Electrochemistry

Direct Redox Reactions

 How much for that necklace?

Gold and silver are widely used metals for making jewelry. One of the reasons these metals are employed for this purpose is that they are very unreactive. They do not react in contact with most other metals, so they are more likely to stay intact under challenging conditions. Who wants their favorite piece of jewelry to fall apart on them?

Direct Redox Reactions

When a strip of zinc metal is placed into a blue solution of copper(II) sulfate ( Figure below ), 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 due to oxidation to zinc ions. At the same time, the copper(II) ions from the solution are reduced to copper metal (see Figure below ), which 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.

Why does this reaction occur spontaneously? The activity series is a listing 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. This series also lists 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. The 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
Lithium	Li(s) → Li+(aq) + e–	Most active or most easily oxidized
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–
Gold	Au(s) → Au3+(aq) + 3e–	Least active or most difficult to oxidize

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. These types of reactions are called direct redox reactions because the electrons flow directly from the atoms of one metal to the cations of the other metal. 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.

Summary

• The activity series of metal reactivities is given.
• Parameters for spontaneous reactions between metals are described.

Practice

Questions

Watch the video at the link below and answer the following questions:

Click on the image above for more content

http://www.youtube.com/watch?v=2MawIDT5DFU

1. What happened when Mg and Zn were placed in the Pb ²⁺ solution?
2. Did the Zn strip react in the Mg ²⁺ solution?
3. How was Ag shown to be least reactive?

Review

Questions

1. What metals are high in the activity series?
2. What metals are low in the activity series?
3. Is tin easier to oxidize than magnesium?

Electrochemical Reaction

What happened to that sculpture?

Metal exposed to the outside elements will usually corrode if not protected. The corrosion process is a series of redox reactions involving the metal of the sculpture. In some situations, the metals are deliberately left unprotected so that the surface will undergo changes that may enhance the esthetic value of the work.

Electrochemical Reactions

Chemical reactions either absorb or release energy, which 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. All sorts of batteries, from those used to power a flashlight to a calculator to an automobile, rely on chemical reactions to generate electricity. Electricity is used to plate objects with decorative metals like gold or chromium. Electrochemistry is important in the transmission of nerve impulses in biological systems. Redox chemistry, the transfer of electrons, is behind all electrochemical processes.

The reaction of zinc metal with copper(II) ions 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 Cu ²⁺ ions in the area of the solution right next to the zinc strip. Electricity on the other hand, requires the passage of electrons through a conducting medium, such as a wire, in order to do work. This work could be lighting a light bulb or powering a refrigerator or heating 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 move from one place to another in between. 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.

There are three components that make up an electrochemical reaction. There must be a solution where redox reactions can occur. These reactions generally take place in water to facilitate electron and ion movement. A conductor must exist for electrons to be transferred. This conductor is usually some kind of wire so that electrons can move from one site to another. Ions also must be able to move through some form of salt bridge that facilitates ion migration.

Summary

• Electrochemistry is defined.
• A description of an electrochemical cell is given.
• Components of an electrochemical reaction are listed.

Practice

Questions

Read the material at the link below and answer the following questions:

http://bouman.chem.georgetown.edu/S02/lect25/lect25.htm

1. Spontaneous reactions occur in what type of system?
2. Nonspontaneous reactions occur in what type of system?
3. What is potential?
4. How is potential measured?

Review

Questions

1. What is an electrochemical reaction?
2. What type of chemical reaction is involved?
3. What needs to be able to move in an electrochemical reaction?

Voltaic Cells

What made it twitch?

Luigi Galvani (1737-1798) was an Italian physician and scientist who did research on nerve conduction in animals. His accidental observation of the twitching of frog legs when they were in contact with an iron scalpel while the legs hung on copper hooks led to studies on electrical conductivity in muscles and nerves. He believed that animal tissues contained an “animal electricity” similar to the natural electricity that caused lightning to form.

