A) Experiment with inert carbon electrodes.

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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.

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.

Equation: CuSO₄ => Cu²⁺ + SO₄^2-

H₂O => H⁺ + OH^-

At anode : OH^- –e = OH°

4OH° → 2H₂O + O₂(g)

At cathode: Cu²⁺(aq) + 2e^- → Cu(s)

In solution: 2H⁺ + SO₄^2- = H₂SO₄

Overall equation: 2CuSO₄ + 2H₂O = 2Cu + O₂ + 2H₂SO₄

Notes: With inert 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.

B) Experiment with active (copper) electrodes 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.

3) Quantitative electrolysis of aqueous copper (II) sulfate. This demonstration experiment shows how to find the value of the Faraday constant using electrolysis.

This demonstration is designed to find the value of the Faraday constant – the amount of electric charge carried by one mole of electrons - from the electrolysis of aqueous copper (II) sulfate solution, using weighed copper electrodes.

4) This is a relatively straightforward demonstration to set up and carry out, but there is not much to see while it is taking place. This provides an ideal opportunity to ask students to prepare a suitable results sheet and to explain how the masses of the electrodes, the current passed and the time elapsed can be used to calculate a value for the Faraday constant. The experiment will take around 40 minutes.

5) Students can be told that the calculations are relevant for industrial applications of electrolysis, for example in the extraction of elements such as sodium and chlorine, and in the electroplating industry.

Procedure:

These should be 2 - 3 cm wide and long enough to reach from the bottom of a 250 cm³beaker to the rim, where they can be folded over and gripped by the crocodile clips on the leads.

a) Clean the copper electrodes with emery paper, rinse under the tap and dry thoroughly using paper towels and a hot-air blower.

b) Mark the electrodes as + or – at one end, and weigh them separately. Record the masses.

c) Set up the circuit as shown in the diagram, clamping the electrodes to the rim of the beaker using the crocodile clips. Make sure the electrodes do not touch.

d) Fill the beaker with copper (II) sulfate solution to just below the crocodile clips.

e) Start the stop clock and switch on the current, setting the rheostat so that a current of 0.50 A passes through the solution.

f) The current alters slightly throughout the electrolysis, so it is important to make continual adjustments to the rheostat to keep the current constant.

g) After about 30 min switch off the current and stop the clock.

h) Remove the electrodes from the electrolyte, wash them carefully under running water, then rinse them in a beaker of propanone in a fume cupboard. Finally dry them by allowing the propanone to evaporate in a well-ventilated laboratory, and away from any naked flames.

i) Re-weigh the dry electrodes.

The results should be entered in a table, for example:

Mass of cathode (-) before /g
Mass of cathode (-) after /g
Change in cathode's mass/g
Mass of anode (+) before /g
Mass of anode (+) after /g
Change in anode's mass /g
Current used /A
Time for electrolysis /s

Notes: It is difficult to avoid fluctuations in current throughout the electrolysis, and thus, obtain an accurate value to use in the calculations.

Be careful to avoid any movement of the electrodes during the electrolysis.

Thorough washing and drying of the electrodes is vital. A tiny amount of moisture left on the electrodes causes a considerable error in the mass calculations. The electrodes should NOT be heated in a flame to dry them as the copper will oxidise on the surface to form black copper (II) oxide. After initial drying by evaporation of the propanone, a hot-air blower could be used.

The student questions work through the process of calculating the Faraday constant. Students could also be asked to use the official value of 96500 Coulombs per mole to calculate the percentage error.

The calculated value for the Faraday constant should be correct to within about 10%. The major error comes from the value for the current, which is hard to keep constant to two significant figures.

The reaction at the cathode is: Cu²⁺(aq) + 2e^– → Cu(s)

The concentration of Cu²⁺ ions in solution remains constant. In a given time as many of these ions are produced at the anode as are being removed at the cathode.

The coating of solid copper on the cathode tends to be ‘spongy’ and some can be lost when electrode is washed or dried. Some may also have fallen off the cathode during electrolysis. Thus, the mass gain of the cathode is an unreliable measure of the quantity of copper actually deposited there.

3^rd Experiment. The electrolysis of tin (II) chloride solution

Tin 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 tin chloride should be carried out in a fume cupboard. The chlorine produced at the positive electrode is toxic and dangerous for the environment.

Tin (II) chloride, sncl₂(s) is corrosive, dangerous for the environment.

Chlorine, Cl₂(g), is toxic, dangerous for the environment. Chlorine is a product of the electrolysis.

Procedure:

Set up apparatus as shown in diagram as in 1^st experiment.

