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School Chemistry: Aspects of Industrial Chemistry: LIMESTONE - its chemistry, uses

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1. Limestone, its chemistry, uses and the thermal decomposition of carbonates

(Suitable for AQA, Edexcel and OCR GCSE chemistry students)

The chemistry of carbonates

including their reaction with acids and thermal decomposition

Sub-index for this page

(a) Limestone in and from the landscape

(b) Limestone - useful material from part of the Carbon Cycle

(c) The Limestone Cycle

(d) The thermal decomposition of other carbonates

(e) Uses of products made from limestone and ceramic materials like clay

(f) Some other uses of calcium oxide and calcium hydroxide

(g) Issues with limestone quarrying and associated industries

(h) More on the formulae of calcium and magnesium compounds and more reactions including the reaction of oxides, hydroxides and carbonates with acids

(i) Thermal decomposition of hydroxides and nitrates

limestone cliffsuseful limestone


(a) Limestone in and from the landscape

This page describes the origin and chemical nature of limestone and marble (both mainly calcium carbonate) and the conversion of limestone to quicklime and slaked lime and their uses in industry and agriculture. The thermal decomposition of carbonates, hydroxides and nitrates is also described and discussed. Some quizzes are listed at the end of the page.

EQUATION NOTE: The equations are often written three times: (i) word equation, (ii) balanced symbol equation without state symbols, and, (iii) with the state symbols (g), (l), (s) or (aq) to give the complete balanced symbol equation.

There are several limestone quarries in the limestone country of the Yorkshire Dales. The quarrying does present a scar on the landscape BUT limestone is a very useful mineral and used as a building stone, in iron extraction in the blast furnace and for making lime for agriculture and kitchen garden. There often has to be a compromise somewhere along the line since many of a countries important mineral resources and rocks are in some of the most beautiful scenic parts of the country!

Craven Villages Walk

A mineral train of limestone or lime filled wagons from Swinden Quarry in the Yorkshire Dales. The local stone walls, barns and houses are also made of limestone - a useful naturally occurring resource which has been exploited for at least 2000 years in the UK starting with the Romans (if not before?).

York Minster is built of limestone, but much of the stone has to be replaced due to weathering.

limestone cliffs (b) Limestone - a very useful material and part of the Carbon Cycle useful limestone
  • Limestone, is a sedimentary rock formed by the mineral and 'shelly' remains of marine organisms, including coral, that once lived in warm shallow fertile seas.

    • The formation of limestone and conversion to lime and cement are parts of the carbon cycle.

      For more details see the diagram above and ...

      Carbon cycle  gcse biology revision notes and

      FOSSIL FUELS, coal, oil & natural gas and the Carbon Cycle gcse chemistry

    • It is chemically mainly calcium carbonate and is a useful material that is quarried and used directly as a building material. The issues of limestone quarrying are discussed in a separate section.

    • Limestone is very useful building material and many great civic buildings, cathedrals and churches are built from limestone from a local quarries.

      • It is also used in carving statues and ornamental decoration on churches, but the medieval buildings have suffered much erosion due to acid rain from the fossil fuel burning of the industrial revolution.

      • Crushed limestone is used to make chippings for road surfaces.

        • These are direct uses of limestone, but it can be processed with other materials to make equally useful products.

        • See more on uses of limestone further down the page.

    • It reacts with acids - 'fizzing' due to carbon dioxide formation - test with 'limewater' - milky white precipitate.

    • Marble is also made of calcium carbonate and is a metamorphic rock formed by the action of heat and pressure on limestone in the Earth's crust. It is a much harder rock than limestone and is used to make highly polished and finely carved stone sculptures, statues etc.

  • Chemically, limestone mainly consists of calcium carbonate, CaCO3, and is a valuable natural mineral resource, quarried in large quantities in many countries. (See uses of limestone)

  • Other uses of limestone rock are outlined below and it is an important raw material in the manufacture of cement, concrete, glass and iron.

  • Powdered limestone can be used to neutralise acidity in lakes and soils. (neutralisation chemistry).

  • Like lime, limestone is a safe agri-chemical to use on the land and does NOT produce any controversial side effects of artificial fertilisers, herbicides and pesticides etc.

  • What happens when limestone is strongly heated?

    • Most metal carbonates you will come across will break down on strong heating (thermal decomposition) to give a metal oxide and carbon dioxide gas.

  • When limestone is heated in a kiln at over 825-900oC, it breaks down into quicklime (calcium oxide) and carbon dioxide. Both are useful products.

  • This type of reaction is endothermic (heat absorbed in the chemical change)

  • It is also an example of thermal decomposition (and other carbonates behave in a similar way).

