compounds of metals

Preparation of Oxides of Some Metals by Direct and Indirect Methods

Prepare 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

Classify 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

Demonstrate 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.

Metal	How it reacts	Product
Barium	burns with a green flame	white solid (barium (II) oxide, BaO)
Calcium	burns with a brick-red flame	white solid (calcium oxide, CaO)
Sodium	burns with a yellow flame	white solid (sodium oxide, Na2O)
Potassium	burns with a purple flame	white solid (potassium oxide, K2O)
Magnesium	burns with a white flame	white solid (magnesium oxide, MgO)
Iron	burns with yellow sparks	Blue-black solid (iron (II) oxide, FeO)
Copper	doesnot burn, turns black	black solid (copper (II) oxide, CuO)

Activity

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

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.Mortar is made by mixing slaked lime, sand and water, and is used in sticking bricks together and in forming smooth surfaces on walls of buildings.
2. Calcium oxide is used for making whitewash, 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: Cement is made by heating together lime or limestone and clay. The product is a mixture of calcium silicates and aluminates. Clay is hydrated aluminium silicate. A mixture of cement, sand, stones and water gives concrete, 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 of glass: heating a mixture of sand, sodium carbonate and lime or limestone gives 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.

Hydroxides

Preparation of Hydroxides of some Metals by Direct and Indirect Methods

Prepare hydroxides of some metals by direct and indirect methods

Metal hydroxides are electrovalent compounds, composed of metallic ions, which are positively charged, and hydroxyl ions, OH-, which are negatively charged. The nature of the hydroxides of the metals varies according to the position of the metal in the reactivity series, as shown below:

Metal hydroxides can be prepared in the laboratory by two methods:

1. Direct method
2. Indirect method

Direct method

This method involves such processes like adding metals directly in water. The method is suitable for preparation of soluble hydroxides (alkalis).

1. 2Na(s)+ 2H2O(l)→ 2NaOH(aq)+ H2(g)
2. 2K(s)+ 2H2O(l)→ 2KOH(aq)+ H2(g)
3. Ca(s)+ 2H2O(l)→ Ca(OH)2(aq)+ H2(g)

Indirect method

This method involves the preparation of insoluble hydroxides by reacting aqueous solutions of sodium or potassium hydroxide with aqueous salts of metals. These are called precipitation reactions. Only NH4OH, KOH and NaOH are completely soluble in water. Calcium hydroxide is sparingly soluble in water (0.173g/100 ml at 20°C). All other hydroxides are insoluble in water and they can be prepared by this method:

• CuCl2(aq)+ 2NaOH(aq)→ Cu(OH)2(s)+ 2NaCl(aq)
• Ionically:Cu2+(aq)+ 2OH-(aq)→ Cu(OH)2(s)
• FeSO4(aq)+ 2KOH(aq)→ Fe(OH)2(s)+ K2SO4(aq)
• Ionically:Fe2+(aq)+ 2OH-(aq)→ Fe(OH)2(s)
• FeCl3(aq)+ 3NaOH(aq)→ Fe(OH)3(s)+ 3NaCl(aq)
• Ionically:Fe3+(aq)+ 3OH-(aq)→ Fe(OH)3(s)

Metal Hydroxides

Classify metal hydroxides

Just like metal oxides, metal hydroxides can be classified based on their solubility in water as either soluble or insoluble hydroxides. They can also be classified as basic or amphoteric hydroxides, based on their reaction with acids and bases.

The classification of hydroxides as either soluble or insoluble is as shown on the previous page.

Based on their reaction with acids and bases, metal hydroxides are classified into basic and amphoteric hydroxides.

A basic hydroxide is a metal hydroxide that contains hydroxyl ions, OH-, and will react with an acid to form a salt and water only. An amphoteric hydroxide is a hydroxide that shows both acidic and basic properties, that is, will react with both an acid and a base.Examples of amphoteric hydroxides are Al(OH)3, Zn(OH)2and Pb(OH)2.

Amphoteric hydroxides behave as bases and as acids under different conditions. These hydroxides are only weakly basic, but like other bases, they still combine with acids to form salt and water:

• Zn(OH)2(s)+ H2SO4(aq)→ ZnSO4(aq)+ 2H2O(l)
• Pb(OH)2(s)+ 2HNO3(aq)→ Pb(NO3)2(aq)+ 2H2O(l)
• Al(OH)3(s)+ 3HCl(aq)→ AlCl3(aq)+ 3H2O(l)

However, in the presence of strong alkalis, these hydroxides behave like acids, and they combine with these strong alkalis to yield a salt and water:

The Chemical Properties of Metal Hydroxides

Explain the chemical properties of metal hydroxides

In this particular case, the chemical properties of only a few common hydroxides will be discussed.

