Category: Basic Chemistry

Electrochemical cells are batteries. One of the most common forms of battery is the alkaline battery (alkaline because potassium hydroxide is the electrolyte).

At the anode:

Unlike electrochemical cells mentioned in an earlier article, the anode is not a piece of zinc but zinc particles suspended in a KOH electrolyte (Zn–KOH paste). Zinc is oxidised to Zn²⁺ as follows:

Zn (s) + 2OH^– (aq) → ZnO (s) + 2H₂O (l) + 2e^–

The free electrons are collected at the anode collector and flow out of the battery via the anode cap when the battery is connected to an external circuit to power an electronic device (see diagram below).

At the cathode:

Returning electrons enter the battery via the cathode cap and flow along the sides to the cathode collector and then to the cathode where they participate in the following reduction reaction:

2MnO₂ (s) + H₂O (l) + 2e^–→ Mn₂O₃ (s) + 2OH^– (aq)

Like the anode, the cathode is not a piece of manganese (IV) oxide but a paste of manganese (IV) oxide and carbon (carbon is added to improve conductivity). The porous cellulose separator prevents the electrode materials from mixing but allows hydroxide ions produced in the cathode to migrate across it to the anode to maintain electrical neutrality.

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A fuel cell converts chemical energy from the fuel (usually H₂) to electrical energy, which is used to power engines. With reference to the diagram below, H₂ is fed into the left compartment, where it is oxidised to H⁺ at the anode, which is porous and impregnated with a Pt catalyst. The protons then migrate across the electrolyte (H₃PO₄) contained in a polymer exchange membrane that only allows the passage of H⁺.

At the cathode, O₂ in the air fed into the right compartment reacts with H⁺ and is reduced to form water. The cathode is again porous and impregnated with a Ni catalyst. The overall redox reaction is:

2H₂ (g) + O₂ (g) → 2H₂O (l)

 

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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 when 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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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

1. • Finding the possible reactants in the electrochemical cell: Zn, Cu, H⁺, SO₄^2– 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.

 

 

 

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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 SO₄^2– 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 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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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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The stability of metal oxides, metal hydroxides, metal carbonates and metal nitrates are summarised in the table below.

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

The displacement occurs because the overall energy of the system is lowered, making the forward reaction energetically favourable. Essentially, the more reactive metal has a greater ability to donate electrons, which drives the displacement reaction. For example, the thermite reaction (also known as the Goldschmidt process) is used to prepare small quantities of metallic iron:

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

 

 

 

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Some of the reactions of metals with water and acid are summarised in the table below.

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