    • A thermal decomposition is when a compound is heated until it breaks down into at least two products.

    • The thermal decomposition of calcium carbonate

    • calcium carbonate (limestone) ==> calcium oxide (quicklime) + carbon dioxide

    • CaCO3 (c) doc b CaO + CO2

    • CaCO3(s) (c) doc b CaO(s) + CO2(g)   (symbol equation with state symbols)

      • Note State symbols in equations: (g) = gas, (l) = liquid, (s) = solid, (aq) = aqueous solution in water

      • If you heat lumps of limestone in a crucible to as high a temperature as can get with a bunsen burner, the thermal decomposition results in a fine powder reside of quicklime (lime).

        • When cooled, if you carefully add water, you get a very exothermic reaction as calcium hydroxide is formed.

        • quicklime + water ==> slaked lime

        • calcium oxide + water ==> calcium hydroxide

        • CaO + H2O ==> Ca(OH)2

        • CaO(s) + H2O(l) ==> Ca(OH)2(aq)

      • Quicklime (calcium oxide) is used in steel making and spread on land to reduce soil acidity. (See uses).

      • Calcium oxide or quicklime is also just known as burnt lime.

      • When heated to a very high temperature, calcium oxide emits intense white light and this property was used in Victorian lamps in the street and buildings.

      • The origin of the phrase 'in the limelight' derives from the use of this phenomena in the theatre.

    • This is a reversible endothermic reaction. To ensure the change is to favour the right hand side, a high temperature of over 900oC is needed as well as the continual removal of the carbon dioxide.

    • The high temperature needed is produced by mixing the limestone with coal/coke (a fuel of mainly carbon) and blowing hot air into the ignited mixture in a rotating kiln for a continuous production line (raw materials in at one end, lime out the other!)  ....

      • C(s) + O2(g) ==> CO2(g) is very exothermic - heat releasing!

    • Note on heating other carbonates you get a similar thermal decomposition.

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(c) The Limestone Cycle

  • Quicklime reacts very exothermically with water to produce slaked lime (solid calcium hydroxide)

    • calcium oxide (quicklime) + water ==> calcium hydroxide (slaked lime)

      • this is a very exothermic reaction, the quicklime 'puffs' up and steam is produced!

        • CaO + H2O ==> Ca(OH)2

        • CaO(s) + H2O(l) ==> Ca(OH)2(s)   (symbol equation with state symbols)

      • With excess water followed by filtration you get calcium hydroxide solution or limewater.

      • An exothermic reaction is one that gives out heat energy as the chemical change takes place.

      • Slaked lime or calcium hydroxide is also known as hydrated lime, builders lime, slack lime.

diagram of the Limestone Cycle (c) doc b

The chemical reactions involved in the LIMESTONE CYCLE

The symbol equations have state symbols too

Ca(OH)2(aq) + CO2(g) ==> CaCO3(s) + H2O(l) calcium carbonate


CaCO3(s) ==> CaO(s) + CO2(g)
bubble carbon dioxide through limewater to give the milky white precipitate of calcium carbonate   heat limestone to over 900oC to form white calcium oxide powder and carbon dioxide gas
calcium hydroxide solution


LIMESTONE CYCLE calcium oxide


dissolving calcium hydroxide in water - the resulting aqueous solution is known as limewater   add water to calcium oxide, exothermic reaction to give white powder of calcium hydroxide
Ca(OH)2(s) + aq ==> Ca(OH)2(aq) calcium hydroxide

(solid slaked lime)

CaO(s) + H2O(l) ==> Ca(OH)2(s)

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(d) The thermal decomposition of other carbonates

  • Other carbonates show a similar thermal decomposition to calcium carbonate e.g. using the simple apparatus illustrated below. Unlike limestone, some carbonates can be decomposed in the school/college laboratory using a bunsen burner and pyrex test/boiling tube. You can also use a crucible, but you couldn't collect the gas to test for carbon dioxide with limewater.

  • The carbonate is heated strongly in a pyrex boiling tube with a bunsen burner and look out for colour changes from the solid reactant to the solid residue product.

  • If the gases given off are VERY carefully bubbled into limewater, the formation of a white precipitate shows that carbon dioxide was formed - make sure the gases do NOT suck back into the hot glass test tube!.

  • The residues an oxide, and if all the carbonate has decomposed, the addition of acid should not produce any fizzing due to carbon dioxide formation, however, the oxide might well dissolve in the acid to give a salt solution.

  • You need to take care that the limewater does not suck back into the hot pyrex tube or it will crack it, no need for a nasty little accident to happen! You can further minimise this risk before starting to heat the carbonate, by tilting the pyrex tube down slightly towards the test tube of limewater.

