Metal Ions Lab

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. Experiments with Metals and Ions of Metals Introduction
Metals are similar in their physical properties in general, but they are not identical. Most of the metals are solids; few of them are liquids, such as mercury and cesium. Density of metals is not similar also. For example, sodium has density of 0.97g/cm3 while lead has density of 11.4g/cm3. Melting point of sodium is 98.0oC while for lead it is 327.6oC.
Metals have the capability to lose electrons when they react with non-metals such as O2, halogens, water, acids and other metal cations.
Metals react with non-metals but each to a different extent. The most reactive metals are alkali metals (group1A, where group is a vertical
column in the periodic table), which include lithium, sodium, potassium, rubidium, cesium and francium. All these metals have one electron in the valence electrons' shell; this electron can be removed relatively easily, which means that this electron has a very small ionization energy.
If we look closely at the elements of this group, we can see that they are not reactive to the same extent. This is because electrons have a negative charge and they are attracted to the nucleus, which has a positive charge. Valence shell electrons are responsible for the reactivity of an element; as the distance between valence electrons and the nucleus increases, the attractive force decreases and so the energy needed to separate the valence electron from an atom (ionization energy) decreases. So we can say that as the size of a metal atom increases, its ionization energy decreases and its reactivity increases (the most reactive alkali metal is francium and the least reactive one is lithium).
An example of the reactivity of alkali metal is their reaction with water. For example:
Na (s) + 2H2O → 2Na+(aq) + 2OH -(aq) + H2 (g) + heat
We can identify this reaction by the following
1- H2 gas emerges
2- Disappearance of metals
3- Heat produced
4- The appearance of hydroxide ions (OH-(aq)) makes the solution basic, which can be identified by using phenolphthalein as the indicator. We can compare between the metals by the rate of production of H2(g) as well as by the heat produced from this reaction. Alkaline earth elements (group IIA) are active metals but less than alkali metals. Since calcium is larger than the magnesium atom, its ionization energy is smaller than that of magnesium. Thus, calcium is more reactive than magnesium. The other metals as aluminum, zinc, iron, lead and copper do not react with water, so we can’t use water to distinguish between their reactivity. HCl(aq) is used to put them in order according to their reactivity. We can put the most common metal in a series according to their reactivity (from the most to the least reactive one).
“Potassium, sodium, lithium, calcium, magnesium, aluminum, zinc, iron, lead, copper, silver……”. When a metal in the upper part of the reactivity series reacts with aqueous solution of other metal ions lower in the series, the more reactive metal loses its electron more easily than the less reactive metal. As a result, the more reactive metal transfers its electrons to the less reactive metal. According to this reactivity series, we can predict what will happen when a piece of zinc metal is put in CuSO4 (aq) solution or in MgSO4 (aq) solution.
Since Zn is more reactive than Cu, when we put Zn in CuSO4(aq), a reaction will occur as follows: Zn (s) + Cu2+(aq) + SO42- (aq) → Zn2+(aq) + SO42- (aq) + Cu (s)
However, Since Zn is less reactive than Mg, if we put Zn in MgSO4(aq) there will be no reaction at all: Zn (s) + Mg2+(aq) + SO42- (aq) → No reaction.