Voltaic Cells

A voltaic cell is an electrochemical cell that uses a spontaneous redox reaction to produce electrical energy.

The voltaic cell (see Figure above ) 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 left half-cell is a strip of zinc metal in a solution of zinc sulfate. The right half-cell is a strip of copper metal in a solution of copper(II) sulfate. 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 electrolyte solutions in which the electrodes are placed. A metal wire connects the two electrodes. A switch opens or closes the circuit. A porous membrane is placed between the two half-cells to complete the circuit.

The various electrochemical processes that occur 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 due to the loss of zinc metal. The zinc ion concentration in the half-cell increases. Because of the production of electrons 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 membrane to maintain electrical neutrality in the cell. In the cell illustrated above, sulfate ions will move from the copper side to the zinc side to compensate for the decrease in Cu ²⁺ and the increase in Zn ²⁺ .

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

Summary

• The structure of a voltaic cell is described.
• The reactions producing electron flow are given.

Practice

Questions

Read the material at the link below and answer the following questions:

Electrochemical Cells

1. What is the difference between an electrolytic cell and a voltaic cell?
2. Where does the oxidation reaction take place in a voltaic cell?
3. Where does the reduction reaction take place?
4. List some examples of voltaic cells that are of commercial importance.

Review

Questions

1. What does a voltaic cell do?
2. Why are the two electrodes physically separated?
3. What is the purpose of the porous membrane?

Electrical Potential

How many volts is that?

The voltmeter doesn’t measure volts directly; it measures electric current flow. But don’t worry – current flow and voltage can be directly related to one another. The first meters were called galvanometers and they used basic laws of electricity to determine voltage. They were heavy and hard to work with, but got the job done. The first multimeters were developed in the 1920s, but true portability had to wait until printed circuits and transistors replaced the cumbersome wires and vacuum tubes.

Electrical Potential

Electrical potential is a measurement of the ability of a voltaic cell to produce an electric current. Electrical potential is typically measured in volts (V). The voltage that is produced by a given voltaic cell is the electrical potential difference between the two half-cells. It is not possible to measure the electrical potential of an isolated half-cell. For example, if only a zinc half-cell were constructed, no complete redox reaction can occur and so no electrical potential can be measured. It is only when another half-cell is combined with the zinc half-cell that an electrical potential difference, or voltage, can be measured.

The electrical potential of a cell results from a competition for electrons. In a zinc-copper voltaic cell, it is the copper(II) ions that will be reduced to copper metal. That is because the Cu ²⁺ ions have a greater attraction for electrons than the Zn ²⁺ ions in the other half-cell. Instead, the zinc metal is oxidized. The reduction potential is a measure of the tendency of a given half-reaction to occur as a reduction in an electrochemical cell. In a given voltaic cell, the half-cell that has the greater reduction potential is the one in which reduction will occur. In the half-cell with the lower reduction potential, oxidation will occur. The cell potential (E _cell )  is the difference in reduction potential between the two half-cells in an electrochemical cell.

Summary

• Definitions for type of electrical potential are given.

Practice

Questions

Read the material at the link below and answer the following questions:

http://www.nobelprize.org/nobel_prizes/chemistry/laureates/1920/nernst-bio.html

1. Where was Nernst born?
2. What theory did he develop in 1889?
3. What musical instrument did he develop that musicians did not like?

Review

Questions

1. Why can’t we measure the electrical potential of an isolated half-cell?
2. What does the reduction potential tell us?
3. What is the cell potential?

Standard Hydrogen Electrode

What is a standard?

We all compare ourselves to someone. Can I run faster than you? Am I taller than my dad? These are relative comparisons that don’t give a lot of useful data. When we use a standard for our comparisons, everybody can tell how one thing compares to another. One meter is the same distance everywhere in the world, so a 100 meter track in one country is exactly the same distance as a 100 meter track in another country. We now have a universal basis for comparison.

