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23.1: Electrochemical Cells

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Lesson Objectives

  • Use the activity series to identify elements that are more easily oxidized than others and write oxidation and reduction half-reactions.
  • Describe the parts of a voltaic cell and explain how redox reactions are used to generate an electric current.
  • Describe the general features of a dry cell, a lead storage battery, and a fuel cell.

Lesson Vocabulary

  • anode
  • battery
  • cathode
  • electrochemical cell
  • electrochemistry
  • electrode
  • fuel cell
  • half-cell
  • salt bridge
  • voltaic cell

Check Your Understanding

Recalling Prior Knowledge

  • What is the activity series, and how is it used?
  • What are the features of oxidation and reduction half-reactions?

Batteries are used in a great many devices in the modern world. Batteries are electrochemical cells that take advantage of redox chemical reactions to generate an electric current. In this lesson, you will be introduced to electrochemistry and some of its applications in several different types of electrochemical cells.

Electrochemical Reactions

Chemical reactions either absorb or release energy, and when they are set up in certain ways, that energy can be in the form of electricity. Electrochemistry is a branch of chemistry that deals with the interconversion of chemical energy and electrical energy. Electrochemistry has many common applications in everyday life. Batteries of all sorts, including those used to power a flashlight, a calculator, or an automobile, rely on chemical reactions to generate electricity. Electricity can also be used to plate objects with decorative metals like gold or chromium. Electrochemistry is also relevant to the transmission of nerve impulses in biological systems. Redox chemistry, the transfer of electrons, is the underlying force behind all electrochemical processes.

Direct Redox Processes

When a strip of zinc metal is placed into a blue solution of copper(II) sulfate, a reaction immediately begins as the zinc strip begins to darken. If left in the solution for a longer period of time, the zinc will gradually decay as it is oxidized to zinc ions, which enter the solution. Meanwhile, the copper(II) ions from the solution are reduced to copper metal, which eventually causes the blue copper(II) sulfate solution to become colorless.

The process that occurs in this redox reaction is shown below as two separate half-reactions, which can then be combined into the full redox reaction.

 &\text{Oxidation:} && \text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2\text{e}^- \\&\text{Reduction:} && \text{Cu}^{2+}(aq) + 2\text{e}^- \rightarrow \text{Cu}(s) \\\hline&\text{Full Reaction:} && \text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)

As you know, the oxidation and reduction processes occur simultaneously. Breaking the process apart into separate oxidation and reduction half-reactions is helpful for analyzing the overall reaction.

Why does this reaction occur spontaneously? In the chapter Chemical Reactions, you learned about the activity series, which is a list of elements in descending order of reactivity. An element that is higher in the activity series is capable of displacing an element that is lower on the series in a single-replacement reaction. Now that you have learned about oxidation and reduction, we can look at the activity series in another way. It is a listing of elements in order of ease of oxidation. The elements at the top are the easiest to oxidize, while those at the bottom are the most difficult to oxidize. Table below shows the activity series together with each element’s oxidation half-reaction.

Activity Series of Metals (in Order of Reactivity)
Element Oxidation Half Reaction
Most active or most easily oxidized Lithium Li(s) → Li+(aq) + e-
Potassium K(s) → K+(aq) + e-
Barium Ba(s) → Ba2+(aq) + 2e-
Calcium Ca(s) → Ca2+(aq) + 2e-
Sodium Na(s) → Na+(aq) + e-
Magnesium Mg(s) → Mg2+(aq) + 2e-
Aluminum Al(s) → Al3+(aq) + 3e-
Zinc Zn(s) → Zn2+(aq) + 2e-
Iron Fe(s) → Fe2+(aq) + 2e-
Nickel Ni(s) → Ni2+(aq) + 2e-
Tin Sn(s) → Sn2+(aq) + 2e-
Lead Pb(s) → Pb2+(aq) + 2e-
Hydrogen H2(g) → 2H+(aq) + 2e-
Copper Cu(s) → Cu2+(aq) + 2e-
Mercury Hg(l) → Hg2+(aq) + 2e-
Silver Ag(s) → Ag+(aq) + e-
Platinum Pt(s) → Pt2+(aq) + 2e-
Least active or most difficult to oxidize Gold Au(s) → Au3+(aq) + 3e-

Notice that zinc is listed above copper on the activity series, which means that zinc is more easily oxidized than copper. That is why copper(II) ions can act as an oxidizing agent when put into contact with zinc metal. Ions of any metal that is below zinc, such as lead or silver, would oxidize the zinc in a similar reaction. However, no reaction will occur if a strip of copper metal is placed into a solution of zinc ions, because the zinc ions are not able to oxidize the copper. In other words, such a reaction is nonspontaneous.

