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    COMPOUNDS OF METALS

    Thursday, April 02, 2020





    Metals combine with other substances to form compounds. There is a diversity of metal compounds known. One can make an endless list of the compounds of the metals. In this chapter, we are going to concentrate our efforts on the following compounds of the metals: oxides, hydroxides, carbonates and hydrogen carbonates, nitrates, chlorides and sulphates.
    Preparation of Oxides of Some Metals by Direct and Indirect Methods
    All elements except helium, neon and argon form compounds with oxygen. This is because oxygen is quite reactive. Binary compounds of oxygen are known as oxides. Therefore, a metal oxide is a binary compound of oxygen and a metal.

    Metal Oxides
    Chemical properties of metal oxides
    1. Reaction with water:The oxides of potassium, sodium and calcium are very soluble in water. They will react vigorously with cold water to produce the corresponding hydroxides. The oxides of metals below calcium in the reactivity series are all insoluble in water.
    2. Reaction with acids:The oxides of metals above hydrogen in the reactivity series react with dilute acids to produce salt and water.
    3. Reaction with alkalis:Some oxides react with alkalis to produce salt and water. The oxides of this nature include ZnO, Al2O3, PbO and SnO.
    The Reactions of Metal Oxides with Water and Dilute Acids
    Preparation of metal oxides
    Metal oxides can be prepared by:
    1. direct methods. This involves heating metals directly in air.
    2. indirect methods. This involves such methods like heating carbonates and hydrogencarbonates of certain metals in air, and reacting certain metals with certain acids.
    Preparation of metal oxides by direct combination(Direct method)
    In this method, oxides can be prepared by direct combination of metals with oxygen. This involves heating a metal in air. When some metals are burned in the air, they react with oxygen of the air to form metal oxides. However, this method is not intensively used because some metals tend to form a protective layer of an oxide on the surface of metal and prevent further attack by oxygen. The best example of such metals is aluminium which, when heated in the air, forms a protective layer of aluminium (III) oxide (Al2O3) on the surface of a metal whichprevents further attack by oxygen.

    Table 9.1 shows the products formed when certain metals are burned in the air.
    Table 9.1: The reaction of metals with oxygen
    MetalHow it reactsProduct
    Bariumburns with a green flamewhite solid (barium (II) oxide, BaO)
    Calciumburns with a brick-red flamewhite solid (calcium oxide, CaO)
    Sodiumburns with a yellow flamewhite solid (sodium oxide, Na2O)
    Potassiumburns with a purple flamewhite solid (potassium oxide, K2O)
    Magnesiumburns with a white flamewhite solid (magnesium oxide, MgO)
    Ironburns with yellow sparksBlue-black solid (iron (II) oxide, FeO)
    Copperdoesnot burn, turns blackblack solid (copper (II) oxide, CuO)
    Activity 1
    Preparation of oxides by direct combination methods
    Aim: to prepare oxides of metals
    Materials: magnesium ribbon, aluminium foil, iron filings, gas jar, Bunsen burner and coal tong.
    Discussion
    1. What changes did you observe when each of the metals above was heated in a Bunsen flame?
    2. What was the function of oxygen in the reaction?
    3. Write well balanced equations for the reactions that took place.
    4. Write and balance the equations for the reactions when the following metals burn in oxygen:(i)sodium (ii)potassium (iii) iron
    Procedure:
    • Lower a piece of burning magnesium ribbon, by means of tongs, into a gas jar of oxygen. Observe and record what happens.
    • Heat the aluminium foil strongly on a Bunsen flame. Observe and record what happens.
    • Perform the same experiment with iron filings. Also, record what happens.
    Preparation of metal oxides by indirect methods
    This method involves thermal decomposition of salts. When some salts are heated, they decompose into oxides and other products as well. If the anion part of the salt heated contains some oxygen, a portion of this oxygen may remain bonded to the central metal atom.
    However, this method is limited only to those compounds of metals below sodium in the electrochemical series. For instance, potassium oxide or sodium oxide cannot be prepared by action of heat on their carbonates. Thermal decomposition of some metals is as shown by the following equations:
    • CuCO3(s)→CuO(s)+ O2(g)
    • CaCO3(s)→CaO(s)+ CO2(g)
    • ZnCO3(s)→ ZnO(s)+ CO2(g)
    Hydroxides also behave in a similar manner. When hydroxides of metals below sodium in the electrochemical series are heated, they decompose into respective oxides, giving off water in the form of steam.