Standard Hydrogen Electrode

The activity series allows us to predict the relative reactivities of different materials when used in oxidation-reduction processes. We also know we can create electric current by a combination of chemical processes. But how do we predict the expected amount of current that will flow through the system? We measure this flow as voltage (an electromotive force or potential difference).

In order to do this, we need some way of comparing the extent of electron flow in the various chemical systems. The best way to do this is to have a baseline that we use – a standard that everything can be measured against. For determination of half-reaction current flows and voltages, we use the standard hydrogen electrode . The Figure below illustrates this electrode. A platinum wire conducts the electricity through the circuit. The wire is immersed in a 1.0 M strong acid solution and H ₂ gas is bubbled in at a pressure of one atmosphere and a temperature of 25°C. The half-reaction at this electrode is .

Under these conditions, the potential for the hydrogen reduction is defined as exactly zero. We call this , the standard reduction potential.

We can then use this system to measure the potentials of other electrodes in the half-cell. A metal and one of its salts (sulfate is often used) is in the second half-cell. We will use zinc as our example (see Figure below ).

As we observe the reaction, we notice that the mass of solid zinc decreases during the course of the reaction. This suggests that the reaction occurring in that half-cell is

So, we have the following process occurring in the cell:

and the measured cell voltage is 0.76 volts (abbreviated as v).

We define the standard emf (electromotive force) of the cell as:

We can do the same determination with a copper cell ( Figure below ).

As we run the reaction, we see that the mass of the copper increases, so we write the half-reaction:

This makes the copper electrode the cathode. We now have the two half-reactions:

and we determine the  for the system to be 0.34 v.

Again, 

Now we want to build a system in which both zinc and copper are involved. We know from the activity series that zinc will be oxidized and cooper reduced, so we can use the values at hand:

Summary

• The structure of the standard hydrogen electrode is described.
• Examples of using this electrode to determine reduction potentials are given.

Practice

Questions

Watch the video at the link below and answer the following questions:

Click on the image above for more content

http://www.youtube.com/watch?v=mrOm6xZip6k

1. Why does a cation move through the salt bridge to the hydrogen side?
2. Why is the zinc half-cell the anode?
3. How is the standard potential defined?

Review

Questions

1. What is the defined potential of the hydrogen electrode?
2. What is the chemical composition of this electrode?
3. What are the standard conditions for the other half-cell?

Calculating Standard Cell Potentials

Keeping rust away

When exposed to moisture, steel will begin to rust fairly quickly. This creates a significant problem for items like nails that are exposed to the atmosphere. The nails can be protected by coated them with zinc metal, making a galvanized nail. The zinc is more likely to oxidize than the iron in the steel, so it prevents rust from developing on the nail.

Calculating Standard Cell Potentials

In order to function, any electrochemical cell must consist of two half-cells. The Table below can be used to determine the reactions that will occur and the standard cell potential for any combination of two half-cells without actually constructing the cell. The half-cell with the higher reduction potential according to the table will undergo reduction within the cell. The half-cell with the lower reduction potential will undergo oxidation within the cell. If those specifications are followed, the overall cell potential will be a positive value. The cell potential must be positive in order for redox reaction of the cell to be spontaneous. If a negative cell potential were to be calculated, that reaction would be spontaneous in the reverse direction.