Bubbles of gas (oxygen) are formed at the anode, bubbles of gas (hydrogen) are formed at the cathode and precipitate of tin hydroxide is formed in solution.

Equation: SnCl₂ => Sn²⁺ + 2Cl^-

H₂O = > H⁺ + OH^-

At anode (+): 2Cl^– → Cl₂ + 2e^–

At cathode (–): 2H⁺ + 2e^– → H₂

In solution: Sn²⁺ + 2OH^- = Sn(OH)₂

Overall equation: SnCl₂ + 2H₂O = H₂ + Cl₂ + Sn(OH)₂

4^th Experiment. The electrolysis of sodium sulfate solution

The electrolysis of an aqueous solution of sodium sulphate using inert electrodes produces hydrogen at the cathode and oxygen at the anode and a neutral solution of sodium sulphate remains unaltered by the electrolysis.

Equation: Na₂SO₄ => 2Na⁺ + SO₄^2-

H₂O => H⁺ + OH^-

At anode : OH^- –e = OH°

4OH° → 2H₂O + O₂(g)

At cathode: 2H⁺ + 2e^- → H₂

In solution: 2Na⁺ + SO₄^2- = Na₂SO₄

The overall cell reaction is: Na₂SO₄ + 2H₂O = 2H₂ + O₂ + Na₂SO₄

If the reaction is carried out in a Hofmann Voltammeter, with some universal indicator in the solution, it will be noticed that around the cathode the solution becomes alkaline and around that anode the solution becomes acidic. This is explained as follows:

At the cathode: Hydrogen ions are being removed from solution, thereby leaving an excess of hydroxyl ions which makes the solution alkaline, and

At the anode: Hydroxyl ions are being removed, so leaving an excess of hydrogen ions which makes the solution acidic.

Procedure:

1) Prepare a very concentrated (it need not be saturated) solution of sodium sulfate in water. Add a generous portion of bromothymol blue solution to the aqueous sodium sulfate solution. If it is necessary add either some sulfuric acid solution or some sodium hydroxide solution in order to keep the electrolyte light green in color. Sodium sulfate solution should color the indicator green (pH = 7). The addition of acid or base may be necessary to counteract local water and/or air conditions.

2) Any DC power supply will work for this experiment.

3) The electrodes for this experiment must be platinum. Many of the commercially available platinum electrodes look black. The black coating is a layer of finely divided platinum (platinum black) deposited on the electrodes in order to increase the surface area of the platinum and thus increase the current that can flow through the system.

4) For this experiment the glassware must be very clean. I keep the Hoffman apparatus filled with distilled water until just before use. After the electrolysis has continued for sufficient time that measurable quantities of gas are produced, stop the reaction. Be careful when discussing the observations that you do not use terms that have not yet been defined in class.

5) Interpret the observations and on the blackboard write the balanced half-reactions for the processes that took place at each of the electrodes.

6) On the basis of the observations and the equations that have been written by the students, ask them to predict the color that will be produced when the cell is inverted into the beaker that was originally used for filling the cell. The students will look at the two equations, and because they have not taken account of the different volumes of hydrogen and oxygen that are produced, they will predict that the resultant solution will be acidic and the indicator will turn yellow. Based on what has been written on the board this is the only prediction that can be made.

7) Pour the contents of the Hoffman apparatus back into the beaker. You must be very careful to get every drop of electrolyte back into the beaker. The reaction is very sensitive and the loss of only a drop of the electrolyte may have a very negative effect on the reaction results. The students are amazed that the resultant solution is green. Use the results of this reaction to define terms such as anode, cathode, oxidation and reduction. The idea that the number of electrons lost at the cathode and gained at the anode are equal will "fall" out of the experiment.

8) Graphite (carbon) electrodes cannot be used in this experiment because the oxygen produced is adsorbed on the surface of the electrode (anode) and the net results will not show the quantitative 2:1 ratio of hydrogen to oxygen that is expected.

9) The choice of electrolytes and acid-base indicators is very limited. The pH of the solution must be 7.0, and the indicator must change through three colors, the middle one at pH 7.0.

5^th Experiment. Preferential discharge of cations during electrolysis

This experiment shows the order in which cations are discharged during electrolysis.

This experiment is designed to show that metal cations are preferentially discharged, in relation to the position of the metal in the reactivity series.

There is probably insufficient time for each group to do all six electrolyses. It is probably best to assign two metal ions and one mixture to each group, and pool the results from all the groups.

