|
 
Doc
Brown's Chemistry Clinic
Extra
KS4 Science GCSE-IGCSE revision-information chemistry notes on
Water and Aqueous Chemistry
including water as
a resource, colloids, acid-base theory, more on salt preparations and salt solubility
See also the KS3-GCSE notes page
on pH, acids, bases and salts
1.
Water: cycle, resource, treatment, pollution,
colloids (sol, foam, emulsion), hard/soft water * 2.
Gas and salt solubility * 3. Acid-base
proton-hydrogen ion (H+) theory of
weak/strong acids/bases (alkalis) * 4. Methods of making salts-
direct synthesis of elements and precipitation reactions *
5.
Calculation of water of crystallisation *
Basic stuff on pH,
Acids, Bases, Neutralisation and Salts * Selected salt solubility data and simple
calculations *
EMAIL
comment?query
1.
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
various ways.
1a Water
Cycle and Resources *
1b Water
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 cycle.
- 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.
1b Water
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
colloidal clay, microscopic organisms,
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)
+ H2SO4(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
below),
- (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
below).
- 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'
taste.
- Cooked vegetables turn brown (complex
compounds with phenols).
- Colloidal clay: A
colloid
consists of one substance (or mixture of substances) very finely dispersed
in
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
phase.
- 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
properties below).
- Examples of colloids
that is the fine dispersion of one substance in another without a new
solution forming:
- 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 in another liquid e.g. (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) and (iii)
margarines
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.
- Colloidal particles may be electrically
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
solution.
- 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
power: Al3+ > Mg2+ > Na+
or [Fe(CN)6]3- > SO42-
> Cl-
- and this explains why aluminium
sulphate Al2(SO4)3 is used to
precipitate (coagulate) colloidal clay in water treatment.
- Other onsite references to water
pollution:
1c Hard and
Soft Water
- 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
with soap
(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.
- Most hardness is due to water containing dissolved calcium or
magnesium compounds. 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)
- or Gypsum rock deposits, which are
mainly calcium sulphate CaSO4,
- and magnesium sulphate 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 and magnesium sulphates are washed
out of rock formations.
- Calcium and magnesium carbonates dissolve
in acid rainwater to hydrogencarbonates
e.g. naturally carbonated water
(dissolved carbon dioxide makes water acidic so it reacts with the
carbonate) ...
- calcium carbonate + water + carbon
dioxide ==> 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.
- * 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.
- It is a simple and effective way of
comparing the 'hardness' in water samples.
- The
apparatus is the same as that used in salt preparation method
(a), but no indicator is used, the end-point is detected by the
appearance of decent froth!
- 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.
Hard water
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
magnesium salts
of the soap are formed ... 'example of a precipitation
reaction' ..
- CaSO4(aq)
+ 2C17H35COONa(aq)
==> (C17H35COO)2Ca(s for scum!) + Na2SO4(aq)
- or more simply ionically: Ca2+(aq)
+ 2C17H35COO-(aq)
==> (C17H35COO-)2Ca2+(s)
- A
precipitation reaction is generally defined as 'the formation of an
insoluble solid on mixing two solutions or a gas bubbled into a
solution'.
- 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)
+ CO2(g)
- 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
any acid
(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
heating element.
- calcium carbonate + ethanoic
acid ==> calcium ethanoate + water + carbon dioxide
- CaCO3(aq)
+ 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
- CaCO3(aq)
+ 2HCl(aq) ==> CaCl2(aq)
+ H2O(l) + CO2(aq)
- or showing the ions involved
- CaCO3(aq)
+ 2H+Cl-(aq) ==> Ca2+(Cl-)2(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
- CaSO4(aq)
+ Na2CO3(aq)
==> CaCO3(s) + Na2SO4(aq)
- or more simply ionically: Ca2+(aq)
+ CO32-(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)
+ 2H+(aq)
- 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-
is
potentially harmful) and sulphate ions (SO42-)e.g.