Standard Reduction Potentials at 25°C
Half Reaction	Eo (V)
F2 + 2e– → 2F−	+2.87
PbO2 + 4H++ SO42− + 2e– → PbSO4 + 2H2O	+1.70
MnO4− + 8H++ 5e– → Mn2+ + 4H2O	+1.51
Au3+ + 3e– → Au	+1.50
Cl2 + 2e– → 2Cl−	+1.36
Cr2O72− + 14H++ 6e– → 2Cr3+ + 7H2O	+1.33
O2 + 4H++ 4e– → 2H2O	+1.23
Br2 + 2e– → 2Br−	+1.07
NO3− + 4H++ 3e– → NO + 2H2O	+0.96
2Hg2+ + 2e– → Hg2 2+	+0.92
Hg2+ + 2e– → Hg	+0.85
Ag+ + e– → Ag	+0.80
Fe3+ + e– → Fe2+	+0.77
I2 + 2e– → 2I−	+0.53
Cu+ + e– → Cu	+0.52
O2 + 2H2O + 4e– → 4OH−	+0.40
Cu2+ + 2e– → Cu	+0.34
Sn4+ + 2e– → Sn2+	+0.13
2H++ 2e– → H2	0.00
Pb2+ + 2e– → Pb	−0.13
Sn2+ + 2e– → Sn	−0.14
Ni2+ + 2e– → Ni	−0.25
Co2+ + 2e– → Co	−0.28
PbSO4 + 2e– → Pb + SO42−	−0.31
Cd2+ + 2e– → Cd	−0.40
Fe2+ + 2e– → Fe	−0.44
Cr3+ + 3e– → Cr	−0.74
Zn2+ + 2e– → Zn	−0.76
2H2O + 2e– → H2 + 2OH−	−0.83
Mn2+ + 2e– → Mn	−1.18
Al3+ + 3e– → Al	−1.66
Be2+ + 2e– → Be	−1.70
Mg2+ + 2e– → Mg	−2.37
Na+ + e– → Na	−2.71
Ca2+ + 2e– → Ca	−2.87
Sr2+ + 2e– → Sr	−2.89
Ba2+ + 2e– → Ba	−2.90
Rb+ + e– → Rb	−2.92
K+ + e– → K	−2.92
Cs+ + e– → Cs	−2.92
Li+ + e– → Li	−3.05

Sample Problem: Calculating Standard Cell Potentials

Calculate the standard cell potential of a voltaic cell that uses the Ag/Ag ⁺ and Sn/Sn ²⁺ half-cell reactions. Write the balanced equation for the overall cell reaction that occurs. Identify the anode and the cathode.

Step 1: List the known values and plan the problem.

Known

• 
• 

Unknown

• 

The silver half-cell will undergo reduction because its standard reduction potential is higher. The tin half-cell will undergo oxidation. The overall cell potential can be calculated by using the equation .

Step 2: Solve.

Before adding the two reactions together, the number of electrons lost in the oxidation must equal the number of electrons gained in the reduction. The silver half-cell reaction must be multiplied by two. After doing that and adding to the tin half-cell reaction, the overall equation is obtained.

The cell potential is calculated.

Step 3: Think about your result.

The standard cell potential is positive, so the reaction is spontaneous as written. Tin is oxidized at the anode, while silver ion is reduced at the cathode. Note that the voltage for the silver ion reduction is not doubled even though the reduction half-reaction had to be doubled to balance the overall redox equation.

Oxidizing and Reducing Agents

A substance which is capable of being reduced very easily is a strong oxidizing agent. Conversely, a substance which is capable of being oxidized very easily is a strong reducing agent. According to the standard cell potential table, fluorine (F ₂ ) is the strongest oxidizing agent. It will oxidize any substance below on the table. For example, fluorine will oxidize gold metal according to the following reaction.

Lithium metal (Li) is the strongest reducing agent. It is capable of reducing any substance above on the table. For example, lithium will reduce water according to this reaction.

Using the Table above will allow you to predict whether reactions will occur or not. For example, nickel metal is capable of reducing copper(II) ions, but is not capable of reducing zinc ions. This is because nickel (Ni) is below Cu ²⁺ , but is above Zn ²⁺ in the table.

Summary

• Standard cell potential calculations are described.
• Guidelines for making predictions of reaction possibilities using standard cell potentials are given.

Practice

Read the material at the link below and answer the questions at the end:

http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Voltaic_Cells/The_Cell_Potential#Problems

Review

Questions

1. What type of reaction will the half-cell with the higher reduction potential undergo?
2. What sign must the overall cell potential be in order for a reaction to be spontaneous?
3. Is Zn ²⁺ a stronger or weaker reducing agent than Mg ²⁺ ?

