Importance of Redox Reactions

Oxidation-reduction reactions are vital for biochemical reactions and industrial processes. The electron transfer system in cells and oxidation of glucose in the human body are examples of redox reactions. Redox reactions are used to reduce ores to obtain metals, to produce electrochemical cells, to convert ammonia into nitric acid for fertilizers, and to coat compact discs.

Oxidation-Reduction or Redox Reaction In a redox reaction the oxidation numbers of atoms are changed. Redox reactions may involve the transfer of electrons between chemical species. The reaction that occurs when In which I₂ is reduced to I^- and S₂O₃^2- (thiosulfate anion) is oxidized to S₄O₆^2- provides an example of a redox reaction:

2 S₂O₃²⁻(aq) + I₂(aq) → S₄O₆²⁻(aq) + 2 I⁻(aq)

 

Oxidation-reduction or redox reactions take place in electrochemical cells. There are two types of electrochemical cells. Spontaneous reactions occur in galvanic (voltaic) cells; nonspontaneous reactions occur in electrolytic cells. Both types of cells containelectrodes where the oxidation and reduction reactions occur. Oxidation occurs at the electrode termed the anode and reduction occurs at the electrode called the cathode.

Electrodes & Charge

The anode of an electrolytic cell is positive (cathode is negative), since the anode attracts anions from the solution. However, the anode of a galvanic cell is negatively charged, since the spontaneous oxidation at the anode is the source of the cell's electrons or negative charge. The cathode of a galvanic cell is its positive terminal. In both galvanic and electrolytic cells, oxidation takes place at the anode and electrons flow from the anode to the cathode.

Galvanic or Voltaic Cells

The redox reaction in a galvanic cell is a spontaneous reaction. For this reason, galvanic cells are commonly used as batteries. Galvanic cell reactions supply energy which is used to perform work. The energy is harnessed by situating the oxidation and reduction reactions in separate containers, joined by an apparatus that allows electrons to flow. A common galvanic cell is the Daniell cell, shown below.

One of the more useful calculations in redox reactions is the Nernst Equation. This equation allows us to calculate the electric potential of a redox reaction in "non-standard" situations. There exist tables of how much voltage, or potential, a r eaction is capable of producing or consuming. These tables, known as standard potential tables, are created by measuring potential at "standard" conditions, with a temperature of 298 degrees Kelvin (or 25 degrees Celsius, or room temperature) and with a concentration of 1.0 moles per liter for each of the products. This standard potential, or E⁰, can be corrected by a factor that includes the actual temperature of the reaction, the number of moles of electrons being transferred, and the conce ntrations of the redox reactants and products. The equation is:

 

Perhaps the best way of understanding this equation is through an example. Suppose we have this reaction:

Fe (s) + Cd²⁺ (aq) ------> Fe²⁺ (aq) + Cd (s)

 

In this reaction iron (Fe) is being oxidized to iron(II) ion, while cadmium (Cd) in aqueous solution is being reduced to cadmium solid. The question is: how does this reaction behave in "non-standard" conditions?

The first thing to answer is how does it behave in standard conditions? We need to look at the standard potential for each half-reaction, then combine them to get a net potential for the reaction. The two (2) half-reactions are:

Fe²⁺ (aq) + 2 e^- ------> Fe (s), E⁰ = -0.44 volts
Cd²⁺ (aq) +2 e^- ------> Cd (s), E⁰ = -0.40 volts

 

Notice that both half-reactions are shown as reductions -- the species gains electrons, and is changed to a new form. But in the complete reaction above, Fe is oxidized, so the half-reaction needs to be reversed. Quite simply, the potential for the half -reaction of iron is now 0.44 volts. To get the potential for the entire reaction, we add up the two (2) half-reactions to get 0.04 volts for the standard potential.

The question now is: what is the total potential (in volts) for a nonstandard reaction? Suppose again that we have the same reaction, except now we have 0.0100 M Fe²⁺ instead of the standard 1.0 M? We need to use the Nernst equation to help us calculate that value. If you go to the Redox Half-Reaction Calculator, you should notice that the reaction is selected and the appropriate values are entered into the boxes. Since we don't have any species "B" or "D", we have entered zero for their concentrations. The concentration of the solid Fe is 1.0 M (actually, concentrations of solids and solvents (liquids) don't enter into the Nernst equation, but we set them to 1.0 so that the mathematics works out). If you click on the "Evaluate" button, you should learn that the standard potential is -0.44 volts, while the nonstandard potential is -0.5 volts. If you scroll down on the calculator, you can enter 0.5 as the first half-reaction. We again change the sign since we're actually reversing the Fe reaction

Using the calculator again, we calculate the nonstandard potential of the Cd reaction. Suppose we now have a concentration of Cd²⁺ of 0.005 M, what is its potential? The calculator should return a standard potential of -0.4 and a nonstandard potential of -0.47 volts. Place this value in the box for the second half-reaction,then click on "Evaluate". You should learn that the net nonstandard potential is 0.03 volts, slightly less than the value of the net standard potential. Since this value is less than the net standard potential of 0.04 volts, there is less of a tendency for this reaction to transfer electrons from reactants to products. In other words, less iron will be oxidized and less cadmium will be reduced than at standard conditions.

