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''' What are half cells? '''

Half cells delve into a branch of chemistry known as “Electrochemistry”. Electrochemical reactions involve oxidation-reduction reactions in which an electron is transferred from one chemical species to another. The simplest type of “Half Cell” reaction is a simple voltaic cell. In the figure above, a simple voltaic cell is demonstrated. A zinc strip (labeled - on the left) is contained in an aqueous solution of ZnSO4 (the blue solution), and the Cu strip (on the right) is contained in aqueous CuSO4. In the figure, Zn and Cu are electrodes. It is important to realize that a half cell is technically theoretical. In the diagram, the Zinc electrode is considered the anode, which is the electrode where oxidation takes place. The Cu strip is the cathode, which is where reduction takes place.

Half cells are considered theoretical because the half cell reactions Zn → Zn2+ + 2e- and Cu2+ + 2e- → Cu cannot technically exist by themselves. As one species is oxidized, the other species is reduced. The total reaction, then, is Zn + Cu2+ → Zn2+ + Cu. Half cells are commonly abbreviated with respect to their potentials for reduction and oxidation. For instance, The Zn half cell would be abbreviated Zn/Zn2+ and the copper half cell would be Cu/Cu2+. Each half cell combines to form an electrochemical cell. An electrochemical cell is a system in which chemical reactions produce electric current, or an electrical current produces a chemical change. Electrochemical cells are abbreviated by the notation: Cathode|Anode. The above example would be abbreviated Cu|Zn. Electrochemical cells are either voltaic or electrolytic. If the redox reaction within the electrochemical cell occurs naturally and produces electrical energy, it is considered a voltaic cell.

 Reduction Potentials and examples of half cell reactions 

In order to predict the reactivity of the metals contained within the cell, one must have knowledge of electrode potentials. The energy required to remove electrons from metal atoms is called electrode potential energy.

A useful table of reduction potentials can be found at the following link: http://www.chem.wisc.edu/deptfiles/genchem/sstutorial/Text14/Tx144/tx144.html (This table assumes standard voltage conditions; 1M, 1 atm, 25 degrees Celsius)

This is a table of reduction potentials, it measures the volts necessary for atoms to gain or lose electrons. As you move up the table you have an increasing tendency for atoms to lose electrons and become positively charged ions, and as you move down the table, you have an increasing tendency for atoms to gain electrons and form atoms of the metal. If you are concerned with whether or not two chemicals will spontaneously react in an electrochemical cell, the reduction potentials are the key. In our example above, we had Zn and Cu reacting. We observe the table and note that since zinc is above copper, it will be oxidized and lose it’s electrons. Zn → Zn2+ + 2e- That means copper must be reduced and gain electrons: Cu2+ + 2e- → Cu

 Spontaneity of the reaction 

Now we are concerned with whether or not this reaction is spontaneous. The table gives the relative energies for each transformation, but it is important to note that the table is giving us the REDUCTION energy (the energy needed to reduce each molecule). The reaction   Cu2+ + 2e- → Cu   is given in the table exactly, and has an E value of .34 volts. However, the table only gives us the reduction reaction potential for Zn. It shows us:   Zn2+ + 2e- → Zn  = -.76 volts     and not     Zn → Zn2+ + 2e- Since the reverse reaction is shown, the negative sign must be changed into a positive in order to obtain the correct value for the oxidation of zinc = .76 volts. Now add the two values together: .34 + .76 =  + 1.1 Since the value is positive, it will proceed spontaneously (without the input of energy).

 Another example 

Suppose a student proposes that a copper solution immersed in silver ions can be described by the equation: Cu + 2Ag+ → Cu2+ + 2Ag     and he wonders whether or not the reaction is spontaneous. Anode = species oxidized = lost electrons =  Cu → Cu2+ + 2e- Cathode = species reduced = gained electrons =  2Ag + 2e- → 2Ag Coppers voltage must be reversed since the table only gives its reduction value and here it is being oxidized. Cu = -.34 volts From the table, silver is being reduced and it’s voltage is given as Ag = .80 volts .80 - .34 = +.46V = spontaneous.

 Applications and uses of Half-cells 

Two half-cells make up a galvanic or voltaic cell which contains an anode in one electrolyte solution and a cathode in another electrolyte solution which allows for the flow of electrons. A voltaic cells is a basic battery. As stated above two half reactions consisting of oxidation of the anode and the reduction of cathode results in the movement of electrons from the anode to the cathode. The electron movement can be harvested as an electricity, which is how a battery works.

 Lithium ion batteries 

Basic half-cell reactions are used in batteries, and power your phones, calculator, and cars. Examples are lithium ion batteries which power most rechargeable electronic devices. Lithium ion batteries utilize a lithium cobalt oxide and a special crystallized carbon compound and the lithium ions move from the carbon to the lithium cobalt oxide producing a charge ,

More complex half cells can be used in the inorganic and electrochemistry laboratories to allow for the reduction of carbon molecules for alternative fuel sources. The two electron of reduction of CO2 results in the formation of carbon monoxide and water with a standard potential of -0.10 V, and yields potential energy from the reduction. The released energy, -0.10 V, can used to activate further reactions or capture as electricity.

Equation: CO2 + 2e- + 2H+ = CO + H2O   E= -0.10 V

 Fuel Cells 

Fuel cells are an example of the use of a half cells for the production of electricity. Fuel cells are a combination of an anode and cathode in an electrolyte solution in which the conversion of chemical energy to electrical energy can be harvested as electricity or storage in a battery. A fuel cell is similar to a battery, however battery cannot be recharged and once the chemical reduction is complete the flow of electrons stop and not power is created. However, fuel cells can be refilled with new reactants which allow for the continuous flow of electrons and production of power and have greater power capacity per volume compared to batteries. Polymer electrolyte membrane (PEM) fuel cells, reduce hydrogen gas to water producing electricity. PEM fuel cells are currently being tested as alternative power sources for the automotive industry with Honda already testing an hydrogen fuel cell car, Honda FCX in select markets and it is the only fuel cell only powered car on the market.