chapter19


 * __Chapter 19 Electrochemistry__**

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=**19.1 Redox Reactions**=

**__Definition:__** Electrochemistry uses processes such as oxidation – reduction reactions to convert chemical energy to electrical energy or vice – versa.

The reason why redox reactions are so common in this process is because in this type of reaction electrons are transferred from one substance to another. It is this flow of electrons which creates an electric current and subsequently, electrical energy. For example, consider the following reaction:

The numbers on top are called the oxidation numbers. As it can be seen, Mg loses two electrons and each H atom gains one electron.

=19.2 Galvanic Cells=

Now consider the next equation:

This occurs when a piece of zinc is placed in a CuSO4 solution. Zinc is oxidized and copper is reduced, thus there is a transfer of electrons. However, for the transfer of electrons to form an electrical current, we need to separate the reducing and oxidizing agents and connect them with a material that allows the flow of electrons between them. This creates an electrical circuit, which can then be used as a source of electricity. (See below)

The apparatus set up above is called a galvanic or voltaic cell. The copper and zinc rods are called electrodes. The anode is the electrode in which oxidation occurs, while the cathode is the electrode in which reduction occurs. The salt bridge is used to complete the circuit between the electrodes (essentially forming a battery). Another way of representing this is through the cell diagram, as shown below.

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

=19.3 Standard Reduction Potentials=

However, an interesting question arises from the previous problem. What is the voltage (V) across the battery? For standard conditions, (meaning that the solutions are 1M and gases are 1atm), this can be done through a chart like the one below. ** Standard Reduction Potentials (in volts), 298K ** For example, consider the reaction shown by the following cell diagram
 * ** Reaction ** || ** Eo **  ||
 * F2 + 2e- ---> 2F- || +2.87 ||
 * Co3+ + e- ---> Co2+ || +1.80 ||
 * PbO2 + 4H+ + SO42- + 2e- ---> PbSO4(s) + 2H2O || +1.69 ||
 * MnO4- + 8H+ + 5e- ---> Mn2+ + 4H2O || +1.49 ||
 * PbO2 + 4H+ + 2e- ---> Pb2+ + 2H2O || +1.46 ||
 * 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 ||
 * Hg2+ + 2e- ---> Hg || +0.85 ||
 * Ag+ + e- ---> Ag || +0.80 ||
 * Fe3+ + e- ---> Fe2+ || +0.77 ||
 * I2 + 2e- ---> 2I- || +0.54 ||
 * Cu+ + e- ---> Cu || +0.52 ||
 * Fe(CN)63- + e- ---> Fe(CN)64- || +0.36 ||
 * Cu2+ + 2e- ---> Cu || +0.34 ||
 * Cu2+ + e- ---> Cu+ || +0.15 ||
 * Sn4+ + 2e- ---> Sn2+ || +0.15 ||
 * 2H+ + 2e- ---> H2 || 0.00 ||
 * Fe3+ + 3e- ---> Fe || -0.04 ||
 * Pb2+ + 2e- ---> Pb || -0.13 ||
 * Sn2+ + 2e- ---> Sn || -0.14 ||
 * Ni2+ + 2e- ---> Ni || -0.25 ||
 * Co2+ + 2e- ---> Co || -0.29 ||
 * PbSO4 + 2e- ---> Pb + SO42- || -0.359 ||
 * PbI2 + 2e- ---> Pb + 2I- || -0.365 ||
 * Cr3+ + e- ---> Cr2+ || -0.40 ||
 * Cd2+ + 2e- ---> Cd || -0.40 ||
 * Fe2+ + 2e- ---> Fe || -0.41 ||
 * Cr3+ + 3e- ---> Cr || -0.74 ||
 * Zn2+ + 2e- ---> Zn || -0.76 ||
 * 2H2O + 2e- ---> H2(g) + 2OH- || -0.83 ||
 * V2+ + 2e- ---> V || -1.18 ||
 * Mn2+ + 2e- ---> Mn || -1.18 ||
 * Al3+ + 3e- ---> Al || -1.66 ||
 * Mg2+ + 2e- ---> Mg || -2.37 ||
 * Mg2+ + 2e- ---> Mg || -2.37 ||

Pt(s) | H2(1 atm) | H+ (1 M) || Cu2+ (1M) |Cu(s)

The voltage of the cell is calculated by the following:

Note that this works only when the concentrations are both 1 M.

