Wednesday, June 3, 2015

Electrolysis Of Zinc Chloride

This demonstration shows that an ionic salt will conduct electricity when molten but not when solid. Zinc chloride is used - this will melt at Bunsen burner temperatures.
To view a video clip of this demonstration experiment.

Zinc chloride offers a safer alternative to lead bromide for demonstrating the electrolysis of molten salts. Lead bromide decomposes to its elements just by heating without the need for electricity. The electrolysis of lead bromide must be carried out in a fume cupboard.
The electrolysis of zinc chloride should be carried out in a fume cupboard. The chlorine produced at the positive electrode is TOXIC and DANGEROUS FOR THE ENVIRONMENT.
There are quite long periods of waiting, including at least 15 minutes for the electrolysis to take place so if you have access to a webcam, or video camera and a data projector, this would enable students to see what is going on inside the crucible.
If not, bring students up in groups of 2 or 3 to view the experiment. They should note at which electrode bubbles are forming but must avoid smelling the bleachy smell ( be aware that many students are asthmatic). They should be able to see crystals of zinc around the negative electrode.
It is worth having another, related, activity for the class to be getting on with.
Solid zinc chloride (CORROSIVE, DANGER TO THE ENVIRONMENT)
Distilled water
Refer to Health & Safety and Technical notes section below for additional information.
Eye protection
Fume cupboard
Low voltage (0-12 V) powerpack and electrical leads
Graphite electrodes, 2, supported in an electrode holder or bung
Ammeter and/or bulb (in holder)
Circuit tester (optional)
Bunsen burner, tripod and heat resistant mat
Pipeclay triangle
Crucible
Clamp and stand
Metal spatula
Tongs
Plastic beaker
Filter paper and funnel
Indicator paper and/or starch-iodide paper 

Work in a fume cupboard. Wear eye protection. 
Zinc chloride, ZnCl2(s), (CORROSIVE, DANGEROUS FOR THE ENVIRONMENT) - see CLEAPSS Hazcard.
Chlorine, Cl2(g), (TOXIC, DANGEROUS FOR THE ENVIRONMENT) - see CLEAPSS Hazcard. Chlorine is a product of the electrolysis.
1 Set up apparatus as shown in diagram below.
Apparatus set-up

Setting up the electrolysis
a Set up a heat-resistant mat, tripod, Bunsen burner and pipeclay triangle. Put the crucible onto the pipeclay triangle, ensuring that it is sitting firmly and is in no danger of falling through.
b Set up the electric circuit with the power pack, ammeter and/or bulb and electrodes in series. Short the circuit at the electrodes with a key or the metal spatula. This is to satisfy yourself and the students that it is working.
c Clamp the electrodes so that they almost touch the bottom of the crucible but do not touch each other.
d Fill the crucible to within about 5 mm of the top with the powdered zinc chloride. As it melts the solid will shrink in volume as air escapes and it is important that the level of the molten salt does not drop below the level of the bottom of the electrodes. Ensure that the leads are well out of the way of the Bunsen flame. Using long electrodes can help with this.
Showing that the solid zinc chloride does not conduct electricity
a Begin to heat the crucible with a low to medium Bunsen flame. Watch the leads, and the bung if you are using one, to ensure that you are not over-heating them.
b The zinc chloride takes about 3 or 4 minutes to melt. It may be tempting to use a roaring Bunsen flame to speed up the melting, but if you do so the zinc chloride can form a crust over the top. This will prevent students from seeing what is going on, and the liquid salt may boil.
c As the salt melts, the bulb will light up and/or the ammeter will give a reading. Turn the Bunsen down a bit at this point. There will be some heating effect from the electric current which may be enough on its own to keep the zinc chloride molten (as in the industrial electrolysis of aluminium oxide.)
d Bubbles of gas will be seen at the positive electrode. The gas can be confirmed as chlorine by holding moist indicator paper close to the bubbles - it will go red and the edges may start to bleach. A more convincing test is to use moist starch iodide paper which will go black. It is also possible to see crystals of zinc forming on the negative electrode. These can form a bridge across the electrodes, effectively shorting them.
e Electrolyse the molten salt for about 15 minutes, with the current adjusted to about 0.5 A. Check every few minutes that the current remains roughly constant as there is a tendency for it to slowly increase.
f After 15 minutes, turn off the power pack and Bunsen burner and remove the electrodes from the crucible. If this is not done while the salt is still molten the electrodes will stick.
g Leave the crucible to cool for about 10 minutes. You may be able to see zinc crystals on both the electrode and on the surface of the mixture in the crucible. You could stop at this point, but to convince students that a metal really has been made you can separate the zinc from the remaining zinc chloride.
Separating the zinc
a When the crucible is cool to the touch, put it into a beaker of distilled water. (If the water is at all basic like most hard tap waters, the zinc ions will flocculate forming large particles which are far harder to remove from the zinc metal.) The zinc chloride will dissolve (which may take some time) and can be decanted off. Swirl the beaker which will cause the zinc metal to concentrate in the centre of the beaker and decant off most of the liquid.
b Filter the remainder and show students the shiny pieces of metal left on the filter paper. Dry the pieces of metal carefully between further sheets of filter paper and then test with a circuit tester to prove that you have a metallic product.
Given that the starting material was zinc chloride and you have made chlorine, most students will have little difficulty in accepting that the metal is zinc.