Reagents:Copper (II) sulfate, about 0.5 M, 200 cm³; Iron (II) sulfate, about 0.5 M, 200 cm³; Zinc sulfate, about 0.5 M, 200 cm³ (IRRITANT, DANGEROUS FOR THE ENVIRONMENT); Nitric acid, about 4 M (CORROSIVE), 20 cm³; Aqueous ammonia, about 4 M (IRRITANT), 10 cm³; Optional: Solution of mercury (II) chloride (VERY TOXIC) mixed with ammonium thiocyanate (HARMFUL)

Materials :Beaker (250 cm³); Beaker (100 cm³), 2 Boiling tube, Test-tubes, 2 or 3, Teat pipette, Bunsen burner, Heat resistant mat, Tripod and gauze, Platinum electrodes, 1 cm square, with platinum leads sealed through glass tubes, both supported in a rubber bung or cork so that the electrodes are about 2 cm apart, DC power pack for supplying about 3–4 V; Light bulb (5 V) and holder; Several lengths of connecting wire, including two fitted with crocodile clips; Emery paper.

Procedure:

1) Platinum electrodes work best – alternatives include graphite.

2) If it is required, to test for zinc in the presence of copper (see teaching notes), the solution containing mercury (II) chloride and ammonium thiocyanate can be prepared by dissolving 2.7 g of mercury (II) chloride and 3 g of ammonium thiocyanate in 100 cm³ of water. The mixed solution at this concentration should be labelled TOXIC.

a) Half-fill the 250 cm³ beaker with water, and heat with a Bunsen burner until boiling. Stop heating and carefully transfer the beaker onto a heat resistant mat.

b) Half-fill the 100 cm³ beaker with the copper (II) sulfate solution.

c) Clamp the electrodes so that they dip into the copper (II) sulfate solution and connect up the rest of the circuit as shown in the diagram, with the power supply switched off.

d) Switch on the current. The bulb should light up. Set the voltage to about 3–4 V, and electrolyse the solution for about 2 min.

e) In the meantime, use the teat pipette to transfer no more than 3–4 cm³ of nitric acid into the boiling tube, and then place this in the beaker of hot water.

f) Switch off the current and remove the electrodes.

g) Place the electrodes under running water, ensuring that all parts that have been in contact with the solution have been washed thoroughly.

h) Dip the electrodes one at a time into the hot nitric acid, leaving them in the acid for about 10 seconds. If the electrodes do not reach into the acid or are difficult to remove from the bung, pour the hot acid into a small beaker first.

i) Test the solution of nitric acid for the presence of copper ions, using the test described at the end of the procedure.

j) Wash the electrodes, the contents of the boiling tube and the 100 cm³ beakers using tap water.

k) Repeat steps b–j using the iron (II) sulfate solution, and then the zinc sulfate solution, using the tests for iron (II) ions and zinc ions in each case.

l) Finally use mixtures of two electrolytes and do the tests for the appropriate metal ions to find out which ion is discharged in preference to the other. Suitable combinations are:

· copper (II) sulfate and iron (II) sulfate

· iron (II) sulfate and zinc sulfate

· copper (II) sulfate and zinc sulfate

Tests for metal ions

Copper ions (Cu²⁺): Using the teat pipette, transfer five drops of the solution into a test-tube. Wash out the teat pipette and then use it to add the aqueous ammonia drop by drop. Shake the test-tube thoroughly throughout the addition. Once all the nitric acid has been neutralised, a pale blue precipitate of copper (II) hydroxide forms initially, but with excess ammonia, this precipitate dissolves and is replaced by a dark blue solution:

CuSO₄ + 2NH₄OH = Cu(OH)₂¯ + (NH₄)₂SO₄

blue precipitate

Cu(OH)₂ + 4NH₄OH = [Cu(NH₃)₄](OH)₂+4H₂O

dark blue complex

Iron (II) ions (Fe²⁺): Using the teat pipette, transfer five drops of the solution into a test-tube. Wash out the teat pipette and then use it to add the aqueous ammonia drop by drop. Shake the test-tube thoroughly throughout the addition. Once all the nitric acid is neutralised, a dirty, dark green precipitate of iron (II) hydroxide eventually forms.

FeSO₄ + 2NH₄OH = Fe(OH)₂¯ + (NH₄)₂SO₄

dark green precipitate

Zinc ions (Zn²⁺): Using the teat pipette, transfer five drops of the solution into a test-tube. Wash out the teat pipette and then use it to add the aqueous ammonia drop by drop. Shake the test-tube thoroughly throughout the addition. Once all the nitric acid is neutralised a white precipitate of zinc hydroxide forms, but this then disappears and is replaced by a colourless solution. If the ammonia is added too quickly the white precipitate will not be seen.