=**19.4 Spontaneity of Redox Reactioins**= =**19.5 The Effect of Concentration on Cell Emf**=

In order to find the emf of a galvanic cell in non - standard conditions, we use the Nernst equatioin:

In the equation above, Eo is the emf of a galvanic cell under standard - state conditions (concentrations 1M). R and F are constants. T is temperature in Kelvin, and n is the moles of electrons transferred per mole of reaction. Finally Q is the ratio of the concentrations of products over reactants (becomes the equilibrium constant at equilibrium).

Note that if the E calculated from the equation above is negative, then the reaction is not spontaneous and vice - versa. Sample Problem:

=**19.6 Batteries**=

A battery is a galvanic cell, or a series of galvanic cells, that can be used as a source of direct electric current at a constant voltage. Main Types of Batteries:

1. Dry Cell Battery Overall reaction for Dry Cell Battery The diagram and the overall reaction for a dry cell battery is shown above. The anode of the battery is a zinc inner case that is in contact with a paste like mixture of manganese dioxide, ammonium chloride, carbon, and water. Its pastelike properties help prevent leakages. A carbon rod is placed in the center of the cell and serves as the cathode.

2. Mercury Battery

Overall reaction for Mercury Batter

The diagram and the overall reaction for a mercury battery is shown above. The anode is a zinc inner case (like the dry cell), surrounded by stainless steel. The anode is then put in contact with an electrolyte containing zinc oxide and mercury(II) oxide. One of its main differences with the dry cell is that it provides a more constant voltage and a considerably longer life. This makes the mercury battery ideal for medicinal and electronic industries.

3. Lead Storage Battery

Overall reaction for lead storage battery: An interesting property of lead storage batteries is their ability to deliver large amounts of current for a short period of time, making them ideal for powering the ignition system in a car engine. Also, unlike the previous two forms of batteries, this type of battery can actually be recharged.

4. Solid - State Lithium Battery

This type of battery utilizes a solid, instead of an aqueous electrolytee to connect the electrodes. The battery has the advantage of having lithium, the element with the most negative Eo value. Lithium is also light weight, so one needs a relatively small amount of Lithium to produce a considerable amount of flowing electrons.

5. Fuel Cells

Fuel cells use combustion reactions (which are a type of redox reactions) to generate electricity. The overall reaction for one type of fuel cell is shown below:

=19.7 Corrosion= Corrosion occurs when a metal deteriorates as a result of an electrochemical process. Examples of this including rusting of iron, tarnish of silver, and green patina on copper. By far the most common is the rusting of iron. It first starts with the oxidation at the surface of the metal as shown below: The Iron ion is then further oxidized by Oxygen in the following reaction: The iron(III) oxide in the products side is what is known as rust.

The most common way to prevent this is by coating the metal with paint or other similar material. However, this has to be done continuosly, since even a small scratch on the paint will cause the metal underneath to corrode. Another way to prevent the oxidation of the metal, is to put another metal on top of it, which oxidizes more readily than the metal itself. However, this other metal also has to be replaced with the passage of time.

=19.8 Electrolysis= Electrolysis is the process in which electrical energy is used to cause a non - spontaneous chemical reaction to occur. Thus, in a way it is the opposite of a galvanic cell.

The most common type of electrolysis is the electrolysis of water, shown below: The ∆G for this reaction is very high (near 475 kJ/mol). Thus, the reaction would not occur spontaneously. It would only occur when an electrical current is run through it, like in the diagram below:



=Lab= Virtual Lab

This lab gives a demonstration of how a galvanic cell works. The solutions and electrodes are selected by the user and the reaction is shown. The user is welcome to calculate the voltage of the galvanic cell that he creates and then compare it with what the virtual lab gives. []

Other labs to consider (if one can make them work) []


 * References**


 * AP Chemistry Book from New Providence High School
 * []
 * http://en.wikipedia.org/wiki/Electrochemistry
 * http://tutors4you.com/electrochemicalcell.jpg
 * http://library.tedankara.k12.tr/chemistry/vol4/Batteries/z176.jpg
 * http://wpcontent.answers.com/wikipedia/en/thumb/f/fd/Mercurybattery2.PNG/180px-Mercurybattery2.PNG
 * http://media-2.web.britannica.com/eb-media/38/238-004-5BE91850.gif
 * http://www.rsc.org/images/FEATURE-batteries-320_tcm18-86070.jpg
 * http://www.saskschools.ca/curr_content/chem30_05/graphics/6_graphics/electrolysis_water.gif