  • You can't do this experiment with limestone because calcium carbonate has much too high a decomposition temperature (> 900oC).

  • The thermal decomposition of copper carbonate

  • copper(II) carbonate(green)  ==> copper(II) oxide(black)  + carbon dioxide

    • CuCO3 ==> CuO + CO2

    • CuCO3(s) ==> CuO(s) + CO2(g)   (symbol equation with state symbols)

    • The colour change from the dark green copper carbonate to the jet black copper oxide is clearly observed.

  • The thermal decomposition of zinc carbonate

  • zinc carbonate(white)  ==> zinc oxide(yellow hot, white cold)  + carbon dioxide

    • ZnCO3 ==> ZnO + CO2

    • ZnCO3(s) ==> ZnO(s) + CO2(g)   (symbol equation with state symbols)

    • Both the zinc carbonate and zinc oxide are white, but zinc oxide turns yellow when very hot, on cooling at the end of the experiment in turns white.

    • Zinc carbonate occurs as the mineral ores calamine/Smithsonite and the resulting zinc oxide can be used to extract zinc metal and zinc oxide itself is used as a whitening agent' in cosmetics and in 'calamine lotion' a mild antiseptic and antipruritic (anti-itching agent) for treating skin irritations.

  • The thermal decomposition of magnesium carbonate

  • magnesium carbonate(white)  ==> magnesium oxide(white)  + carbon dioxide

    • MgCO3 ==> MgO + CO2

    • MgCO3(s) ==> MgO(s) + CO2(g)   (symbol equation with state symbols)

    • Both magnesium carbonate and magnesium oxide are white.

  • FeCO3 , PbCO3 and MnCO3 behave in a similar way

    • The thermal decomposition of iron carbonate

    • iron(II) carbonate(s, dark green)  ==> iron(II) oxide(s, black)  + carbon dioxide

      • FeCO3 ==> FeO + CO2

      • FeCO3(s) ==> FeO(s) + CO2(g)   (symbol equation with state symbols)

      • The colour change is from a dark green reactant to a black solid residue.

    • The thermal decomposition of lead carbonate

    • lead(II) carbonate(s, white)  ==> lead(II) oxide(s, yellow-orange)  + carbon dioxide

      • PbCO3 ==> PbO + CO2

      • PbCO3(s) ==> PbO(s) + CO2(g)   (symbol equation with state symbols)

      • The colour change is from white to yellow-orange.

    • The thermal decomposition of manganese carbonate

    • manganese(II) carbonate(s, pale pink)  ==> manganese(II) oxide(s, white)  + carbon dioxide

      • MnCO3 ==> MnO + CO2

      • MnCO3(s) ==> MnO(s) + CO2(g)   (symbol equation with state symbols)

      • The colour change is from a pale pink solid to a white solid residue.

  • Sodium hydrogen carbonate is used in baking powder because on heating it thermally decomposes releasing carbon dioxide gas that gives the 'rising' action in baking.

    • The thermal decomposition of sodium hydrogencarbonate

    • sodium hydrogencarbonate ==> sodium carbonate + water + carbon dioxide

      • 2NaHCO3 ==> Na2CO3 + H2O + CO2

      • 2NaHCO3(s) ==> Na2CO3(s) + H2O(g/l) + CO2(g)   (symbol equation with state symbols)

    • This is just one of many chemical process that occur when food is cooked.

  • Note that the Group 1alkali metal carbonates like sodium carbonate Na2CO3, are so thermally stable that a bunsen flame temperature isn't high enough to decompose it. At very high temperatures sodium carbonate would break down into sodium oxide and carbon dioxide.

    Often, but not always, the more reactive a metal, the more thermally stable is the metal carbonate i.e.

    reactivity trend with water and acids: sodium > calcium > copper,

    thermal stability of carbonate: Na2CO3 > CaCO3 > CuCO3

BUT many carbonates do not fit into such a sequence, so you can't regards this as a general rule

(e) Uses of products made from limestone and ceramic materials like clay

  • What do we use limestone for? Quite a lot things actually! more than you may think!

  • Cement is produced by roasting a mixture of powdered limestone with powdered clay** in a high temperature rotary kiln at 1400oC.

    • Cement is a mixture of calcium silicate and aluminium silicate.

    • When cement is mixed with water a slow chemical reaction takes place and the cement sets to form a hard rock like material.

    • ** Clay is also used directly to make pottery and other ceramics

    • Clay is mineral formed from weathered and decomposed rock, when mixed with water it is soft and malleable and easy to mould into pottery, ceramic tiles for walls or flooring or bricks for building construction.

