconceptions. Correct answers are not important, only engagement. This is also an opportunity for you to gauge students’ answers as indicators of their prior knowledge and conceptions.

investigate

part a: Solutions that Conduct electricity

1.

You may not need a resistor (see Teacher Preparation section) but if in doubt, use a 100-ohm resistor in series with the other components. Instruct students to use electrical tape to hold the tester together as well as hold the connections in place.

3. the salt bridge is unnecessary.

This preconception is incorrect because a circuit must be complete in order for electricity to flow. Electricity cannot flow in an electrochemical cell unless charge balance can be maintained in the solutions. Solutions cannot build up extra positive or negative charges, due to the oxidation and reduction half-reactions taking place separately in the two half-cells. In the oxidation half-cell (or anode), a metal is oxidized and the resulting electrons move through the circuit, leaving behind positively charged metal ions in the solution.

In order to balance the overall charge, negative ions must be able to flow into the solution. These negative ions travel through the salt bridge from the reduction half-cell (or cathode). Similarly, positive ions travel through the salt bridge toward the reduction half-cell, to balance the charge in that solution.

StuDentS’ prior ConCeptionS

What Do You Think?a ChemiSt’S reSponSe

A battery uses a spontaneous oxidation-reduction reaction separated into half-cells so that electrons are not transferred directly from atom to ion but are forced to flow through a circuit which generates a potential difference (voltage). A battery eventually dies because the reactants for the reaction are consumed.

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2.a)

The LED will not light, or only very weakly. If it does light, then the water is contaminated. Repeat the test with pure water without ions.

2.b)

Make sure students understand it is necessary to test distilled water to establish that it is not the substance that is conducting electricity. The testing of the distilled water serves as a control for the testing of other substances.

3.

Students should have a good sense of which compounds are ionic in nature. They may be less familiar with the two sugars, which are organic compounds.

4.a)

The solutions will be identical in appearance clear, colorless liquids.

4.b)

Students should be able to predict that the substances most likely to create charged ions are the NaCl and KNO3, which are the ones with metal ions.

5.

The solutions that should conduct electricity are NaCl and KNO3. Solutions that should not conduct electricity are sucrose and fructose. If the sucrose and fructose solutions do conduct, either the students’ samples have been contaminated by insufficient washing of the leads between uses, or the stock solution is contaminated.

6.

Tap water should conduct electricity, because it has dissolved electrolytes in it (such as sodium fluoride and sodium chloride).

7.

All solutions can be disposed of by pouring down a drain. The beakers should be rinsed well.

teaChinG tip

Students should be able to complete Part A in a 45 minute class period. If you are running short of time, you could provide students with commercial conductivity testers. However, constructing their own device will provide them with better insight that will help them complete the challenge.

part B: making a Battery

1.

To save time, you may want to clean the electrodes before the lab. If students do it, be sure they use the steel wool on a durable or protected surface. This would be a good time to discuss why it is necessary to clean the electrodes and identify the coatings (ZnO and CuO). This section provides a review of what students learned previously about the Metal Activity Series. They should be able to predict that the Zn2+(aq)/Cu(s) reaction will not occur, but that there should be evidence of the occurrence of a reaction in the case of Cu2+(aq)/Zn(s).

NOTES


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2.

If this is the students’ first exposure to the terminology and concept of solution concentration, it might be helpful to show students a volumetric flask and talk about the preparation procedure of making a 1 molar solution.

2.a)

Nothing should happen to the copper electrode. Students may observe some air bubbles on the surface of the metal but this is not a sign of a reaction.

3.a)

Nothing should happen to the zinc electrode. If a reaction is observed, the students might have placed the zinc electrode into the copper nitrate (blue) solution.

3.b)

Since nothing happened to the electrodes, the students do not have much information available to answer this question. However, if they understand the Metal Activity Series, they may correctly infer that the two metal electrodes need to be in solutions of their own ions in order for the battery to work properly. If the zinc electrode is placed in the copper solution, the oxidation reaction and the reduction reaction will occur at the zinc electrode and no electricity will pass through the circuit. In short, if you put the zinc in the copper solution, a reaction will happen that will produce Zn2+ and Cu in the same beaker.

5.

Students should use tweezers to handle the glass wool (or the cotton ball). This serves as a barrier to slow the migration of Na+ ions. A strip of paper towel soaked in the salt solution may also serve as the salt bridge.

6.a)

No change should occur.

4.a)

Zinc atoms lose electrons and copper atoms gain electrons. The two half-reactions are:

At the anode: oxidation Zn Zn e2+→ + −2

At the cathode: reduction Cu e Cu2+ −+ →2

4.b)

Cu Zn Zn Cu2+ 2++ → + and the electrons cancel out.


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check all connections. The zinc electrode is the anode and the copper electrode is the cathode.

9.

