Chemistry KS4 science GCSE/IGCSE/O Level Chemistry Revision Notes
Water Cycle - Water as
a resource, hard & soft water, colloids, emulsions, salt solubility
& water of crystallisation
1. The components of the
water cycle are described and explained and the use of water as an important
resource, its treatment to make it safe to drink, pollution problems.
Colloids (e.g. sol, foam, emulsion) are described with examples. The
difference between hard water and soft water is explained and the causes and
treatment of hard water * 2. Gas and salt solubility and solubility curves
are considered and finally 3. The explanation and calculation of water of
cycle, treatment, pollution,
colloids (sol, foam, emulsion) hard/soft water
2. Gas and salt solubility
in water and solubility curves
3. Calculation of water of crystallisation
What happens to water on the Earth's Surface?
The water on the Earth's surface is continually
being re-cycled. As it falls, rain water contains only
dissolved gases but once it reaches the ground water becomes contaminated in
Cycle and Resources *
Treatment-pollution-colloids * 1c Hard & Soft Water
1a The Water
Cycle and Water as a Resource
- Water is the most abundant substance on the
surface of our planet and is essential for all life. Water in rivers,
lakes and the oceans is evaporated by the heat of the Sun (endothermic). The water
vapour formed rises into the atmosphere, cools and forms clouds of
condensation (exothermic). Eventually this gives rain and snow
'precipitation' which on melting returns to the rivers, seas and oceans. This is known as the
- Water is an important raw material and has
many uses. It is used as a solvent and as a coolant both in the home and
in industry. It is used in many important industrial processes including the manufacture of sulphuric acid.
- Seawater/brine is a valuable resource e.g. large scale evaporation in 'salt pans' (using fuel burning or solar
energy) to produce 'sea salt' sodium chloride NaCl, the water also
contains lots of other salts including bromides from which the element
bromine is extracted.
Treatment and pollution - domestic and industrial contexts
- There are various undesirable materials that need
to be removed from water before it is fit for domestic consumption. They include
chemicals which cause tastes or
odours and Acidic substances:
- Drinking water is made fit for
domestic home consumption by
- (i) allowing sedimentation to occur, where
larger insoluble particles settle out,
- (ii) passing it through sand filter beds to remove
finer solid particles,
- (iii) treating
with chlorine to kill bacteria,
- (iv) adding small amounts of
sulphur dioxide to remove excess toxic chlorine
- the molecular equation is SO2(aq)
+ Cl2(aq) + 2H2O(l) ==> 2HCl(aq)
- the ionic equation is SO2(aq)
+ Cl2(aq) + 2H2O(l) ==> 2Cl-(aq)
+ SO42-(aq) + 4H+(aq)
- (v) aluminium sulphate is added to
coagulate colloidal clay (see colloids
- (vi) carbon slurry absorbs
molecules causing 'tastes' and 'odours'.
- (vii) adding lime slurry to
neutralise the water if it is too acid.
- The use of artificial fertilisers results
in many natural waters being contaminated with dissolved nitrate and
ammonium ions. Dissolved nitrate ions can have harmful effects on babies
and so the levels of nitrate are carefully monitored. Nitrates may be
carcinogenic. The ions from this pollution are not easy to remove on a
large cost-efficient scale.
- An ion-exchange filter can remove
these and other ions which can cause problems e.g. calcium and magnesium
which cause hardness in water and iron compounds (see
- Iron in water
is a non-harmful but an aesthetic nuisance impurity:
- readily soluble iron(II) when exposed
to air form rusty brown insoluble iron(III) hydroxide or hydrated
iron(III) oxide compounds. These stain yellow/orange/brown washing
laundry and white plumbing facilities!
- The iron(III) ions also form inky black
compounds with the tannic acids in tea and giving it a 'metallic'
- Cooked vegetables turn brown (complex
compounds with phenols).
- Colloidal clay: A
consists of one substance (or mixture of substances) very finely dispersed
another substance (or a mixture of substances) without a new true
solution forming. So a
colloid is a mixture of a dispersed phase
and a continuous phase (disperse
medium) BUT the dispersed phase is NOT dissolved in the continuous
- A colloid is NOT a solution, although the
colloid particles are not usually seen under a microscope, they are much
bigger than molecules, and much bigger than the molecules of the
continuous phase (disperse medium e.g. water).