- [resin]+OH-(s)
+ Cl-(aq) ==> [resin]+Cl-(s)
+ OH-(aq)
- [resin]+OH-(s)
+ NO3-(aq) ==> [resin]+NO3-(s)
+ OH-(aq)
- 2[resin]+OH-(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
water.
- H+(aq)
+ OH-(aq) ==> H2O(l)
- However, it will not remove
non-ionic substances like organic pesticides etc.
- Permanently hard water means the
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
hydrogencarbonate.
- 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.
2.
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:
-
solute: the
material which is to be dissolved in a solvent.
-
solvent: the
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.
-
solution: the
result of dissolving something in a liquid (solute + solvent =>
solution).
-
solubility: to
what extent a solute material will dissolve.
-
soluble: the
material will dissolve in a particular liquid solvent.
-
saturated:
means that no more of a substance (the solute) will dissolve in its
solution i.e. maximum solubility achieved at a particular
temperature.
-
insoluble: not
soluble, will not dissolve in a particular liquid (don't assume it
means will not dissolve in anything).
-
hydration:
means the addition of water to a material.
-
dehydration:
means to remove water from a substance.
-
Factors affecting
rates of dissolving.
-
heat: heating
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.
-
surface area:
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.
-
stirring:
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
concentrated
-
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
increases.
-
Carbonated water
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.
-
Thermal Pollution:
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
assumptions!).
-
The
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
(crystallisation). The
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.
-
Solubility curves:
Excel data file/graph of
selected solubility ... original excel file ... web
page version
-
General rules
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
-
All nitrate
salts are soluble
e.g. NaNO3,
Mg(NO3)2, Al(NO3)3, NH4NO3
-
Some ethanoate
salts are soluble
e.g. CH3COONa
-
Common chloride
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.
-
soluble:
Na2SO4,
MgSO4, Al2(SO4)3
-
insoluble: PbSO4, BaSO4, CaSO4
is slightly soluble.
-
Common oxides,
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:
-
soluble:
K2O,
KOH, NaOH, NH4OH actually NH3(aq), Na2CO3,
(NH4)2CO3
-
insoluble:
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.
-
See section 4. for
summary of salt preparation methods and details of some methods on
Acids,
Bases Salts page.
3.
Why do some substances
produce acidic or alkaline solutions?
Most general aspects of Acids,
Bases, Salts and pH are covered on a separate web page which should be
studied first. There is more on salt preparations in
section 4.
-
Some compounds react will water to produce
acidic or alkaline solutions.
-
Water must be present for a substance to
act as an acid or as a base (usually at gcse level!).
-
Acids in aqueous solution produce
hydrogen H+
ions. The H+ ion is a proton. In water this proton is hydrated
(associated with water and more correctly expressed as H3O+(aq))
but H+(aq) is adequate here. The greater the
concentration of hydrogen ions the more acid the solution and the lower the pH.
-
Alkalis in aqueous solution produce OH-(aq)
hydroxide ions. The greater the concentration of
hydroxide ions the more alkaline the solution and the higher the pH.
-
When alkalis and acids
react,
the 'general word' and e.g. 'molecular formula' neutralisation equation might be ...
-
ACID
+ ALKALI ==> SALT
+ WATER ... e.g.
-
hydrochloric
acid + sodium hydroxide ==> sodium
chloride + water
-
HCl(aq)
+ NaOH(aq)
==> NaCl(aq) + H2O(l)
-
BUT
the ionic equation for ANY neutralisation is
-
H+(aq)
+ OH-(aq) ==> H2O(l)
-
and the remaining ions
e.g. Na+(aq) and Cl-(aq) become
the salt crystals NaCl(s) on evaporating the water.
-
Acids can be defined as proton donors.
A base can be defined as a proton acceptor (Bronsted-Lowry theory).
-
e.g. here the hydroxide ion is the base
and accepts a proton from an acid.
-
or here the hydrogen chloride is the
acid and the ammonia is the base when ammonium chloride is formed when
the two gases are mixed. The acid hydrogen chloride donates a proton
to the base ammonia. (note: no water present!)
