(b) Limestone - a very useful material
and part of the Carbon Cycle
-
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?
-
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
CaO + CO2
-
CaCO3(s)
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!) ....
-
Note
on heating other carbonates you
get a similar
thermal decomposition.
TOP OF PAGE and
sub-index
(c)
The Limestone Cycle
|
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
(limestone) |
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
(limewater) |
LIMESTONE CYCLE |
calcium oxide
(quicklime) |
 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) |
TOP OF PAGE and
sub-index
(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
-
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
-
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
-
The thermal
decomposition of lead carbonate
-
lead(II) carbonate(s,
white)
==> lead(II) oxide(s, yellow-orange)
+ carbon dioxide
-
The thermal
decomposition of manganese carbonate
-
manganese(II) carbonate(s,
pale pink)
==> manganese(II) oxide(s, white)
+ carbon dioxide
-
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
-
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.
-
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.
-
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!
-
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!).
Summary of limestone chemistry and uses
TOP OF PAGE and
sub-index
(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.
TOP OF PAGE and
sub-index
(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.
TOP OF PAGE and
sub-index
(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).
-
 Equation 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)
-
 This
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
-
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
-
zinc carbonate + nitric acid ==>
zinc nitrate + water + carbon dioxide
-
copper(II) carbonate + sulfuric acid ==>
copper(II) sulfate + water + carbon dioxide
-
The equations are
slightly different for sodium carbonate because of its different
carbonate formula
-
sodium carbonate +
hydrochloric acid ==> sodium chloride + water + carbon dioxide
-
sodium carbonate +
nitric acid ==> sodium nitrate + water + carbon dioxide
-
sodium carbonate +
sulfuric acid ==> sodium sulfate + water + carbon dioxide
-
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
|
|
TOP OF PAGE and
sub-index
(i) Thermal decomposition
of hydroxides and nitrates (not
needed by some syllabuses-specifications)
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See also s-block Group I/II detailed
notes on Alkaline Earth Metals for A level chemistry students
Index of
selected pages describing industrial processes:
Limestone, lime
- uses, thermal decomposition of carbonates, hydroxides and nitrates
Enzymes and
Biotechnology
Contact Process, the importance of sulfuric acid
How can
metals be made more useful? (alloys of Al, Fe, steel etc.)
Instrumental Methods of Chemical Analysis
Chemical & Pharmaceutical Industry Economics & Sustainability
- Life Cycle Assessment
Products of the
Chemical & Pharmaceutical Industries & impact on us
The Principles & Practice of Chemical
Production - Synthesising Molecules
Ammonia
synthesis/uses/fertilisers
Oil Products
Extraction of Metals
Halogens
- sodium
chloride electrolysis
Transition
Metals
Extra Electrochemistry
- electrolysis and cells
Quizzes and other web pages
of industrial chemistry
GCSE/IGCSE Multiple choice QUIZ
on Limestone and its uses:
easier
GCSE/IGCSE Foundation
on limestone chemistry
or harder
GCSE/IGCSE higher
on limestone chemistry
and a
GCSE/IGCSE multi-word fill exercise
on limestone chemistry and uses
2nd Word-fill quiz
"Limestone chemistry"
Word-fill quiz
"More on carbonate chemistry"
See also s-block Group I/II detailed
notes on Alkaline Earth Metals for A level chemistry students
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