- In a solution the solvent or solute particles are
usually of comparable size and completely mixed at the 'individual
particle level' i.e. completely homogeneous in the same phase.
- A colloid can be thought of as
intermediate between a true solution and a mixture of e.g. a liquid and
an insoluble solid. No filtration separation is possible with solutions
and filtration is easy and effective with an insoluble solid.
Similarly, most colloid particles
are too small to be filtered, but separation from truly dissolved
substances is possible with a membrane.
- The colloidal particles of the disperse
phase are equivalent to the solute of a solution and the continuous
phase is equivalent to the solvent. The mixture is sometimes referred
to as the 'colloidal solution'. These descriptors can be somewhat
'blurred' by the intermediate particulate nature of colloidal systems!
- The particles in a colloid are so small
that they remain 'suspended' (the mixture is called a 'suspension') in the disperse medium (e.g. colloidal clay
particles in water) with little tendency to settle out. However the
colloidal particles are big enough for their surface area properties
to be significant (see electrical
- Examples of colloids
that is the fine dispersion of one substance in another without a new
- A sol is a solid dispersed in a
liquid e.g. tiny particles of clay in water.
- A foam is a gas dispersed in
a liquid e.g. a well shaken soap solution or shaving cream foam.
- An emulsion is a liquid
dispersed or suspended in another liquid ...
- and is a mixture of two immiscible
liquids like oil and water.
- Emulsions are thicker than either
liquid eg the emulsion 'French dressing', is thicker than olive oil
- With time, the two layers settle out, so
the less dense oil floats on top of the aqueous/water layer.
- One way to inhibit the two layers
settling out is to use an emulsifier.
- An emulsifying agent stabilises an
- Two of the most commonly used
emulsifiers are lecithin (E322) and the mono- and
di-glycerides of fatty acids (E471), and are classified as food
additives in the E number system.
- Egg yolk acts as an emulsifying agent
(it contains lecithin).
- examples of emulsions.
- (i) milk (aqueous solution + insoluble,
but dispersed fats)
- (ii) French dressing in salads (based on vinegar +
olive oil, but these do reform the oil and aqueous layers quite
easily which is why they are shaken before use)
- (iii) Mayonnaise is a mixture of
oil, water, emulsifier and other ingredients.
contain emulsifiers to stop the salty water from separating out and mayonnaise also contains an emulsifier to stop the oil and aqueous
based components separating out.
- Emulsifier molecules have
a 'water loving'/'oil hating' (hydrophilic) part and a 'water
hating'/'oil loving' part (hydrophobic). Therefore they can interact
with the different components and keep the different types of molecules
dispersed in each other.
- Diagram A: This diagram represents a true solution where the black dots represent the dissolved individual
molecules - they do NOT clump together.
- Diagram B: This diagram represents an
emulsion of oil droplets dispersed in water (oil in water emulsion).
- Each oil droplet will have
millions of oil molecules in it.
- The oil is the disperse phase and the
water is the continuous phase.
- This is NOT a true solution.
- Semi-skimmed or full fat milk is like this, droplets
of fat (~1-3% oil) are dispersed in water.
- Single cream (~18% oil), double cream
(~50% oil) are oil in
- Whipped cream and ice cream are oil in
- Air is whipped or whisked into cream to
give it a soft frothy texture to use as a topping.
- Whipping air into ice cream gives it a
softer texture so you can scoop out portions easily.
- Mayonnaise is an emulsion of sunflower
oil or olive oil with vinegar, and these mixtures are used in salad
dressings and sauces. A salad dressing coats the salad materials better
than either the olive oil or vinegar.
- Some non-food examples of oil in water
emulsions include moisturising creams and other cosmetic
- Diagram C: This diagram represents an
emulsion of water droplets dispersed in an oil (water in oil emulsion).
- Each water droplet will
have millions of water molecules in it.
- The water is the disperse phase and the
oil is the continuous phase.
- This is NOT a true solution.
