# 4.4 Applications  (Page 2/2)

 Page 2 / 2

The basic electrolytic reactions involved are as follows: At the cathode :

$\begin{array}{ccc}\hfill A{l}^{+3}+3{e}^{-}& \to & Al\left(s\right)\phantom{\rule{4pt}{0ex}}\phantom{\rule{4pt}{0ex}}\phantom{\rule{4pt}{0ex}}\phantom{\rule{4pt}{0ex}}\phantom{\rule{4pt}{0ex}}\left(99%\mathrm{purity}\right)\hfill \end{array}$

At the anode :

$\begin{array}{ccc}\hfill 2{O}^{2-}& \to & {O}_{2}\left(g\right)+4{e}^{-}\hfill \end{array}$

The overall reaction is as follows:

$\begin{array}{ccc}\hfill 2A{l}_{2}{O}_{3}& \to & 4Al+3{O}_{2}\hfill \end{array}$

The only problem with this process is that the reaction is endothermic and large amounts of electricity are needed to drive the reaction. The process is therefore very expensive.

## Summary

• An electrochemical reaction is one where either a chemical reaction produces an external voltage, or where an external voltage causes a chemical reaction to take place.
• In a galvanic cell a chemical reaction produces a current in the external circuit. An example is the zinc-copper cell.
• A galvanic cell has a number of components . It consists of two electrodes , each of which is placed in a separate beaker in an electrolyte solution. The two electrolytes are connected by a salt bridge . The electrodes are connected two each other by an external circuit wire.
• One of the electrodes is the anode , where oxidation takes place. The cathode is the electrode where reduction takes place.
• In a galvanic cell, the build up of electrons at the anode sets up a potential difference between the two electrodes, and this causes a current to flow in the external circuit.
• A galvanic cell is therefore an electrochemical cell that uses a chemical reaction between two dissimilar electrodes dipped in an electrolyte to generate an electric current.
• The standard notation for a galvanic cell such as the zinc-copper cell is as follows:
$Zn|Z{n}^{2+}||C{u}^{2+}|Cu$
where
$\begin{array}{ccc}\hfill |& =& \mathrm{a}\phantom{\rule{4pt}{0ex}}\mathrm{phase}\phantom{\rule{4pt}{0ex}}\mathrm{boundary}\phantom{\rule{4pt}{0ex}}\left(\mathrm{solid}/\mathrm{aqueous}\right)\hfill \\ \hfill ||& =& \mathrm{the}\phantom{\rule{4pt}{0ex}}\mathrm{salt}\phantom{\rule{4pt}{0ex}}\mathrm{bridge}\hfill \end{array}$
• The galvanic cell is used in batteries and in electroplating .
• An electrolytic cell is an electrochemical cell that uses electricity to drive a non-spontaneous reaction. In an electrolytic cell, electrolysis occurs, which is a process of separating elements and compounds using an electric current.
• One example of an electrolytic cell is the electrolysis of copper sulphate to produce copper and sulphate ions.
• Different metals have different reaction potentials . The reaction potential of metals (in other words, their ability to ionise), is recorded in a standard table of electrode potential . The more negative the value, the greater the tendency of the metal to be oxidised. The more positive the value, the greater the tendency of the metal to be reduced.
• The values on the standard table of electrode potentials are measured relative to the standard hydrogen electrode .
• The emf of a cell can be calculated using one of the following equations: E ${}_{\left(cell\right)}^{0}$ = E ${}^{0}$ (right) - E ${}^{0}$ (left) E ${}_{\left(cell\right)}^{0}$ = E ${}^{0}$ (reduction half reaction) - E ${}^{0}$ (oxidation half reaction) E ${}_{\left(cell\right)}^{0}$ = E ${}^{0}$ (oxidising agent) - E ${}^{0}$ (reducing agent) E ${}_{\left(cell\right)}^{0}$ = E ${}^{0}$ (cathode) - E ${}^{0}$ (anode)
• It is possible to predict whether a reaction is spontaneous or not, either by looking at the sign of the cell's emf or by comparing the electrode potentials of the two half cells.
• It is possible to balance redox equations using the half-reactions that take place.
• There are a number of important applications of electrochemistry. These include electroplating , the production of chlorine and the extraction of aluminium .