If the bulb does not light when testing the circuit, and the electrodes are mounted in a bung, check that the electrodes are not cracked.
The boiling point of zinc chloride is about 730°C. This can easily be reached by the combination of the heat from the Bunsen burner and the electric current. If the zinc chloride does begin to boil, it can boil over from the crucible and will also produce fumes of zinc chloride in the air. These rapidly turn back to the solid, forming a fine powder.
It is possible to confuse the boiling bubbles with those of chlorine gas being formed. Therefore do not heat the molten zinc chloride too strongly.
Do not try to remove the Bunsen and cool the salt while still electrolysing it, in order to show that the salt only conducts when molten. The heating effect of the electric current will keep the salt molten for several minutes, and when it does cool, a crust forms which is very difficult to re-melt.
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Electrolysis Of Copper (II) Sulphate Solution

This experiment enables students to carry out the electrolysis of copper(II) sulfate solution and to link their findings with the industrial electrolytic refining of copper.
 

This class experiment can be done by students working either in pairs or threes.
CHEMICALS
Aqueous copper(II) sulfate, about 0.5 M, 200 cm3
Copper strips (optional), 2 (Note 3)
Small pieces of emery paper
Refer to Health & Safety and Technical notes section below for additional information.
APPARATUS
Eye protection
Each group of students will require:
Beaker (250 cm3)
Graphite electrodes (about 5 mm diameter), 2
Retort stand and clamp to hold electrodes (Note 1)
DC power supply (6 volt)
Light bulb (small, 6 volt, 5 watt) - optional
(Note 2)
Leads and crocodile clips 

Wear eye protection. Students must wash their hands at the end of all practical work. 
Copper(II) sulfate solution, CuSO4(aq) - see CLEAPSS Hazcard and CLEAPSS Recipe Book. At the suggested concentrations, the copper(II) sulfate solution is LOW HAZARD If the concentrations are increased, the solutions must be labelled with the correct hazard warnings. Copper(II) sulphate solution is HARMFUL if concentration is equal to or greater than 1M.
1 There are several ways of securing the graphite electrodes. Using a retort stand and clamp is probably the most convenient. They can also be fixed using Blutac on to a small strip of wood resting on the top of the beaker.
2 A bulb can be included in the circuit to indicate that there is a flow of current.
3 As an extension to the basic experiment, strips of copper can be used in place of the graphite rods.

a Ask the students to set up the cell as shown. They should watch for any activity on each of the electrodes, and write down their observations.
The cathodes can be cleaned using emery paper.
Apparatus set-up

Students should see a deposit of copper forming on the cathode. This will often be powdery and uneven. You should explain that, if the current used is much lower, then the solid coating is shiny, impermeable and very difficult to rub off; this process forms the basis of electroplating.
Bubbles of gas (oxygen) are formed at the anode.
Cathode reaction: Cu2+(aq) + 2e- → Cu(s)
Anode reaction: 2H2O(l) → O2(g) + 4H+(aq) + 4e-
With carbon (graphite) electrodes, the oxygen usually reacts with the anode to form CO2. If copper is used for the electrodes, the copper anode dissolves. The reaction is the reverse of the cathode reaction.
The results of this experiment can lead to a discussion about electroplating and the electrolytic refining of copper.
It can be instructive to allow students to copper-plate metal objects supplied by the school and previously tested for their suitability. Personal items should not be used. In many cases, an alternative redox reaction often takes place before any current is actually passed. This happens for instance in items made of metals above copper in the reactivity series. It is wise not to complicate electrolytic deposition with chemical displacement - valued articles can be effectively ruined.
Extension experiments for copper refining
1 After doing the electrolysis as described above, the electrodes can be interchanged. Students can then see the copper disappearing from the surface of the copper-coated anode
Cu(s) → Cu2+(aq) + 2e-
This leads to a discussion as to why, during electrolytic refining:
  • the anode consists of an unrefined sample of the metal
  • the cathode is made of pure copper or a support metal such as stainless steel.
     