- Melted margarine is a water in oil
- One of the problems with useful
emulsions is that the two main components, the two immiscible liquids,
tend to separate out rendering the emulsion useless for its designed
- The way round this is to use an
emulsifying agent (emulsifier) which inhibits the separation of the
emulsion back into two layers.
- D: This diagram represents the
effect of mixing an oil in water emulsion with an emulsifying agent like
soap (edible substances like lecithin are used in processed food!).
- This diagram illustrates the mechanism
by which soaps wash oily/greasy clothes or surfaces.
- The washing process is described and
- Diagrams E1, E2 and S3: Emulsifying molecules like
soap have a negative ionic hydrophilic 'head' ('water liking'/'oil
hating' end of molecule) and a
hydrophobic 'tail' ('water hating'/'oil liking' end of molecule').
- eg the stearate ion from the soap sodium
stearate shown above.
- When you shake soap with an oily/greasy
material (washing clothes or scrubbing a surface), the oil/grease breaks
up into tiny droplets or globules. Why? ...
- The hydrocarbon hydrophobic tail of the
soap dissolves in the oil or grease globule and the negative head is on
the surface of the globules/droplets.
- The hydrophobic tail can only
interact with oil/grease ie is attracted to oil and grease.
- The hydrophilic head can only
interact with water ie is attracted to water.
- Two hydrophilic heads cannot interact
with each other and tend to repel each other especially if the
hydrophilic head carries a negative charge.
- In effect, the globules of oil/fat get a
surface coating of the emulsifier - a general name for these emulsifying
molecules is surfactants and includes soaps, detergents and naturally
occurring molecules like lecithin found in egg yolk..
- Because of the negative head of the
soap ion on the oil/grease droplet surface, you get repulsion
between the globules and an emulsion is formed.
- The ionic end also strongly interacts
with water, again this prevents the oil/grease particles flocking
- So, the oil and grease particles cannot
re-clump together to form a separate layer on the clothes or surface
being cleaned, and so the emulsion is stabilised.
- Therefore the oil/grease remains
dispersed in the soapy washing water and hence washed away.
- In other contexts eg food, you use a
soap like molecule, but harmless and edible!, to do exactly the same
effect, that is, emulsifying the mixture to make a stable emulsion which
doesn't separate into two layers.
- In the food industry emulsifiers are
very important for stopping recipe components separating out from
emulsions and give processed foods greater stability and longer
shelf-life and helps to produce less fatty food and still retain
acceptable texture for the consumer. There can be some diet
restrictions for some people eg if you are allergic to eggs then any
processed food using egg yolk as an emulsifying agent is a no go
area! As with any processed food, if you have a sensitive
constitution, you must carefully check the ingredients.
- Incidentally, the emulsifier molecule
does not have to be an ionic compound like soap.
- It can be a non-ionic neutral molecule
like lecithin BUT the molecule must have a hydrophilic head that bonds with water and a hydrophobic
tail that bonds with oil/grease.
- The bonds formed are intermolecular
bonds (from intermolecular forces of attraction) and NOT chemical bonds
like ionic or covalent bonds.
- Detergents are also emulsifiers,
and not just used for washing in the home, they are also used to help
disperse oil spilt from tankers into rivers, seas and oceans. Much of
the oil spill can be contained by booms and pumped off the surface of
the water - but not all unfortunately. Dispersed oil droplets break down
(biodegrade) more quickly than large patches of oil, but the process is
very slow. Rescued seabirds coated in oil can be washed with detergent
to clean them BUT their own natural protective oils are also washed away
so their lives are still in danger and the birds need care and
- See also ...
- Colloidal particles may be
charged. (Note: So far the discussion has
been confined to hydrophobic ('water hating') colloids which do NOT interact strongly with
the continuous phase.
- In contrast 'gels' for example, are hydrophilic
('water liking') colloids, in which the colloid particles are very solvated* and
stabilised by the continuous phase).
- Solvated means the particle is weakly attracted to layers
of surrounding 'solvent' molecules of the dispersal medium e.g. water.
- Colloidal particles of a sol absorb ions, but not in
electrically balanced proportions. Depending on which ion(s) are preferentially
absorbed from the water,
the net charge on the colloid particle can be positive or
negative. The situation is complicated further because the
charged colloid particles attract a sheath of oppositely charged ions
around them. This is called the electrical double layer effect.