From Part A of this Investigate, students may recall that the LED will light when the short lead on the LED is connected to the negative terminal of the battery. The Cu/Zn cell only produces about 1–1.1 volts, so the LED probably will not light using the Cu/Zn cell alone. Nor will it light using a D-cell battery, which only produces 1.5 volts. However,

7.a)

Drawings should look like the picture in their textbook without the voltmeter. The parts that should be labeled are: Cu electrode, Zn electrode, copper (II) nitrate solution, zinc nitrate solution, salt bridge.

8.a)

The voltmeter should record the voltage in the positive direction. If it appears to be negative, switch the wires attached to it. It should read 1.0–1.1 volts. If the voltage is appreciably less,

with the two connected in series, the 2.5–2.6 volts available should be enough to light the LED. In the diagram in the Student Edition, the LED is deliberately shown with two leads of equal length, so that students will have to use their knowledge of how to connect the LED.

10.a)

With the battery in operation, students should begin to see copper metal coating the Cu electrode, or dropping into the Cu2+ solution beneath the electrode. It is difficult to see the decrease in mass of the Zn electrode, however. Students should note that the Cu electrode is gaining mass as the Cu2+ions become Cu atoms. They might say that the Zn atoms are becoming Zn2+ ions and dissolving in the solution although they cannot observe that change. They could also say that the zinc electrode is not disappearing.

11.a)

The LED light should go out when the salt bridge is removed from the circuit.

11.b)

The LED light should go on again when the salt bridge is replaced.

12.a)

Students may predict that the LED will not light with distilled water in the salt bridge because distilled water does not conduct electricity. On the other hand, they may predict that the LED will light because there is a connection between the two cells.

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12.b)

When the NaNO3 (or KNO3) solution in the salt bridge is replaced with distilled water, the cell will not operate and the LED will not light up.

Chem talk

In the first section, students are introduced to the requirements for a conductive solution. The Path of Electricity in a Battery defines and analyzes the design of a battery. In the final section, students learn how the Metal Activity Series is related to the selection of materials for a battery.


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Checking up

1.

A salt solution can conduct electricity because there are electrically charged particles (the Na+ and Cl− ions) that can move through the solution. A sucrose solution cannot conduct electricity because there are no charged particles to conduct the electricity.

2.

An electrolyte is a substance that has charged particles (ions) when dissolved in water, forming a solution that conducts electricity. A non-electrolyte is a substance that dissolves in water but does not form charged particles, meaning that the solution does not conduct electricity.

3.

In the zinc-copper battery, the electrons come from the oxidation half-reaction (when Zn metal gives up two electrons to become zinc ions). The electrons go through the wire and arrive at the copper electrode, where they are taken up by Cu2+ ions in the solution through a reduction half-reaction. The copper ions then become copper metal atoms at the cathode.

4.

The anode of a battery is the terminal from which the electrons in an electrical circuit flow. It is where the oxidation half-reaction takes place.

5.

The cathode of a battery is the terminal where the reduction half-reaction takes place.

What Do you think now?

This would be a good time to have the students go back to the What Do You See? illustration and review their initial observations. Some aspects of the illustration should be more significant to them now.

how does a battery work?

Students’ answers will vary but they should make reference to two electrodes and two different metals. Electrons arise from the oxidation taking place at the anode and they are consumed by the reduction process taking place at the cathode. The difference in activity of the two metals should be mentioned as the driving force for the oxidation-reduction reaction.


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Why does a battery eventually “die”?

As the reactions proceed at each terminal of the battery, the reactants are used up. The voltage produced drops off slowly and finally reaches a level too low to operate a light or motor.

You may want to provide students with the answer found in A Chemist’s Response at the beginning of this section and open a discussion to compare that response with their answers.

NOTES

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Chem essential Questions What does it mean?maCro — Descriptions of the construction of a battery should be similar to the procedure outlined in the Investigate section.

nano — Electrons flow from the zinc anode through a wire, the LED, a second wire, and to the copper cathode where they react with Cu2+ ions in the cathode solution, forming Cu atoms. The negatively charged nitrate ions in solution migrate toward the positively charged anode to balance the change in positive ions at each electrode.

SymBoliC — The drawing should be similar to the one shown in Investigate, Part B, Step 10, in the Student Edition.

how do you know?The voltmeter and the lighting of the LED provide evidence that it is a battery.

Why do you believe?Students’ answers will vary. Suggested items might range from MP3 Players, CD Players and cameras to remote controls, calculators and flashlights.

Why should you care?Three ways to change the design of a battery system to produce more voltage:

1. Using metals that are farther apart on the activity series would increase voltage because the difference in the metals’ abilities to lose electrons would increase, resulting in greater voltage.

