9.2.D - Electrochemistry


The term oxidation was originally used to describe reactions in which an element combines with oxygen. After electrons were discovered, chemists extended the definition of oxidationto involve the transfer of electrons from one atom to another. Because electrons cannot be created or destroyed, the species losing electrons was said to be oxides and the species gaining electrons was said to be reduced.

If the electrons from an oxidation reaction can be made to flow around a circuit to the site of reduction, we have a galvanic cell or battery. These cells convert chemical energy to electrical energy and rely on oxidation-reduction reactions.

In this unit you will learn to:

  • Recognise oxidation-reduction reactions
  • Assigning oxidation states to species to determine if oxidation or reducation has occurred
  • Use metal displacement reactions to put metals into an order or activity
  • Describe and construct galvanic cells and calcaulte their theoretical cell potentials
  • Compare different types of galvanic cells


Oxidation-Reduction Reactions

Redox reactions, or oxidation-reduction reactions are a class of reactions that involve the transfer of electrons from one chemical species to another. If oxidation in occuring in one place, reduction will be occuring at the same time. This is because oxidation refers to the loss of electrons, while reduction refers to the gain of those same electrons. Each reaction by itself is described as a half-equation simply because we need two half-equations to form an overall or net equation. In writing half equations, chemists typically include the electrons in the equation.

The oxidation half-equation above describes the process of a copper atom (with no charge) being oxidized (losing electrons) to form a copper ion with a 2+ charge. Notice that we have a balance of charge and mass on both sides of the equation. We have one copper atom on both sides and the charges balance as well.

Cu (s) → Cu2+ (aq) + 2e-

You should note that metals are almost always oxidised (lose electrons) because they have valence electrons that are removed to create a stable electron configuration.

In the reduction half-equation below, two silver ions are being reduced through the addition of the two electrons lost from the copper atom above to form a silver atom.

2Ag+ (aq) + 2e-→ 2Ag (s)


Since oxidation and reduction occur simultaneously, these two equations can be added together to give a net or overall equation.

Cu (s) + 2Ag+ (aq) + 2e- → Cu2+ (aq) + 2e-+ 2Ag (s)

Cu (s) + 2Ag+ (aq) → Cu2+ (aq) + 2Ag (s)


Some things to commit to memory:

A species that is oxidised is called the reductant or reducing agent.

A species that is reduced is is called the oxidant or the oxidising agent.

In the above reactions, the copper atoms (oxidised) are the reductants and the silver ions (reduced) are the oxidants.

 

Metal Displacement Reactions

A single metal displacement reaction is one where a metal is placed into a solution conatining another metal ion and an oxidation-reduction reaction occurs. If a piece of copper metal was placed into a solution of silver nitrate, a metal displacement reaction would occur. The copper atoms and oxidised to copper ions and the silver ions are reduced to copper atoms according to the equations shown above. There is a net transfer of two electrons from one copper atom to two silver ions (not that copper and silver have different valencies here). The transfer happens on the metal surface as newly created copper ions float out into the solution and silver ions in the solution form a deposit of silver atoms on the surface of the copper metal.



A metal single metal displacement reaction between copper and silver. In (a) the reactants are shown as brown
copper atoms and grey silver ions with a positive charge. After the reaction (b) a deposit of silver has formed
on the copper and grey silver atoms and blue copper ions can be seen.


If instead we were to put a piece of silver metal into a solution of copper nitrate, no reaction would occur. This is because copper is more easily oxidised than silver. Said another way, copper is a stronger reducant than silver so the copper will be oxidised in preference to the silver.

Metal displacement reactions can be used to create an activity series for the metals which is a list of metals listed in descending order from the strongest to the weakest reductant. The metals at the top of the activity series are most easily oxidised or the strongest reductants.

The rule for metal displacement reactions goes like this:

A more reactive metal will displace the ion of a less reactive metal.

Metal displacement reactions provide the basis understanding how galvanic cells work. In a metal displacement reaction, the electrons are transferred from the oxidised to the reduced species on the surface of the metal. If we could put a conducting wire between the oxidation reaction and the reduction reaction, we have a way of creating an electric current from a redox reaction and this is exactly what a galvanic cell or battery does.