Cu(s) + Zn2+(aq) → NR

The reaction of zinc metal with copper(II) ions described above is called a direct redox process or reaction. The electrons that are transferred in the reaction go directly from the Zn atoms on the surface of the strip to the Cu2+ ions that are in the solution adjacent to the zinc strip. In this case, no electricity is generated. Electricity requires the passage of electrons through a conducting medium, such as a wire, in order to do work. This work could be used to light a light bulb, power a refrigerator, or heat a house. When the redox reaction is direct, those electrons cannot be made to do work. Instead, we must separate the oxidation process from the reduction process and force the electrons to travel from one place to another in order for the reaction to proceed. That is the key to the structure of the electrochemical cell. An electrochemical cell is any device that converts chemical energy into electrical energy or electrical energy into chemical energy.

Voltaic Cells

In 1800, Italian physicist Alessandro Volta (1745-1827) constructed the first electrochemical cell that was able to generate a direct current (DC). A voltaic cell is an electrochemical cell that uses a spontaneous redox reaction to produce electrical energy. There are other types of electrochemical cells that use an external source of electricity to drive an otherwise nonspontaneous reaction. You will learn about these in a later lesson. The Figure below shows a diagram of a voltaic cell.

Diagram of a voltaic cell consisting of zinc and copper half-cells.

The voltaic cell consists of two separate compartments. A half-cell is one part of a voltaic cell in which either the oxidation or reduction half-reaction takes place. The half-cell on the left consists of a strip of zinc metal immersed in a solution of zinc nitrate. The half-cell on the right consists of a strip of copper metal immersed in a solution of copper(II) nitrate. The strips of metal are called electrodes. An electrode is a conductor in a circuit that is used to carry electrons to a nonmetallic part of the circuit. The nonmetallic part of the circuit is the electrolytic solutions in which the electrodes are placed. A metal wire connects the two electrodes to one another. In the above figure, that wire is equipped with a switch to open or close the circuit and a voltmeter to measure the electrical potential generated by the cell. The half-cells are also connected by a salt bridge, the u-shaped tube in the figure. A salt bridge is a tube containing an inert electrolyte that allows the passage of ions between the two half-cells. Without the salt bridge, the voltaic cell will not function because the circuit will not be complete. The inert electrolyte in the salt bridge is often potassium chloride (KCl) or sodium nitrate (NaNO3).

The various electrochemical processes that take place in a voltaic cell occur simultaneously. It is easiest to describe them in the following steps, using the above zinc-copper cell as an example.

  1. Zinc atoms from the zinc electrode are oxidized to zinc ions. This happens because zinc is higher than copper on the activity series and so is more easily oxidized.The electrode at which oxidation occurs is called the anode. The zinc anode gradually diminishes as the cell operates, because zinc metal is being consumed by the reaction. Since zinc ions are a product of the reaction, the zinc ion concentration in the half-cell increases. Because a surplus of electrons is generated at the anode, it is labeled as the negative electrode.
  2. The electrons that are generated at the zinc anode travel through the external wire and register a reading on the voltmeter. They continue to the copper electrode.
  3. Electrons enter the copper electrode where they combine with the copper(II) ions in the solution, reducing them to copper metal.The electrode at which reduction occurs is called the cathode. The cathode gradually increases in mass because of the production of copper metal. The concentration of copper(II) ions in the half-cell solution decreases. The cathode is the positive electrode.
  4. Ions move through the salt bridge to maintain electrical neutrality in the cell. Negative ions move toward the anode to compensate for the production of positive zinc ions. Positive ions move toward the cathode to compensate for the consumption of positive copper(II) ions.

The two half-reactions can again be added together to provide the overall redox reaction occurring in the voltaic cell.

Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

Notice in the Figure above that the reading on the voltmeter is 1.10 volts (V). This will be the electrical potential (voltage) in a zinc-copper cell when the ion concentrations are both 1.0 M. You will learn how to determine cell voltages in the following lesson.

There is a simple shorthand notation used to illustrate a particular electrochemical cell. The cell notation for the zinc-copper cell is shown below.

Zn(s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu(s)

The single vertical lines represent the phase boundaries between the metal electrodes and the solutions. The double vertical line represents the salt bridge. The anode is conventionally written on the left and the cathode on the right. The molarities of the half-cell solutions are also indicated.