This means neighbouring colloid particles have the same 'outer charge'
and so are repelled, rather than attracted together. The sol
itself is overall electrically neutral like any other
- Colloids are destroyed when the
particles of the disperse phase join together and separate out from the
continuous phase. This process is called coagulation. For sols,
any disturbance of the double layer can cause coagulation to happen.
It can be caused by boiling the sol, the increased random
thermal collisions disturb the electrical balance and allows the
colloid particles to collect together.
- Sols are also very sensitive to the
presence of ions, so any electrolyte ions present can affect the
electrical double layer (the theory is complex but just think of the
ions charge as affecting the stability of the double layer). The more
highly charged the ion, the greater the electrical field
force effect, so the greater its coagulating
power. The ions reduce the repulsion between the colloid
particles and allow coagulation to occur.
- Examples of coagulating
- positive cations: Al3+ > Mg2+ > Na+
- negative anions: [Fe(CN)6]3- > SO42-
- and this explains why aluminium
sulphate Al2(SO4)3 is used to
precipitate (coagulate) colloidal clay in water treatment for domestic
- Other onsite references to water
1c Hard and
- HARD and SOFT WATER: Many compounds dissolve in
water without chemical change but may have a variety of consequences!
- Water which readily gives a lather
(not detergents) is described as soft water.
- Note: Detergents usually give a
good lather with any water.
- Some of
these dissolved substances make the water hard.
- This means the water does
not readily give a good lather with soap and so wastes soap as
well as causing a 'scum'! though it does
not affect soapless detergents.
- The 'scum' is due to the formation of
insoluble calcium and magnesium salts formed by a reaction between the soap
molecules and calcium and magnesium ions.
- Most hardness is due to water containing dissolved calcium or
- The hard water is formed when natural waters flow
over ground or rocks containing calcium or magnesium compounds.
- e.g. Chalk and limestone, mainly calcium
carbonate CaCO3 with some magnesium carbonate too.
- or Gypsum rock deposits, which are
mainly calcium sulphate CaSO4,
- and magnesium sulphate which was
called 'Epsom Salts', formula MgSO4.7H2O, because it crystallised out of evaporated spring
water from Epsom on the chalk downs of southern England.
- Calcium sulphate (slightly soluble) and magnesium sulphate
(very soluble) are washed
out of rock formations.
- Insoluble calcium carbonate (in limestone, chalk) and
insoluble magnesium carbonate both dissolve
in acid rainwater to form soluble hydrogencarbonates
- e.g. naturally carbonated water
(dissolved carbon dioxide makes water acidic so it reacts with the
- insoluble calcium carbonate + water + carbon
dioxide ==> soluble calcium hydrogencarbonate
- CaCO3(s) + H2O(l)
+ CO2(g) ==> Ca(HCO3)2(aq)
- The simplest test
for 'hardness' is to shake the water with an old fashioned 'soapy'* soap.
The term 'soapy soap' is NOT a joke! e.g.
the blocks of 'household' soap based on sodium stearate, sodium palmitate
(from palm oil) or sodium oleate (from olive oil).
- Soft water readily forms a lather with
soap but hard water does not.
- Hard water forms a scum from the
dissolved calcium or magnesium compounds.
- The scum is a precipitate formed from insoluble calcium and magnesium
soap salts, instead of a nice frothy lather (see below).
- Eventually with enough soap,
a lather does form, when all the calcium and magnesium ions have been
precipitated as a 'scum salt'! However, it does mean a lot of
soap is wasted!
- The amount of hardness in water sample
can be estimated by titrating it with soap solution and noting what
volume of soap solution is needed to produce a lather.
- A modern detergent is sometimes called
a 'soapless soap', at least when I was a
student!, or soapless detergent. Its advantage is that no insoluble salt 'scum' is formed,,
because the Ca and Mg salts of it are soluble.
So modern detergents e.g. like 'washing up liquids' give a lather with any
water which is more acceptable for dish washing.