-
or copper(II) oxide (base) +
sulphuric acid (acid) ==> copper(II) sulphate + water
-
CuO(s)
+ H2SO4(aq)
==> CuSO4(aq) + H2O(l)
-
ionically it is: Cu2+O2-(s)
+ 2H+(aq)
==> Cu2+(aq) + H2O(l)
-
Acids are characterised by having at
least one replaceable hydrogen atom in forming a salt, the H is
replaced by a metal ion (Na+, Mg2+ etc.) or the
ammonium ion (NH4+):
-
Several scientists have made contributions
to ionic and acid-base theory e.g.
-
Arrhenius (1887), was one of the first
scientists to suggest that substances could split into free positive
and negative ions when
dissolved in water, the so called 'electrolytic dissociation'
giving rise to electrically conducting solutions. His theory was
considered a bit revolutionary, and he was given a low rating for his PhD
at Paris at first! - however the 'professors' recanted when
other scientists decided it was a good idea and in 1903 he was awarded
the Nobel Prize for his ionic theory work!
-
Lowry and Bronsted (1923) took
further the work of Arrhenius and applied ionic theory to the concept
of acids and bases - that is, that acids and bases are proton donors
and acceptors (see above). It should be noted that the work of Arrhenius took much longer to be accepted than the work of
Lowry and Bronsted because there was no pre-existing (and proven) theory of
ion formation.
-
Acids and alkalis are
further classified by the
extent of their ionisation in water.
-
They are described as strong or weak
depending on their degree of ionisation in water.
-
Do not confuse
the terms weak
and strong about how far the 'molecules' become ionised in
water with the terms dilute and
concentrated, they mean different things!
-
Dilute and
concentrated refer to the concentration of the acid or
alkali in terms of how much (i.e. a little or a lot) of the original material is dissolved
in water as measured by concentration e.g. molarity.
-
A strong acid or alkali is one that
is that is nearly or completely 100% ionised in water
(not an equilibrium situation)
-
examples of strong acids are hydrochloric,
sulphuric and nitric acids.
-
e.g. the maximum (or nearly) hydrogen
ion concentration results in the lowest pH ...
-
nitric acid is: HNO3(l)
+ aq ==> H+(aq) +
NO3-(aq)
-
and sulphuric acid is: H2SO4(l)
+ aq ==> 2H+(aq) + SO42-(aq)
-
The greater
the concentration of hydrogen ions the lower the pH, so strong
acids make the most acidic solutions.
-
examples of strong alkalis
(soluble strong bases) are sodium hydroxide or potassium hydroxide
etc. (usually Group 1 or 2 hydroxides).
-
A weak acid or alkali is only partially
ionised in water.
-
examples of weak acids are ethanoic, citric and carbonic
acids.
-
e.g. for ethanoic about 2% ionises
(forward reaction to the right), the
equilibrium lies mainly to the un-ionised form on the left and for the
weaker carbonic acid even less is ionised. So only a relatively low
concentration of free hydrogen ions form giving a less acidic higher pH
solution than
strong acids (but pH still less than 7) ...
-
An example of a weak alkali/base
(weak soluble base) is ammonia
solution, about 2% changes to the ionic forms on the right. So only a
relatively low concentration of free hydroxide ions form giving a less
alkaline solution, so the pH is less than a strong base/alkali (but
pH still over 7) ...
-
NH3(aq) + H2O(l) NH4+(aq)
+ OH-(aq)
-
or sodium carbonate: CO32-
+ H2O(l)
HCO3-(aq) + OH-(aq)
-
both of which, when
dissolved in water, produce hydroxide ions giving an alkaline solution, despite the fact that OH doesn't appear in their
formulae!
-
You can distinguish between strong and weak acids of the same concentration by
using the pH scale and observations from a variety of experiments
support the low or high of ionisation theory.
-
e.g. by the rate of reaction with metals.
-
If you put magnesium ribbon into 1 molar
solutions of hydrochloric acid (strong, high % ionisation so high H+(aq)
concentration) and ethanoic acid (weak, low percentage ionization so much lower H+(aq)
concentration), you can
see the difference in the fast and slow 'fizzing' rates!
