SAT Subject Test Chemistry

PART 2

REVIEW OF MAJOR TOPICS

CHAPTER 12

Oxidation-Reduction and Electrochemistry

These skills are usually tested on the SAT Subject Test in Chemistry. You should be able to …

• Explain the reactions in a voltaic cell.

• Use electrode potentials to determine if a reaction will occur.

• Describe electroplating and how to determine the quantity of product plated.

• Explain the differences between voltaic and electrolytic cells.

• Explain how ordinary batteries work.

• Balance redox equations using oxidation numbers.

This chapter will review and strengthen these skills. Be sure to do the Practice Exercises at the end of the chapter.

In the early 1830s, Michael Faraday discovered that water solutions of certain substances conduct an electric current. He called these substances electrolytes. Our definition today of an electrolyte is much the same. It is a substance that dissolves in water to form a solution that will conduct an electric current.

The usual apparatus to test for this conductivity is a lightbulb placed in series with two prongs that are immersed in the solution tested, as shown in Figure 35.

Figure 35. Test for Conductivity of Electrolytes

Using this type of apparatus, we can classify solutions as good, moderate, or poor electrolytes. If they do not conduct at all, they are called nonelectrolytes. The table below gives the classifications of some common substances.

The reason that these substances conduct with varying degrees of efficiency is related to the number of ions in solution.

The ionic lattice substances like sodium chloride are dissociated by the water molecules so that the individual positive and negative ions are dispersed throughout the solution. In the case of a covalent bonded substance the degree of polarity determines the extent to which it will be ionized. The water molecules, which are polar themselves, can help weaken and finally break the polar covalent bonds by clustering around the substance. When the ions are formed in this manner, the process is called ionization; when the ionic lattice comes apart, the process is dissociation. Substances that are nonelectrolytes are usually bonded so that the molecule is a nonpolar molecule. The polar water molecule cannot orient itself around the molecule and cause its ionization.

Let us see how the current was carried through the apparatus in Figure 35. The electricity causes one electrode to become positively charged and one negatively charged. If the solution contains ions, they will be attracted to the electrode with the charge opposite to their own. This means that the positive ions migrate to the negative pole, and they are referred to as cations. The negative ions migrate toward the positive pole, and they are called anions. When these ions arrive at the respective electrodes, the negative ions give up electrons and the positive ions accept electrons, and we have a completed path for the electric current. The more highly ionized the substance, the more current flows and the brighter the lightbulb glows. It was the Swedish chemist Arrhenius who proposed a theory to explain the behavior of electrolytes in aqueous solutions.

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REMEMBER Cations are positively charged ions. Anions are negatively charged ions.

OXIDATION-REDUCTION AND ELECTROCHEMISTRY

The branch of chemistry that deals with electricity-related applications of oxidation-reduction reactions is called electrochemistryOxidation-reduction reactions involve a transfer of electrons from the substance oxidized to the substance reduced. If the two substances are in contact with each other, a transfer of heat energy accompanies the electron transfer. This can be shown when a zinc strip is in contact with a copper(II) sulfate solution in a beaker, as shown in the accompanying diagram. The zinc strip loses electrons to the copper(II) ions in solution. The copper(II) ions accept the electrons and fall out of solution as copper atoms. As electrons are transferred between zinc atoms and copper(II) ions, energy is released in the form of heat, as indicated by a rise in temperature.

TIP 

REMEMBER Leo the lion says Ger” stands for Loss of Electrons is Oxidation and Gain of Electrons is Reduction

Voltaic Cells (or Galvanic Cells)

In another example of an oxidation-reduction reaction, the substance that is oxidized during the reaction is separated from the substance that is reduced during the reaction. The electron transfer is accompanied by a transfer of electrical energy instead of heat. One means of separating oxidation and reduction half-reactions is with a porous barrier, which prevents the metal atoms of one half-reaction from mixing with the ions of the other half-reaction. Ions in the two solutions, however, can move through the porous barrier. Electrons can be transferred from one side to the other through an external connecting wire. Electric current moves in a closed loop path, or circuit, so this movement of electrons through the wire is balanced by the movement of ions in the solution, as shown in the figure below.

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REMEMBER: “An Ox (anode is oxidation) and a Red Cat” (reduction at cathode)

The zinc strip is in an aqueous solution of ZnSO4; the copper strip, in an aqueous solution of CuSO4. Both solutions conduct electricity, so they are classified as electrolytes. An electrode is a conductor used to establish electrical contact with a nonmetallic part of a circuit, such as an electrolyte. In the figure, zinc and copper are electrodes. A single electrode immersed in a solution of its ions is a half-cell. The zinc strip in aqueous ZnSO4 is an anode, the electrode where oxidation takes place. The copper strip in CuSO4 is a cathode, the electrode where reduction takes place. The copper half-cell can be written as Cu2+/Cu, and the zinc half-cell as Zn2+/Zn. Notice that the electrode is identified as the anode if it is the electrode at which oxidation takes place and as the cathode if it is the site at which reduction takes place. Both the anode and the cathode may be the positive or the negative electrode of a cell. These terms are not interchangeable. The determining factor in the identification is whether oxidation or reduction is occurring at the site.

