Electro chemical cells

In electrochemical cells, electricity is produced from chemical reactions i.e. chemical energy is converted to electric energy. It consists of two half cells or electrodes and at each electrode/half cell, an element is in contact with a solution of its ions.
For example, consider zinc plate dipped into a solution containing zinc ions and copper plate dipped in a solution of copper ions as below.

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The metal surface there fore loses electrons and becomes positively charged. There will be a potential difference between the metal surface and its solution.

The above reactions depend on the reactivity of the metals. Metals which are more reactive e.g. zinc have a greater tendency to dissolve as ions and a smaller tendency to be deposited as metal. While metals that are less reactive like copper have greater tendencies to be deposited as metal and small tendency to dissolve and form ions. Zinc there fore acquires a more negative potential than copper.
The tendency of a metal to dissolve when in solution of its own ions can be measured by joining two half cells to from an electrochemical cell/galvanic or voltaic cell.

Connection between the two half cells
i) The two half cells can be connected by a porous partition which allows conduction of electricity by movement of ions across without physically mixing the solutions.
For example the Daniel cell (a primary cell) which consists of zinc and copper metals in solutions of their ions combined as shown below.

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The zinc electrode is on the left hand side as it has a more negative potential and copper on the right hand side as it has a more positive potential. Current flows from the positive electrode to the negative electrode and electrons flow from negative to positive electrode.
Reaction at the electrodes
At Zinc (negative) electrode
Zinc dissolves from the zinc electrode forming zinc ions and electrons are released.

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At copper (positive) electrode
Copper ions gain electrons and copper is deposited on the electrode.

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When PbO2 and Pb have been converted into PbSO4, there is no difference between the plates, and the cell can no longer give a current. The battery must be charged.
The battery is charged by passing a direct electric current through it. This reverses the electrode processes in the cells to produce lead and lead (IV) oxide. The battery can again give a current. When the vehicle is in motion, it drives a generator which charges the battery. If there is too much stopping and starting, the battery loses its charge and becomes flat until it is recharged.
During discharge, sulphuric acid is formed. During charge, it is used up. Since sulphuric acid is much denser than water, the state of the battery can be assessed by measuring the density of the battery acid with a hydrometer.
Part of a lead-acid battery

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REDOX REACTIONS AND IONIC EQUATIONS
These are reactions in which both reductions and oxidations take place simultaneously.


Oxidation
Can be defined as;
i) Addition of oxygen to a substance
ii) Removal of hydrogen from a substance
iii) Loss of electrons from a substance
NB. Oxidizing agents are electron acceptors. Therefore non metals are oxidizing agents.

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Exercise
Calculate or determine the oxidation numbers of stated elements in the following;
i) Sulphur in Sulphite ion
ii) Hydrogen in Hydrochloric acid
iii) Phosphorous in phosphate
iv) Carbon in Carbon dioide
v) Carbon in cabon monoxide
vi) Sulphur in Sulphur dioxide
vii) Oxygen in Hydroxyl ion
viii) Chromium in Dichromate ion
ix) Copper in Copper (II) oxide
x) Manganese in Permanganate (VII) ions

Examples

State what is taking place and write half reaction equation in each case.
NB. Oxidation is an increase in oxidation number and reduction is a decrease in the oxidation number.
i) The conversion of hydrogen ions (H+) to hydrogen molecules (H2)

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  1. Displacement reactions as redox reactions
    (a) Metals

    A displacement reaction is one in which one element takes the place of another element in a chemical reaction. Example, metals high up in the reactivity series [more reactive metals] tend to display those below them from their solutions [compounds]. In this case, the more reactive metals act as reducing agents [electron donors] and their reducing power decreases as reactivity decreases.
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b) Displacement reactions of halogens
Halogens are displaced from their solutions by other halogens which are higher in the electro chemical series [more reactive halogens]. The order of reactivity of halogens is Flourine >Chlorine>Bromine>Iodine>Astatine.
For example
When chlorine gas is passed through an aqueous colourless solution of potassium iodide, a brown colouration develops immediately. The brown colouration is due to iodine liberated as iodide ions are displaced by chlorine from solution example;

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Chlorine in the above case acts as the oxidizing agent and the oxidizing power of the halogens decreases with their decreasing reactivity.


IONIC EQUATIONS
The ionic equation is used to describe the chemical reaction while also clearly indicating which of the reactants and/or products exist primarily as ions in aqueous solution.
Rules in writing ionic equations

  1. Start with a balanced molecular equation.
    A molecular equation is one that shows the chemical formulas of all reactants and products but does not expressly indicate their ionic nature.
    For example,
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  1. Break all soluble strong electrolytes (compounds with (aq) beside them) into their ions
    o indicate the correct formula and charge of each ion
    o indicate the correct number of each ion
    o write (aq) after each ion
  2. Bring down all compounds with (s), (l), or (g) unchanged
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