    • Ca(OH)2(s)→ CaO(s)+ H2O(g)
    • Mg(OH)2(s)→ MgO(s)+ H2O(g)
    The oxides can also prepared by heating some nitrates and sulphates:
    • 2Pb(NO3)2(s)→ 2PbO(s)+ 4NO2(g)+ O2(g)
    • 2Cu(NO3)2(s)→ 2CuO(s)+ 4NO2(g)+ O2(g)
    Nitrates of silver and mercury are not suitable for preparation of oxides by thermal decomposition because they decompose to metals directly when heated. The bond between oxygen and a metal atom is not strong enough to withstand the thermal energy:
    • 2AgNO3(s)→ 2Ag(s)+ 2NO2(g)+ O2(g)
    Activity 2
    Preparation of oxides by thermal decomposition of carbonates
    Discussion
    1. (a) What was the colour of the copper carbonate before heating? (b) What was the colour of the residue in the test tube after cooling? What substance is this? (c) Which gas was evolved on heating the carbonate?
    2. (a) What happened to lead carbonate whenit was heated? (b) What gas was evolved? (c) What colour was the residue after cooling?
    3. Write well balanced chemical equations for the two experiments performed.
    4. Explain why metal oxides cannot be prepared by thermal decomposition of either potassium or sodium carbonate.
    Aim:To prepare metal oxides by thermal decomposition of carbonates
    Procedure:
    1. Put a sample of copper carbonate in a test tube.
    2. Place the test tube on a Bunsen flame and heat slowly then strongly.
    3. Observe and record any changes, including testing the gases evolved.
    4. When no further changes take place in the test tube, cool down the contents.
    5. Perform the same experiment with lead carbonate.
    The Uses of Metal Oxides
    Explain the uses of metal oxides
    Metal oxides find a wide range of uses. The following are the uses of the most common oxides:
    Uses of calcium oxide (CaO)
    1. Making mortar:Calcium oxide reacts with water to form the hydroxide, Ca(OH)2, known as slaked lime.Mortaris made by mixingslaked lime,sandandwater, and is used in sticking bricks together and in forming smooth surfaces on walls of buildings.
    2. Calcium oxide is used for makingwhitewash, which is used in marking sport’s fields, roads and is brushed on walls of buildings to give them the white colour. Whitewash is a suspension of slaked lime in water.
    3. Cement and concrete:Cementis made by heating togetherlimeorlimestoneandclay. The product is a mixture of calcium silicates and aluminates. Clay is hydrated aluminium silicate. A mixture ofcement,sand,stonesandwatergivesconcrete, which on setting becomes extremely hard. It is the materials used for making foundations of buildings, pillars, roads, paths, bridges, etc
    4. Soil treatment:In agriculture,quicklime is used to neutralize soil acidity, and it also adds mineral nutrients (Ca2+) to the soil.
    5. Calcium oxide is dissolved in water to make slaked lime, which is used in the softening of water.
    6. Drying agent: Calcium oxide is used for drying ammonia and ethanol.
    7. It is used in the manufacture of bleaching powder, CaOCl2.
    8. Manufacture ofglass:heating a mixture ofsand,sodium carbonateandlimeorlimestonegives glass.
    9. Lining of furnaces: It is mixed with magnesium oxide to form the basic lining of the furnaces to remove acidic impurities in the form of slag.
    10. it is used in the blast furnace to remove impurities fromiron ore which is removedin the formof slag.
    11. Preparation calcium carbide: Calcium carbide (CaC2) is manufactured in an electric furnace at 2000oC. CaO(s)+ 3C(s)→ CaC2(s)+ CO(g)
    Uses of magnesium oxide (MgO)
    1. The oxide is used as a lining material in refractory furnaces, owing its high melting point, which is around 2900oC. It is also used as a refractory agent in the construction of crucibles.
    2. The oxide in its solution form (magnesium hydroxide) is commonly used as an antacid. This works because magnesium hydroxide is a basic substance, which means that, it will neutralize excess acidity and end up indigestion, caused by too much hydrochloric acid in the stomach.
    3. It is used to manufacture the common chemical reagents in the laboratory such a magnesium chloride, magnesium sulphate and magnesium hydroxide.
    4. Magnesium oxide is a popular drying agent. In its powder form, it is hydroscopic in nature. This makes it suitable for drying different substances.
    5. Insulation: Due to its heat resistance properties, magnesium oxide powder makes an excellent insulator.
    6. Dietary supplement: Since it is a good source of magnesium, the oxide is used as or in dietary supplements forhumans and animals.