Chemical properties of sodium hydroxide (NaOH) and potassium hydroxide (KOH)

When heated together with aluminium, zinc and lead metals, concentrated sodium hydroxide dissolves these metals to form sodium aluminate, zincate and plumbate respectively.

Caustic alkalis react with soluble salts of certain metals like copper, lead, zinc and iron to form their insoluble hydroxides by double decomposition.

• 3NaOH(aq)+ FeCl3(aq)→ Fe(OH)3(s)+ 3NaCl(aq)
• 2NaOH(aq)+ ZnSO4(aq)→ Zn(OH)2(s)+ Na2SO4(aq)

Sodium hydroxide liberates ammonia gas from ammonia salts when both are warmed together. This is the standard test for any ammonium salt.

• NH4Cl(aq)+ NaOH(aq)→ NaCl(aq)+ H2O(l)+ NH3(g)
• NH4NO3(aq)+ NaOH(aq)→ NaNO3(aq)+ H2O(l)+ NH3(g)

When carbon dioxide is bubbled through aqueous solutions of the caustic alkalis, the carbonates are formed. With excess of the gas, the hydrogencarbonates are formed.

• 2NaOH(aq)+ CO2(g)→ Na2CO3(aq)+ H2O(l)
• Na2CO3(aq)+ H2O(l)+ CO2(g)→ 2NaHCO3(aq)

Chlorine reacts with excess of cold dilute caustic alkalis to form the hypochlorite, (NaClO or KClO).

• 2NaOH(aq)+ Cl2(g)→ NaCl(aq)+ NaClO(aq)+ H2O(l)
• 2KOH(aq)+ Cl2(g)→ KCl(aq)+ KClO(aq)+ H2O(l)

If excess chlorine is bubbled through hot concentrated solutions of caustic alkalis, the chlorates are formed, (NaClO3 or KClO3).

• 6NaOH(aq)+ 3Cl2(g)→ 5NaCl(aq)+ NaClO3(aq)+ 3H2O(l)
• 6KOH(aq)+ 3Cl2(g)→ 5KCl(aq)+ KClO3(aq)+ 3H2O(l)

Caustic alkalis are strong bases, which react with acids by neutralization reactions:

• NaOH(aq)+ HCl(aq)→ NaCl(aq)+ H2O(l)
• 2KOH(aq)+ H2SO4(aq)→ K2SO4(aq)+ 2H2O(l)

Chemical properties of Ca(OH)2

Action of heat

All hydroxides except potassium hydroxide (KOH) and sodium hydroxide (NaOH) decompose under the action heat to the oxide and steam. Calcium hydroxide undergoes thermal decomposition to calcium oxide and steam. Ca(OH)2(aq)→CaO(s)+ H2O(g)

Reaction with acids

Calcium hydroxide dissolves readily in dilute hydrochloric acid and nitric acid to form the corresponding calcium salts: Ca(OH)2(aq)+ 2HCl(aq)→CaCl2(aq)+ H2O(l) ;Ca(OH)2(aq)+ 2HNO3(aq)→Ca(NO3)2(aq)+ 2H2O(l)

Its reaction with dilute sulphuric acid is unsatisfactory due to the formation of calcium sulphate, which tends to precipitate on any dissolved lime.

Reaction with carbon dioxide

A solution of calcium hydroxide in water is called limewater. Carbon dioxide turns limewater “milky” due to the precipitation of white particles of calcium carbonate. Ca(OH)2(aq)+ CO2(g)→CaCO3(s)+ H2O(l)

Reaction with ammonium salts

Any ammonium salt will release ammonia gas when heated with calcium hydroxide solution. This is the common laboratory method for preparation of ammonia gas. Ca(OH)2(aq)+ 2NH4Cl(s)→2NH3(g)+ CaCl2(aq)+ 2H2O(l); Ionically; NH+4(s)+ OH-(s)→NH3(g)+ H2O(l)