Materials:
10 mL Cu2+(aq) solution 0.5M
10 mL Ag+(aq) solution 0.5M
10 mL Mg2+(aq) solution 0.5M
10 mL Zn2+(aq) saturated solution
Pieces of the metals: Mg(s), Zn(s), Ag(s), Cu(s) 1cm x 1cm each. Equipments:
16 test tubs
Emery cloth
Tweezers

Experimental procedure
In the current laboratory session you will explore some of the characteristics of certain elements. You will start by comparing the reactivity of Zn, Mg, Cu and Ag metals by their ability to release electrons. a. Wipe the metals Zn, Mg, Cu and Ag using the emery cloth and put each of them into a clean test tube using a tweezers. b. Fill each of the test tubes with about 2 cm of the Cu2+(aq) solution. 1. Watch what happens and write down your observations. Wait for 2-3 minutes to determine whether a chemical reaction occurred in each test tube, and if so, write a balanced chemical reaction. Zn in Cu2+(aq) :
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____________ Mg in Cu2+(aq) :
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____________ Cu in Cu2+(aq) :
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____________ Ag in Cu2+(aq) :
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____________ c. Wipe the metals Zn, Mg, Cu and Ag using the emery cloth and put each of them into a clean test tube. d. Fill each of the test tubes with about 2 cm of the Mg2+(aq) solution. 2. Watch what happens and write down you observations. Wait for 2-3 minutes to determine whether a chemical reaction occurred in the test tube, and if so, write a balanced chemical reaction. Zn in Mg2+(aq) :
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_______________ Mg in Mg2+(aq) :
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____________ Cu in Mg2+(aq) :
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____________ Ag in Mg2+(aq) :
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____________ e. Wipe the metals Zn, Mg, Cu and Ag using the emery cloth and put each of them into a clean test tube. f. Fill each of the test tubes with about 2 cm of the Zn2+(aq) solution. 3. Watch what happens and write down you observations. Wait for 2-3 minutes to determine whether a chemical reaction occurred in the test tube, and if so, write a balanced chemical reaction. Zn in Zn2+(aq) :
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____________ Mg in Zn2+(aq) :
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____________ Cu in Zn2+(aq) :
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____________ Ag in Zn2+(aq) :
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____________ g. Wipe the metals Zn, Mg, Cu and Ag using the emery cloth and put each of them into a clean test tube. h. Fill each of the test tubes with about 2 cm of the Ag+(aq) solution. 4. Watch what happens and write down you observations. Wait for 2-3 minutes to determine whether a chemical reaction occurred in the test tube, and if so, write a balanced chemical reaction. Zn in Ag+(aq) :
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____________ Mg in Ag+(aq) :
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____________ Cu in Ag+(aq) :
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____________ Ag in Ag+(aq) :
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____________ 5. Sum your observations in the following table (color, precipitance, other changes):

6. List the metals in order of their ability to release electrons.
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___________ i. Receive an unknown solution from your teacher and determine what is the metal ion in the solution 7. Describe the tests you did in order to find the unknown ion in the solution and write what it was. ____________________________________________________________

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Reactions of chlorine, bromine and iodine with aluminium
In this demonstration experiment, aluminium reacts with each of three halogens, chlorine, bromine and iodine.
Spectacular exothermic redox reactions occur, accompanied by flames and coloured ‘smoke’, forming the solid aluminium halides:
2Al + 3X2 → 2AlX3 (X = Cl, Br and I)
Read our standard health & safety guidance
Lesson organisation
This is a class demonstration that shows the spectacular reactivity of three non-metals from Group 7 with a metal.
These experiments must be done in a fume cupboard as both the reactants and products are hazardous. Teachers attempting this demonstration for the first time are strongly advised to do a trial run before doing it in front of a class.
Each experiment should take about 5
minutes.