A demonstration showing two different metal displacement reactions between
(1) aluminium metal and copper ions and (2) magnesium metal and copper ions.

 

Galvanic Cells

To make a functioning electrochemical cell, the oxidiation and reduction reactions are separated into two half-cells with a voltmeter between them to measure the energy output of the cell. A typical electrochemical cell has each metal electrode immersed in a solution containing its ions. By definition, the electrode where oxidation occurs is called the anode and the electrode where reduction occurs is the cathode. Since electrons flow from the anode to the cathode spontaneously, they must be flowing away from a negative electrode towards a positive one. This makes the anode negative and the cathode positive in an electrochemical cell.

A salt bridge connects the two half-cells to help neutralise the charge that builds up as a result of oxidation and reduction. The oxidation of zinc in the half-cell below is increasing the positive charge in the solution due to the production of zinc ions so chloride ions flow out of the salt bridge in this direction to reset the charge to zero in the solution. In the reduction half-cell, copper ions are being reduced to atoms so the positive charge is decreasing and the negative charge increasing. Sodium ions flow out of the salt bridge to reset the negative charge back to zero. If the salt bridge were not present, the build-up of charge in each half-cell would prevent the flow of electrons through the external circuit.



A typical electromchemical cell with a zinc anode and a copper cathode. Note the presence of zinc and
copper ions in each respective half-cell and the flow of electrons from the anode to the cathode. As the
reaction proceeds, the zinc electrode loses mass, the copper electrode gains mass and the copper
catholyte solution fades due to the reduction in concentration of the blue copper ions.


A simpler notation sometimes used to describe
electrochemical cells.

You could be asked to label an electrochemical cell where the only information you are given is the metals to be used for the anode and cathode. In this case you should follow the steps below:

  1. Anode and Cathode
    First identify all species available for oxidation (Zn and Cu above). Use the table of standard potentials to decide which species is most easily oxidised (see below). This species will be the anode (Zn) and the other will be the cathode (Cu).

  2. Electron Flow
    Label the electron flow from the anode to the cathode.

  3. Solutions
    Each species in the oxidation and reduction reaction must be present in the respective half-cell. If the electrode is made of zinc, the solution must contain zinc ions and the same goes for the other electrode. Select a suitable soluiton for both half-cells making sure that they are soluble salts. Sometimes the anode solution is called the annolyte and the cathode solution the catholyte.

  4. Salt Bridge
    Select a suitable ionic salt for the salt bridge. Make sure that its soluble and that it doesn't form any inoluble precipitates with the ions in each half-cell. Potassium nitrate is a good choice because all nitrates are soluble and all potassium salts are soluble.



The table of standard potentials which can be found on the data sheet in the supplementary materials for the
HSC examination. When working out which species is oxidised or reduced, use the middle column which
contains mostly metals. This list is a kind of activity series and shows the most easily oxidised at the top
and the least easily oxidised at the bottom.



An elecrochemical cell containing an inert platinum cathode for the reduction
of hydrogen ions to hydrogen gas. In this case, zinc is more easily oxidised
than hydorgen gas so zinc becomes the anode.

Some cells will use a platinum or graphite electrode because one of the species being oxidised or reduced cannot be made into an electrode.

For example, if hydrogen gas is being reduced to hydrogen ions, an electrode cannot be made from hydrogen ions or hydrogen gas. In this case an inert electrode such as platinum or graphite is used.

Both species in the reduction reaction still need to be present in the half-cell, so hydrogen gas must be pumped in and the catholyte must contain hydrogen ions (usually as an acid). Otherwise, the cell and its labelling occur exactly as described in the sequence of steps above.

 

Calculating the Cell Potential

The cell potential represents the theoretical maximum voltage output of a cell. It can be calculated as follows.