    Uses of aluminium (III) oxide (Al2O3)
    1. Aluminium (III) oxide, in the form of bauxite, is used as a source of aluminium.
    2. Owing its rough surface, the oxide is used as an abrasive, i.e. it is used to rub and clean other surfaces.
    3. It is used as an adsorbent in chromatography.
    4. It is used in the lining of furnaces as a refractory material because it has a high melting point (2040oC).
    Uses of zinc oxide (ZnO)
    1. Zinc oxide is chiefly used in the manufacture of paints and pigments. In addition, the oxide is used to manufacture anti-corrosive coatings, lubricants, adhesive batteries, fire retardants, plastic, cement, glass and ceramics (as a component of glazes).
    2. Manufacture of rubber: It is mainly used to activate vulcanization, which aims at improving the strength and elasticity of rubber.
    3. Manufacture of cigarette filter: As a cigarette filter, zinc oxide helps to remove certain harmful compounds from the tobacco smoke, without altering its flavour.
    4. Making concrete: It helps to make the concrete more resistant to water, besides improving the processing time required.
    5. Medical uses: Zinc oxide has anti-bacterial properties, for which it is extensively used to treat a number of skin conditions. It is topically applied to provide relief in skin irritation, diaper rash, minor burns and cuts, and for dry and chapped skin. It is added to baby powder, anti-dandruff shampoos as well as antiseptic creams and surgical tapes due to its medicinal properties. In addition, together with iron oxide, it is used to make calamine solution.
    6. Cosmetic uses: the most important use of zinc oxide in the cosmetic industry is in the preparation of sunscreen lotions and creams. Zinc oxide can absorb ultraviolet (UV) radiation of the sun and thereby protect the skin from sunburn and other damaging effects of UV radiation.
    Preparation of Soluble and Insoluble Sulphates
    Prepare soluble and insoluble sulphates
    Metal sulphates are generally soluble in water, except for the commonly known insoluble sulphates of lead (PbSO4) and barium (BaSO4). Calcium sulphate is sparingly soluble.The sulphate of alkali metals and those of the alkaline earth metals (Mg, Ca, Sr and Ba) are very stable to heat.
    Soluble sulphates can be prepared in the laboratory by reacting metals, oxides, hydroxides or carbonates with dilute sulphuric acid and then isolating the crystals. The sulphates are isolated by the usual methods of preparing soluble salts.
    1. Fe(s)+ H2SO4(aq)→ FeSO4(aq)+ H2(g)
    2. MgO(s)+ H2SO4(aq)→ MgSO4(aq)+ H2O(l)
    3. 2NaOH(aq)+ H2SO4(aq)→ Na2SO4(aq)+ 2H2O(l)(d)
    4. CaCO3(s)+ H2SO4(aq)→ CaSO4(aq)+ H2O(l)+ CO2(g)
    Preparation of insoluble sulphates
    The best method to prepare insoluble metal sulphates is by precipitation reactions. This is achieved by reacting their soluble salts with dilute sulphuric acid:
    1. BaCl2(aq)+ H2SO4(aq)→ BaSO4(s)+ 2HCl(aq)
    2. Pb(NO3)2(aq)+ H2SO4(aq)→PbSO4(s)+ 2HNO3(aq)
    The sulphates are isolated by the usual methods of preparing insoluble salts.
    Chemical Properties of Sulphates
    Explain chemical properties of sulphates
    Sulphates of all metals are normal salts and have the following chemical properties:
    1. All sulphates give a white precipitate when treated with aqueous salts ofleadandbarium, e.g.BaCl2(aq)+ K2SO4(aq)→ BaSO4(s)+ 2KCl(aq);Na2SO4(aq)+Pb(NO3)2(aq)→ PbSO4(s)+ 2NaNO3(aq)
    2. Barium and lead sulphates are the only two common insoluble sulphates. Calcium sulphate is sparingly soluble. The rest of the sulphates are soluble in water.
    3. Sulphates of the metals in group I and II of the periodic table are very stable to heat. Strong heating decomposes some of the sulphates of the heavier metals.Iron (II) sulphate disproportionates when heated:2FeSO4(s)→ Fe2O3(s)+ SO2(g)+ SO3(g)Iron (III) sulphate gives a good yield of sulphur trioxide gas: Fe2(SO4)3(s)→ Fe2O3(s)+ 3SO3(g)
    4. Hydrated sulphates decompose on heating to form oxides, water and sulphur trioxide.CuSO4.5H2O(s)→ CuO(s)+ 5H2O(g)+ SO3(g);FeSO4.7H2O(s)→ FeO(s)+ 7H2O(g)+ SO3(g).But hydrated sodium sulphates is stable to heat.It only loses its water of crystallization when heated. Na2SO4.10H2O(s)→ Na2SO4(s)+ 10H2O(g)
    Uses of Sulphates
    Describe uses of sulphates
    Sulphates are salts of considerable importance in our everyday lives. The following are the uses of some sulphates:
    Sodium sulphate (Na2SO4.10H2O), also known as Glauber’s salt.