2 The electrolysis can be done using two weighed copper strips. This is to confirm that the mass gained at the cathode is equal to the mass loss at the anode.
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Colourfull Electrolysis

An interesting introduction to the electrolysis of brine (sodium chloride solution). Students use Universal Indicator to help them follow what is happening during the reaction.
 

This experiment works well if students are directed to make detailed observations and then attempt to explain for themselves what they think is happening.
The main issue is likely to be the availability of sufficient U-shaped test tubes.
Chemicals
Sodium chloride (table salt) (Note 1)
Universal Indicator solution (FLAMMABLE)
Distilled water (Note 2)
Refer to Health & Safety and Technical notes section below for additional information.
Aparatus
Eye protection
Each working group will require:
U-shaped test tube
Clamp and clamp stand
Carbon electrodes and electrode holders, 2 of each (Note 3) 
Electrical leads, 2
Power pack (low voltage, d.c.)
Beaker (100 cm3)
Spatula
Stirring rod

Hydrogen, H2(g), (HIGHLY FLAMMABLE) - see CLEAPSS Hazcard.
Chlorine, Cl2(g), (TOXIC, DANGEROUS FOR THE ENVIRONMENT) - see CLEAPSS Hazcard 
Sodium hydroxide, NaOH(aq),  (CORROSIVE) - see CLEAPSS Hazcard 
1 The products of the electrolysis of the salt solution are all more hazardous than the starting materials. Hydrogen is EXTREMELY FLAMMABLE, chlorine is TOXIC and DANGEROUS FOR THE ENVIRONMENT, and sodium hydroxide is CORROSIVE. Ensure that the current is turned off a soon as a trace of chlorine is detected.  Chlorine (TOXIC, DANGEROUS FOR THE ENVIRONMENT) can be a problem for asthmatic pupils. If the directions in the procedure notes are followed then very little chlorine is produced. Sodium hydroxide is CORROSIVE. Ensure that students wear eye protection, especially when they are clearing up the experiment. 
2 If distilled water is a problem, then tap water could be used. But it may affect the colours produced, especially in areas with hard water.
3 If electrode holders are not available, another suitable means of securing the electrodes could be used. Do not use bungs because the products are gases.

apparatus set-upa Put about 75 cm3 distilled water into the beaker. Add about 2 heaped spatulas of sodium chloride.
b Stir until the salt dissolves. Then add several drops of Universal Indicator solution. Stir to mix thoroughly. You need enough indicator to give the water a reasonable depth of green colour.
c Pour coloured salt solution into the U-shaped test tube and clamp it as shown in the diagram.
d Wash the carbon electrodes carefully in distilled water and then fix them so that there is about 3 cm of electrode in each side of the U-tube – see diagram. This is most easily done using electrode holders.
e Attach leads and connect to a power pack set to 10 V (d.c.).
f Turn on the power pack and observe closely what happens. A piece of white paper held behind the U-tube can help. Make sure the U-tube is kept very still during the experiment.
g Turn off the power as soon as you notice any change at the positive electrode, or when you smell a ‘bleachy, swimming pool’ smell. This will probably take less than 5 minutes.

This experiment is an interesting introduction to the electrolysis of brine. It is probably best not used as the first electrolysis that students encounter. They would really struggle to explain for themselves what is going on. It could be followed by the electrolysis of salt solution in industry.
Students should be able to notice bubbles of gas at each electrode. At the positive electrode, the indicator turns red initially, and is then bleached to colourless. This indicates the presence of chlorine. At the negative electrode the indicator turns purple. The remainder of the solution stays green.
The product at the negative electrode is hydrogen. This can be difficult for students to understand.
Some of the water will ionise, that is, turn to hydrogen (H+) and hydroxide (OH-) ions.
When the sodium chloride is dissolved in water, the ions forming the ionic solid separate out. This means that there are actually 4 ions present in the solution: H+, OH-, Na+ and Cl-.
The negative ions are attracted to the positive electrode. The chloride ions are discharged (giving chlorine) in preference to the hydroxide ions. These are left behind in solution.
At the negative electrode, the hydrogen ions are discharged (producing hydrogen gas) in preference to the sodium ions. These are also left behind in solution. Thus sodium hydroxide solution remains. This is the cause of the purple colour of the indicator at the negative electrode.
In time, the green colour of the indicator in the middle would change too, as the ions diffuse through the resulting solution.
Equations:
2H+ + 2e– → H2 [negative electrode, cathode]
2Cl– → Cl2 + 2e– [positive electrode, anode]
H2O → H+ + OH–
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Chemicals From Seawater