ZnSO₄ + 2NH₄OH = Zn(OH)₂¯ + (NH₄)₂SO₄

white precipitate

Zn(OH)₂ + 2NH₄OH = (NH₄)₂ZnO₂+4H₂O

Notes: However tempting might be, electrolysing solutions of other metal ions give poor results in practice. This is because the metals concerned tend not to adhere to the platinum cathode very effectively. However, copper, iron and zinc usually work very well.

It is essential that the washing process described in g is done thoroughly, otherwise the solution to be tested will contain the original ions, in addition to those derived from the electrolysis.

There are grounds for criticising these tests, as there is the possibility of two metals being deposited on the cathode. For example, during the electrolysis of a mixture of copper ions with zinc ions, if there is any zinc deposited on the cathode in addition to copper, the test described for copper will mask that for zinc. More complex test reactions for zinc ions do exist, producing definitive coloured products, but they would not be understood by the students.

In theory, and indeed in practice, no such mixture of ions should be present, because the experiment is designed to show preferential electrolytic discharge.

However, if you want to know how to test for zinc in the presence of copper, you can use the following technique. Transfer a few drops of the solution to a test-tube using a teat pipette. Add a few drops of the mercury(II) chloride-ammonium thiocyanate reagent. A yellow precipitate is formed by copper alone, but if zinc is also present the colour is darker.

Obviously you need to try this out beforehand so that the test can be described and demonstrated with a prior knowledge of the colour changes involved. During the class experiments it may be helpful to have both colours available on display to enable students to compare their results with yours.

Questions to the student:

1) For each of the individual solutions being electrolysed, on which electrode (positive or negative) is a metal deposited?

2) Write an equation for the electrical discharge of each metal ion: Cu²⁺, Fe²⁺, Zn²⁺.

3) When mixtures of electrolytes are used, is a mixture of metals deposited at an electrode, or is it a single metal? If so, which?

4) What conclusion can you make about the priority of metal ions discharged during electrolysis?

EXERCISES:

1) Calculate the number of moles of hydrogen released when 5 amps of current passes for 3000 seconds through a solution of sulphuric acid.

2) Calculate the number of coulombs needed to deposit 6.35g of copper at the cathode in an electrolysis of CuCl₂ molten. Calculate the time that a current of 4 amps must pass to deposit this mass of copper.

3) A current of 0,5 ampere was passed through a solution of AgNO₃ for one hour. Calculate the masses of substances deposited on the electrodes.

4) When one Faraday or 96500 Coulombs of electricity is passed through silver nitrate solution, 108 gram of silver are deposited. Calculate electrochemical equivalent of silver.

5) How many grams of oxygen is liberated by the electrolysis of water after pass in of 0,0565 ampere for 185 sec.

6) Perform test:

№	Question	Variant
1	A device in which electric current is produced at the expense of spontaneous chemical reaction is:	a) voltaic cell b) voltameter c) galvanic cell d) both a and c
2	Which among the following is not true for electrochemical cell?	a) It consists of a battery. b) It consists of generally two electrolytes. c) Anode acquires negative charge. d) It needs a porous partition.
3	Among the following which is not a true statement for Faraday’s laws of electrolysis?	a) The weight of substance deposited is directly proportional to the quantity of electricity passed. b) The weight of substance deposited at the respective electrodes are directly proportional to their equivalent weights. c) Both a and b d) None of these
4	The quantity of electricity required to liberate or deposit 1 gram equivalent of a substance from its solution during electrolysis is known as:	a) Faraday b) Ampere c) Coulomb d) None of these
5	Salt bridge is an inverted U-glass tube containing saturated solution of:	a) KCl b) NH₄NO₃ c) K₂SO₄ d) Any one of these
6	In a redox reaction if the emf is positive then the reaction is:	a) Non-spontaneous b) Spontaneous c) Both a and b d) None of these.
7	For the following reduction reaction how many coulombs are required? 1 mol of Cu²⁺to Cu	a) 193000 C b) 96500 C c) 482500 C d) 386000 C
8	Find the valency of the metal whose atomic weight is 96 and 0.3605g of a metal is deposited on the electrode by passing 1.2 ampere current for 15 minutes:	a) 1 b) 2 c) 3 d) None
9	For a given concentration of a solution with increase in temperature:	a) electrolytic conduction increases, metallic conduction decreases. b) electrolytic conduction decreases, metallic conduction increases. c) both electrolytic and metallic conduction increases. d) Both electrolytic and metallic conduction decreases.
10	When electrolysis of aqueous NaCl solution is carried out, the products are:	a) Na at cathode and Cl₂at anode. b) H₂ at cathode and O₂at anode. c) Na at cathode and O₂at anode. d) H₂ at cathode and Cl₂at anode.
11	…………….. is the amount of electricity required to deposit 1 mol of aluminium from a solution of AlCl₃ .	a) Faraday b) 1 ampere c) 3 Faradays d) 3 Ampere
12	An aqueous solution of Na₂SO₄ is electrolysed using platinum electrodes. The products at the anode and the cathode are:	a) O₂, Na b) SO_4-, H₂ c) SO_4-, Na d) O₂, H₂
13	Due to the presence of …………… NaCl conducts electricity.	a) free ions b) free atoms c) free molecules d) free electrons
14	Which one of the following is a false statement?	a) Salt bridge maintains electrical neutrality b) When salt bridge is removed then the potential of the cell drops to zero. c) Salt bridge increases the emf of the cell. d) Salt bridge connects two half cells.
15	A current of x A flowing for 10 min deposits 3g of the metal which is monovalent. The atomic mass of the metal is 50. Find the value of x?	a) 10.5 b) 9.65 c) 8.75 d) 6.72
16	The chemical equivalent of the metal is 19.3. When 3A of electric current is passed for 50 min. what will be the amount of metal deposited from an electrolyte?	a) 1.8 g b) 0.9 g c) 3.6 g d) 4.2 g
17	NaCl in solid form does not conduct electricity because:	a) It does not ionize b) It ionizes c) It is present in fused form d) None of these.
18	When 10A current is passed for 96.5 sec for electrolysis of Al (NO₃)₃. Find the number of aluminium atoms deposited at cathode?	a) 2 x 10²³ b) 2 x 10²⁴ c) 2 x 10²² d) 2 x 10²¹
19	X ampere current is passed for y seconds through copper and silver voltameters. The metal that is deposited more is:	a) Ag b) Na c) Cu d) Al
20	A current of 2 ampere was passed through solutions of CuSO₄ and AgNO₃. The weight of copper deposited is 0.636 g. What is the weight of silver deposited?	a) 3.16g b) 4.32g c) 2.61g d) 2.16g
21	A certain quantity of electricity was passed through CuSO₄ solution for 6 minutes 20 sec. The amount of Cu deposited is 0.317g. (At.wt.of Cuis 63.5 and Faraday = 96500 coulombs). Find the amount of electricity consumed in ampere?	a) 25.0 b) 27.3 c) 2.54 d) 2.73
22	1 C contains ……………… number of electrons.	a) 6.24 x 10¹⁸ b) 6.03 x 10¹⁶ c) 6.05 x 10²¹ d) 6.24 x 10¹⁷
23	What is the correct efficiency, if an electricity of 10800C decomposes 14.25 gm of AgNO₃?	a) 100% b) 50% c) 25% d) 75%
24	Calculate the time required for a current of 2.5 A to obtain 1% coating of Ag on a tea-pot weighing 0.27 Kg.	a) 965 sec b) 40 sec c) 1930 sec d) 486 sec.

LABORATORY WORK

METALS

Objectives: 1) Students understand the terms like metals, metal displacement reaction, reactivity series, etc.

2) Students acquire skills to perform & visualize the reactions of Al, Zn, Fe and Cu with the following salt solutions:

· Aluminium sulphate

· Zinc sulphate

· Ferrous sulphate

· Copper sulphate

3) Students can analyze the meaning of reactivity series of metals based on the inferences from the experiment.

4) Students acquire skills to perform an experiment to determine the reactivity of metals in salt solutions.

Materials and reagents: strip of metals – Zn, Cu, Al, Fe; solutions – ZnSO₄, CuSO₄, Al₂(SO₄)₃, FeSO₄, four 100 mL beakers, forceps

Precautions:

· Do not touch any chemical.

· Label the beakers properly and put them in sequences. It is better to label the beakers both with the chemical names and also by chemical symbols.

· Till the reaction goes on, keep the beakers at a safe place. Cover them so that insects or dust may not pollute the solutions.

· It is very important that the apparatus you use must be very clean and dry.

· Always clean the metallic strips before use with a sand paper. Use the same strips of metals of same size, weight and number.

Metals are elements and are good conductors of heat and electricity. Most metals are electropositive in nature and the metal atoms lose electrons in chemical reactions to form cations. The more reactive a metal, the greater tendency it has to form a positive ion in a chemical reaction. For example:

How are metals arranged in a periodic table? Metals occupy the bulk of the periodic table. The alkali metals are a series of chemical elements in the periodic table.

· Alkali metals comprise group 1 in the periodic table along with hydrogen. Alkali metals are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs) and francium (Fr).