-
Since stronger/weak
acid solutions (or alkalis) contain more/less hydrogen ions, they are
better/poorer conductors of electricity.
-
e.g. If you carry out
electrolysis experiments with the same two solutions, you get a much greater
volume of hydrogen collected at the cathode from the hydrochloric acid
compared to the ethanoic acid.
-
You must use solutions
of the same concentration and electrolysed them for the same time before
measuring the gas volumes (Electrolysis
methods 1a and 1b).
-
Remember that its
the H+ ion that is the active chemical species in acid
solutions NOT a 'HCl' or a 'H2SO4' or a 'CH3COOH'
molecule.
-
More on pH scale and
indicators on Acids,
Bases, Salts page.
-
The pH is dependent on the relative concentrations of the H+(aq)
and the OH-(aq) concentrations.
-
a high H+(aq)
concentration means a low pH and low OH-(aq)
concentration, usually strong acid
-
lower H+(aq)
concentration means higher pH and higher OH-(aq)
concentration, less acid
-
a high OH-(aq)
concentration means a high pH and low H+(aq)
concentration, usually strong base/alkali
-
lower OH-(aq)
concentration means lower pH and higher H+(aq)
concentration, less alkaline
-
In general: pH 1-2 strong
acids, pH 3-6 weak acids, pH
7 neutral, pH 8-11 weak base/alkali, pH 12-14 strong base/alkali
-
Neutralisation ionically is: H+(aq)
+ OH-(aq)
==> H2O(l)
(exothermic)
-
The pH of a solution, or determining
the neutralisation point, can be measured with
-
When mixing an acid and alkali the
neutralisation end-point can also be determined by
-
Further work-study
links:
4.
What different ways are
there of making salts?
-
Most general aspects of Acids,
Bases, Salts and pH are covered on a separate web page and there is more
on acid-base theory and neutralisation in section 3. above.
-
Neutralising an acidic solution with an
alkaline solution is one way of making of making salts but salts can also be made
using several other methods. All the methods (1) to (5) are listed below, but
methods (1) to (3) are described on the Acids,
Bases, Salts and pH page.
-
(1) Reaction of an acid with a metal, see method
(b) on Acids, Bases and Salts , the metal is effectively a water insoluble
material that dissolves in acid to form the salt.
-
(2) Reaction of an acid with an insoluble base using water insoluble, oxide, hydroxide or carbonate) with an acid, see method
(b) on Acids, Bases and Salts page.
-
(3) Reaction of a soluble base (alkali) with an
acid, see method
(a) on Acids, Bases and Salts page. Usually an Alkali metal hydroxide or ammonia.
-
(4) An insoluble salt
can be made by
mixing two solutions of soluble salts
in a process is called precipitation.
One solution contains the 1st required ion, and the other solution
contains the 2nd required ion. The
precipitated salt can then be filtered off with a filter funnel and paper. The
collected solid is washed with distilled water to
remove any remaining soluble salt impurities and removed from the
filter paper to be dried. Examples ...
-
(i) Silver chloride is made by
mixing solutions of solutions of silver nitrate and sodium chloride.
-
silver nitrate + sodium chloride ==>
silver chloride + sodium nitrate
-
AgNO3(aq) + NaCl(aq)
==> AgCl(s) + NaNO3(aq)
-
in terms of
ions it could be written as
-
Ag+NO3-(aq)
+ Na+Cl-(aq) ==> AgCl(s) +
Na+NO3-(aq)
-
or: Ag+(aq)
+ NO3-(aq)
+ Na+(aq)
+ Cl-(aq) ==> AgCl(s) +
Na+(aq)
+ NO3-(aq)
-
but the
spectator
ions are
nitrate NO3- and
sodium Na+
which do not change at all,
-
so the ionic
equation is simply: Ag+(aq)
+ Cl-(aq) ==> AgCl(s)
-
Note that ionic
equations omit ions that do not change there chemical or physical
state.