 This experiment is a simple one to carry out and is designed to show that seawater contains a mixture of different salts. It can be used in conjunction with Earth Science topics linked to the conveyance of mineral salts into the sea via rivers.
This experiment could be carried out early in a science course while using scientific apparatus is still a novelty, because it provides practice at using a variety of equipment and also introduces the concept of a mixture and some simple chemical tests. It is also an opportunity to reinforce safety messages eg wearing eye protection throughout and activity and the need to keep standing up.
 

The experiment can be carried by groups of two or three and will take about one hour to complete.
Seawater, 200 cm3 (Note 1)
Access to hydrochloric acid, 1 M 
Refer to Health & Safety and Technical notes section below for additional information.
Eye protection
Beaker (250 cm3)
Beaker (100 cm3)
Conical flask (100 cm3)
Filter funnel
Filter paper
Bunsen burner
Heat resistant mat
Tripod
Gauze
Teat pipette

Wear eye protection throughout. Take care with hot apparatus and solutions. 
Hydrochloric acid, HCl(aq) - see CLEAPSS Hazcard and CLEAPSS Recipe Book.
Limewater, Ca(OH)2(aq), (treat as IRRITANT) - see CLEAPSS Hazcard and CLEAPSS Recipe Book. 
Calcium sulfate, hydrated, CaSO4.2H2O(s) - see CLEAPSS Hazcard.
Sodium chloride, NaCl(s) - see CLEAPSS Hazcard. 
1 Although it is tempting to use genuine seawater if available, this experiment is usually more successful if the seawater is generated artificially. Genuine seawater will not always yield a full range of solids in sufficient quantities to be detected. Artificial seawater can be generated by adopting the following procedure:
  • Bubble carbon dioxide through a mixture of 250 cm3 limewater plus 750 cm3 deionised water for about 20 minutes or until the cloudy precipitate disappears completely.
  • Filter.
  • Add as much solid hydrated calcium sulfate as will dissolve.
  • Add about 15 g of sodium chloride.
  • Stir until all the solid has dissolved, leave to settle and then decant the liquid if necessary.

a Place 200 cm3 of seawater in a 250 cm3 beaker.
b Heat and boil the seawater.
Apparatus set-up
c Stop heating when about 60-70 cm3 of liquid remains. Solid will be precipitated during this evaporation process.
d Allow to cool and for any solids to settle.
e Pour the clear liquid into the 100 cm3 beaker, leaving the solids behind.
f Add a few drops of hydrochloric acid to the solids left behind and observe what happens.
g Put the 100 cm3 beaker on the tripod and gauze and heat the liquid until another solid appears. This will occur when about 30-40 cm3 of liquid remains.
h Carefully filter the liquid into the conical flask.
i Wash out the 100 cm3 beaker and pour the filtrate into the beaker.
j Boil the liquid yet again until there is almost none left.
k Let it cool and note what is observed.

Encourage the students to write down what they observe at each stage.
Teachers may wish to mark 250 cm3 beakers at the 60 cm3 level if there are no gradations already present.
The artificial seawater contains calcium hydrogencarbonate owing to the reaction of limewater with excess carbon dioxide:
Ca(OH)2 (aq) + 2CO2(g) → Ca(HCO3)2(aq)
When this solution is boiled it soon precipitates calcium carbonate:
Ca(HCO3)2(aq) → CaCO3(s) + CO2(g) + H2O(l)
This is the identity of the predominant solid first left behind when the liquid is boiled. However, some calcium sulfate will also be present.
When hydrochloric acid is added to this solid, students should observe effervescence (fizzing), since the calcium carbonate is producing carbon dioxide gas:
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)
The solid which continuously crystallises out on further evaporation is sparingly soluble calcium sulfate, which is the predominant solid filtered off when 30 cm3 of sea water remains.
The more soluble sodium chloride precipitates out during the final stages of evaporation.
Student questions
Here are some questions for students.
a What evidence is there that seawater is a mixture of salts?
b What gas is likely to have been evolved when hydrochloric acid is added to the solids first collected?
c What does this tell you about the identity of these solids?
d Research the web to try to find information about the solubilities of sodium chloride and calcium sulfate – two common compounds present in seawater. Use this information to predict the possible identity of the very last solid left at the end of your experiment.
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