- The chemistry of 'scum' formation.
contains dissolved compounds that react with soap to form scum. e.g. with
soaps made from the sodium salts of fatty acids, insoluble calcium or
of the soap are formed ... 'example of a precipitation
- calcium sulfate + sodium stearate
(a soap) ==>calcium stearate (scum ppt.) + sodium sulfate
==> (C17H35COO)2Ca(s for scum!) + Na2SO4(aq)
- or more simply ionically:
- calcium ion + stearate ion ===> calcium
+ 2C17H35COO-(aq) ==> (C17H35COO-)2Ca2+(s)
precipitation reaction is generally defined as 'the formation of an
insoluble solid on mixing two solutions or a gas bubbled into a
- Below are some diagrams of the
organic molecules or ions involved
- Diagram S1: The stearic acid molecule C17H35COOH
- Diagram S2: The salt sodium stearate C17H35COO-Na+,
formed when stearic acid is neutralised with sodium hydroxide
- Diagrams S3 and E2: The negative stearate anion C17H35COO-,
its structure is important in understanding how it forms the calcium
salt precipitate, calcium stearate AND explaining
how emulsifiers work.
- Using hard
water can increase costs
because more soap is needed to make a useful
'washing lather' and hard water often
leads to deposits (lime scale) forming in heating systems and kettles which
require cleaning at times.
- The 'lime scale'
is usually caused by the thermal decomposition of the dissolved hydrogencarbonates producing insoluble calcium carbonate (so it does
remove some of the temporary hardness before washing! and chemically, it is the opposite of the carbonated water dissolving
action above) ...
- Ca(HCO3)2(aq) ==>
CaCO3(s) + H2O(l)
- However there
is a plus side to the deposition!
The coating on the inner
surface of the pipe work prevents corrosion
and the dissolving of potentially poisonous
salts of copper or lead into the water supply.
- The lime scale can be removed by
(hydrogen ion solution) treatment which dissolves the calcium carbonate.
- ionically this is: CaCO3(aq)
+ 2H+(aq) ==> Ca2+(aq)
+ H2O(l) + CO2(aq)
- e.g. vinegar contains
the weak organic acid ethanoic acid and will dissolve lime
scale in kettles but shouldn't react with the steel container or
- calcium carbonate + ethanoic
acid ==> calcium ethanoate + water + carbon dioxide
+ 2CH3COOH(aq) ==> Ca2+(CH3COO-)2(aq)
+ H2O(l) + CO2(aq)
- In the school lab. you will
doubt at some point you add the 'strong' hydrochloric acid to marble
chips, which is essentially a very similar reaction to the one
dissolving limescale above. The reaction is faster if the
vinegar is hot because all reactions are speeded by higher
temperatures because of the increased kinetic energy of the reactant
particles (see rates of
reaction for more details) and maybe also because calcium
ethanoate is not that soluble in cold water and dissolves more in
hot water (not sure of the importance of this 2nd factor?).
- calcium carbonate +
hydrochloric acid ==> calcium chloride + water + carbon dioxide
+ 2HCl(aq) ==> CaCl2(aq)
+ H2O(l) + CO2(aq)
- or showing the ions involved
+ 2H+Cl-(aq) ==> Ca2+(Cl-)2(aq)
+ H2O(l) + CO2(aq)
- more simply and the more correct ionic
+ 2H+(aq) ==> Ca2+(aq)
+ H2O(l) + CO2(aq)
- Hard water can be made soft by removing the
dissolved calcium and magnesium ions.
- If due to calcium/magnesium hydrogencarbonates it is
removed by boiling (see above).
- Adding enough 'soapy' soap, see above,
but the water is best treated before the washing!, so its not the
desired solution with the scum and all that!
- The addition of sodium carbonate
(as 'washing soda' crystals),
which dissolves and precipitates out the calcium or magnesium
ions as their insoluble carbonates(s) formed.
- calcium sulphate + sodium carbonate
==> calcium carbonate + sodium sulphate
==> CaCO3(s) + Na2SO4(aq)
- or more simply ionically: Ca2+(aq)
==> Ca2+CO32-(s) (called
an 'ionic equation')
- Packs of ion
exchange resins can hold or release ions in an ion exchange process.