-
In this case the
nitrate, NO3-(aq) and sodium
Na+(aq)
ions do not change physically or chemically and are called
spectator ions,
-
BUT the aqueous
silver ion, Ag+(aq), combines with the aqueous
chloride ion, Cl-(aq), to form the insoluble
salt silver chloride, AgCl(s), thereby changing their
states both chemically and physically.
-
More Ionic equations
explained with all spectator ions
indicated.
-
If you use
barium chloride the word and symbol equations are ...
-
barium
chloride + silver nitrate ==> silver chloride + barium
nitrate
-
BaCl2(aq)
+ 2AgNO3(aq) ==> 2AgCl(s) +
Ba(NO3)2(aq)
-
which can be
written as
-
Ba2+(aq)
+ 2Cl-(aq) + 2Ag+(aq) +
2NO3-(aq) ==> 2AgCl(s)
+ Ba2+(aq) + 2NO3-(aq)
-
the spectator
ions are Ba2+ and
NO3-
-
so the ionic
equation is: Ag+(aq)
+ Cl-(aq) ==> AgCl(s)
-
(ii) Lead(II) iodide,
a yellow precipitate (insoluble in water!) can be made
by mixing lead(II) nitrate solution with e.g. potassium iodide solution.
-
lead(II) nitrate + potassium iodide
==> lead(II) iodide + potassium nitrate
-
Pb(NO3)2(aq) +
2KI(aq)
==> PbI2(s) + 2KNO3(aq)
-
which can be
written as
-
Pb2+(aq)
+ 2NO3-(aq) + 2K+(aq)
+ 2I-(aq) ==> PbI2(s) +
2K+(aq) + 2NO3-(aq)
-
the ionic equation is:
Pb2+(aq)
+ 2I-(aq) ==> PbI2(s)
-
because the
spectator ions are nitrate NO3- and
potassium K+.
-
In a similar
way you can make lead(II) chloride by e.g. using dilute
hydrochloric acid
-
lead(II) nitrate +
hydrochloric acid
==> lead(II) chloride + nitric acid
-
Pb(NO3)2(aq)
+ 2HCl(aq)
==> PbCl2(s) + 2HNO3(aq)
-
Pb2+(aq)
+ 2NO3-(aq) + 2H+(aq)
+ 2Cl-(aq) ==> PbCl2(s)
+ 2H+(aq) + 2NO3-(aq)
-
the ionic equation is:
Pb2+(aq)
+ 2Cl-(aq) ==> PbCl2(s)
-
because the
spectator ions are
nitrate NO3- and
hydrogen H+.
-
and you can
make lead(II) bromide by e.g. using sodium bromide
-
lead(II) nitrate +
sodium bromide
==> lead(II) bromide + sodium nitrate
-
Pb(NO3)2(aq)
+ 2NaBr(aq)
==> PbBr2(s) + 2NaNO3(aq)
-
Pb2+(aq)
+ 2NO3-(aq) + 2Na+(aq)
+ 2Br-(aq) ==> PbBr2(s)
+ 2Na+(aq) + 2NO3-(aq)
-
the ionic equation is:
Pb2+(aq)
+ 2Br-(aq) ==> PbBr2(s)
-
because the
spectator ions are
nitrate NO3- and
sodium Na+.
-
(iii) Calcium carbonate,
a white precipitate, forms on
e.g. mixing calcium chloride and sodium carbonate solutions ...
-
calcium chloride + sodium carbonate
==> calcium carbonate + sodium chloride
-
CaCl2(aq) +
Na2CO3(aq)
==> CaCO3(s) + 2NaCl(aq)
-
Ca2+(aq)
+ 2Cl-(aq) + 2Na+(aq) +
CO32-(aq)
==> CaCO3(s) + 2Na+(aq)
+ 2Cl-(aq)
-
ionically: Ca2+(aq)
+ CO32-(aq) ==> CaCO3(s)
-
because the
spectator ions are chloride Cl-
and sodium Na+.
-
(iv) Barium sulphate,
a white precipitate, forms on
mixing e.g. barium chloride and dilute sulphuric acid ...