    It is used as a mild purgative in medicine. The anhydrous salt, Na2SO4,is used as a laxative. It also finds its use in the manufacture of glass.
    Calcium sulphate
    In the form of Plaster of Paris (CaSO4.½H2O), is used to make plaster casts that are used in hospitals for the repair of broken limbs. When in the form of gypsum (CaSO4.2H2O), it is used in the manufacture of cement, moulds, wall plasters and wallboards, and inexpensive art objects. Among the many other uses of calcium sulphate are its uses as a pigment in white paints, coating agent in papers, in the manufacture of sulphuric acid and sulphur, and as a drying agent in many laboratory and commercial purposes.
    Some uses of aluminium sulphate are as follows:
    • It is used in paper making where it binds the paper fibres together.
    • It is also used in the manufacture of aluminium hydroxide, which is used for mordant dyeing. The aluminium hydroxide formed by the hydrolysis of the sulphate is deposited on the cloth fibres, where it helps the dye to stick well on the fibre. The salt is also used in paper sizing (i.e. giving it body and strength), and waterproofing the cloth.
    • Aluminium sulphate is an important chemical in the treatment of urban water. It precipitates colloidal matter from water. Microorganisms e.g. bacteria and algae are also captured during the coagulation process and precipitated with mud.
    • Aluminium sulphate is used in the “foam” type fire extinguishers. The sulphate is mixed with sodium carbonate or hydrogencarbonate to produce carbon dioxide and aluminium (III) hydroxide, Al(OH)3, which mix together to formthe foam. The foam is effective in excluding air from oil fires hence helping to put the fire off.
    • Aluminium sulphate is used in the tanning of leather. It is also used as a fertilizer.
    Iron (II) sulphate is used:
    • in the manufacture of ink and dye
    • in tanning leather (iron-tanning)
    • to make tablets prescribed to patients with iron deficiency.
    • to prepare a reddish-brown iron (III) oxide (‘red oxide’) which is used as a pigment.
    • as a weed killer ( herbicide) and as a fungicide.
    • for treating sewage and water.
    • to coagulate ( bind together) blood in slaughterhouses.
    Copper (II) sulphate finds many uses which include:
    • Manufacture of copper fungicides (CuSO4.5H2O) such as red and blue copper, which is sprayed on crops to prevent certain species of fungi.
    • Manufacture of certain green pigments.
    • Copper (II) sulphates is used for makingwashes such as a “Bordeaux mixture”, used in sprayingvines and potatoes to kill moulds which would otherwise injure the plants.
    • Manufacture of insecticides such as copper arsenite and Paris green for control of fungus diseases.
    • Correction of copper deficiency in soils and in animals
    • It is used as a growth stimulant in pigs and broiler chickens.
    • It is also used as a molluscide for the destruction of slugs and snails, particularly the snail, which is a host of liver fluke.
    • Copper sulphate is used as a timber preservative for the prevention of wood rot.
    • It is used as an electrolyte in copper-plating and as a catalyst in preparation of ethanol.
    Zinc sulphate has the following uses:
    • it is used as an emetic and for treatment of certain skin diseases.
    • Hydrated zinc sulphate (ZnSO4.7H2O) is used as an agent in printing and textile dyeing, as an antiseptic and in preserving wood and hides. It is also in zinc-plating by electrolysis.
    • It is used as an antibacterial treatment for sewage, a miticide and an herbicide.
    • It is used as a component of cosmetics (such as skin fresheners) and an ingredient in some deodorants.
    • Diluted “White Vitriol”, (ZnSO4.7H2O), is used in medicine in the preparation of eye lotions and mouth washes. It is also used to assists the healing of wounds.
    • Ammonium sulphate is used as a fertilizer.
    • Magnesium sulphate, in the form of Epsom salt (MgSO4.7H2O), is used as a mild purgative.
    • Barium sulphate (BaSO4) is used in the manufacture of white pigments, in white paints.
    • The alums are used in dyes and in leather industry. Alums are double salts of general formula, X2SO4.Y2(SO4)3.24H2O, whereX is Na, K or NH4and Y is Fe(III), Al or Cr. The two commonest alums are: Potash alum, K2SO4.Al2(SO4)3.24H2O (colourless); and; Iron (III) alum, (NH4)2SO4.Fe2(SO4)3.24H2O(purple
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