2. Linking batteries together in a series increases voltage.

3. Decreasing the concentration of the cathode cell and/or increasing the concentration of the anode cell would change the reaction from standard conditions and thereby generate higher voltage.

NOTES


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reflecting on the Section and the Challenge

Students should read this section for a specific, direct connection between the section and the Chapter Challenge. While students do not answer any questions in this section, it will provide them with valuable direction in the Chapter Challenge. You may want to provide some class time for students to read this paragraph silently or aloud.

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Chem to Go

1.

The solutes that were electrolytes were KNO3 and NaCl. The solutions that conducted electricity could be distinguished by the light coming on, indicating a complete circuit was formed.

2.

a) and d): KF and Na2SO4 should make solutions that would conduct electricity because they are ionic compounds and produce ions when dissolved in water.

3.a)

It would be dangerous to use an electrical appliance while bathing because tap water (bath water) conducts electricity. If you were to drop the appliance into the water, you could be electrocuted.

3.b)

If you took a bath in distilled water, you would still not be safe to operate an electrical appliance, because the salts from your body would dissolve in the bath water, making the bath water conductive.

4.a)

Yes, a sports drink will conduct electricity because it contains electrolytes.

4.b)

The experiment design should include an indication of how concentrated the electrolytes are in solution. One way to do this would be to continually dilute the sports drink solution and test the dilutions for conductivity until the conductivity falls

below a detectable level. The drink requiring the most dilutions before ceasing to be conductive would be the one with the greatest concentration of electrolytes in the original sports drink. Testing the sports drink for voltage (i.e., attaching a voltmeter to the conductivity tester) would not work, because the voltmeter would simply measure the battery’s voltage, regardless of the concentration of the solution that completes the circuit.

4.c)

You would need a control in the experiment, to prove that there is no conductivity in a solution that has no electrolytes. Distilled water should be used as the control.

5.

Students should choose a metal farther up on the activity series (i.e., more reactive) because those atoms lose electrons more easily, which would increase voltage. Tin, aluminum, and magnesium would be the realistic choices.


chapter 4


6.a)

The voltage should decrease to zero because when the salt bridge is removed, the ions cannot flow through it to keep the anode and cathode electrically neutral.

6.b)

The voltage should decrease to zero because distilled water cannot conduct the electrical current between the anode and cathode and therefore cannot maintain their electrical neutrality.

8.

c) Au e Au3+ + →−3

For reduction, the electrons will always be on the left side of an equation.

9.

b) The number lost is always equal to the number gained.

10.

Preparing for the Chapter Challenge

Students may use the information in this section to design a battery for the Chapter Challenge. Because of the difficulty in producing adequate voltage, students may elect to use a commercial battery. This will require them to somehow complete a circuit to light the LED. The first part of the section in which students build a conductivity tester may give them some good ideas about how to design such a device.

inquiring further

factors affecting voltage generated

Be sure to check students’ proposed procedures for safety, efficacy, and practicality before approving.

6.c)

The voltage should decrease to zero because electrons cannot flow out of the anode.

6.d)

The voltage should decrease because fewer atoms and ions will be left to react since they have already reacted over the preceding hour.

7.

d) concentration of ions decreases

chemical dominoes

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Section 6 – Quiz anSWerS

❶ a) Charged particles must be present and able to move around.

❷ b) Fe Fe e2+ 3+→ + −

❸ It serves as the control. Since all of the substances are dissolved in distilled water, you have to test distilled water to make sure that it is not responsible for the conductivity of a solution.

❹ a) The circuit cannot be complete without mobile charged particles that can move inside the salt bridge.

b) The voltage will decrease because the ion concentration in the cathode cell will decrease, which makes the reaction run more slowly.

SeCtion 6 – Quiz

4-6a Blackline master

1. For a solution to conduct electricity, which of the following conditions must be met?

a) Charged particles must be present and able to move around.

b) Charged particles must be present but not able to move around.

c) Charged particles cannot be present and molecules must be able to move around.

d) Charged particles cannot be present and molecules must not be able to move around.

2. Which half-reaction correctly represents oxidation?

a) 2 2Cl e Cl2

− −+ → b) Fe Fe e2+ 3+→ + − c) 2e Sn Sn4+ 2+− + → d) Na e Na1+ + →−

3. Why was it necessary to test distilled water for conductivity before you tested the solutions for conductivity? Explain why the experiment would not be valid if you didn’t test distilled water.

4. The following picture shows a typical electrochemical cell, similar to ones that you built in the activity. Refer to the picture to give explanations for the following questions.

Salt bridgeElectron flow

Anode —

Cathode+

a) Why is it necessary to fill the salt bridge with a solution containing dissolved ions?

b) If you connect a lamp in the circuit and let the electrochemical cell run for a while, the light will become dimmer. Why?


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NOTES

electrochemical cells and half-reactions

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Transcript of electrochemical cells and half-reactions