· The group 2 elements in the periodic table are called alkaline earth metals. Alkaline earth metals are beryllium (Be), magnesium (Mg) calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra). Group1 and group 2 elements are together called s-block elements.

· Elements of group 3 to 12 are transition elements. These are also metallic in nature and are called transition metals. They are also called d-block elements.

· Non-metallic elements occupy the right side of the periodic table. A diagonal line from boron to polonium separates metals from non-metals.

Physical properties of metals:

· Good conductors of heat and electricity: In metals, the positive ions are surrounded by a sea of electrons and these are responsible for their conductivity.

· High melting point and boiling point.

· They are malleable and ductile, so they can be bent and stretched without breaking.

· In most metals, the atoms are highly close packed, so they have high density.

· They have a shiny appearance.

· Metals exist in solid state at room temperature except for mercury, which is in a liquid state at room temperature.

Chemical properties of metals:

A more reactive metal readily reacts with other elements.

· The most reactive metals will react with even water, while the least reactive metals will not react even with acid. Most metals, on reacting with water produce hydroxide.

Example: If we put a small piece of sodium metal in water, sodium reacts exothermically with water producing hydrogen and metal hydroxide.

Magnesium reacts mildly with water but vigorously with steam. Zinc and iron react mildly with steam. Copper, gold and silver do not react with water at all.

· Metals react with oxygen in the air to form oxides.

Examples:

· Some metals react with acid and replace hydrogen from the acid.

· Most metals corrode when they are exposed to atmosphere. For example, the iron gets rusty after sometime if it is not painted. Titanium is highly resistant to corrosion.

· Single displacement reactions: A single displacement reaction is an important type of chemical reaction. It is also called substitution reaction. In these reactions, a free element displaces another element from its compound, producing a new compound. The reaction is usually written as:

Single displacement reactions are all oxidation-reduction reactions. For example,

Definition of Displacement reaction: The chemical reaction by which one element takes the position or place of another element in a compound .

Displacement reactions are very common in metals. They can be used to find out the relative reactivities of metals. In a displacement reaction, a more reactive metal can displace a less reactive metal from its salt solution. The reaction is often known as metal displacement reaction.

Some of the commonly used metals have been arranged in the decreasing order of reactivity. This is known as the reactivity series or activity series. The activity series of metals is an important concept in chemistry. The activity series of metals is an important tool for predicting the products of displacement reactions and the reactivity of metals in other reactions. Potassium is the most reactive metal, while platinum is the least reactive.

The higher of metal in the series, the more reactive it is and the more vigorously it reacts with water, oxygen and acid. A metal in the activity series can displace any metal below it in the series from its compound. The elements potassium, sodium, lithium and calcium are very reactive and they react with cold water to produce hydroxides and hydrogen gas. The elements magnesium, aluminium and iron are also considered as active metals and they react with steam to produce oxides and hydrogen gas. The metals above hydrogen are more reactive than hydrogen. These metals can displace hydrogen from acids or water and liberate hydrogen gas. The metals like copper, silver, gold and platinum are less reactive than hydrogen and they do not replace hydrogen from water or acid.

Examples for metal-displacement reactions:

· Reaction with acids:

a) Potassium, sodium, lithium and calcium react violently with dilute H₂SO₄ and dilute HCl, forming the metal salt (either sulphate or chloride) and hydrogen gas:

b) Zinc with dilute sulphuric acid is often used for the laboratory preparation of hydrogen. The reaction is slow at room temperature, but its rate can be increased by the addition of a little copper (II) sulphate. Zinc displaces copper metal, which acts as a catalyst.

c) Metals below hydrogen (copper, silver, gold and platinum), will not react with dilute acids. They cannot displace hydrogen from the non-metal anion.

· Reaction with concentrated acids as HNO₃ and H₂SO₄:

Hydrogen gas is not evolved when metals react with nitric acid (HNO₃) because it is a strong oxidising agent and it oxidises the H₂ produced to water and is itself reduced to nitrogen dioxide.

1) With active metals:

Mg + HNO_3(dilut) = Mg(NO₃)₂ + H₂O + NH₃ (NH₄NO₃)

Mg + 6HNO_3(conc) = Mg(NO₃)₂ + 4H₂O + 2N₂O

2) With passive metals:

Cu + HNO_3(dilut) = Cu(NO₃)₂ + H₂O + NO

3Cu + 8HNO_3(conc) = 3Cu(NO₃)₂ + 4H₂O + 2NO₂

3) Reaction with concentrated sulfuric acid:

Me + H₂SO₄ (conc) = MeSO₄ + H₂O + (H₂S, S, SO₂)

4) Fe and Al will not react with conc H₂SO₄ acid, they are passivated.