Apparatus and chemicals
Eye protection
Thick chemically-resistant gloves such as marigold industrial blue nitrile
Access to a fume cupboard
The teacher will require:
Mortar and pestle
Heat resistant mat, 30 x 30 cm
Watch glasses, about 10 cm diameter, 2
Reduction tube (see note 1)
Test-tubes, 3
Test-tube rack
Teat pipette
Filter paper
Spatula or wooden splint
Bosses, clamps and stands
Chemicals for one demonstration:
Aluminium foil, a few cm2
Aluminium powder (Highly flammable, Contact with water may liberate hydrogen), 0.3 g
Liquid bromine (Corrosive, Very toxic), 1 cm3 (see note 2)
Solid iodine (Harmful), 2 g
Silver nitrate solution, about 0.1 mol dm3 (Low hazard but stains skin and clothing at this concentration), about 10 cm3
A little deionised water in a washbottle
Chlorine generator (Toxic, Irritant) Generating gases - scroll down for chlorine
Sodium chlorate(I) solution (14% (w/v) available chlorine) (Corrosive), about 100cm3
Hydrochloric acid, 5 mol dm-3 (Corrosive), about 50 cm3
Technical notes
Aluminium powder (Highly flammable, Contact with water may liberate hydrogen) Refer to CLEAPSS Hazcard 1
Liquid bromine (Corrosive, Very toxic) Refer to CLEAPSS Hazcard 15A, SRA04 The aluminium/bromine reaction
Solid iodine (Harmful) Refer to CLEAPSS Hazcard 54A
Silver nitrate solution (Low hazard at concentration used) Refer to CLEAPSS Hazcard 87 and Recipe card 58
Chlorine (Toxic, Irritant) Refer to CLEAPSS Hazcard 22A and Recipe card 26
Sodium chlorate(I) solution (Corrosive) Refer to CLEAPSS Hazcard 89
Hydrochloric acid (Corrosive) Refer to CLEAPSS Hazcard 47A and Recipe card 31
1 The reduction tube should be fitted with a one-holed rubber stopper fitted with short length of glass tubing and attached to the chlorine generator. Alternatively an 8–10 cm length of wide bore glass tubing with a stopper fitted with a short length of glass tubing at each end could be used – see diagram.
2 Wear suitable protective gloves (thick, chemically resistant) when handling liquid bromine. Have 500 cm3 of 1 mol dm-3 solution of sodium thiosulfate available to treat any spillages on the skin.

Procedure
HEALTH & SAFETY: Wear eye protection and gloves
Reaction of Al + Cl2 a Set up the chlorine generator in a fume cupboard. Make sure it is securely clamped. b Loosely crumple a piece of aluminium foil (10 x 5 cm) so that it will just fit inside the reduction tube and push it into the tube. Attach the tube to the generator with a short length of rubber tubing and clamp it in position at the end nearest to the generator, so that the aluminium foil can easily be heated using a Bunsen burner – see diagram. c Generate a gentle stream of chlorine by dripping the hydrochloric acid slowly on to the sodium chlorate(I) solution, and allow it to pass over the aluminium. When the green colour of the chlorine gas fills the reduction tube, start heating the aluminium foil with a Bunsen burner at the end nearest to the chlorine supply. Once the aluminium is hot, a bright glow will suddenly appear where it starts to react with chlorine. d Remove the heat. The bright glow should spread along the aluminium. If not, heat again, and increase the flow of chlorine gas. A lot of white ‘smoke’ – aluminium chloride – is produced, some of it condensing as a white powder on the walls of the reduction tube and the rest streaming out of the hole in the end of the tube. When the reaction is over, stop the chlorine supply and remove the heat. e When the reduction tube has cooled down, disconnect it and, still in the fume cupboard, scrape a little of the white powder into a test-tube. Add a little deionised/distilled water to the solid to dissolve it. Care: the reaction between anhydrous aluminium chloride and water can be quite vigorous – an audible hiss can often be heard - see Teaching Notes. f Test a drop of the solution with Universal indicator paper. It is strongly acidic. Test the remainder with a little silver nitrate solution. A white precipitate forms, showing the presence of chloride ions.
Reaction of Al + Br2 a Tear or cut some aluminium foil into several small pieces about 2 x 2 mm in size. Carefully pour 1 cm3 of liquid bromine onto a watchglass on a heat resistant mat in a fume cupboard. Sprinkle a few pieces of aluminium foil on to the surface of the bromine. Within a minute there are flashes of flame and a white ‘smoke’ of aluminium bromide is formed, together the orange vapour of bromine. Carefully hold another watchglass over the reaction to condense some of the ‘smoke’ on to its surface as a solid. b Wash any aluminium bromide collected in this way off the watchglass into a test-tube using a little deionised water (Care: see Teaching Notes below). Test the solution with indicator paper and silver nitrate solution as above. The solution is acidic and a cream precipitate of silver bromide is formed.