Types of Voltaic Cells

Several variations on the basic voltaic cell presented above are in common use today. A few of those will be described, including the dry cell, the lead storage battery, and the fuel cell.

Dry Cells

Many common batteries, such as those used in a flashlight or a remote control, are voltaic dry cells. There are several kinds of dry cells in common usage that differ by the substances that are undergoing redox reactions. Figure below shows a zinc-carbon dry cell.

A dry cell is commonly known as a battery, like those used in a flashlight. They are relatively inexpensive, but they do not last a long time and are generally not rechargeable.

These batteries are called dry cells because the electrolyte is a paste instead of an aqueous solution. In a zinc-carbon dry cell, the anode is the zinc container, while the cathode is a carbon rod through the center of the cell. The paste is made of manganese(IV) oxide (MnO2), ammonium chloride (NH4Cl), and zinc chloride (ZnCl2), plus just enough water to allow current to flow. The half-reactions for this dry cell are:

&\text{Anode (oxidation):} && \text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2\text{e}^- \\&\text{Cathode (reduction):} && \text{2MnO}_2(s) + 2\text{NH}_4^+(aq) + 2\text{e}^- \rightarrow \text{Mn}_2\text{O}_3(s) + 2\text{NH}_3(aq) + \text{H}_2\text{O}(l)

The paste prevents the contents of the dry cell from freely mixing, so a salt bridge is not needed. The carbon rod serves only as a conductor and does not participate in the actual reaction. The voltage produced by a fresh dry cell battery is 1.5 V, but this value decreases somewhat over time as the battery is used up.

An alkaline battery is a variation on the zinc-carbon dry cell. The alkaline battery has no carbon rod and uses a paste of zinc metal and potassium hydroxide instead of a solid metal anode. The cathode half-reaction is the same, but the anode half-reaction is different.

&\text{Anode (oxidation):} && \text{Zn}(s) + 2\text{OH}^-(aq) \rightarrow \text{Zn(OH)}_2(s) + 2\text{e}^- \\&\text{Cathode (reduction):} && \text{2MnO}_2(s) + 2\text{NH}_4^+(aq) + 2\text{e}^- \rightarrow \text{Mn}_2\text{O}_3(s) + 2\text{NH}_3(aq) + \text{H}_2\text{O}(l)

Alkaline batteries tend to have a longer shelf life, and its voltage does not decrease as much over time.

Lead Storage Batteries

A battery is a group of electrochemical cells combined together as a source of direct electric current at a constant voltage. Technically, dry cells are not true batteries, since they consist of only one cell. The lead storage battery is commonly used as the power source in cars and other vehicles. It consists of six identical cells joined together, each of which has a lead anode and a cathode made of lead(IV) oxide (PbO2) packed on a metal plate (Figure below).

A lead storage battery, such as those used in cars, consists of six identical electrochemical cells and is rechargeable.

The cathode and anode are both immersed in an aqueous solution of sulfuric acid, which acts as the electrolyte. The cell reactions are:

&\text{Anode (oxidation):} && \text{Pb}(s) + \text{SO}_4^{2-}(aq) \rightarrow \text{PbSO}_4(s) + 2\text{e}^- \\&\text{Cathode (reduction):} && \text{PbO}_2(s) + 4\text{H}^+(aq) + \text{SO}_4^{2-}(aq) + 2\text{e}^- \rightarrow \text{PbSO}_4(s) + 2\text{H}_2\text{O}(l) \\\hline&\text{Overall:} && \text{Pb}(s) + \text{PbO}_2(s) + 4\text{H}^+(aq) + 2\text{SO}_4^{2-}(aq) \rightarrow 2\text{PbSO}_4(s) + 2\text{H}_2\text{O}(l)

Each cell in a lead storage battery produces 2 V, so a total of 12 V is generated by the entire battery. This electrical potential is used to start a car or power a different type of electrical system.

Unlike a dry cell, the lead storage battery is designed to be rechargeable. Note that the forward redox reaction generates solid lead(II) sulfate, which slowly builds up on the plates. Additionally, the concentration of sulfuric acid decreases. When the car is running normally, its generator recharges the battery by forcing the above reactions to run in the opposite, or nonspontaneous direction.

2\text{PbSO}_4(s) + 2\text{H}_2\text{O}(l) \rightarrow \text{Pb}(s) + \text{PbO}_2(s) + 4\text{H}^+(aq) + 2\text{SO}_4^{2-}(aq)

This regenerates the lead, lead(IV) oxide, and sulfuric acid needed for the battery to function properly. Theoretically, a lead storage battery should last forever. In practice, the recharging is not 100% efficient, because some of the lead(II) sulfate falls from the electrodes and collects on the bottom of the cells.

