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Other forms of electrochemical cells

The electrochemical cell below contains two electrolytes, aqueous zinc sulphate and aqueous copper sulphate, that are separated by a semi-permeable membrane, which in this case, is an anion-exchange membrane (AEM). The AEM, made with positively charged polymers, allows only anions in the electrolyte to pass through it, with the cations remaining in their respective compartments. It serves to maintain electrical neutrality in both compartments and to complete the electrochemical circuit.

To determine the reactions in the cell, we follow the steps outline in the previous article, which result in the following overall redox reaction equation:

Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s)

As the reaction progresses, the blue copper sulphate solution decolourises as Cu2+ ions are reduced to Cu.

Another way to construct an electrochemical cell is shown in the diagram below. The electrolytes are housed in separate vessels that are linked by a salt bridge, which is a gel-filled glass tube containing a suitable electrolyte. The preferred electrolyte in the salt bridge is one that does not react with the electrolytes in the vessels and with the electrodes.

Like the semi-permeable membrane in the earlier setup, the salt bridge maintains electrical neutrality in both vessels and completes the electrochemical circuit by allowing the flow of anions from one vessel to the other.

 

 

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Analysing reactions in an electrochemical cell

To analyse reactions in an electrochemical cell, for example the Zn–Cu cell, we follow these steps:

Step 1:  Determine the direction of flow of electrons (to identify the oxidation reaction) by

    • Finding the possible reactants in the electrochemical cell: Zn, Cu, H+, SO42– and H2O (note that H2O is always assumed to be the species instead of OH– in non-hydroxide aqueous solutions, as the concentration of OH– from H2O is very small).
    • Distinguishing the most reactive species according to the electrochemical series: Zn
    • Writing the ionic half-reaction equation for the most reactive species, which undergoes oxidation:

Zn (s) → Zn2+ (aq) + 2e–

Zinc, being most reactive amongst all the species in the electrochemical cell, is oxidised with free electrons flowing from the anode (zinc electrode) to the cathode (copper electrode). The electrons, upon reaching the cathode, preferentially reduce one of the species at the electrode’s surface.

Step 2:  Identify the reduction reaction at the cathode by

    • Finding all possible reducible species (usually cations): Zn2+, H2O and H+.
    • Distinguishing the species that is most readily reduced, if there is more than one competing species, by:
      • comparing the species on the electrochemical series: H+.
      • comparing the concentrations of the species if their positions in electrochemical series are relatively close to one another: Zn2+ and H2O are relatively far apart in the series from H+, so the result is still H+.
    • Writing the ionic half-reaction equation for the reduction reaction:

2H+ (aq) + 2e– → H2 (g)

Combining the two half reaction equations, the overall redox reaction in the electrochemical cell is:

Zn (s) + 2H+ (aq) → Zn2+ (aq) + H2 (g)

A potential difference is established between the two electrodes as a result of the redox reaction, and its magnitude (voltage) is measured by the voltmeter attached to the wire.

If we perceive the electrolyte as the chemical in a battery, the electrodes as terminals of the battery and the voltmeter as the external load, we can label the anode as negative and the cathode as positive since current (opposite direction to the flow of electron) flows from the positive terminal of a battery to the negative terminal. Note that the electrochemical circuit is complete by the flow of electrons from the anode to the cathode via the wire and a net flux of negative ions towards the anode and positive ions towards the cathode in the electrolyte.

 

Question

Why is the voltmeter not connected parallel to a load in the circuit?

Answer

By connecting the voltmeter in series, we are measuring the potential difference of an open circuit with almost no current flowing. This is an accurate way to measure the potential difference, as it allows the concentration of the electrolytes to remain unchanged. To understand more about the effects of the flow of a current on the potential of an electrochemical cell, read the article on overpotential in the advanced section.

 

 

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Electrochemical cell

An electrochemical cell (also known as a simple cell) is a device that generates electrical energy from chemical reactions. The earliest forms of electrochemical cells were invented by Luigi Galvani, Alessandro Volta and John Frederic Daniell, and hence the terms Galvanic cell, Voltaic cell and Daniell cell respectively.