5) One part of HNO₃ and three parts of HCl are mixed together, the mixture wich obtained is called Aqua Regia. Noble metals such as gold, platinum are not dissolved either of the acid alone but in Aqua Regia they easily dissolve due to liberation of atomic Cl which is highly reactive:

HNO₃ + 3HCl NOCl + 2H₂O + 2Cl^-

· Zinc can displace copper from copper sulphate solution and iron from ferrous sulphate solution. So zinc is more reactive than iron and copper.

· If a piece of zinc metal is added to aluminium sulphate solution, no reaction takes place. So aluminium is more reactive than zinc.

· If small pieces of copper are added into solutions of zinc sulphate, aluminium sulphate and ferrous sulphate, no reactions take place. So, copper is less reactive than zinc, aluminium and iron.

· Aluminium can displace Zn from zinc sulphate solution, Cu from copper sulphate solution and Fe from ferrous sulphate solution.

· Iron displaces Cu from copper sulphate solution, but it does not displace Al and Zn from their salt solutions.

· Not only metals but also non-metals can take part in displacement reactions.

For example, chlorine gas displaces bromine from potassium bromide solution. The solution of potassium bromide acquires a yellowish and orange color of the liberated bromine gas.

Occurrence of Metals. The earth’s crust is the major source of metals. Seawater also contains some soluble salts such as sodium chloride, magnesium chloride, etc. The elements or compounds, which occur naturally in the earth’s crust, are known as minerals. At some places, minerals contain a very high percentage of a particular metal and the metal can be profitably extracted from it. These minerals are called ores.

Extraction of Metals: Some metals are found in the earth’s crust in the free state. Some are found in the form of their compounds.

The metals at the bottom of the activity series are the least reactive. They are often found in a free state. For example, gold, silver, platinum and copper are found in the free state. Copper and silver are also found in the combined state as their sulphide or oxide ores.

The metals at the top of the activity series (K, Na, Ca, Mg and Al) are so reactive that they are never found in nature as free elements.

The metals in the middle of the activity series (Zn, Fe, Pb, etc.) are moderately reactive. They are found in the earth’s crust mainly as oxides, sulphides or carbonates. The ores of many metals are oxides. This is because oxygen is a very reactive element and is very abundant on the earth.

Thus on the basis of reactivity, we can group the metals into the following three categories:

1 Metals of low reactivity (after H atom);

2 Metals of medium reactivity (between Al and Pb);

3 Metals of high reactivity (above Mg metal atom).

Different techniques are to be used for obtaining the metals falling in each category. Metallurgy as Per Reactivity:

· Electrolysis of molten is used for highly active metals such as K, Na, Ca, Mg and Al.

· Reduction with carbon is used for medium active metals such as Zn, Fe, Pb and Cu.

· Least reactive metals such as Ag and Au are found in free state in nature.

Refining of Metals: The metals produced by various reduction processes described above are not very pure. They contain impurities, which must be removed to obtain pure metals. The most widely used method for refining impure metals is electrolytic refining.

Electrolytic Refining: Many metals, such as copper, zinc, tin, nickel, silver, gold, etc., are refined electrolytically. In this process, the impure metal is made the anode and a thin strip of pure metal is made the cathode. A solution of the metal salt is used as an electrolyte. On passing the current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte is deposited on the cathode. The soluble impurities go into the solution, whereas, the insoluble impurities settle down at the bottom of the anode and are known as anode mud.

1^st Experiment. Reaction of metals with water

We have the following metals powder in laboratory: Cu, Fe, Zn, Mg, Na, Al, K, Ca, Pb. Please, make up a series of activity of these metals in accordance with their standard potentials from active to passive:

Me
E°, V

PROCEDURE:

1) Primarily we carry out experiments with alkali metals: for the alkali metals prepare porcelain cups with water and metal tweezers. Take a small piece of each metal by tweezers and gently lower it into the water in cup. Be careful, the reaction proceeds vigorously. After the metal is completely reacted with the water, drip 1-2 drops phenolphthalein indicator in cup and note the color of the solution. What products are formed? What metal is the most active? Write the equation redox reactions of alkali metals with water:

K+ H₂O =>

Na + H₂O =>

Ca + H₂O = >

2) For other metals, prepare the six test tubes and label them. Put a few of the metal powder in each tube and add a small amount of water and heat the tube until the reaction starts. Then, add 1 drop of phenolphthalein indicator in each tube and note the change in colors of solutions in test tubes. What products are formed? What is the metal most active? What metals do not react with water? Write the equation redox reactions of metals with water.