Batteries

Ouch, that hurts

Alessandro Volta developed the first “voltaic cell” in 1800 (pictured above). This battery consisted of alternating disks of zinc and silver with pieces of cardboard soaked in brine separating the disks. Since there were no voltmeters at the time (and no idea that the electric current was due to electron flow), Volta had to rely on another measure of battery strength: the amount of shock produced (it’s never a good idea to test things on yourself). He found that the intensity of the shock increased with the number of metal plates in the system. Devices with twenty plates produced a shock that was quite painful. It’s a good thing we have voltmeters today to measure electric current instead of the “stick your finger on this and tell me what you feel” method.

Batteries

Two variations on the basic voltaic cell are the dry cell and the lead storage battery.

Dry Cells

Many common batteries, such as those used in a flashlight or remote control, are voltaic dry cells. These batteries are called dry cells because the electrolyte is a paste. They are relatively inexpensive, but do not last a long time and are not rechargeable.

In the zinc-carbon dry cell, the anode is a zinc container, while the cathode is a carbon rod through the center of the cell. The paste is made of manganese(IV) oxide (MnO ₂ ), ammonium chloride (NH ₄ Cl), and zinc chloride (ZnCl ₂ ) in water. The half-reactions for this dry cell are:

Anode (oxidation):

Cathode (reduction):

The paste prevents the contents of the dry cell from freely mixing, so a salt bridge is not needed. The carbon rod is a conductor only and does not undergo reduction. The voltage produced by a fresh dry cell is 1.5 V, but decreases during use.

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.

Anode (oxidation):

Advantages of the alkaline battery are that it has a longer shelf life and the voltage does not decrease during use.

Lead Storage Batteries

A battery is a group of electrochemical cells combined together as a source of direct electric current at a constant voltage. Dry cells are not true batteries since they are 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 (PbO ₂ ) packed on a metal plate.

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

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

Unlike a dry cell, the lead storage battery is 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.

This reaction 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.

Summary

• Construction of a dry cell and a battery are given.
• Chemical reactions for both types are described.

Practice

Questions

Read the material at the link below and answer the following questions:

http://www.fueleconomy.gov/feg/fuelcell.shtml

also click on the fuel cell stack link highlighted on the page.

1. Where does hydrogen enter the fuel cell?
2. How are electrons produced?
3. Where do the electrons go?
4. What is the product of the fuel cell reaction?

Review

Questions

1. What purpose does the carbon rod serve in a dry cell?
2. Where does an alkaline battery get its name?
3. Why is recharging a car battery not 100% efficient?

Electrolytic Cells

Do we have heat yet?

In 1989, two scientists announced that they had achieved “cold fusion”, the process of fusing together elements at essentially room temperature to achieve energy production. The hypothesis was that the fusion would produce more energy than was required to cause the process to occur. Their process involved the electrolysis of heavy water (water molecules containing some deuterium instead of normal hydrogen) on a palladium electrode. The experiments could not be reproduced and their scientific reputations were pretty well shot. However, in more recent years, both industry and government researchers are taking another look at this process. The device illustrated above is part of a government project, and NASA is completing some studies on the topic as well. Cold fusion may not be so “cold” after all.

Electrolytic Cells

A voltaic cell uses a spontaneous redox reaction to generate an electric current. It is also possible to do the opposite. When an external source of direct current is applied to an electrochemical cell, a reaction that is normally nonspontaneous can be made to proceed. Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur. Electrolysis is responsible for the appearance of many everyday objects such as gold-plated or silver-plated jewelry and chrome-plated car bumpers.

An electrolytic cell is the apparatus used for carrying out an electrolysis reaction. In an electrolytic cell, electric current is applied to provide a source of electrons for driving the reaction in a nonspontaneous direction. In a voltaic cell, the reaction goes in a direction that releases electrons spontaneously. In an electrolytic cell, the input of electrons from an external source forces the reaction to go in the opposite direction.