Fuel Cells

The burning of fossil fuels to generate electricity is an inherently inefficient process, and it is harmful to the environment as well. The upper limit for a power plant is to convert about 40% of the chemical energy into electricity. Fuel cells offer an alternative approach to extracting chemical energy in a useful form. A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning. A diagram for a hydrogen-oxygen fuel cell is shown in Figure below.

A hydrogen-oxygen fuel cell is a clean source of power that generates only water as a product.

A typical hydrogen-oxygen fuel cell uses an electrolyte solution containing hot, concentrated potassium hydroxide. The inert electrodes are made of porous carbon. Hydrogen gas is fed into the anode compartment while oxygen gas is fed into the cathode compartment. The gases slowly diffuse through the electrodes, and the following reactions take place.

&\text{Anode (oxidation):} && 2\text{H}_2(g) + 4\text{OH}^-(aq) \rightarrow 4\text{H}_2\text{O}(l) + 4\text{e}^- \\&\text{Cathode (reduction):} && \text{O}_2(g) + 2\text{H}_2\text{O}(l) + 4\text{e}^- \rightarrow 4\text{OH}^-(aq) \\\hline&\text{Overall:} && 2\text{H}_2(g) + \text{O}_2(g) \rightarrow 2\text{H}_2\text{O}(l)

The standard voltage from a hydrogen-oxygen fuel cell is 1.23 V.

A number of other fuels have been developed for fuel cells, including methane, propane, and ammonia. Fuel cells are more efficient than other engines and have been used for many years on space missions. Another advantage to fuel cells is that they produce fewer pollutants, particularly in the case of the hydrogen-oxygen fuel cell, where the only product of the reaction is water. The primary drawback to current fuel cell technology is that they are very expensive to build and maintain.

Lesson Summary

  • Electrochemistry is the interconversion of chemical energy and electrical energy. Electrochemical reactions are redox reactions.
  • The elements at the top of an activity series are the most easily oxidized, while the lowest elements are the most difficult to oxidize.
  • A direct redox reaction cannot be used to generate an electric current. The oxidation and reduction half-reactions must be separated, as in a voltaic cell. Voltaic cells use spontaneous redox reactions to generate a current.
  • Dry cells, lead storage batteries, and fuel cells are three modern devices that take advantage of electrochemical reactions to produce energy.

Lesson Review Questions

Reviewing Concepts

  1. What type of reaction drives any electrochemical process?
  2. Manganese metal is more active than cadmium. Which element is more easily oxidized?
  3. What half-reaction occurs at the cathode of a voltaic cell? At the anode?
  4. What substance is oxidized in a typical dry cell? What substance is reduced?
  5. What is the function of the salt bridge of an electrochemical cell? Why should an inert electrolyte be used in the salt bridge?

Problems

  1. Predict whether a reaction will occur when the metals listed below are immersed into the given solutions. Explain.
    1. tin into nickel(II) nitrate
    2. magnesium into lead(II) nitrate
    3. lead into silver nitrate
    4. silver into copper(II) chloride
  2. For those experiments in question 6 in which a reaction occurs, write the following.
    1. the balanced molecular equation
    2. the balanced net ionic equation
    3. the oxidation and reduction half-reactions
  3. A voltaic cell is constructed with a strip of aluminum metal immersed in a 1 M solution of aluminum nitrate as one half-cell and a strip of tin metal immersed in a 1 M solution of tin(II) nitrate as the other half-cell. The half-cells are connected by a conducting wire and a salt bridge.
    1. Write the oxidation half-reaction that will occur when the cell is operating.
    2. Write the reduction half-reaction that will occur when the cell is operating.
    3. Write the balanced overall redox reaction.
  4. Depict the voltaic cell from question 8 in shorthand cell notation.
  5. What are the primary advantages and disadvantages of fuel cells compared to conventional power plants?

Further Reading / Supplemental Links

Points to Consider

The electrical potential of an electrochemical cell is the voltage that the cell produces. It is dependent on the particular oxidation and reduction reactions that take place in the cell.

  • What is the standard for measuring half-cell electrical potentials?
  • How can the electrochemical cell potential be calculated for any cell?

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Date Created:

Mar 29, 2013

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Sep 09, 2014
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CK.SCI.ENG.SE.1.Chemistry-Intermediate.23.1

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