Question 16: Which among the following is not a usage of standard hydrogen electrode?

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a. Used for pH measurement of a given solution.
b. Used in acid-base titrations.
c. Used as a primary reference electrode.
d. None of these.

Question 17: In calomel electrode the bottom of tube contains:

a. Hg
b. KCl solution
c. HCl solution
d. Both b and c

Question 18: In high energy density batteries Lithium is used as an electrode because:

a) it is reactive.
b) it does not corrode easily.
c) it has high negative reduction potential.
d) it is the lightest metal.

Question 19: A standard hydrogen electrode has zero electrode potential because:

a) hydrogen is the lightest element.
b) hydrogen is easiest to oxidize.
c) this electrode potential is assumed to be zero.
d) hydrogen has only one electron.

Question 20: The constitute of a Danlell cell is:

a) Cu - Ag cell.
b) Zn – Ag cell
c) Zn – Cu cell
d) None of these.

Question 20: Which is a secondary cell?

a) Mercury cell.
b) Dry cell.
c) H₂ – O₂ cell.
d) Ni – Cd cell.

Question 21: If the electrode potentials of both the electrodes become equal in magnitude but opposite in sign then:

a) an electrochemical cell stops working.
b) an electrochemical cell starts working.
c) an electrochemical cell is reversed.
d) none of these.

Question 22: …………. cell directly converts the free energy of a chemical reaction into electricity.

a) Lead storage battery.
b) Fuel cell.
c) Leclanche cell
d) Concentration cell.

Question 23: In an electrochemical cell:

a) potential energy changes into electrical energy.
b) chemical energy changes into electrical energy.
c) kinetic energy decreases.
d) potential energy decreases.

LABORATORY WORK

ELECTROLYSIS

Objective: to dentify the products of electrolysis . This experiment allows students to investigate the products formed by electrolysing various solutions.

This experiment enables students to carry out the electrolysis of various solutions and to investigate the identity of the products formed at the electrodes. They should be able to link their practical experiences with theory and learn how to construct simple ionic equations.

Laboratory conditions: It should be carried out in a well-ventilated laboratory as significant amounts of toxic chlorine, bromine and iodine can be produced in some cases, as well as highly flammable hydrogen.

Reagents: Access to the following aqueous solutions (all approx.
0.5 M concentration) potassium iodide KI, sodium sulfate Na₂SO₄, tin chloride SnCl₂, copper sulfate CuSO₄, 1% starch solution

Materials:

· Electrolysis apparatus (see diagram)

· Graphite electrodes (about 5 mm diameter),

· Large rubber bung to fit electrolysis cell, with holes to carry the graphite electrodes

· Small test-tubes (to fit over the electrodes),

· DC power supply (6 V)

· Small light bulb in holder (6 V, 5 W)

· Leads and crocodile clips

· Wooden spills

· Small pieces of emery paper

· Strips of Universal indicator paper

· Diposable plastic gloves

· Clamp and stand

Recommendations on the implementation of the laboratory experience and the use of electrolysis

Notes to the student :

1. The electrolysis apparatus shown below can be purchased ready-made. Alternatively, it can be made from thick glass tubing of 8 - 10 cm diameter, professionally cut into lengths of about 12 cm (fig. 1). A suitable glass bottle, with a wide-necked top and its base cut off, could be used instead (as shown in the diagram). The graphite rods should be well sealed into the holes, 2 - 3 cm apart, of the rubber bung, otherwise the electrolyte may leak onto the external wiring, causing it to corrode. Once made, this apparatus should last for several sessions, but the graphite rods tend to erode away quite quickly, particularly if students use larger than recommended voltages. The rods do eventually become thin and snap fairly easily, but they are cheap enough to replace.

2. Once copper (II) sulfate has been electrolysed (preferably last), a deposit of copper will have formed on the cathode. This has to be removed before the cells can be used again. Immersing the plated part of the electrode in a small quantity of 50% concentrated nitric acid (CORROSIVE) in a small beaker can be used to do this. Gloves and eye protection should be worn and the cleaning done in a fume cupboard by a suitably qualified person.