Reaction of Al + I2 a Weigh out 2 g of iodine, which should be dry, and grind it to a powder in a fume cupboard, using a mortar and pestle. Place the powdered iodine on a filter paper on a dry heat resistant mat and add 0.3 g of aluminium powder to it. Mix the two solids together in the fume cupboard using a wooden splint – do not grind them together. When they are thoroughly mixed, pour the mixture into a heap on the heat resistant mat or in a metal tray, such as a tin lid, positioned in the middle of the fume cupboard. b To start the reaction, use a teat pipette to place a few drops of water on the mixture. After a time lag, the water begins to steam and clouds of purple iodine vapour are given off, indicating that an exothermic reaction has started. After a few more seconds sparks are given off and the mixture bursts into flame. When the reaction subsides, a white residue of aluminium iodide remains. Scrape a little of this into a test-tube (Care: see Teaching Notes below), add some deionised water and filter if necessary. Test the solution with indicator paper and silver nitrate solution as above. The solution is acidic and a yellow precipitate indicates the presence of iodide ions.
Teaching notes
These reactions make quite spectacular demonstrations in themselves, the bromine + aluminium reaction even more so in a partly darkened room. Classroom management in semi-darkness (Practical Physics advice)
The demonstrations can be used to show the reaction between reactive non-metallic elements and a fairly reactive metal to form compounds, or as part of the study of the reactions of the Group 7 elements. Here the apparent order of reactivity is not that predicted from their position in the Group (that is chlorine → bromine → iodine). This is because of the different physical states of the three halogens, and the different surface area of the aluminium as a powder or foil. This can be used to make an important point about ‘fair’ comparisons of reactivity.
These reactions also serve to show that aluminium is in fact a more reactive metal than it appears in everyday use. The protective oxide layer of aluminium has to be penetrated by the halogens before the reactions can start, hence the delays, and the need for water to assist the two solid elements getting into contact, in the case of aluminium and iodine.
The clouds of iodine vapour released when aluminium and iodine react can stain the inside of a fume cupboard. Teachers may prefer to demonstrate this reaction outdoors, if possible.
The anhydrous aluminium halides are vigorously hydrolysed (sometimes violently if freshly prepared and hot, as here) by water, giving off fumes of a hydrogen halide and a forming an acidic solution of the aluminium salt. To dispose of the solid residues, allow them to cool completely before adding in small amounts to 1 mol dm-3 sodium carbonate solution in a fume cupboard. Wait until the reaction has subsided before adding more solid. Dispose of the resultant slurry with plenty of water.

Allotropes of sulfur
Sulfur is heated slowly and steadily from room temperature, so that all the changes in colour and consistency as it melts and eventually reaches boiling point, can be observed. A fresh sample of sulfur is heated to just above the melting point, then allowed to cool and crystallise slowly as monoclinic sulfur. A further sample is heated to boiling point, and the liquid rapidly chilled in cold water to form plastic sulfur.
A separate sample of sulfur is dissolved in a warm solvent, and the solution allowed to cool and evaporate, leaving crystals of rhombic sulfur.
All the observed changes in properties can be related to the different molecular structures of the three solid forms of sulfur, and to the changes in structure as the temperature of liquid sulfur is gradually raised.
Read our standard health & safety guidance
Lesson organisation
This practical is described here as a demonstration. However, some teachers may wish to consider whether certain parts could be used as class practicals with appropriately skilful and reliable classes.
A demonstration, without any accompanying discussion about the possible reasons for the changes in properties in terms of structure, would take up to 45 minutes. However, to derive maximum benefit from the experiment, more time needs to be allowed for such discussion.
Apparatus and chemicals
The teacher will require:
Eye protection
Heat resistant gloves
Access to a fume cupboard
Flexicam or similar camera, digital microscope, digital projector and screen or other method of projecting images of small crystals to the class (as available).