## Summary exercise

1. For each of the following, say whether the statement is true or false . If it is false, re-write the statement correctly.
1. The anode in an electrolytic cell has a negative charge.
2. The reaction $2{\mathrm{KClO}}_{3}\to 2\mathrm{KCl}+3{\mathrm{O}}_{2}$ is an example of a redox reaction.
3. Lead is a stronger oxidising agent than nickel.
2. For each of the following questions, choose the one correct answer.
1. Which one of the following reactions is a redox reaction?
1. $HCl+NaOH\to NaCl+{H}_{2}O$
2. $AgN{O}_{3}+NaI\to AgI+NaN{O}_{3}$
3. $2FeC{l}_{3}+2{H}_{2}O+S{O}_{2}\to {H}_{2}S{O}_{4}+2HCl+2FeC{l}_{2}$
4. $BaC{l}_{2}+MgS{O}_{4}\to MgC{l}_{2}+BaS{O}_{4}$
(IEB Paper 2, 2003)
2. Consider the reaction represented by the following equation: $B{r}_{2\left(l\right)}+2{I}_{aq}^{-}\to 2B{r}_{aq}^{-}+{I}_{2\left(s\right)}$ Which one of the following statements about this reaction is correct?
1. bromine is oxidised
2. bromine acts as a reducing agent
3. the iodide ions are oxidised
4. iodine acts as a reducing agent
(IEB Paper 2, 2002)
3. The following equations represent two hypothetical half-reactions: ${X}_{2}+2{e}^{-}⇌2{X}^{-}$ (+1.09 V) and ${Y}^{+}+{e}^{-}⇌Y$ (-2.80 V) Which one of the following substances from these half-reactions has the greatest tendency to donate electrons?
1. X ${}^{-}$
2. X ${}_{2}$
3. Y
4. Y ${}^{+}$
4. Which one of the following redox reactions will not occur spontaneously at room temperature?
1. $Mn+C{u}^{2+}\to M{n}^{2+}+Cu$
2. $Zn+S{O}_{4}^{2-}+4{H}^{+}\to Z{n}^{2+}+S{O}_{2}+2{H}_{2}O$
3. $F{e}^{3+}+3N{O}_{2}+3{H}_{2}O\to Fe+3N{O}_{3}^{-}+6{H}^{+}$
4. $5{H}_{2}S+2Mn{O}_{4}^{-}+6{H}^{+}\to 5S+2M{n}^{2+}+8{H}_{2}O$
5. Which statement is CORRECT for a Zn-Cu galvanic cell that operates under standard conditions?
1. The concentration of the Zn ${}^{2+}$ ions in the zinc half-cell gradually decreases.
2. The concentration of the Cu ${}^{2+}$ ions in the copper half-cell gradually increases.
3. Negative ions migrate from the zinc half-cell to the copper half-cell.
4. The intensity of the colour of the electrolyte in the copper half-cell gradually decreases.
(DoE Exemplar Paper 2, 2008)
3. In order to investigate the rate at which a reaction proceeds, a learner places a beaker containing concentrated nitric acid on a sensitive balance. A few pieces of copper metal are dropped into the nitric acid.
1. Use the relevant half-reactions from the table of Standard Reduction Potentials to derive the balanced nett ionic equation for the reaction that takes place in the beaker.
2. What chemical property of nitric acid is illustrated by this reaction?
3. List three observations that this learner would make during the investigation.
(IEB Paper 2, 2005)
4. The following reaction takes place in an electrochemical cell:
$Cu\left(s\right)+2AgN{O}_{3}\left(aq\right)\to Cu{\left(N{O}_{3}\right)}_{2}\left(aq\right)+2Ag\left(s\right)$
1. Give an equation for the oxidation half reaction.
2. Which metal is used as the anode?
3. Determine the emf of the cell under standard conditions.
(IEB Paper 2, 2003)
5. The nickel-cadmium (NiCad) battery is small and light and is made in a sealed unit. It is used in portable appliances such as calculators and electric razors. The following two half reactions occur when electrical energy is produced by the cell. Half reaction 1: $\mathrm{Cd}\left(\mathrm{s}\right)+2{\mathrm{OH}}^{-}\left(\mathrm{aq}\right)\to \mathrm{Cd}{\left(\mathrm{OH}\right)}_{2}\left(\mathrm{s}\right)+2{\mathrm{e}}^{-}$ Half reaction 2: $\mathrm{NiO}\left(\mathrm{OH}\right)\left(\mathrm{s}\right)+{\mathrm{H}}_{2}\mathrm{O}\left(\mathrm{l}\right)+{\mathrm{e}}^{-}\to \mathrm{Ni}{\left(\mathrm{OH}\right)}_{2}\left(\mathrm{s}\right)+{\mathrm{OH}}^{-}\left(\mathrm{aq}\right)$
1. Which half reaction (1 or 2) occurs at the anode? Give a reason for your answer.
2. Which substance is oxidised?
3. Derive a balanced ionic equation for the overall cell reaction for the discharging process.
4. Use your result above to state in which direction the cell reaction will proceed (forward or reverse) when the cell is being charged.
(IEB Paper 2, 2001)
6. An electrochemical cell is constructed by placing a lead rod in a porous pot containing a solution of lead nitrate (see sketch). The porous pot is then placed in a large aluminium container filled with a solution of aluminium sulphate. The lead rod is then connected to the aluminium container by a copper wire and voltmeter as shown.
1. Define the term reduction .
2. In which direction do electrons flow in the copper wire? (Al to Pb or Pb to Al)
3. Write balanced equations for the reactions that take place at...
1. the anode
2. the cathode
4. Write a balanced nett ionic equation for the reaction which takes place in this cell.
5. What are the two functions of the porous pot?
6. Calculate the emf of this cell under standard conditions.
(IEB Paper 2, 2005)

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