The two half-cells together constitute an electrochemical cell. An electrochemical cell is a system of electrodes and electrolytes in which either chemical reactions produce electrical energy or an electric current produces chemical change. An electrochemical cell may be represented by the following notation: cathode | anode. For example, the cell made up of zinc and copper can be written as Cu | Zn. There are two types of electrochemical cells: voltaic (also called galvanic) and electrolytic.

If the redox reaction in an electrochemical cell occurs naturally and produces electrical energy as is shown in the drawing immediately above, the cell is a voltaic cell or galvanic cell. Cations in the solution are reduced when they gain electrons at the surface of the cathode to become metal atoms. This half-reaction for the voltaic cell shown in the drawing is as follows:

Cu2+(aq) + 2e → Cu(s)

The half-reaction occurring at the negative electrode or anode is:

Zn(s) → Zn2+(aq) + 2e

Electrons given up at the anode pass along the external connecting wire to the cathode.

The movement of electrons through the wire must be balanced by a movement of ions in the solution. Anions move toward the anode to replace the negatively charged electrons that are moving away. Cations move toward the cathode as positive charge is lost through reduction. Thus, sulfate ions in the CuSO4 solution can move through the barrier into the ZnSO4 solution. Likewise, the Zn2+ ions pass through the barrier into the CuSO4 solution. This porous barrier can be replaced by a “salt bridge” as shown in Figure 44.

It is important to note that if a battery is connected so that the positive terminal contacts the copper electrode and the negative terminal contacts the zinc electrode, the electrons flow in the opposite direction. The battery forces the cell to reverse its reaction; the zinc electrode becomes the cathode, and the copper electrode becomes the anode. The half-reaction at the anode, in which copper metal is oxidized, can be written as follows:

The reduction half-reaction of zinc at the cathode is written as follows:

Electrode Potentials

In the discussion of acids and metals reacting to produce hydrogen, the activities of metals compared with the activity of hydrogen are shown. A more inclusive presentation of this activity of metals, called the electromotive series, is shown in Table 10.

From this chart you notice that zinc is above copper and, therefore, more active. This means zinc can displace copper ions in a solution of copper sulfate:

Zn0(s) + Cu2+(aq) + SO42−(aq) → Cu0(s) + Zn2+(aq) + SO42−(aq)

The zinc atom must have lost 2 electrons to become Zn2+ ions:

Zn0(s) → Zn2+(aq) + 2e(electrons)

At the same time, the Cu2+ must have gained 2 electrons to become the Cu0 atom:

Cu2+(aq) + 2e (electrons) → Cu0(s)

These two equations are called half-reactions. The loss of electrons by the zinc is called oxidation; the gain of electrons by the copper ion, reduction. It is important to remember that the gain of electrons is reduction and the loss of electrons is oxidation. A way to remember this is the statement “LEO the lion says GER.” LEO translates into “Loss of Electrons is Oxidation,” and GER into “Gain of Electrons is Reduction.” The fact that the zinc in this reaction is oxidized by giving up electrons makes it possible for electrons to be gained by the copper, which is acting as a reducing agent. Similarly, because the copper in this reaction is being reduced by gaining electrons, electrons can be lost by the zinc, which is acting as an oxidizing agent.

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REMEMBER The reducing agent is oxidized. The oxidizing agent is reduced.

The metal elements that lose electrons easily and become positive ions are placed high in the electromotive series. The metal elements that lose electrons with more difficulty are placed lower in Table 10.

The energy required to remove electrons from metallic atoms can be assigned numerical values called electrode potentials. The energies, E0, required for the reduction of common elements are shown in Tables 10 and 11.

* A measure in volts of the tendency of atoms to gain or lose electrons.

These voltages depend on the nature of the reaction, the concentrations of reactants and products, and the temperature. Throughout this discussion, we use standard concentrations, that is, all ions or molecules in aqueous solution are at a concentration of 1 molar. Furthermore, all gases taking part in the reactions are at 1 atmosphere pressure, and the temperature is 25°C. The voltage measured under these conditions is called standard voltage.

The currently accepted convention is to give the potentials of half-reactions as reduction processes. For example:

2H+(aq) + 2e → H2(g)

The E0 values corresponding to these half-reactions are called standard reduction potentials.

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If the E0cell is positive, the reaction will occur.

Some reduction electrode potentials are shown in the last column of Tables 10 and 11. Notice that hydrogen is used as the standard with an electrode potential of zero. These values help you to predict what reactions will occur and how readily they will occur.

The following examples will clarify the use of electrode potentials.

If magnesium reacts with chlorine, we can write the equation

Mg(s) + Cl2(g) → MgCl2(s)

The two half-reactions with the electrode potentials are:

     

In the net reaction, E0 is a positive number. This indicates that the reaction occurs spontaneously.

TIP 

REMEMBER When the E0cell is a positive value, the reaction will occur spontaneously.