      • Chemically, clay consists mainly of aluminium silicates which on heating in the oven form a strongly bonded giant 3D lattice ceramic materials.

    • When these moulded clay objects are heated to a high temperature, i.e. firing in an oven they become really hard, strong and durable (if a brittle!), due to the formation of strong ionic and covalent bond networks.

    • Ceramic are non-metallic solids with high melting points and are not based on organic molecules e.g. from oil, and there are huge deposits of clay in most countries. Most ceramics are made from clay.

    • Both clay ceramics and glass are both heat and electrical insulators (very low electrical conductivity and low thermal conductivity).

      • Environmental note:

        • When limestone thermally decomposes in the production of cement carbon dioxide is produced.

        • CaCO3(s) ===> CaO(s) + CO2(g)

        •  In doing so, the production of cement contributes 8% to the World's carbon dioxide emissions

        • The cement industry is the third biggest emitter of carbon dioxide behind China and the US.

        • BUT, what is the alternative to the use of cement in the construction industry?

        • Research is going on into 'lower carbon' production methods of making cement.

        • Another possibility is using a mixture of sand and bacteria to grow bio-concrete blocks at room temperature, considerably reducing the energy needs of the production process.

          • No need for all the energy needed to heat the raw materials to 1400oC !!!

          • BUT, is it as strong? Can it withstand weathering? Can it be produced on a large enough scale to meet the demands of developed or developing countries as the World's population increases?

        • -

  • Mixing cement with sand and water makes mortar, the material you use to hold bricks together in walls.

    • You can also add calcium hydroxide to mortar to slow the setting down so the solid mortar doesn't crack as easily.

  • When cement is mixed with water, sand and crushed rock (gravel), a slow chemical reaction produces a hard, stone-like building material called concrete.

    • Concrete is a kind of composite mixture, the cement is the binder and the reinforcement is the gravel, which can also be supplemented by using steel rods too - reinforced concrete.

    • A water and gravel mixture is referred to as aggregate.

    • Concrete is a very useful cheap building material that is strong and durable against the weather, so, concrete is widely used in the construction industry from tower blocks of offices and flats to bridges, road surfaces etc.

    • Concrete is actually quite brittle, so for many concrete applications like bridges and buildings, steel rods are laid in the concrete to produce the much more durable reinforced concrete. You then get the combination of the hardness of concrete with strength of steel which does allow a minute amount of flexibility e.g. when heavy road vehicles go over a concrete bridge they will cause vibration. The concrete sections could easily fracture if it wasn't for the steel rods inside them.

  • Glass is made by heating together a mixture of limestone (calcium carbonate CaCO3), sand (mainly silica = silicon dioxide = SiO2) and 'soda' (sodium carbonate, Na2CO3).

    • limestone + sand + sodium carbonate == heat ==> soda-lime glass

    • When all three are mixed they fuse together at high temperatures and then cooled to produce the transparent glass-like solid we call glass!

    • More expensive borosilicate glass (e.g. pyrex), doesn't use limestone and is made by heating and fusing together sand and boron trioxide, and melts at higher temperatures than soda-lime glass, so very good for chemical laboratory apparatus!

      • Pyrex has better properties - its stronger than soda-glass, has a higher melting point and withstands heat changes better e.g. on heating or cooling because it has a lower thermal expansion coefficient.

    • Technically, glass is another ceramic material, but unlike clay ceramics, it is transparent and on heating readily moulded into any desired shape from bottles to glass drinking vessels. It is readily coloured by using transition metal compound pigments mixed into the hot glass. However, like fine ceramics, thin glass is brittle and easily fractures and shatters.

  • Limestone is used to remove acidic oxide impurities in the extraction of iron and in making steel.

  • Calcium oxide and calcium hydroxide also react with acids to form calcium salts.

  • You will find details of this kind of reaction on the Acids and Bases pages

  • Limestone and hard/soft water are covered on the Extra Aqueous Chemistry page.

  • Lime (calcium oxide) and slaked lime (calcium hydroxide) are both used to reduce the acidity of soil on land, they are both faster and stronger acting than limestone powder.

    • Plants grow best within a certain pH range, over acidity can affect plant growth, adding powdered limestone or lime to the soil neutralises the acid and raises the pH to improve crop growth.

    • They are also used to reduce acidity in lakes and rivers due to acid rain.

    • They are also used to neutralise potentially harmful industrial acid waste including sulphur dioxide in the flue gases of power stations. The process is called flue gas desulfurization.

      • Power stations can be fitted acid gas scrubbers eg removing the acidic sulfur dioxide with an alkaline mixture of water mixed with powdered lime/limestone.