Boiling tubes, 4 (see note 4)
Test-tube holders, 2
Test-tube rack
Stands and clamps, 2
Conical flask, 250 cm3
Cork, to fit conical flask
Beaker (250 cm3), 2
Beaker, 1 dm3 (see note 5)
Thermometer, 0 – 250 °C
Petri dishes or watchglasses, 4 (or more)
Bunsen burner, tripod and gauze or Electric hotplates, 2 (optional, if available)
Heat resistant mats, 2
Filter paper, about 18 - 20 cm diameter
Spatula
Paper clips
Damp cloth (to extinguish small sulfur fires)
Sulfur, powdered roll (Low hazard), 100 g (see note 1)
Dimethylbenzene (xylene), (Harmful), 100 cm3 (see note 2)
Cooking oil (Low Hazard), 700 cm3 (see note 3)
Technical notes
Dimethylbenzene (xylene) (Harmful) Refer to CLEAPSS Hazcard 46
Sulfur (Low hazard) Refer to CLEAPSS Hazcard 96A
Cooking oil (Low hazard)
1 The sulfur used must be roll sulfur, crushed to a powder. To crush the rolls of sulfur, place in a strong plastic bag on a hard surface. Use a hammer or a vice to break up the roll sulfur into small pieces, then crush to a powder in a mortar and pestle. ‘Flowers of sulfur’ is not suitable because it contains a lot of insoluble amorphous sulfur.
During the experiments sulfur may catch fire, releasing sulfur dioxide (Toxic – refer to CLEAPSS Hazcard 97), which may cause breathing difficulties to some students. If this happens, extinguish quickly by placing a damp cloth over the mouth of the test-tube. If the combustion cannot be extinguished quickly, the test-tube should be placed in fume cupboard, and the fan left running.
2 Although other hydrocarbon solvents, such as methylbenzene, can be used to dissolve sulfur and form monoclinic sulfur, dimethylbenzene (xylene) is the least hazardous.
3 If suitable cooking oil is not available, other clear, high-boiling oils may be used, e.g. paraffin oil (Refer to CLEAPSS Hazcard 45B).
4 These are large (150 x 25 mm) test-tubes, and should be clean and dry. The test-tubes in which sulfur has been heated can be difficult to clean for general use. It may be worth keeping a set of such tubes from year-to-year for this experiment.
5 The large beaker containing the cooking oil functions as an oil-bath for heating the sulfur slowly and uniformly, while allowing students to see clearly what is happening to the sulfur. Other containers may be preferred for the oil-bath, provided the visibility is maintained, for example by use of a webcam and digital projector.
Procedure
HEALTH & SAFETY: Wear eye protection.
Before the demonstration: a Pre-heat the oil-bath to about 130 °C, and maintain this temperature. b Clamp one of the sulfur-containing tubes in the oil bath, so that the sulfur is below the level of the oil in the bath. c Half fill the 250 cm3 beaker with cold water. d In the fume cupboard, put about 10 g of powdered roll sulfur into the conical flask and add about 100 cm3 of dimethylbenzene. e Prepare filter paper cone held together by a paper clip and supported in a beaker, as shown below:

The demonstration: a Two-thirds fill two test-tubes with powdered roll sulfur (about 20 g in each tube) and place in the oil bath. The sulfur will melt to a transparent, amber, mobile liquid in about 15 minutes. b Remove one tube from the oil-bath and pour the molten sulfur into the filter paper cone. Allow the sulfur to cool slowly and solidify, forming a crust. c Break the crust with a spatula and, handling the filter paper cone with heat resistant gloves, tilt it so that any remaining liquid flows out of the cone of solidifying sulfur on to a piece of scrap paper or card (for disposal). Needle-shaped crystals of monoclinic sulfur will be seen inside the hollow cone. When cool, the cone can be passed around the class. It may be necessary to break the cone open to see the crystals more easily. d Over the next day or two, look carefully at the needle crystals from time to time. They will slowly go cloudy, yet retain their needle shape, as the monoclinic form slowly turns back to the more stable rhombic sulfur – each needle becomes a mass of tiny rhombic crystals.