- Negative polymer resin columns hold hydrogen ions or sodium ions. These
can be replaced by calcium and magnesium ions when hard water passes down the column.
The calcium or magnesium ions are held on the negatively charged
resin. The freed hydrogen or sodium ions do not form a scum with soap.
- e.g. 2[resin]-H+(s)
+ Ca2+(aq) ==> [resin]-Ca2+[resin]-(s)
- or 2[resin]-Na+(s)
+ Mg2+(aq) ==> [resin]-Mg2+[resin]-(s)
+ 2Na+(aq) etc.
- Extra Note on water
purification: You can also use an ion-exchange resin to replace
negative ions by using a positively charged resin initially holding
hydroxide ions e.g. to remove chloride (Cl-), nitrate (NO3-
potentially harmful) and sulphate ions (SO42-)e.g.
+ Cl-(aq) ==> [resin]+Cl-(s)
+ NO3-(aq) ==> [resin]+NO3-(s)
+ SO42-(aq) ==> [resin]+SO42-[resin]+(s)
+ 2OH-(aq) etc.
- Now, by using both a positive
and negatively charged resin, you can completely de-ionise water
because the released hydrogen ions and hydroxide ions combine to form pure
+ OH-(aq) ==> H2O(l)
- However, it will not remove
non-ionic substances like organic pesticides etc.
- Permanently hard water
hardness cannot be removed by boiling e.g. when caused by dissolved
magnesium or calcium sulphate.
- Temporary hard water means it is
softened by boiling e.g. when caused by magnesium hydrogencarbonate or calcium
- HOWEVER, a plus
point! Hard water contains dissolved compounds
that are good for health. Hard water often provides calcium compounds that
help the development of strong bones and teeth and help to reduce heart
illnesses and also traces of other essential elements like iron and iodine.
How well do different
gases and solids dissolve in water?
First, some definitions
of words you may encounter in talking about solubility and
other water related situations:
material which is to be dissolved in a solvent.
liquid which dissolves the material (the solute). You will come
across water more than any other liquid solvent BUT lots of
important organic solvents like hexane (petrol like), ethanol
(alcohol) and propanone (acetone) are in common laboratory use.
result of dissolving something in a liquid (solute + solvent =>
what extent a solute material will dissolve.
material will dissolve in a particular liquid solvent.
means that no more of a substance (the solute) will dissolve in its
solution i.e. maximum solubility achieved at a particular
soluble, will not dissolve in a particular liquid (don't assume it
means will not dissolve in anything).
means the addition of water to a material.
means to remove water from a substance.
rates of dissolving.
the mixture to raise the temperature will increase the rate of a
substance dissolving - the energy of all the particles involved is
increased - increased rate of more energetic collisions between solute
and solvent particles speeding up the dissolving process.
if a solid is broken up and crushed into smaller pieces or a powder it
will dissolve faster. This breaking down of a solid increases the
surface are for the solvent to 'attack' and dissolve the solid.
this increases the rate of dissolving because it prevents 'local'
saturation of the solution which will inhibit dissolving.
volume of solvent:
adding more solvent increases the speed of dissolving, the less
These factors are
similar with those affecting the rates of chemical reactions except
there is no catalyst that speed up dissolving as far as I know?
Also, increasing the volume of the solvent will decrease the rate of
reaction because concentrations are reduced.
Some gases and solid substances are more
soluble in water than others and some are hardly
soluble at all.
The solubility of gases and solids in water also depends on the
temperature of the water:
Many gases are soluble in water and the
solubility increases as the temperature decreases and as the pressure
is produced by dissolving
carbon dioxide under high pressure. When the pressure is released the gas
bubbles out of the solution. Carbonated water is used to give fizzy drinks
a 'tang' to the taste.
Dissolved oxygen is essential for aquatic
life and the colder the water, the more of it dissolves. Hot water from power stations may be discharged into rivers or
lakes. This discharge reduces the amount of oxygen dissolved in the water
and this can damage aquatic life and disrupt the natural eco-systems.
Chlorine water is made by dissolving Chlorine
gas in water and can be a useful chemical reagent, both in the laboratory
and industry (e.g. displaces iodine from sea water).