      • An alkaline slurry (mixture of solid + powdered solid) of calcium hydroxide (calcium oxide + water) is sprayed into the flue gases from the power station furnaces.

      • In a neutralisation reaction, the sulfur dioxide reacts with the calcium hydroxide to make the neutral salt calcium sulfite - a waste product, but much of the acidic sulphur dioxide is removed, so less acid rain damage to the environment.

  • In the test for carbon dioxide, calcium hydroxide solution (limewater) forms a white milky precipitate of calcium carbonate (back to where you started!). 

    • calcium hydroxide  + carbon dioxide ==> calcium carbonate + water

      • Ca(OH)2 + CO2 ==> CaCO3 + H2O

      • Ca(OH)2(aq) + CO2(g) ==> CaCO3(s) + H2O(l)   (symbol equation with state symbols)

Limestone Chemistry and Uses (c) doc b

Summary of limestone chemistry and uses

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(f) Some other uses of calcium oxide and calcium hydroxide

Some other uses of calcium oxide (lime/quicklime)

  • Read first CaO uses above

  • As well as a major ingredient in cement manufacture (see CaO uses above) it is used in making paper pulp.

  • Much of the calcium oxide produced, apart that used in cement,  is converted to calcium hydroxide.

Some other uses of calcium hydroxide, Ca(OH)2, slaked lime

  • Read first Ca(OH)2 uses above

  • In life support systems as a carbon dioxide scrubber, particularly in closed-circuit diving re-breathers, being an alkaline substance, it absorbs the weakly acidic gas carbon dioxide, a product of respiration.

  • An ingredient in whitewash and plaster as well as mortar.

  • It has been suggested that it is added to sea water to reduce atmospheric CO2 to reduce the greenhouse effect - rather a lot needed?

  • In the production of metals, limewater is injected into the waste gas stream to neutralize acids, such as fluorides and chlorides prior to being released to atmosphere.

  • In Bordeaux mixture to neutralize the solution and form a long-lasting fungicide with copper sulfate solution.

  • In the chemical industry for manufacture of calcium stearate.

  • For preparation of dry mixes for painting and decorating.

  • Calcium hydroxide is used to clear a brine of carbonates of calcium and magnesium in the manufacture of salt for food and pharmaceutical uses.

  • To enrich or fortify (Ca supplement) fruit drinks, such as orange juice, and infant formula as a calcium supplement.

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(g) Issues with limestone quarrying and associated industries

Exploitation of natural resources issues - 'pros and cons'!

Why is limestone quarried (mined)? Why is it so useful?

  • Whenever mineral extraction takes place, whether its to obtain stone, sand a metal ore etc. there are bound to be issues of concern in the 21st century.

  • Exploiting any natural resource on a large scale will always have to be a balance (compromise) between economic need and usefulness of the resource and environmental and ecological factors.

  • In the past centuries (eg 18th and 19th centuries, less so in 20th and 21st centuries) little consideration was given to the environment and the conditions of workers, but things are different now in many '1st world countries', sadly, not so in many parts of the world eg parts of Europe, China, South America and 3rd world countries in Africa.

  • The following 'discussion' focuses on limestone, but many of the points could be made in the context of sandstone quarrying, mining such as open cast coal mines or iron ore mines.

  • Advantages - the plus points for limestone quarrying - why do we quarry limestone?

    • Limestone is very useful for house building and road building and ...

      • unlike wood doesn't burn (fire-resistant),

      • it doesn't rot by being attacked by fungi, no treatment required,

      • it isn't attacked by insects, no prevention chemicals required,

      • it doesn't corrode as fast as most metals.

    • Quarrying provides jobs for local people and adds to the local economy directly by increasing the wealth of the local population and indirectly with eg local businesses like shops and better transport links, road improvement, health provision and recreation facilities.

    • Limestone is quite widely available and readily quarried with dynamite and diggers!

    • Limestone, like sandstone, is cheaper than eg marble or granite and is quite easy to cut into blocks.

    • It isn't as hard wearing as granite or marble but it still looks attractive and will last for hundreds of years. It is readily eroded by acid rain, but this effect is decreasing as we switch away from fossil fuel power.

    • Products involving the direct use or indirect use of limestone include cement, mortar, concrete and steel.

    • They may not be attractive, but products such as concrete on initial preparation, can be poured into moulds or cast into panels for quick and efficient building construction techniques.

    • Limestone and lime are used to increase the fertility of soil by reducing acidity and there is much pressure on food production around the world.

    • Limestone and lime are used to decrease the acidity of waste gases from fossil fuel power stations by neutralising the harmful and acidic gas sulfur dioxide, which causes acid rain - which damages the environment.