2^nd Experiment. Reaction of metals with diluted acids as HCl and H₂SO_4.

Hazards: Dilute hydrochloric acid irritates the eyes and magnesium ribbon is highly flammable. When a metal reacts with acid, an acid mist is formed which irritates the eyes and throat. Hydrogen gas is produced in the reaction and it is highly flammable.

Procedure:

1.	Add dilute hydrochloric acid to the beaker until it is half full.
2.	Put five test tubes in the test tube rack. Pour some of the hydrochloric acid into the first test tube to a depth of about 4 cm. Pour the same volume of acid into the other four test tubes.
3.	Add a piece of zinc to the first test tube. Add a piece of magnesium to the second test tube. Add a piece of copper to the third test tube. Add a piece of iron to the fourth test tube Add a piece of aluminium to the fifth test tube
4.	Watch carefully what happens in each test tube. • the name of each metal • whether bubbles of gas were given off or not • the speed at which the bubbles were given off • write the equation redox reactions of metals with water.
5	Repeat the experience with diluted 0.01Msulfuric acid H₂SO₄. What products are formed? What metal is the most active? What metals do not react with acid? Write the equation redox reactions of metals with water.

3^rd Experiment. Reaction of metals with nitric acid HNO₃.

Nitric acid is a strong oxidizing agent as shown by its large positive E^◦ values:

NO₃^- _(aq) + 2H⁺ _(aq) +1e^- → NO₂ _(g) + H₂O	E°= 0.79 V
NO₃^- _(aq) + 4H⁺_(aq) + 3e^- → NO _(g) + 2H₂	E°= 0.96 V

Being a powerful oxidizing agent, nitric acid reacts violently with many non-metallic compounds and the reactions may be explosive. Depending on the acid concentration, temperature and the reducing agent involved, the end products can be variable. Reaction takes place with all metals except the precious metal series and certain alloys. As a general rule, oxidizing reactions occur primarily with the concentrated acid, favoring the formation of nitrogen dioxide (NO₂).

Nitric acid dissolves most metals including iron, copper, and silver, with generally the liberation of lower oxides of nitrogen rather than hydrogen:

Cu + 4HNO_{3 (conc)} → Cu(NO₃)₂ + 2NO₂ + 2H₂O

The acidic properties tend to dominate with dilute acid, coupled with the preferential formation of nitrogen oxide (NO):

3Cu + 8HNO_{3 (dil)} → 3Cu(NO₃)₂ + 2NO + 4H₂O

Since nitric acid is an oxidizing agent, hydrogen (H) is rarely formed. Only magnesium (Mg) and calcium (Ca) react with cold, dilute nitric acid to give hydrogen:

Mg_(s) + 2HNO_{3 (aq)} → Mg(NO₃)_{2 (aq)} + H_{2 (g)}

4Mg+10HNO_{3 (conc)} ® Mg(NO₃)₂ + 3H₂O + NH₄NO₃

Procedure:

1) Safety: This acid is dangerous because it is highly corrosive and has oxidizing properties and as such it must be handled with great care. When handling nitric acid, one must wear protective gears especially to protect the eyes. The skin must also be protected because when nitric acid comes in contact with the skin, it results in a yellow discolouration. Perform this reaction in the fume hood.

2) Prepare the seven test tubes and label them.

3) Place 20-25 drops of 16 M nitric acid (concentrated) in all test tubes and support the test tubes in a beakers in the hood.

4) Add a small amount of each metals powder to the test tubes. Allow the metals approximately five minutes to dissolve. The brown gas produced is toxic nitrogen dioxide (NO₂). When all of the metal has dissolved, fill the test tube approximately ¼ full with distilled water.

5) What metal is the most active? Write the equation redox reactions of metals with nitric acid.

6) Repeat this experiment with dilute nitric acid. What products are formed? What metal is the most active? Write the equation redox reactions of metals with acid.

4^nd Experiment. Single displacement reaction of metals with salt.

In this experiment you will be ranking some of the metals according to their activities, from most to least active. One way to do this is to observe the relative vigor of their reactions (tendency to react) with nonmetals such as oxygen or chlorine. Another method, which differentiates more clearly, is to note whether one metal can replace another in a chemical compound. The general rule is that the more active metal replaces the less active one. This is easily observed in reactions between metals and metal ions (cations) in solution.

The general reaction for the replacement of a metal ion by another metal may be written as follows (We are arbitrarily showing only one electron per atom being transferred):

M_A+ M_B⁺X^- ® M_A+ X^-+ M_B

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