Chlorine water is used to bleach materials
and kill bacteria.
Many ionic compounds are soluble in water
and many covalent compounds are insoluble in water (but don't make
solubility of a solute
in water, or any other solvent, is usually given in grams of solute per
100 grams of solvent (e.g. water) at that temperature.
The solubility of most solid solutes increases as
the temperature increases (opposite of gases, but the ambient air
pressure has no effect).
A saturated solution is one in which no
more solute will dissolve at that temperature giving the maximum
solubility at that particular temperature.
When a hot saturated
solution cools some of the solute will separate from the solution
crystals form because the solubility is lower at the lower temperature.
From solubility graphs-data you can calculate how much will dissolve at a
given temperature and how much will crystallise out on cooling.
which describe the solubility of common types of compounds in water:
All common sodium, potassium and
ammonium salts are soluble e.g. NaCl, K2SO4, NH4NO3
salts are soluble
Mg(NO3)2, Al(NO3)3, NH4NO3
salts are soluble
salts are soluble except
those of silver and lead e.g.
Common sulfates are soluble except
those of lead, barium and calcium: soluble e.g.
insoluble: PbSO4, BaSO4, CaSO4
is slightly soluble.
hydroxides and carbonates are usually insoluble (e.g. Group 2 and Transition Metals)
except those of the Group
1 Alkali Metals sodium, potassium etc. and ammonium:
KOH, NaOH, NH4OH actually NH3(aq), Na2CO3,
MgO, CuO, ZnO, Mg(OH)2, Fe(OH)2,
Cu(OH)2, CuCO3, ZnCO3, CaCO3
Knowledge of salt solubility is
important in deciding which method of salt preparation is employed.
salt preparation methods and details of
2b. Solubility curves for selected
Interpretation of graph eg
Reading graph: at 38oC the
solubility of copper sulphate, CuSO4, is 28g of anhydrous
salt per 100g of water.
Reading graph: at 84oC the
solubility of potassium sulphate, K2SO4, is 22g
per 100g of water.
Ex Q1: How much potassium
nitrate will dissolve in 20g of water at 34oC?
At 34oC the
solubility is 52g per 100g of water,
so scaling down, 52 x
20 / 100 = 10.4g will dissolve in 20g of water
Ex Q2: At 25oC
6.9g of copper sulphate dissolved in 30g of water, what is its
solubility in g/100cm3 of water?
Ex Q3: 200 cm3
of saturated copper solution was prepared at a temperature of 90oC.
What mass of copper sulphate crystals form if the solution was cooled to
Solubility of copper
sulphate at 90oC is 67g/100g water, and 21g/100g water at 20oC.
Therefore for mass
of crystals formed = 67 - 21 = 46g (for 100 cm3 of solution).
However, 200 cm3 of
solution was prepared,
so total mass
of copper sulphate crystallised = 2 x 46 = 92g
Note: The density of
water is close to 1.0g/cm3 or ml, so for approximate purposes.
the volume in cm3 or ml of just the water is numerically close to
the value in g, i.e. 100 cm3 of water or solution is about 100g
SALT SOLUBILITY DATA
g salt / 100g water
* multiply by 1.562 for hydrated crystals CuSO4.5H2O
Water of crystallisation calculations
Solubility graphs and data
are covered in section 2.
How to calculate the
theoretical % of water in a hydrated salt
eg magnesium sulphate MgSO4.7H2O
Relative atomic masses:
Mg = 24, S = 32, O = 16 and H = 1
relative formula mass =
24 + 32 + (4 x 16) + [7 x (1 + 1 + 16)] = 246
7 x 18 = 126 is the mass
so % water = 126
x 100 / 246 = 51.2%
calculation of salt formula containing 'water of
when crystallised from aqueous solution, incorporate water molecules
into the structure. This is known as 'water of crystallisation', and the
'hydrated' form of the compound.
e.g. magnesium sulphate MgSO4.7H2O.
The formula can be determined by a simple experiment (see the copper
sulphate example below).
A known mass of the hydrated salt is gently
heated in a crucible until no further water is driven off and the weight
remains constant despite further heating. The mass of the anhydrous salt left
The original mass of hydrated salt and the mass of the anhydrous salt
residue can be worked out from the various weighings.