    • It is possible to convert and landscape old limestone quarries into nature reserves or parks for local use, and so partly redressing the damage done by quarrying.

      • You can do a lot by landscaping and putting some effort into restoration projects which may end up as an attractive and useful local amenity that's good for local people as well as plants and wildlife.

      • Even if the deeper parts of a quarry filled with water will provide a habitat for aquatic life.

    • Limestone or lime is used in chemical processes that produce paint mixtures, dyestuffs for colouring materials and drugs/medicines.

    • So, all in all, its a pretty useful material and relatively cheap to produce.

  • Disadvantages - the downside negative points for limestone quarrying - what are the problems and issues?

    • Large quarries disfigure the landscape, big holes can be ugly, and limestone country is often part of some of the most beautiful landscapes we have.

    • The holes left can fill with water producing a habit.

    • Limestone is usually transported by road giving rise to noise and pollution from lorries and road wear damage.

    • Blasting the rock face of a limestone quarry is periodically a bit noisy!

    • There may be waste from the process that lies around in tips.

    • Wildlife habitats are destroyed for birds and animals.

    • Cement factories produce dust which is harmful to lungs, particularly people with respiratory problems.

    • Making lime from limestone uses lots of energy eg from fossil fuels.

    • It is quite brittle and cracks easily, but it can be reinforced.

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(h) More on the formulae of calcium and magnesium compounds and more reactions including the reaction of oxides, hydroxides and carbonates with acids

  • Formulae of magnesium and calcium compounds (M = metal = Mg or Ca, same group 2, same formula!)

    • IONS: The metal ion in aqueous solution or solid compounds is M2+, which combines with other ions such as: oxide O2-, hydroxide OH-, carbonate CO32-, hydrogencarbonate HCO3-, chloride Cl-, sulfate SO42-, nitrate NO3- to form the calcium or magnesium compounds.

    • COMPOUND FORMULAE: oxide MO, hydroxide M(OH)2, carbonate MCO3, hydrogencarbonate M(HCO3)2, chloride MCl2, sulfate MSO4, nitrate M(NO3)2

  • The oxides and hydroxides readily react with acids to form salts

    • general word equation: oxide or hydroxide  +  acid ==>  salt  +  water

      • examples ...

      • calcium oxide + hydrochloric acid ==> calcium chloride + water

      • magnesium hydroxide + nitric acid ==> magnesium nitrate + water

      • calcium hydroxide + sulfuric acid ==> calcium sulfate + water

    • since hydrochloric acid gives a chloride salt, nitric acid  gives a nitrate salt, sulfuric acid a sulfate salt ... the symbol equations are ... where M = Mg or Ca (or any other Group 2 metal)

      • MO + 2HCl ==> MCl2 + H2O

      • MO(s) + 2HCl(aq) ==> MCl2(aq) + H2O(l)    (state symbol equation)

      • MO + 2HNO3 ==> M(NO3)2 + H2O

      • MO(s) + 2HNO3(aq) ==> M(NO3)2(aq) + H2O(l)    (state symbol equation)

      • MO + H2SO4 ==> MSO4 + H2O

      • MO(s) + H2SO4(aq) ==> MSO4(aq/s*)  + H2O(l)    (state symbol equation)

      • if M(OH)2 involved, there is a 2H2O at the end NOT a single H2O to balance the symbol equation

      • M(OH)2 + 2HCl ==> MCl2 + 2H2O

      • M(OH)2(s) + 2HCl(aq) ==> MCl2(aq) + 2H2O(l)    (state symbol equation)

      • M(OH)2 + 2HNO3 ==> M(NO3)2 + 2H2O

      • M(OH)2(s) + 2HNO3(aq) ==> M(NO3)2(aq) + 2H2O(l)    (state symbol equation)

      • M(OH)2 + H2SO4 ==> MSO4 + 2H2O

      • M(OH)2(s) + H2SO4(aq) ==> MSO4(aq/s*)  + 2H2O(l)    (state symbol equation)

      • * the sulfates of e.g. calcium and barium are not very soluble and this slows the reaction down!

    • For more notes see 'Reactions of acids with oxides, hydroxides and carbonates'

  • Solubility of calcium compounds and reactions (and the chemically similar magnesium):

    • Magnesium and calcium oxides or hydroxides are slightly soluble in water forming alkaline solutions. They readily react and dissolve in most acids (see above).

    • Magnesium and calcium carbonate are insoluble in water but readily dissolve in most dilute acids like hydrochloric, nitric and sulfuric (see below).