Reaction with chlorine

• When chlorine gas is passed over moist cold solid calcium hydroxide, bleaching powder is formed. The formula of the bleaching powder is complex, and it is known with certainty but it probably contains Ca2+, Cl-and OH-ions and water. The reaction equation for its formation is usually written in “approximate” form: Ca(OH)2(s)+ Cl2(g)→ CaOCl2(s)+ H2O(l).The accepted but highly approximate formula of bleaching powder is Ca(OCl)2. It is known to contain a mixture of calcium hypochlorite, Ca(OCl)2, and basic calcium chloride, CaCl2.Ca(OH)2.H2O.
• When chlorine gas is bubbled in a cold milk of lime, a mixture of calcium chloride and calcium hypochlorite, Ca(OCl)2, is formed. 2Ca(OH)2(aq)+ 2Cl2(g)→ CaCl2(aq)+ Ca(OCl)2(aq)+ 2H2O(l)
• When chlorine is bubbled through hot milk of lime, calcium chloride and calcium chlorate are formed. 6Ca(OH)2(aq)+ 6Cl2(g)→ 5CaCl2(aq)+ Ca(ClO3)2(aq)+ 6H2O(l)

Reaction with hydrogencarbonates

Calcium hydroxide removes temporary hardness of water by precipitating the carbonate from the hydrogencarbonate. Ca(HCO3)2(aq)+ Ca(OH)2(aq)→2CaCO3(s)+ 2H2O(l)

The Uses of Metal Hydroxides

Describe the uses of metal hydroxides

Uses of NaOH

1. Sodium hydroxide is used in the industrial manufacture of sodium, soaps and in the extraction of aluminium from bauxite.
2. It is used in the manufacture of paper, dyes and bleach.
3. Sodium hydroxide is a very common laboratory chemical. It is used in the qualitative as well as quantitative analysis.
4. It is used in the preparation of insoluble metal hydroxides.
5. It is used in the manufacture of artificial textile fibres.
6. It is used to produce mineral salts by neutralization reactions with mineral acids.

Uses of KOH

1. Potassium hydroxide is used in the manufacture of soft soap.
2. Use as an electrolyte: aqueous potassium hydroxide is used as an electrolyte in alkaline batteries.
3. Manufacture of biodiesel: potassium hydroxide works well in the manufacture of biodiesel by saponification of the fats in vegetable oil.4. Preparation of salts: many potassium salts are prepared by neutralization reactions involving potassium hydroxide.

Uses of Ca(OH)2

1. Treatment of acid soils: Calcium hydroxide is a cheap alkali which can be used in large quantities to treat acid soils. The alkali neutralizes the excessive soil acidity, just like any alkali reacts with an acid, to form salt and water.
2. Softening of hard water: Calcium hydroxide, in precisely calculated quantities, is used in the softening of temporary hard water as discussed early in chapter three. Ca(OH)2(aq)+ Ca(HCO3)2(aq)→ 2CaCO3(s)+ 2H2O(l)
3. Preparation of mortar and whitewash: Mortar is prepared by mixing 1 part of slaked lime to 3 parts of sand into a paste with water. Slaked lime mixed with sand and water is used to stick bricks together.Whitewash is a thick suspension of slaked lime in water. It is smeared on the walls of buildings to give them a smooth protective finish.
4. Manufacture of bleaching powder: Bleaching powder is manufactured by passing chlorine gas over moist calcium hydroxide as discussed previously in this chapter.
5. Manufacture of paperFirst, a suspension of calcium hydroxide in water (milk of lime) is treated with sulphur dioxide gas to form calcium hydrogensulphite. Ca(OH)2(aq)+ 2SO2(g)→ Ca(HSO3)2(aq)Then, a solution of calcium hydrogen sulphite is used to remove the lignin from wood, leaving cellulose, which is used in paper manufacture. The action of removing lignin from wood is called bleaching of pulp.
6. Manufacture of paints:Calcium hydroxide is used in the manufacture of undercoat paints which are applied as the first coat on plaster walls or on wood before applying the final gloss paint.
7. Extraction of metals:Sodium hydroxide is used in the extraction of aluminium from bauxite ore.
8. Liberation of ammonia: Calcium hydroxide is used in the Solvary Process to produce ammonia from ammonium chloride.
9. Calcium hydroxide is a very important reagent inqualitative and quantitative analysis. It is used to determine the concentrations of acids in volumetric analysis, and in detection of ions present in unknown solutions in qualitative analysis.

