Sodium Hydroxide

Explain the difference between galvanic cells and electrolytic cells in terms of energy requirements

The difference between galvanic cells and electrolytic cells causes endless grief for students if they attempt to think of it in terms of positive and negative anodes and cathodes. Rather, keep thinking in terms of AnOx (Anode Oxidation).

In galvanic cells, a spontaneous chemical reaction converts chemical energy into electrical energy, whereas in an electrolytic cell, electrical energy is converted into chemical energy so as to bring about a chemical reaction.

In both the galvanic and electrolytic cells, oxidation still occurs at the anode and reduction occurs at the cathode.

Remember- A galvanic cell converts chemical energy into electrical energy. An electrolytic cell con- verts electrical energy into chemical energy.

Identify data, plan and perform a first-hand investigation to identify the prod- ucts of the electrolysis of sodium chloride

This experiment is relatively simple, and simply confirms the fact that sodium ions do not reduce in the production of sodium hydroxide as readily as water. A concentrated sodium chloride solution is used as a molten sodium chloride solution is not as practical.

Procedure:

  1. Prepare one beaker of dilute sodium chloride solution and one beaker of concentrated sodium chloride
  2. Set up two electrolytic cells using the two A current should be run through the beakers on at a time.
  3. Place a test tube at both electrodes, taking care to trap any gas formed during the
  1. Place a pH meter at both the electrodes of the dilute sodium chloride
  2. Light a splint, gently blow it out, and place near each of two electrodes in the dilute sodium chloride solution beaker in quick
  3. Repeat steps 3 to 5 for the beaker containing the concentrated sodium chloride
  4. Waft (Ideally, conduct the experiment in a fume cupboard) any gases formed during the ex- periment carefully towards yourself to identify its
  5. Place a piece of litmus paper each at both electrodes of the concentrated sodium chloride solution.

Expected Results:

  • The pH was below 7 at the anode, and above 7 at the This indicates the self-ionisation of water brought about by the unresponsiveness of sodium ions to reduction where water is present.
  • A re-lighted splint indicates the presence of
  • A popping sound produced when gas caught in a test tube is ignited indicates the presence of hydrogen
  • Chlorine gas has a distinct odour which is immediately
  • The bleaching of litmus paper at the anode of the concentrated sodium chloride solution confirms the presence of chlorine

Analyse information from secondary sources to predict and explain the different products of the electrolysis of aqueous and molten sodium chloride

In an aqueous sodium chloride solution:

At the anode:

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At the cathode:

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Full equation:

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The sodium ion is more stable than water, and as such, water is reduced instead.

In a molten sodium chloride solution:

At the anode:

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At the cathode:

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Full equation:

 

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Although water is less stable than a sodium ion and thus more likely to be reduced, there is no water in molten sodium chloride! As such, the sodium ions must be reduced.

Remember- Where water exists, it will be reduced more readily than a sodium ion. Where no water exists, the sodium ions must be reduced. Chlorine ions are always oxidised in the formation of sodium hydroxide.

Outline the steps in the industrial production of sodium hydroxide from sodium chloride solution and describe the reaction in terms of net ionic and full formulae equations

In the industrial production of sodium hydroxide, a solution known as brine (really just concentrated NaCl solution) is electrolysed.

At the anode, chlorine ions are oxidised (and therefore lose electrons becoming more positive) to form chlorine gas.

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At the cathode, sodium ions are not reduced as would be expected. Rather, because the sodium ions are simply too stable to undergo this process, the water in the solution is reduced instead, forming hydroxide ions and hydrogen gas.

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Through these processes, sodium hydroxide is formed.

Net ionic equation:

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Full equation:

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Remember- Chlorine ions are oxidised at the ions and water, not sodium ions, are reduced at the cathode to form sodium hydroxide from brine.

 

Distinguish between the three electrolysis methods used to extract sodium hydroxide: The mercury process, diaphragm process, and membrane process by describing each process and analysing the technical and environmental difficulties involved in each process

After examining the general gist of how electrolysis occurs and then how sodium hydroxide is pro- duced, this dotpoint focuses specifically on three of the more common processes by which these reactions are brought about. Take the time to understand each of these equations, as they may prove important in your exams.

Mercury Process

In this process, the anode is a segmented titanium plate and the cathode is mercury. At the anode, chlorine ions are oxidised to produce chlorine gas.

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At the cathode, sodium ions are reduced to produce sodium, which forms a mixture with the mercury in the cell.

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This mixture is frequently referred to as an ‘amalgam’, which literally means an alloy of mercury and another metal, or metals.

Net ionic equation:

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The amalgam of sodium and mercury then flows into the next area of the cell, where the sodium and water combine to form sodium hydroxide

Full equation:

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Difficulties:

  • Technical: Although achieving a high degree of purity, this process is extremely costly, partic- ularly when taking into account the cost of replacing the mercury lost in the
  • Environmental: Mercury, if mishandled, clearly poses both environmental and general safety problems. Mercury can enter the food chain and continue its journey through organisms through a process known as

 

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Diaphragm Process

In this process, the anode is once again titanium, yet the cathode is an iron mesh.

At the anode, brine enters the cell from the side closest to the anode, where the chlorine ions are attracted to the anode (which is positive in an electrolytic reaction), where they become oxidised to form chlorine gas.

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At the cathode, sodium ions are drawn towards the cathode (which is negative in an electrolytic reaction), travelling through the asbestos diaphragm. However, given its relatively stable nature, these sodium ions are not reduced. Instead, water is reduced to form hydroxide ions and hydrogen gas.

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Net ionic equation:

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Full equation:

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Given the soluble nature of sodium hydroxide, water can be removed to crystallise out any remaining chlorine ions. Hydrogen and chlorine gas are tapped off separately.

Difficulties:

  • Technical: Purity is affected as it is extremely difficult to remove all traces of chlorine ions from the sodium hydroxide In contrast, in the mercury process, chlorine ions are separated before the sodium hydroxide is formed, therefore resulting in a product that is still relatively pure.
  • Technical: If the hydroxide ions come into contact with chlorine ions, hydrochlorous acid (HOCl) may be As such, care must be taken to prevent this.
  • Environmental: Asbestos is equally as problematic, if perhaps more controversial than, If asbestos fibers come into contact with humans, mesothelioma may result amongst other diseases.
  • Environmental: One potential product, the hypochlorite ion (ClO – ), is an extremely potent oxidant which can have devastating repercussions upon the

Membrane Process

This process is virtually identical to the diaphragm process. The one distinguishing feature is that where the diaphragm process makes use of an asbestos diaphragm, the membrane process relies upon a polytetrafluoroethylene (PTFE) membrane.

At the anode:

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At the cathode:

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Net ionic equation:

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Full equation:

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Difficulties:

  • Technical: As with the mercury process, the membrane process is rather costly given its reliance upon membranes formed by expensive petrochemical
  • Environmental: The petrochemicals used to produce the membrane can have an adverse effect upon the environment by way of the greenhouse

Remember- While the cathode half-equations may change, the reaction at the anode is always the same: Chlorine ions are oxidised to form chlorine gas.