The simplest electrochemical cell consists of two dissimilar metals (preferably far apart from each other on the metal reactivity series or electrochemical series) dipped in a conducting solution and connected by a wire. For example, the diagram above has Zn and Cu (electrodes) dipped in a H2SO4 solution (electrolyte), which consists of H+ and SO42– from sulphuric acid, and H+ and OH– from water.

Due to the different reactivity of the two metals making up the electrodes, and the ions in the electrolyte, chemical reactions occur. The electrode where oxidation reactions occur is called the anode, while the electrode where reduction reactions take place is known as the cathode.

 

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The electrochemical series

The electrochemical series is a tabulation of the half-cell standard electrode potentials of various chemical species.

Just as the metal reactivity series compares the reactivity of metals, the electrochemical series lists the reactivity of metals and ions in an electrochemical cell. A summary of the electrochemical series is as follows:-

Since potassium is at the top of the cation series, its elemental form K is the most reactive and, therefore, most easily oxidised to its ionic form at the anode, compared to the other elements in the series. This implies that its ionic form K+ is least easily reduced at the cathode compared to the other ions in the series. The same logic applies to anions. Note that the oxidation of the sulphate (and nitrate) ion at the anode is relatively thermodynamically infeasible, as the anion is quite stable. The two series can be combined into one, in the form of reduction potentials (i.e. oxidised-to-reduced from left to right) as follows:

 

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Basic electrochemistry: overview

Electrochemistry is the study of the relationship between electricity and chemical reactions. Such a relationship is manifested in an electrochemical process that relies on the principle of oxidation and reduction to enable the flow of electrons between chemical species. In this section, we shall focus on two electrochemical devices, namely, the electrochemical cell and the electrolytic cell. We want to understand the chemistry behind these devices, investigate how they work and learn about their applications.

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Metal displacement reactions

A metal displacement reaction is one where a less reactive metal in a compound is replaced by a more reactive metal.

In general, a more reactive metal can displace a less reactive metal from its compound. This concept is used in the thermite reaction (also known as the Goldschmidt process) for preparing small quantities of metallic iron.

2Al (s) + Fe2O3 (s) → Al2O3 (s) + 2Fe (s)

 

Question

What do you observe when you place a strip of magnesium in a beaker containing dilute aqueous copper sulphate?

Answer

The blue solution turns colourless with pink deposits at the bottom of the beaker.

Mg (s) + CuSO4 (aq) → MgSO4 (aq) + Cu (s)

 

 

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Reactivity series of metals

The reactivity of metals broadly refers to the relative extent of change that metals undergo with respect to certain reactions. Some of these reactions are:

    1. Metal reaction with water
    2. Metal reaction with acid
    3. Metal displacement reactions
    4. Extraction of metals from metal oxides (or other ores)
    5. Heating of metal hydroxides, metal carbonates and metal nitrates

In general, the greater the extent of change of a metal from its elemental form to a compound is, the more reactive it is. For example, sodium is more reactive than copper, as elemental sodium reacts violently with cold water, while elemental copper does not react with water and steam. The reactivities of metals for reactions 1 to 5 are summarised in the following table that is known as the metal reactivity series:

The two non-metals, C and H are included for reference.

 

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Effects of pH on redox reactions

What are the effects of pH on redox reactions?

pH affects the strength of an oxidising agent. For example, the oxidising strength of potassium manganate (VII), as measured by its standard electrode potential, is as follows:

Acidic medium

MnO4– + 8H++ 5e– → Mn2+ + 4H2O               Eored  = +1.51 V

Neutral/slightly basic medium

MnO4– + 2H2O + 3e– → MnO2 + 4OH–              Eored  = +1.23 V

Strongly basic medium

MnO4– + e– → MnO42-               Eored  = +0.56 V

Hence, potassium manganate (VII) is most strongly oxidising in an acidic medium and least oxidising in a strongly basic medium. Note that the manganate (VI) ion is stable only in very basic conditions. It disproportionates in acidic, neutral and slightly basic conditions as follows:-

3MnO42- + 4H+ → 2MnO4– + MnO2 + 2H2O

 

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