3. Depending on the volume of the electrolysis apparatus, each group of students needs enough solution to cover the electrodes plus about 2 cm to enable the full test-tubes of liquid to be inverted over the electrodes.

a) Clamp the electrolysis cell and pour in enough of the first electrolyte so that the tops of the electrodes are covered with about 1–2 cm of liquid. Fill the two test-tubes with the same electrolyte. Wearing gloves, close the end of each test-tube in turn with a finger and invert it over an electrode, so that no air is allowed to enter (see diagram). During electrolysis it may be necessary to lift the test-tubes slightly to ensure that the electrodes are not completely enclosed, preventing the flow of current.

b) Connect the circuit, and mark the polarity of each electrode on the bung. The circuit should be checked before being switched on.

c) Switch on the circuit, then:

· observe whether or not the lamp lights up;

· look for the substances produced at each electrode – ie gaseous, solid or in solution;

· write down results after each observation, not when all the experiments are finished.

d) Only carry out the electrolysis for long enough to make the necessary observations. Prolonging the electrolyses unnecessarily causes toxic gases such as chlorine and bromine to be produced in unacceptably hazardous quantities.

e) After each electrolysis switch off the current and remove the test-tubes from the cell to test any gases present by lifting them slowly in turn to let any remaining solution drain out before closing the end with a finger. Carry out the tests on the gases as instructed.

f) (Optional) After removing the test-tubes from the cell, quickly pour the liquid down the sink with plenty of water. Wipe a piece of Universal indicator paper over each electrode and note any colour changes.

g) Wash the cell with plenty of water and dry the outside with a paper towel before fixing it back into position and re-connecting the power supply. It is important to connect the leads according to the polarities marked on the bung.

h) Repeat the experiment with each of the other four solutions, trying to keep to the order given in the table. Zinc chloride and copper nitrate should be the last electrolytes tested. This is because they deposit solids on the cathode. If zinc chloride is electrolysed first, the solid deposit on the cathode can be easily removed with a piece of emery paper or dipping the end of the electrode in some dilute hydrochloric acid in a beaker.

Notes to the teacher: The electrolysis of aqueous solutions, rather than molten salts, is easier and safer for students to do for themselves, Unfortunately the theory is more complicated, because the presence of water complicates what students may decide are the products formed at the electrodes.

Ensure that students do not attempt to smell directly any of the halogen fumes produced. It is important that you are aware of any students who are asthmatic or who might have an allergic reaction to these toxic gases. In this context do not allow the electrolyses of the halide solutions to proceed any longer than is absolutely necessary.

When testing for hydrogen or oxygen, the mouth of the test-tube can be closed with a gloved finger, and the test-tube transported to a central area, where a single naked flame has been set up, well away from the experiments. A supply of spills can also be kept in this area for the tests.

For the hydrogen test, students may well ask why there is little or no ‘pop’ or ‘squeak’. Explain that pure hydrogen – rather than a mixture of hydrogen and air – is being tested if the test-tube was full of gas before it was removed.

For the oxygen test, care should be taken that the dampness at the mouth of the test-tube does not extinguish the ‘glow’, causing the test to fail.

Once the electrolysis of zinc chloride or copper nitrate has been done, a deposit of metal will have formed on the cathode. This will have to be cleaned before the cell can be used again. These metal deposits can be removed using emery paper. Alternatively, small quantity of 50% concentrated nitric acid (CORROSIVE) in a small beaker can be used to remove the copper, providing gloves are worn and the operation is done in a fume cupboard by a suitably qualified person. Similarly dilute hydrochloric acid will remove the zinc.

4. Results and conclusions:

Experimental results must be tabulated:

Lamp lights?