Uses of Mg(OH)2

1. Magnesium hydroxide is used as an antacid to neutralize stomach acid, and as a laxative. It is sold for medical use as chewable tablets, capsules, and as liquids having various added flavours. It is primarily used to alleviate constipation, but also relieve indigestion and heartburn.
2. Magnesium hydroxide is also used as an antiperspirant and as armpit deodorant.
3. Milk of magnesia (a suspension of magnesium hydroxide in water) is also applied and massaged onto the scalp, a few minutes before washing, to relieve symptoms ofseborrhoeaanddandruff. An additional use is for the treatment of acne or oily skin by applying topically, allowing to dry, and then washing it off the face (or other body part). It is also said to be used for seborrhoeaic dermatitis, which is a drying and flaking of the skin similar to dandruff but often occurring on the face.
4. Magnesium hydroxide powder is used industrially as a non-hazardous alkali toneutralizeacidic waste waters.
5. Solid magnesium hydroxide is used as fire and smoke retardant.

Carbonates and Hydrogen Carbonates

Preparation of Metal Carbonates and Hydrogen Carbonates by Different Methods

Prepare metal carbonate and hydrogen carbonates by different methods

here are many types and forms of metal carbonates in the earth’s crust. The most common and important carbonate is calcium carbonate which occurs naturally in the form of chalk, limestone or marble. Carbonates of other metals such as iron, copper, manganese, lead and zinc also occur naturally.Shells of snail, tortoise, fish and eggs are made of a great deal of carbonates.

There are four solid hydrogencarbonates. These are potassium, sodium, lithium and ammonium hydrogencarbonates. The hydrogencarbonates of calcium and magnesium occur in solution forms. Hydrogencarbonates are bases just like other ordinary carbonates.All metal hydrogencarbonates are soluble in water. Aluminium, zinc, iron, lead and copper hydrogencarbonates do not exist.

The method used to prepare carbonates depends on whether the carbonate is soluble in water or not. Sodium, potassium and ammonium carbonates are the only soluble metal carbonates.Solublecarbonatesare prepared by passing carbon dioxide gas to the alkali. For example, sodium carbonate is formed when carbon dioxide gas is blown through sodium hydroxide solution.

CO2(g)+ 2NaOH(aq)→Na2CO3(aq)+ H2O(l)

If more carbon dioxide is bubbled through the solution, a second reaction occurs and sodium hydrogencarbonate is formed.Na2CO3(aq)+ H2O(l)+ CO2(g)→2NaHCO3(aq)

This is the easiest and convenient way for preparing hydrogencarbonates. Insoluble carbonates can be prepared by precipitation reactions. This involves adding a soluble carbonate to a solution of a salt of heavy metal. For example, when a solution of zinc carbonate is mixed with sodium carbonate solution, a precipitate of zinc carbonate is formed. ZnSO4(aq)+ Na2CO3(aq)→ZnCO3(s)+ Na2SO4(aq)

Likewise, copper carbonate can be precipitated by mixing copper (II) chloride solution with potassium carbonate solution.CuCl2(aq)+ K2CO3(aq)→CuCO3(s)+ 2KCl(aq)

Classification ofMetal Carbonates

Classify metal carbonates

Metal carbonates are classified based on their solubility in water. Classified on this basis, we have soluble and insoluble carbonates. Potassium, sodium and ammonium carbonates are soluble in water. All other carbonates are insoluble in water.

The Chemical Properties of Metal Carbonates

Analyse the chemical properties of metal carbonates

Reaction with acids

Carbonates and hydrogencarbonates react with dilute acids to produce carbon dioxide, salt and water. For example:

CaCO3(s)+ 2HCl(aq)→CO2(g)+ CaCl2(aq)+ H2O(l)

2NaHCO3(aq)+ H2SO4(aq)→ 2CO2(g)+ Na2SO4(aq)+ 2H2O

Action of heat

Excluding sodium and potassium carbonates, all other carbonates decompose on heating to form oxides of corresponding metals and carbon dioxide. For example:PbCO3(s)→PbO(s)+ CO2(g)

A hydrogencarbonate decomposes to give a carbonate, carbon dioxide and water.2NaHCO3(aq)+ Na2CO3(s)→H2O(l)+CO2(g)