The spontaneous direction for the reaction between Zn and Cu is for the Zn metal to be oxidized to Zn ²⁺ ions, while the Cu ²⁺ ions are reduced to Cu metal. This makes the zinc electrode the anode and the copper electrode the cathode. When the same half-cells are connected to a battery via the external wire, the reaction is forced to run in the opposite direction. The zinc electrode is now the cathode and the copper electrode is the anode.

The standard cell potential is negative, indicating a nonspontaneous reaction. The battery must be capable of delivering at least 1.10 V of direct current in order for the reaction to occur. Another difference between a voltaic cell and an electrolytic cell is the signs of the electrodes. In a voltaic cell, the anode is negative and the cathode is positive. In an electrolytic cell, the anode is positive because it is connected to the positive terminal of the battery. The cathode is therefore negative. Electrons still flow through the cell form the anode to the cathode.

Summary

• The function of an electrolytic cell is described.
• Reactions illustrating electrolysis are given.

Practice

Questions

Watch the video at the link below and answer the following questions:

Click on the image above for more content

http://www.youtube.com/watch?v=y4yYF8gSHdA

1. What was the source of electricity?
2. What was the purpose of the steel attached to an electrode?
3. What is used to help carry the electric current?

Review

Questions

1. What would be the products of a spontaneous reaction between Zn/Zn ²⁺ and Cu/Cu ²⁺ ?
2. How do we know that the reaction forming Cu ²⁺ is not spontaneous?
3. What would be the voltage for the reaction where Zn metal forms Zn ²⁺ ?

Electrolysis of Water

More energy from the sun?

With fossil fuels becoming more expensive and less available, scientists are looking for other energy sources. Hydrogen has long been considered an ideal source, since it does not pollute when it burns. The problem has been finding ways to generate hydrogen economically. One new approach that is being studied is photoelectrolysis – the generation of electricity using photovoltaic cells to split water molecules. This technique is still in the research stage, but appears to be a very promising source of power in the future.

Electrolysis of Water

The electrolysis of water produces hydrogen and oxygen gases. The electrolytic cell consists of a pair of platinum electrodes immersed in water to which a small amount of an electrolyte such as H ₂ SO ₄ has been added. The electrolyte is necessary because pure water will not carry enough charge due to the lack of ions. At the anode, water is oxidized to oxygen gas and hydrogen ions. At the cathode, water is reduced to hydrogen gas and hydroxide ions.

oxidation (anode)	[latex]2\text{H}_2\text{O}(l)\to\text{O}_2(g)+4\text{H}^+(aq)+4e^-[/latex]	[latex]E^0=-1.23\text{ V}[/latex]
reduction (cathode)	[latex]2\text{H}_2\text{O}(l)+2e^-\to\text{H}_2(g)+2\text{OH}^-(aq)[/latex]	[latex]E^0=-0.83\text{ V}[/latex]
overall reaction	[latex]2\text{H}_2\text{O}(l)\to\text{O}(g)+2\text{H}_2(g)[/latex]	[latex]E^0_{\text{cell}}=-2.06\text{ V}[/latex]

In order to obtain the overall reaction, the reduction half-reaction was multiplied by two to equalize the electrons. The hydrogen ion and hydroxide ions produced in each reaction combine to form water. The H₂ SO₄ is not consumed in the reaction.

Summary

• The electrolysis of water is described.

Practice

Questions

Watch the video at the link below and answer the following questions:

Click on the image above for more content

www.youtube.com/watch?v=HQ9Fhd7P_HA

1. What are the electrodes?
2. What is the power source?
3. What is put in the water to facilitate flow of electricity?
4. Which test tube contains hydrogen gas?