Test your use of the redox calculator by calculating the net standard potential for this reaction:

2 Ag⁺ (aq, 0.80 M) + Hg (l)------> 2 Ag (s) + Hg²⁺ (aq, 0.0010M)

Answer: 0.025 volts. Since the value is positive, the reaction will work to form the products indicated. Negative values of the potential indicate that the reaction tends to stay as reactants and not form the products. The net standard potential for t his reaction is 0.01 volts -- since the nonstandard potential is higher, this reaction will form more products than the standard reaction.

Free energy and the standard potential can also be related through the following equation:

 

Where:   delta G = change in free energy n = number of moles

If a reaction is spontaneous, it will have a positive E^o, and negative delta G, and a large K value (where K is the equilibrium constant-this will be discussed more in the kinetics section).

The energy released in any spontaneous redox reaction can be used to perform electrical work using an electrochemical cell (a device where electron transfer is forced to take an external pathway instead of going directly between the reactants. Think of the reaction between zinc and copper. Instead of placing a piece of zinc directly into a solution containing copper, we can form a cell where solid pieces of zinc and copper are placed in two different solutions such as sodium nitrate. The two solids a re called electrodes. The anode is the electrode where oxidation occurs and mass is lost where as the cathode is t he electrode where reduction occurs and mass is gained. The two electrodes are connected by a circuit and the two (2) solutions are connected by a "salt bridge" which allows ions to pass through. The anions are the negative ions and they move towards the anode. The cations are the positive ions and they move towards the cathode.

 Electrolytic Cells

The redox reaction in an electrolytic cell is nonspontaneous. Electrical energy is required to induce the electrolysis reaction. An example of an electrolytic cell is shown below, in which molten NaCl is electrolyzed to form liquid sodium and chlorine gas. The sodium ions migrate toward the cathode, where they are reduced to sodium metal. Similarly, chloride ions migrate to the anode and are oxided to form chlorine gas. This type of cell is used to produce sodium and chlorine. The chlorine gas can be collected surrounding the cell. The sodium metal is less dense than the molten salt and is removed as it floats to the top of the reaction container.

A(+)
K(-)

Electrolysis is the passage of a direct electric current through an ionic substance that is either molten or dissolved in a suitable solvent, resulting in chemical reactions at the electrodes and separation of materials.

The main components required to achieve electrolysis are:

§ An electrolyte: a substance containing free ions which are the carriers of electric current in the electrolyte. If the ions are not mobile, as in a solid salt then electrolysis cannot occur.

§ A direct current (DC) supply: provides the energy necessary to create or discharge the ions in the electrolyte. Electric current is carried byelectrons in the external circuit.

§ Two electrodes: an electrical conductor which provides the physical interface between the electrical circuit providing the energy and theelectrolyte

Electrodes of metal, graphite and semiconductor material are widely used. Choice of suitable electrode depends on chemical reactivity between the electrode and electrolyte and the cost of manufacture.

Process of electrolysis

The key process of electrolysis is the interchange of atoms and ions by the removal or addition of electrons from the external circuit. The desired products of electrolysis are often in a different physical state from the electrolyte and can be removed by some physical processes. For example, in the electrolysis of brine to produce hydrogen and chlorine, the products are gaseous. These gaseous products bubble from the electrolyte and are collected.

2 NaCl + 2 H₂O → 2 NaOH + H₂ + Cl₂

 

A liquid containing mobile ions (electrolyte) is produced by:

§ Solvation or reaction of an ionic compound with a solvent (such as water) to produce mobile ions

§ An ionic compound is melted (fused) by heating

An electrical potential is applied across a pair of electrodes immersed in the electrolyte.

Each electrode attracts ions that are of the opposite charge. Positively charged ions (cations) move towards the electron-providing (negative) cathode, whereas negatively charged ions (anions) move towards the positive anode.

At the electrodes, electrons are absorbed or released by the atoms and ions. Those atoms that gain or lose electrons to become charged ions pass into the electrolyte. Those ions that gain or lose electrons to become uncharged atoms separate from the electrolyte. The formation of uncharged atoms from ions is called discharging.

The energy required to cause the ions to migrate to the electrodes, and the energy to cause the change in ionic state, is provided by the external source of electrical potential.