Na₂SO₄

SnCl₂

CuSO₄

Observations

anode (–)

cathode (+)

Test used for products

anode (–)

cathode (+)

Identity of product formed

anode (–)

cathode (+)

Write oxidation and reduction reactions occurring at the electrodes during electrolysis of these solutions by using rules of electrolysis in solution:

Anode (–): the anion which is stronger reducing agent (low value of standard potential) is liberated first at the anode	Cathode (+): the ion which is stronger oxidizing agent (high value of standard potential) is discharged first at the cathode
1. If anions of halogens (Сl^-, Вr^-and I^-) are present in the solution, halide ion is oxidized at the anode and the halogen is produced: Г^-n – ne = Г₂	1. If the cations of active metals are present in the solution (they stay above to hydrogen in the electrochemical series), at the cathode hydrogen ions are recovered: 2H⁺ +2e = H₂
2. If the solution contains no halide ions, such as NO₃^-, SO₄^2-,CO₃^2-, hydroxide OH^- ion is oxidized at the anode and oxygen is produced: OH^- – e = OH ° 4OH ° = 2H₂O + O₂	2. If in solution present cations of passive metals (they stay after hydrogen in the electrochemical series), at the cathode these cations will be recovered: Me⁺ⁿ +ne = Me °

For examples: a) electrolysis of silver nitrate solution

AgNO₃ => Ag⁺ + NO₃^-

HOH => H⁺ + OH^-

At Cathode (– ):	At Anode (+)
Silver ions, Ag⁺ ions are selectively discharged by receiving electrons to form Ag atom. This is because Ag⁺ ions has a lower position and high E° than H⁺ ions in the electrochemical series.	Hydroxide ions, OH are selectively discharged by donating electrons. This is because OH-ions has a lower position than SO₄^2- ions in the electrochemical series.
Ag⁺¹+1e = Ag ° H⁺	NO₃^- OH^- – e = OH ° 4OH ° = 2H₂O + O₂
· A silvery grey solid is deposited at cathode. · Silver, Ag metal is formed at the cathode.	· Colourless gas bubbles are released at anode. (The gas relights a glowing wooden splinter.) · Oxygen gas, O₂ and water, H₂O are produced at the anode.
4AgNO₃ + 2H₂O = 4Ag + O₂ + 4HNO₃

b) electrolysis of magnesium bromide solution:

MgBr₂ => Mg²⁺ + 2Br^-

HOH => H⁺ + OH^-

Cathode (– ):	Anode (+)
Hydrogen ions, H⁺ ions are selectively discharged by receiving electrons to form H₂. This is because H⁺ ions has a lower position and high E° than Mg²⁺ ions in the electrochemical series.	Bromide ions, Br- are selectively discharged by donating electrons. This is because in the first order halide ions are oxidized in solution
Mg²⁺ 2H⁺ +2e = H₂	2Br^-– 2e = Br₂ OH^-
· Colourless gas bubbles are released at cathode. · Hydrogen gas, H₂ is produced at the cathode.	· Brown vapor is released at anode. The gas turns the blue litmus paper to red. · Bromine gas, Br₂ is produced at the anode.
MgBr₂ + 2H₂O = H₂ + Br₂ + Mg(OH)₂

1^st Experiment. The electrolysis of salt potassium iodide solution performed in a U-shaped test tube with universal indicator.

1) The products of the electrolysis of the salt solution are all more hazardous than the starting materials. Hydrogen is EXTREMELY FLAMMABLE, iodine 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.

If the directions in the procedure notes are followed then very little iodine 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 (заглушки, пробки) because the products are gases.

Procedure:

a) Put about 75 cm³ distilled water into the beaker. Add about 2 heaped spatulas of potassium iodide.

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. This indicates the presence of iodine. 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 KI 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^-, K⁺ and I^-.

The negative ions are attracted to the positive electrode. The iodide ions are discharged (giving iodine) 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: KI => K⁺ + I^-

H₂O = > H⁺ + OH^-

At anode (+): 2I^– → I₂ + 2e^–

At cathode (–): 2H⁺ + 2e^– → H₂

In solution: K⁺ + OH^- = KOH

Overall equation: 2KI + 2H₂O = H₂ + I₂ + 2KOH

2^nd Experiment. The electrolysis of copper sulfate 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.

Copper (II) sulfate solution, CuSO₄(aq) – see CLEAPSS Hazcadr 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.

Procedure:

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