Step
Description
Reaction
Cell Potential
1

Oxidation half equation Reverse the standard reduction equation and the cell potential on the data sheet

Zn → Zn2+ + 2e-

EØ= +0.76 V
2

Reduction half equation

Cu2+ + 2e-→ Cu

EØ= +0.34 V
3 Full equation

Zn + Cu2+ → Zn2+ + 2Cu

EØ= 0.76 + 0.52 = 1.1 V


A positive value for the cell potential indicates that this reaction will occur spontaneously and it is therefore a galvanic cell. A negative value indicates that this reaction will only occur if electrical energy is applied to the system and this would make it an electrolytical cell which we study later in the option.

The standard potentials provided in the table are measure under quite specific conditions. The standard electrode potentials are customarily determined at solute concentrations of 1 mol L-1, gas pressures of 1 atm and a standard temperature which is usually 25°C. In the school laboratory, these conditions are not always the same and as such, the measured cell potential can vary from the theoretical cell potential.


Dry Cell Batteries

Although there are many different types of batteries ranging from the relatively large flashlight batteries to the minaturised versions used for wristwatches or calculators, the most common type in use today isthe Lelanche dry cell battery. Different types of batteries vary widely in composition and form, but they all work on the same principles of oxidation-reduction reactions.

The structure of a typical acidic dry cell will a
zinc anode and graphite (carbon) cathode.

A dry-cell battery is essentially comprised of a metal electrode or graphite rod (elemental carbon) surrounded by a moist electrolyte paste enclosed in a metal cylinder as shown on the right. In the most common type of dry cell battery, the cathode is composed of a form of elemental carbon called graphite, which serves as a conductor of electrons to the reduction half-reaction.

In an acidic dry cell, a thin zinc cylinder serves as the anode and it undergoes oxidation.

Zn → Zn2+ + 2e-

The reduction reaction occurs within the moist paste comprised of ammonium chloride (NH4Cl) and manganese dioxide (MnO2).

2NH4+ + 2MnO2 + 2e- → Mn2O3 + 2NH3 + H2O

This dry cell couple produces about 1.5 volts.

In the alkaline version or alkaline battery, the ammonium chloride is replaced by KOH or NaOH.

The anode reaction is:

Zn + 2OH- → ZnO + H2O + 2e-

The reduction equation is:

2 MnO2 + 2e- + H2O → Mn2O3 + 2 OH-

The alkaline dry cell lasts much longer as the zinc anode corrodes less rapidly under basic conditons than under acidic conditions. Other types of dry cell batteries are the silver battery in which silver metal serves as an inert cathode to support the reduction of silver oxide (Ag2O) and the oxidation of zinc (anode) in a basic medium. The type of battery commonly used for calculators is the mercury cell. In this type of battery, HgO serves as the oxidizing agent (cathode) in a basic medium, while zinc metal serves as the anode. Another type of battery is the nickel/cadmium battery, in which cadmium metal serves as the anode and nickel oxide serves as the cathode in an alkaline medium. Unlike the other types of dry cells described above, the nickel/cadmium cell can be recharged.

 

Oxidation States

In chemical reactions where it is not immediately obvious which species is oxidised and which is reduced, oxidation states or numbers can be used to work it out.

An increase in oxidation state (number) means that a species has been oxidised and a decrease means that it has been reduced.

There are simple rules for assigning oxidation numbers and these are summarised below.

  1. The oxidation state of any element such as Fe, H2, O2, P4, S8 is zero (0).

  2. The oxidation state of oxygen in its compounds is -2, except for peroxides like H2O2, and Na2O2, in which the oxidation state for O is -1.

  3. The oxidation state of hydrogen is +1 in its compounds, except for metal hydrides, such as NaH, LiH, etc., in which the oxidation state for H is -1.

  4. The oxidation states of other elements are then assigned to make the algebraic sum of the oxidation states equal to the net charge on the molecule or ion.

  5. The following elements usually have the same oxidation states in their compounds:
    +1 for alkali metals - Li, Na, K, Rb, Cs;
    +2 for alkaline earth metals - Be, Mg, Ca, Sr, Ba; and
    -1 for halogens except when they form compounds with oxygen or one another.

Other than these, you may simply remember the oxidation states for H and O are +1 and -2 respectively in a coumpound and oxidation of other elements can be asigned by making the algebraic sum of the oxidation states equal to the net charge on the molecule or ion.