You should also note that the total number of electrons lost in oxidation is equal to the total number of electrons gained so that the net reaction (arrived at by adding the two reactions) does not contain any electrons.

For sodium reacting with chlorine the equation is:

2Na(s) + Cl2(g) → 2NaCl(s)

The two half-reactions are:

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This note is important.

Again, the E0 for this reaction is positive and the reaction is spontaneous.

The next example shows a negative E0.

Copper metal placed in an acid solution is shown as follows:

     

Since E0 is negative, we know the reaction will not take place.

Many times the reduction reactions with their E0 values are shown for metals. If you must use the oxidation reactions of these metals, the equations must be reversed, and the sign of E0 changed. For example, when a piece of copper is placed into a solution of silver ions:

     

This reaction will occur spontaneously since E0 is positive.

• SAMPLE PROBLEM:

Calculate the cell voltage of the following reaction:

Zn(s) | Zn2+(0.001 M) || Ag+(0.1 M) | Ag(s)

(Each side of the || represents the half-reaction characters, with the concentration when appropriate.)

The reactions are:

This method is based on balancing the electrons gained and the electrons lost in the two half-reactions.

Since the cell voltage is positive, the reaction will occur.

Electrolytic Cells

In the second type of electrochemical reactions, the redox reactions do not occur spontaneously but can be forced to take place by supplying energy with an external current. These reactions are called electrolytic reactions. Some examples of this type of reaction are electroplating, electrolysis of salt solution, electrolysis of water, and electrolysis of molten salts. An example of the electrolytic setup is shown in Figure 36.

Figure 36. Electrochemical Reactions

If this solution contained Cu2+ cations and Cl anions, the half-reactions would be:

Notice that the E0 for this reaction is negative, so an outside source of energy must be used to make it occur.

In electroplating, where electrolysis is used to coat a material with a layer of metal, the object to be plated is made the cathode in the reaction. A bar of the plating metal is made the anode, and the solution contains ions of the plating metal. See Figure 37 below.

Figure 37. Electroplating a Metal Fork

In this example, silver nitrate, AgNO3, can be used to silver plate. When dissolved in H2O, it forms a solution containing silver ions, Ag+. Assume a metal fork is made the cathode and a bar of silver is made the anode. When the current is switched on, the positive silver ions in the solution are attracted to the fork. When the silver ions make contact, they are reduced and change from ions to atoms of silver. These atoms gradually form a metallic coating on the fork. At the anode, oxidation occurs, and the anode itself is oxidized. These two half reactions are:

     Cathode reaction: Ag+(aq) + e → Ag(s)

     Anode reaction: Ag(s) → Ag+(aq) + e

The silver ions formed at the anode replace those in the solution that are plated onto the fork during the cathode reaction.

Another example is the electrolysis of a water solution of sodium chloride. This water solution contains chloride ions, which are attracted to the anode and set free as chlorine molecules. The cathode reaction is somewhat more complicated. Although the sodium ions are attracted to the cathode, they are not set free as atoms. Remember that water can ionize to some extent, and the electromotive series shows that the hydrogen ion is reduced more easily than the sodium ion. Therefore, hydrogen, not sodium, is set free at the cathode. The reaction can be summarized in this way:

Another example of electrolysis is the decomposition of water by using an apparatus like the one shown in Figure 38.

The solution in this apparatus contains distilled water and a small amount of H2SO4. The reason for adding H2SO4 is to make the solution an electrolyte because distilled water alone will not conduct an electric current. The solution, therefore, contains ions of H+, HSO4, and SO42−.

Figure 38. Setup for the Electrolysis of Water

The hydrogen ions (H+) migrate to the cathode where they are reduced to hydrogen atoms and form hydrogen molecules (H2) in the form of a gas. The SO42− and HSO4 migrate to the anode but are not oxidized since the oxidation of water occurs more readily. These ions then are merely spectator ions. The oxidized water reacts as shown in the following half-reaction.

     

Notice that the equation shows 2 volumes of hydrogen gas are released while only 1 volume of oxygen gas is liberated. This agrees with the discussion of the composition of water in Chapter 7.

There are two important differences between the voltaic cell and the electrolytic cell:

1. The anode and cathode of an electrolytic cell are connected to a battery or other direct-current source, whereas a voltaic cell serves as a source of electrical energy.

2. In electrolytic cells, electrical energy from an external source causes nonspontaneous redox reactions to occur. In voltaic cells, however, spontaneous redox reactions produce electricity. Thus, in an electrolytic cell, electrical energy is converted into chemical energy; in a voltaic cell, chemical energy is converted into electrical energy.

TIP 

Know the differences between the voltaic cell and the electrolytic cell.

Applications of Electrochemical Cells (Commercial Voltaic Cells)

One of the most common voltaic cells is the ordinary “dry cell” used in flashlights. Its makeup is shown in the drawing below, along with anode, cathode, and paste reactions.

The automobile lead storage battery is also a voltaic cell. When it discharges, the reactions are as shown in the drawing below.