Distinction between carbonates and hydrogencarbonates

We can easily distinguish a carbonate from a hydrogencarbonate by the use of magnesium sulphate.When a carbonate solution is added to a magnesium sulphate solution, a white precipitate of magnesium carbonate is formed:Na2CO3(aq)+ MgSO4(aq)→MgCO3(s)+ Na2SO4(aq)

When a similar reaction is performed with a hydrogencarbonate solution, no white precipitate is formed:2NaHCO3(aq)+ MgSO4(aq)→Mg(HCO3)2(aq)+ Na2SO4(aq)

This is because magnesium hydrogencarbonate is soluble in water. However, when the solution is boiled, magnesium hydrogencarbonate decomposes to give a white precipitate of magnesium carbonate.Mg(HCO3)2(aq)→MgCO3(s)+ H2O(l)CO2(g)

The Uses of Carbonates and Hydrogen Carbonates

Describe the uses of carbonates and hydrogen carbonates

For easy understanding of the uses of carbonates and hydrogencarbonates, each of the most common of these compounds will be dealt with separately.

Uses of sodium carbonate, Na2CO3

1. It is widely used in the manufacture of sodium silicate, which is used to manufacture glass (for more details, and the reaction equations involved, refer to the uses of calcium carbonate).
2. It is used to manufacture soap and paper.
3. Manufacture of sodium hydroxide: sodium hydroxide is prepared by adding slaked lime (calcium hydroxide) to sodium carbonate and the mixture continuously stirred. A precipitate of calcium carbonate forms and is filtered off: Na2CO3(aq)+ Ca(OH)2(s)→ 2NaOH(aq)+ CaCO3(s).The filtrate, which is a solution of sodium hydroxide, is then evaporated to obtain crystals of sodium hydroxide.
4. It is used in volumetric analysis to determine the concentration of acids (standardize acids).
5. Sodium carbonate is used for softening hard water.
6. It is used in medicine as an antacid.
7. It is also important inphotographyand the textile industry (where itis used to facilitate the chemical bonding between the dye and the fibre).
8. Manufacture of ‘water glass’: When sodium carbonate is heated together with silicon dioxide (SiO2), sodium silicate (NaSiO3) and carbon dioxide are produced. Na2CO3(s)+ SiO2(s)→ NaSiO3(s)+ CO2(g).A concentrated solution of sodium silicate in water, known as water glass is used as a preservative for eggs. It is also used asan adhesive in paper making and in television tubes.

Uses of sodium hydrogencarbonate, NaHCO3

This is the most important hydrogencarbonate. The following are some of its uses:

1. Health salt: The salt finds a variety of medicinal uses, particularly in stopping diarrhoea and, as anantacidto treat heartburn, indigestion, and other stomach disorders. It is also used to treat various kidney disorders and to increase the effectiveness of sulphonamides.A heath salt is a mixture of sodium hydrogencarbonate, tartaric acid and sodium potassium tartrate.Acting as antacid, sodium hydrogencarbonate neutralizes the hydrochloric acid produced in the stomach. Because it is very soluble, it acts very fast thus providing a quick relief for symptoms caused by excess acid in the stomach.
2. It is used in the removal of grease from clothes and other articles. Most of the cleaning agents that are used for the removal of stains tend to contain a small amount of sodium bicarbonate as a very important ingredient.
3. Use in industry: sodium bicarbonate is used in textile industry for dyeing and printing operations; in leather industry as a neutralizer of dyeing agents in tanning processes; and in rubber and plastic industry as a blowing agent, as it releases carbon dioxide which is used to shape the object that is made from the rubber and plastic.
4. Baking: sodium bicarbonate is widely used to bake breads and cakes. The baking soda contains a large quantity of sodium hydrogencarbonate and is always accompanied by a small quantity of acid phosphate. When baking soda mixed with dough is heated, the sodium hydrogencarbonate decomposes to give bubbles of carbon dioxide which make the dough rise: 2NaHCO3(s)→ Na2CO3(s)+ CO2(g)+ H2O(l).Commercial baking powder is a mixture of baking soda and organic acids such as citric or tartaric acid.

Uses of calcium carbonate, CaCO3

Remember we learned early that calcium carbonate exists in three forms: chalk, limestone and marble. So, in this context, the uses of three forms will be discussed.