  • methods of gas preparation - apparatus, chemicals and equation (c) doc bEquation examples for the reactions of calcium carbonate with acids

    • Apart from copper compounds, in every case the white carbonate solid dissolves to give a colourless solution and effervescence accompanies the dissolving as carbon dioxide gas is given off (test - gives white precipitate if bubbled into limewater). Crystallisation - on evaporation of the resulting solution, the colourless salt will crystallise out. Copper(II) carbonate is a greyish green colour and dissolves to form a blue solution and from it blue copper salts can be crystallised.

    • Here are some examples of the chemical equations describing these acid-carbonate reactions.

    • calcium carbonate + hydrochloric acid ==> calcium chloride + water + carbon dioxide

      • CaCO3 + 2HCl ==> CaCl2 + H2O + CO2

      • CaCO3(s) + 2HCl(aq) ==> CaCl2(aq) + H2O(l) + CO2(g)   (symbol equation with state symbols)

      • methods of gas preparation - apparatus, chemicals and equation (c) doc bThis reaction is a simple way to make a sample of carbon dioxide gas (see the two right-hand diagrams).

    • calcium carbonate + nitric acid ==> calcium nitrate + water + carbon dioxide

      • CaCO3 + 2HNO3 ==> Ca(NO3)2 + H2O + CO2

      • CaCO3(s) + 2HNO3(aq) ==> Ca(NO3)2(aq) + H2O(l) + CO2(g)

    • calcium carbonate + sulfuric acid ==> calcium sulfate + water + carbon dioxide

      • CaCO3 + H2SO4 ==> CaSO4 + H2O + CO2

      • CaCO3(s) + H2SO4(aq) ==> CaSO4(aq,s)  + H2O(l) + CO2(g)   (symbol equation with state symbols)

      • Calcium carbonate reacts slowly in dilute sulfuric acid because calcium sulfate is not very soluble and coats the limestone inhibiting the reaction.

    • The equations are similar for magnesium carbonate, zinc carbonate and copper carbonate ...

    • You simply replace Ca in the above equations with Mg, Zn or Cu

    • eg

    • magnesium carbonate + hydrochloric acid ==> magnesium chloride + water + carbon dioxide

      • MgCO3 + 2HCl ==> MgCl2 + H2O + CO2

      • MgCO3(s) + 2HCl(aq) ==> MgCl2(aq) + H2O(l) + CO2(g)   (symbol equation with state symbols)

    • zinc carbonate + nitric acid ==> zinc nitrate + water + carbon dioxide

      • ZnCO3 + 2HNO3 ==> Zn(NO3)2 + H2O + CO2

      • ZnCO3(s) + 2HNO3(aq) ==> Zn(NO3)2(aq) + H2O(l) + CO2(g)   (symbol equation with state symbols)

    • copper(II) carbonate + sulfuric acid ==> copper(II) sulfate + water + carbon dioxide

      • CuCO3 + H2SO4 ==> CuSO4 + H2O + CO2     (blue solution formed)

      • CuCO3(s) + H2SO4(aq) ==> CuSO4(aq)  + H2O(l) + CO2(g)   (symbol equation with state symbols)

    • The equations are slightly different for sodium carbonate because of its different carbonate formula

      • sodium carbonate + hydrochloric acid ==> sodium chloride + water + carbon dioxide

        • Na2CO3 + 2HCl ==> 2NaCl + H2O + CO2

          • Na2CO3(s) + 2HCl(aq) ==> 2NaCl(aq) + H2O(l) + CO2(g)

      • sodium carbonate + nitric acid ==> sodium nitrate + water + carbon dioxide

        • Na2CO3 + 2HNO3 ==> 2NaNO3 + H2O + CO2

          • Na2CO3(s) + 2HNO3(aq) ==> 2NaNO3(aq) + H2O(l) + CO2(g)

      • sodium carbonate + sulfuric acid ==> sodium sulfate + water + carbon dioxide

        • Na2CO3 + H2SO4 ==> Na2SO4 + H2O + CO2

          • Na2CO3(s) + H2SO4(aq) ==> Na2SO4(aq) + H2O(l) + CO2(g)

    • For more notes see 'Reactions of acids with oxides, hydroxides and carbonates'

  • Magnesium and calcium hydrogencarbonate are soluble in water and cause 'hardness' i.e. scum with 'traditional' non-detergent soaps. Formulae are Mg(HCO3)2 and Ca(HCO3)2

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(i) Thermal decomposition of hydroxides and nitrates (not needed by some syllabuses-specifications)
  • Decomposition of carbonates: see above.

  • Decomposition of metal hydroxides:

    • The Group 1 Alkali Metal hydroxides do not readily decompose on heating even 'up to red heat'.