-
barium chloride + sulphuric acid
==> barium sulphate + hydrochloric acid
-
BaCl2(aq) +
H2SO4(aq)
==> BaSO4(s) + 2HCl(aq)
-
Ba2+(aq)
+ 2Cl-(aq) + 2H+(aq) +
SO42-(aq)
==> BaSO4(s) + 2H+(aq)
+ 2Cl-(aq)
-
ionic
equation: Ba2+(aq)
+ SO42-(aq) ==> BaSO4(s)
-
because the
spectator ions are chloride Cl- and
hydrogen H+.
-
Or you can use
sulphate salts like sodium sulphate, so the word and symbol
equations are ..
-
barium chloride +
sodium sulfate
==> barium sulfate + sodium chloride
-
BaCl2(aq)
+ Na2SO4(aq)
==> BaSO4(s) + 2NaCl(aq)
-
The ionic
equation is the same: Ba2+(aq)
+ SO42-(aq) ==> BaSO4(s)
-
because the
spectator ions are sodium Na+ and
chloride Cl-
-
(v) Lead(II)
sulphate, a white precipitate, forms in a similar way e.g.
-
lead(II) nitrate +
sodium sulphate ==> lead(II) sulphate + sodium nitrate
-
Pb(NO3)2
(aq) + Na2SO4 (aq) ==> PbSO4
(s) + 2NaNO3 (aq)
-
ionically: Pb2+(aq)
+ SO42-(aq) ==> PbSO4(s)
-
because the
spectator ions are sodium
Na+ and
nitrate
NO3-
-
NOTE:
A precipitation reaction is generally defined as 'the formation of
an insoluble solid on mixing two solutions or bubbling a gas into a
solution'.
-
General
guide rules on salt solubility are given above in section 2.
5.
How can we work out
and use the
concentration of solutions?
and water of crystallisation calculations
-
Solubility graphs and data
are covered in section 2.
-
Determination and
calculation of salt formula containing 'water of
crystallisation'.
-
Some salts,
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
is measured.
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
unknown) was
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%
-
[ Ar's
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
'species'
-
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
other 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.
ks4 science examinations gcse-igcse chemistry
revision *
ks4 science examinations-gcse-igcse chemistry revision * ks4 science
examinations-gcse-igcse chemistry revision * ks4 science examinations-gcse-igcse chemistry
revision * ks4 science examinations-gcse-igcse
chemistry revision * ks4 science examinations-gcse-igcse chemistry revision * SITE PURPOSE EDUCATION - online learning
or 'self-private-tuition' using revision notes, quizzes,
practice tests involving GCSE Science CHEMISTRY in the areas of REVISING
only the CHEMISTRY-Earth Science-Radioactivity at Doc Brown's
Chemistry Clinic via HOMEPAGE in secondary school/schools, 6th form college/colleges,
academy/academies or home self-study. Hopefully it will encourage
interest and understanding of Chemistry, Earth Science and Radioactivity
in any country of the world, though the site is written entirely in English. The website is designed to help
and unofficially support students/teachers revise-learn/teach the chemistry for modular
or co-ordinated examination science
courses from UK QCA based AQA, OCR (Oxford and Cambridge) Twenty First (21st) Century and Gateway Science, Edexcel
360Science ,
Nuffield, Salters, Cambridge International (CIE), London International, WJEC, CCEA
exams etc.