1. Sodium carbonate is used in the blast furnace for extraction of iron from its ore.
2. It is used in the manufacture of glass. Heating a mixture of calcium carbonate (CaCO3), sodium carbonate (Na2CO3) and sand (SiO2) very strongly, at a temperature of 1300 to 1400°C, produces a clear melt. When this melt is cooled down, it hardness to form glass.Upon heating these compounds, carbon dioxide is given off, leaving a mixture of sodium silicate and calcium silicate with excess silicon dioxide: CaCO3(s)+ SiO2(s)→ CaSiO3(s)+ CO2(g)Na2CO3(s)+ SiO2(s)→ Na2SiO3(s)+ CO2(g).Glass is a mixture of sodium and calcium silicates.
3. Limestone is used to make cement. Limestone powder and clay are roasted and the product is grinded to form a fine powder called cement.
4. Calcium carbonate is used in the manufacture of lime, which is used as a fertilizer.
5. Marble is used on walls of houses and other buildings to make them look attractive.
6. Limestone is used in road construction. Small pieces of marble are stuck together and then covered with tar to gives a smooth hard surface, which can last for several years.
7. A mixture of limestone and linseed oil is used as putty, which helps in retaining the structure of windows, doors and other wooden structures in position in houses and buildings.
8. Calcium carbonate, in the fine precipitate form, can be used in the manufacture of toothpaste.

Nitrates

Preparation of Metal Nitrates

Prepare metal nitrates

Nitrates are important compounds widely known for industrial purposes. Sodium nitrate has been known for quite a long time. It occurs in nature as saltpetre (NaNO3).

Nitrates are usually prepared by methods which involve crystallization. This can be done by reacting nitric acid with metals, oxides, hydroxides or carbonates.

1. In the laboratory, sodium nitrate may be prepared by neutralizing sodium hydroxide solution with nitric acid.NaOH(aq)+ HNO3(aq)→ NaNO3(aq)+ H2O(l)
2. Calcium nitrate may be made in the laboratory by the action of nitric acid upon calcium carbonate.CaCO3(s)+ 2HNO3(aq)→ Ca(NO3)2(aq)+ H2O(l)+ CO2(g)
3. Ammonium nitrate can be made by neutralization of ammonia solution with nitric acid.NH3(aq)+ HNO3(aq)→ NH4NO3(aq)
4. Lead nitrate can be obtained by the reaction between lead oxide and dilute nitric acid.PbO(s)+ 2HNO3(aq)→ Pb(NO3)2(aq)+ H2O(l)
5. Copper nitrate can be prepared by dissolving copper oxide in dilute nitric acid.CuO(s)+ 2HNO3(aq)→ Cu(NO3)2(aq)+ H2O(l)
6. Nitrates of certain metals can be prepared by reacting the metals with dilute or concentrated nitric acid.3Pb(s)+ 8HNO3(aq)→ 3Pb(NO3)2(aq)+ 4H2O(l)+ 2NO(g);3Cu(s)+ 8HNO3(aq)→ 3Cu(NO3)2(aq)+ 4H2O(l)+ 2NO(g)

However, the nitrates of common heavy metals, except lead nitrate, are very soluble in water and they are deliquescent. This makes it a matter of grater difficult to prepare their crystals. All the crystals are white in colour except those of copper (II) nitrate, which are blue. All nitrates are soluble in water.

The Chemical Properties of Metal Nitrates

Explain the chemical properties of metal nitrates

Action of heat

The chemical properties of metal nitrates vary according to the position of the metal in the reactivity series. Metal nitrates give a variety of products when thermally decomposed. This fact is summarized below:

Ammonium nitrate is decomposed by heat into dinitrogen oxide and water: NH4NO3(s)→N2O(g)+2H2O(l)

The brown ring test

All nitrates undergo the same reaction with iron (II) sulphate and concentrated sulphuric acid and this reaction becomes a test for the soluble nitrates, the brown ring test. This test is carried out by crushing a little potassium nitrate, putting it in a test tube, adding water to a depth of about 2 cm and then shaking the test tube to dissolve the potassium nitrate (note that any metal nitrate could have been used instead of potassium nitrate). This is followed by adding a little sulphuric acid and then two or three crystals of iron (II) sulphate, which have also been crushed. The contents are shaken to dissolve them. Finally, the test tube is held in a slanting position and a slow continuous stream of concentrated sulphuric acid is poured down the side of the test tube. The acid forms a separate layer underneath the aqueous layer and, at the junction of the two, a brown ring will be seen. This brown ring is the characteristic test for all soluble nitrates.

Explanation

• The concentrated sulphuric

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