      • Except for lithium hydroxide which forms lithium oxide and water.

        • lithium hydroxide ==> lithium oxide + water

        • 2LiOH ==> Li2O + H2O

        • 2LiOH(s) ==> Li2O(s) + H2O(l)    (symbol equation with state symbols)

      • These hydroxides are white solids and soluble in water to give an alkaline solution.

    • On heating, the Group II, Lead, Aluminium and Transition Metal hydroxides decompose to form the metal oxide and water vapour.

      • The original hydroxides are usually relatively insoluble solids, white in colour, except copper(II) hydroxide is blue and iron(III) hydroxide is brown.

    • M(OH)2(s) ==> MO(s) + H2O(l)     (symbol equation with state symbols)

      • M = Mg, Ca, Zn giving white oxide MO (ZnO yellow when hot),

      • M = Cu gives black copper(II) oxide CuO, M = Pb gives yellow lead(II) oxide PbO

      • e.g. if M = Zn:

        • zinc hydroxide ==> zinc oxide + water

        • Zn(OH)2 ==> ZnO + H2O

        • Zn(OH)2(s) ==> ZnO(s) + H2O(l)   (symbol equation with state symbols)

        • or

        • Ca(OH)2 ==> CaO + H2O

        • Ca(OH)2(s) ==> CaO(s) + H2O(l) to give calcium oxide (quicklime)

    • and

      • 2M(OH)3 ==> M2O3 + 3H2O

      • 2M(OH)3(s) ==> M2O3(s) + 3H2O(l)     (symbol equation with state symbols)

      • where M = Al to give white aluminium oxide, and M = Fe to give reddish-brown iron(III) oxide.

      • e.g.

      • aluminium hydroxide ==> aluminium oxide + water

      • 2Al(OH)3 ==> Al2O3 + 3H2O

      • 2Al(OH)3(s) ==> Al2O3(s) + 3H2O(l)     (symbol equation with state symbols)

  • Decomposition of nitrate salts:

    • The Group 1 Alkali Metal nitrates (NO3) decompose to form the nitrite (NO2) salt and oxygen gas.

      • 2MNO3 ==> 2MNO2 + O2    (where M = Na or K)

      • 2MNO3(s) ==> 2MNO2(s) + O2(g)   (symbol equation with state symbols)

      • so when M = Na/K:

      • sodium/potassium nitrate ==> sodium/potassium nitrite + oxygen

      • Nitrates are colourless crystals and nitrites are white solids and are all soluble in water giving neutral solutions.

    • Many metal nitrates decompose to form the metal oxide, nasty brown nitrogen dioxide gas (NO2) and oxygen gas (O2) when strongly heated. Note: Nitrogen dioxide is more correctly and systematically names as nitrogen(IV) oxide.

      • These nitrates are all water soluble neutral salts, all colourless crystals except Cu is blue and Fe is pale brown-dark orange.

      • 2M(NO3)2 ==> 2MO + 4NO2 + O2

      • 2M(NO3)2(s) ==> 2MO(s) + 4NO2(g) + O2(g)     (symbol equation with state symbols)

        • where M = Mg, Ca, Zn giving the white oxide MO (ZnO yellow when hot),

        • when M = Cu, it gives the black copper(II) oxide CuO

        • if M = Pb it gives the yellow lead(II) oxide PbO.

        • e.g.

        • lead(II) nitrate ==> lead(II) oxide + nitrogen dioxide + oxygen

          • 2Pg(NO3)2 ==> 2PbO + 4NO2 + O2

          • 2Pg(NO3)2(s) ==> 2PbO(s) + 4NO2(g) + O2(g)     (symbol equation with state symbols)

      • 4M(NO3)3(s) ==> 2M2O3(s) + 12NO2(g) + 3O2(g)     (symbol equation with state symbols)

        • If M = Al, it gives white aluminium oxide, if M = Fe it gives give reddish-brown iron(III) oxide

        • e.g.

        • aluminium nitrate ==> aluminium oxide + nitrogen dioxide + oxygen

        • 4Al(NO3)3 ==> 2Al2O3 + 12NO2 + 3O2

        • 4Al(NO3)3(s) ==> 2Al2O3(s) + 12NO2(g) + 3O2(g) 

        • iron(III) nitrate ==> iron(III) oxide + nitrogen dioxide + oxygen

        • 4Fe(NO3)3 ==> 2Fe2O3 + 12NO2 + 3O2

        • 4Fe(NO3)3(s) ==> 2Fe2O3(s) + 12NO2(g) + 3O2(g)

    • See also (c) doc b gas preparation and collection page for methods.

    • See also s-block Group I/II detailed notes for advanced level chemistry students

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