The % water of
crystallisation and the formula of the salt are calculated as follows:
Suppose 6.25g of blue
hydrated copper(II) sulphate, CuSO4.xH2O, (x
gently heated in a crucible until the mass remaining was 4.00g. This
is the white anhydrous copper(II) sulphate.
The mass of anhydrous
salt = 4.00g, mass of water (of crystallisation) driven off =
6.25-4.00 = 2.25g
The % water of
crystallisation in the crystals is 2.25 x 100 / 6.25 = 36%
Cu=64, S=32, O=16, H=1 ]
The mass ratio of CuSO4
: H2O is 4.00 : 2.25
To convert from mass
ratio to mole ratio, you divide by the molecular mass of each
CuSO4 = 64
+ 32 + (4x18) = 160 and H2O = 1+1+16 = 18
The mole ratio of CuSO4
: H2O is 4.00/160 : 2.25/18
which is 0.025 : 0.125
or 1 : 5, so the formula of the hydrated salt is CuSO4.5H2O
All concentration calculations are covered on the on-line
calculations page, especially sections 7. on molarity, 11. and 12. on molarity and
acid-base (alkali) titrations, section 14.3 on dilutions.
GCE A level advanced notes on the
structure of hydrated salts
Revision KS4 Science GCSE/IGCSE/O level
Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel
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Chemistry (revise courses equal to US grade 8, grade 9
SO2(aq) + Cl2(aq) + 2H2O(l) ==> 2HCl(aq)
+ H2SO4(aq) SO2(aq) + Cl2(aq) + 2H2O(l) ==> 2Cl-(aq) + SO42-(aq) + 4H+(aq)
CaCO3(s) + H2O(l) +
CO2(g) ==> Ca(HCO3)2(aq) CaSO4(aq) + 2C17H35COONa(aq) ==> (C17H35COO)2Ca(s
for scum!) + Na2SO4(aq) Ca2+(aq) + 2C17H35COO-(aq)
==> (C17H35COO-)2Ca2+(s) Ca(HCO3)2(aq) ==> CaCO3(s) + H2O(l) + CO2(g)
CaCO3(aq) + 2H+(aq) ==> Ca2+(aq) + H2O(l) + CO2(aq) CaCO3
(aq) + 2CH3COOH(aq) ==> Ca2+(CH3COO-)2(aq) + H2O(l) + CO2(aq) CaCO3(aq) +
2HCl(aq) ==> CaCl2(aq) + H2O(l) + CO2(aq) CaCO3(aq) +
2H+Cl-(aq) ==> Ca2+(Cl-)2(aq) + H2O(l) + CO2(aq) CaCO3(aq) + 2H+(aq) ==>
Ca2+(aq) + H2O(l) + CO2(aq) CaSO4(aq) + Na2CO3(aq) ==> CaCO3
(s) + Na2SO4(aq) Ca2+(aq) + CO32-(aq) ==> Ca2+CO32-(s) CO2(aq) + H2O(l)
<==> H+(aq) + HCO3-(aq) SO2 + Cl2 + 2H2O ==> 2HCl + H2SO4 SO2 + Cl2 +
2H2O ==> 2Cl- + SO42- + 4H+ CaCO3 + H2O +
CO2 ==> Ca(HCO3)2 CaSO4 + 2C17H35COONa ==> (C17H35COO)2Ca + Na2SO4 Ca2+ +
==> (C17H35COO-)2Ca2+ Ca(HCO3)2 ==> CaCO3 + H2O + CO2 CaCO3 + 2H+ ==> Ca2+
+ H2O + CO2 CaCO3
+ 2CH3COOH ==> Ca2+(CH3COO-)2 + H2O + CO2 CaCO3 + 2HCl ==> CaCl2 + H2O + CO2
2H+Cl- ==> Ca2+(Cl-)2 + H2O + CO2 CaCO3 + 2H+ ==> Ca2+ + H2O + CO2 CaSO4 +
Na2CO3 ==> CaCO3
+ Na2SO4 Ca2+ + CO32- ==> Ca2+CO32- CO2 + H2O <==> H+ + HCO3-
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