Liquid sulfur: a Remove the second tube from the hot oil using a reliable test-tube holder and wipe off any oil using a paper towel. Heat the molten sulfur gently over a small Bunsen flame, keeping the contents moving to prevent local overheating. The liquid gets darker and, fairly suddenly, becomes a viscous, gel-like substance. This occurs at about 200 °C. b The tube can be inverted and the sulfur will remain in it. Show that the mobile liquid re-forms on cooling. c Now heat the sulfur slowly and steadily beyond the gel-like stage. The sulfur liquefies again to a very dark red-brown liquid. Note that during this heating the sulfur may catch fire and sulfur dioxide will be produced. Have a heat resistant mat or damp cloth to hand to place over the mouth of the tube to extinguish the blue flames. d When the sulfur begins to boil (441°C), pour the liquid sulfur in a slow stream into a beaker of cold water. A tangled mass of brown plastic sulfur will form. e Allow this to cool thoroughly. The inside of the plastic sulfur may remain molten after the outside has solidified. f Remove the plastic sulfur from the water and show that it is rubbery – it can be stretched and will return to its original shape. g The shiny surface of the plastic sulfur begins to dull and some of the elasticity is lost within 30 minutes, as it begins to turn back to the more stable rhombic sulfur. h Leave the plastic sulfur until the following lesson to monitor the progress of this change. This will be very noticeable after a week or so but complete change will take a long time. It will become brittle.
Rhombic sulfur: a Gently warm the conical flask containing sulfur and dimethylbenzene to about 50°C (preferably on an electric hotplate) to complete dissolving of the sulfur. Some teachers may prefer to have done this before the demonstration to save time. b Pour a little of the solution into each of a set of petri dishes or watch glasses and leave them in the fume cupboard for the solvent to evaporate. This will take about 10 minutes. c The small crystals of rhombic sulfur formed should be viewed by projection of images onto a screen if possible.
Teaching notes
Some stages of this demonstration are time-consuming, e.g. melting the sulfur in the oil bath, dissolving the sulfur in dimethylbenzene, and evaporating the solvent. Some teachers may prefer to melt some sulfur before the lesson and to prepare rhombic crystals before the lesson to save time. In the latter case, slower evaporation (which can be brought about by covering the petri dish with filter paper with a few holes in) will produce larger crystals. Particularly large and/or well-formed crystals could be retained as examples for future use.
Monoclinic crystals can be formed by allowing a hot solution of sulfur in boiling dimethyl- benzene to cool so that crystallisation starts at above 96 °C.
Carbon disulphide has been use in the past as a better solvent for making rhombic sulfur, However its smell, toxicity and high flammability make it unsuitable for use in schools - see CLEAPSS Hazcard 20.
Very slow heating is essential if all of the changes on heating sulfur are to be seen clearly.
Sulfur is a poor thermal conductor, hence the changes can overlap one another if the heating is too fast. It is difficult to heat slowly enough using a Bunsen burner – hence the use of an oil bath.
Crystalline sulfur consists of puckered S8 rings in the shape of crowns. These can be packed together in two different ways – to form rhombic crystals and to form needle-shaped monoclinic crystals, as shown below:

Below about 96 °C, rhombic sulfur is the more stable allotrope. On melting at about 118 °C, sulfur first forms a mobile, amber liquid containing S8 rings. If this is allowed to cool, monoclinic sulfur forms as crystallisation occurs above 96 °C.
Monoclinic sulfur will turn slowly into the more stable rhombic form on standing below 96 °C.
Further heating of the S8-containing liquid breaks the rings into S8 chains. These may join to form longer chains which tangle, causing an increase in viscosity. At higher temperatures, these chains break into shorter ones, perhaps as short as S2, and the viscosity decreases
again.
Rapid cooling of this liquid traps the resulting solid sulfur in the tangled chain state – this is plastic sulfur. On stretching, the chains uncoil and on releasing the tension they return to the partly coiled state (see scheme below). If solid sulfur is formed below 96 °C by crystallisation from a solution, the stable rhombic form is produced.