Also, national award assessments-examinations for GCSE-IGCSE-KS4-O
level-BTEC-NVQ
applied, additional and chemistry national science courses. Also covers,
mainly via quizzes the UK National KS3 SATs Science-biology/chemistry/physics (SAT revision levels 3-5
or 5-7) and covers much of the revising, learning and teaching chemistry
examinations for the
national curriculum for secondary schools and colleges. The site does
not support the content of England, Wales or Northern Ireland primary
science KS1 or KS2. The notes should also provide some background theory
for a coursework assignment or project. BUT please note that
my on-line revision notes and quizzes are no substitute for good classroom
teaching-lecturing and thorough studying of your own notes and textbooks, practicing past papers
and a copy of the syllabus which are readily downloaded from the
examination board sites, but I hope here and there they will lend a
tutoring hand on some topic, unit, module etc. For final revision you
have to be intellectually honest about what you don't know or follow, YOU have to
take the stuff to pieces, analyse what you do/do not understand
and reconstruct it so it all makes sense in the end. There is no other
way, there are no magic secrets on how to revise and learn, its mainly
down to hard work and just good old fashioned study and employing teach-yourself
strategies without the need for extra tutors and tutoring lessons. I also think
there is too much hit and miss revision using past papers (which I do NOT
supply) and not enough
systematic revision. I also hope it will help teachers in planning
lessons and developing schemes of work for science-chemistry. There are no
lesson plans on the site but there are plenty of quizzes to incorporate into
classroom activities whether photocopied or on electronic whiteboard projector
for use as self-tuition-assessment purposes and a variety of teaching and
learning styles and the images may be used in Microsoft Word documents and powerpoint projections.
The site seems to be used by a large number of home study tutors, particularly
the revision notes. An individual tutor may print out the notes for
science-chemistry learning teaching-tuition purposes and for background material
for assignments and projects. I have no interest or time in producing WORD.doc or xxxx.pdf files
of the notes at the moment. Neither have I time to write up many practical
laboratory experiments ('lab'-'labs') at the moment, but the notes contain lots
of background information of chemical reactions in terms of
observations-balanced equations-reactants-products-theory etc. I also find it
difficult to recommend specific exam websites or syllabus textbooks, it depends exactly on
what you need, what you have time for, and there are so many of them to choose
from and I do not supply past examination papers for classes. The sites
resources include revision notes, quizzes and worksheets which provide support
for home study or tuition for homework and coursework help e.g. science
investigations for any of the key stage courses indicated, but I do not supply
lesson plans. Dr W P Brown gcse
10-11-2007 * ks4 science examinations gcse-igcse chemistry
revision *
ks4 science examinations-gcse-igcse chemistry revision * ks4 science
examinations-gcse-igcse chemistry revision * ks4 science examinations-gcse-igcse chemistry
revision * ks4 science examinations-gcse-igcse
chemistry revision * ks4 science examinations-gcse-igcse chemistry
revision
IGCSE GCSE KS4 Science
Chemistry * IGCSE GCSE KS4 Science Chemistry * IGCSE GCSE
KS4 Science Chemistry * IGCSE GCSE KS4 Science
Chemistry * IGCSE GCSE KS4 Science Chemistry * IGCSE GCSE
KS4 Science Chemistry * IGCSE GCSE KS4 Science
Chemistry * IGCSE GCSE KS4 Science Chemistry * IGCSE GCSE
KS4 Science Chemistry * IGCSE GCSE KS4 Science
Chemistry * IGCSE GCSE KS4 Science Chemistry * IGCSE GCSE
KS4 Science Chemistry * IGCSE GCSE KS4 Science
Chemistry * IGCSE GCSE KS4 Science Chemistry * IGCSE GCSE
KS4 Science Chemistry * IGCSE GCSE KS4 Science
Chemistry * IGCSE GCSE KS4 Science Chemistry * IGCSE GCSE
KS4 Science Chemistry * IGCSE GCSE KS4 Science
Chemistry * IGCSE GCSE KS4 Science Chemistry * IGCSE GCSE
KS4 Science Chemistry * IGCSE GCSE KS4 Science
Chemistry * IGCSE GCSE KS4 Science Chemistry * IGCSE GCSE
KS4 Science Chemistry * IGCSE GCSE KS4 Science
Chemistry * IGCSE GCSE KS4 Science Chemistry * IGCSE GCSE
KS4 Science Chemistry * scientific investigations, educational development, scientific exhibitions,
scientific adventures, science projects, fantasy science, science fiction,
interesting science demonstrations, fascinating science experiments, science
education conferences, scientific expeditions, scientific information and
databases, revision tutoring resources for syllabuses specifications
examinations, chemical physical biological forensic science, scientific
applications, science-chemistry tuition courses
|