Review

Questions

1. What are the electrodes used in the reaction?
2. Why is sulfuric acid used?
3. At which electrode does oxygen appear?

Electrolysis of Molten Salts and Electrolysis of Brine

A big electric bill

Production of NaOH is an important industrial process. Three different methods are employed, all of which involve the use of electricity. When calculating the price of sodium hydroxide a company has to charge in order to make a profit, the cost of electricity has to be factored in. To make a metric ton of NaOH, between 3300-5000 kWh (kilowatt hours) are required. Compare that with the power needed to run an average house. You could power a home for 6-10 months with the same amount of electricity.

Electrolysis of Molten Sodium Chloride

Molten (liquid) sodium chloride can be electrolyzed to produce sodium metal and chlorine gas. The electrolytic cell used in the process is called a Down’s cell (see Figure below ).

In a Down’s cell, the liquid sodium ions are reduced at the cathode to liquid sodium metal. At the anode, liquid chloride ions are oxidized to chlorine gas. The reactions and cell potentials are shown below:

The battery must supply over 4 volts to carry out this electrolysis. This reaction is a major source of production of chlorine gas and is the only way to obtain pure sodium metal. Chlorine gas is widely used in cleaning, disinfecting, and in swimming pools.

Electrolysis of Aqueous Sodium Chloride

It may be logical to assume that the electrolysis of aqueous sodium chloride, called brine , would yield the same result through the same reactions as the process in molten NaCl. However, the reduction reaction that occurs at the cathode does not produce sodium metal because the water is reduced instead. This is because the reduction potential for water is only -0.83 V compared to -2.71 V for the reduction of sodium ions. This makes the reduction of water preferable because its reduction potential is less negative. Chlorine gas is still produced at the anode, just as in the electrolysis of molten NaCl.

Since hydroxide ions are also a product of the net reaction, the important chemical sodium hydroxide (NaOH) is obtained from evaporation of the aqueous solution at the end of the hydrolysis.

Summary

• The reactions involving the electrolysis of molten NaCl are described.
• The reactions involving the electrolysis of brine are described.

Practice

Questions

Read the material at the link below and answer the following questions:

http://www.citycollegiate.com/sblock1.htm

1. How is sodium removed from the cell?
2. Why is CaCl ₂ added to the system?
3. Why doesn’t metallic calcium contaminate the sodium production?

Review

Questions

1. What are the products of the electrolysis of molten NaCl?
2. What are the products of the electrolysis of aqueous NaCl?
3. What spectator ion is not shown in the overall equation for the electrolysis of aqueous NaCl?

Electroplating

Does anybody know where we are?

The astrolabe (pictured above disassembled) was a device used to study the motions of planets and to do surveying. Most astrolabes were made of brass, but this one has been overlaid with gold which is wearing off. Persian mystics also used astrolabes for following the stars and making astrological predictions.

Electroplating

Many decorative objects like jewelry are manufactured with the aid of an electrolytic process. Electroplating is a process in which a metal ion is reduced in an electrolytic cell and the solid metal is deposited onto a surface. The Figure below shows a cell in which copper metal is to be plated onto a second metal.

The cell consists of a solution of copper sulfate and a strip of copper which acts as the anode. The metal (Me) is the cathode. The anode is connected to the positive electrode of a battery, while the metal is connected to the negative electrode.

When the circuit is closed, copper metal from the anode is oxidized, allowing copper ions to enter the solution.

Meanwhile copper ions from the solution are reduced to copper metal on the surface of the cathode (the second metal):

The concentration of copper ions in the solution is effectively constant. This is because the electroplating process transfers metal from the anode to the cathode of the cell. Other metals commonly plated onto objects include chromium, gold, silver, and platinum.

Summary

• The process of electroplating is described.

Practice

Questions

Watch the video at the link below and answer the following questions:

Click on the image above for more content

http://www.youtube.com/watch?v=Q8Xo43sfLgY

1. What is the solution used?
2. How did he test the system?
3. Why are batteries better than wall current for the electrical current?
4. What was the anode?

Review

Questions

1. In an electroplating process using copper, what is the anode?
2. What supplies the electric current?
3. What other metals can be coated onto objects?