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Doc Brown's
chemistry revision notes: GCSE chemistry, IGCSE
chemistry, O level
& ~US grades 9-10 school science courses or equivalent for ~14-16 year old
students of chemistry
3. Soft water, hard water - causes &
treatment
The
difference between hard water and soft water is explained and the causes and
treatment of hard water fully explained
Extra Aqueous Chemistry
Index:
1.
Water
cycle, potable water, water treatment, pollution, tests for ions
2.
Colloids – sols, foam and emulsions
3. Hard & soft water: Causes and treatment (this page)
4.
Gas and salt solubility
in water and solubility curves
5.
Calculation of water of crystallisation
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3.
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
'soapy' 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.
- 'Hard water' means the water does
not readily give a good lather with soap and so wastes soap as
well as causing a 'scum'! though the 'hardness' does
not affect soapless detergents.
- Hard water come in two varieties (or a
mixture of them).
- Temporary hard water can be softened by
boiling the water.
- Temporary hardness is usually caused by the
thermally unstable magnesium hydrogencarbonate and calcium
hydrogencarbonate dissolved in the water from geological formations like
limestone or chalk.
- Permanently hard water cannot be softened
by boiling.
- Permanent hardness is caused by very soluble
magnesium sulfate (from salt deposits underground) and slightly
soluble calcium sulfate (from gypsum deposits).
- In order to soften water you must
remove the calcium ions (Ca2+) or
magnesium ions (Mg2+) ions from it by one means or
another (methods discussed in detail later).
- So, what causes water to be hard?
- The 'scum' is due to the formation of
grey–white insoluble calcium and magnesium compound formed by a reaction between the soap
molecules and calcium (Ca2+) and magnesium ions (Mg2+).
- This is a method of removing the calcium and
magnesium ions, so excess soap does soften the hard water.
- Eventually, if enough soap is added you get
a good soap bubble lather when all the calcium and magnesium ions are
precipitated, so the water is effectively softened, BUT at the expense of
wasted soap AND 'dirty' water from the 'scum'!
- These insoluble salts are seen as a grey-white
precipitate, commonly known as 'scum'!
- The details of chemical reaction between the
soap and the dissolved calcium and magnesium compounds is dealt with later.
- A variety of magnesium and calcium compounds
in water cause it to be hard e.g. magnesium sulfate (forms Epsom Salts),
calcium sulfate (gypsum), magnesium hydrogencarbonate and calcium
hydrogencarbonate from dissolved limestone, but note ....
- the sulfate salts of calcium and magnesium
cause 'permanent hardness', because boiling the water does not remove
the hardness, these salts are unaffected by heat,
- and the hydrogencarbonates of calcium and
magnesium cause 'temporary hardness', because boiling the
water removes the hardness, these compounds are decomposed on heating.
- Despite the problems with hard water e.g.
scum formation, furring-up of kettles etc. there are some healthy aspects
to living in a hard water area.
- The calcium compounds (calcium ions) are
good for teeth and bone growth and maintenance.
- Traces of iodine and iron compounds in hard
water provide other essential minerals for the body.
- People living in hard water areas (e.g.
limestone areas) generally have a lower incidence of heart disease, though
the reasons are not fully understood.
- Most hardness is due to water containing dissolved calcium or
magnesium compounds.
- What is the origin of the compounds that
cause 'hardness' in water?
- 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,
MgCO3, these become dissolved in rain naturally acidified
with carbon dioxide gas, dissolved from the atmosphere..
- Gypsum rock deposits, which are
mainly calcium sulphate CaSO4 (calcium sulfate)
which is slightly soluble in water,
- and magnesium sulphate which was
called 'Epsom Salts', formula MgSO4.7H2O
(hydrated magnesium sulfate), 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
carbonate) ...
- The equations for the formation of temporary hard
water due to calcium hydrogencarbonate and magnesium hydrogencarbonate.
- insoluble calcium carbonate + water + carbon
dioxide ==> soluble calcium hydrogencarbonate
- CaCO3(s) + H2O(l)
+ CO2(g) ==> Ca(HCO3)2(aq)
- The H2O + CO2
is sometimes written as H2CO3 (referred to as
'carbonic acid'), but this does not exist, and what we are dealing with is
slightly 'carbonated water', a very weak acid solution.
- This is enough to very slowly dissolve
limestone and chalk, and if rainwater runs off acidic soil then the effect
is enhanced.
- insoluble magnesium carbonate + water + carbon
dioxide ==> soluble magnesium hydrogencarbonate
- MgCO3(s) + H2O(l)
+ CO2(g) ==> Mg(HCO3)2(aq)
- You find the equations involving magnesium
and calcium very similar because they are next to each other in the same
group, namely Group 2 Alkaline Earth Metals.
- 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
(Palmolive soap from palm oil) or sodium oleate (from olive oil), NOT modern household
washing up detergents etc.
- 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 titration does not involve an indicator is used, the end–point is detected by the
appearance of decent frothy lather!
- This method
to determine the hardness in water is described near the end of
the page.
- 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.
- Modern soap powders contain chemical agents
to stop the scum precipitates forming, effectively incorporating a 'water
softener'.
- 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' ..
- The equations for the formation of the 'scum'
precipitate when soap is added to hard water.
- calcium sulfate + sodium stearate
(a soap) ==>calcium stearate (scum ppt.) + sodium sulfate
- CaSO4(aq)
+ 2C17H35COONa(aq) ==> (C17H35COO)2Ca(s for scum!) + Na2SO4(aq)
- or more simply ionically:
- calcium ion + stearate ion ===> calcium
stearate
- Ca2+(aq)
+ 2C17H35COO–(aq) ==> (C17H35COO–)2Ca2+(s)
- This ionic equations applies to ANY dissolved
calcium compound such as the sulfate or hydrogencarbonate.
- magnesium sulfate + sodium palmitate
(a soap) ==>magnesium palmitate (scum ppt.) + sodium sulfate
- MgSO4(aq)
+ 2C15H35COONa(aq) ==> (C15H35COO)2Ca(s for scum!) + Na2SO4(aq)
- or more simply ionically:
- calcium ion + palmitate ion ===> calcium
palmitate
- Mg2+(aq)
+ 2C15H35COO–(aq) ==> (C15H35COO–)2Mg2+(s)
- Like calcium compounds, this ionic equations
applies to ANY dissolved magnesium compound such as the sulfate or
hydrogencarbonate.
- A
precipitation reaction is generally defined as 'the formation of an
insoluble solid on mixing two solutions or a gas bubbled into a
solution'.
- Below are some diagrams of the
organic molecules or ions involved
-
- Diagram S1: The stearic acid molecule C17H35COOH
or CH3(CH2)16COOH
-
- 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 in solution of the dissolved hydrogencarbonates producing insoluble calcium carbonate (so it does
remove some of the temporary hardness before washing!).
- When heated, hard water can form lime scale
on the insides of boilers, kettles and any hot water pipe.
- Badly scaled-up pipes and boilers will be
less efficient and may need replacing, all adding to the cost of running a
home in a hard water area.
- If the heating element of a kettle becomes
'furred-up' due to scale deposition, the scale acts as an insulator and
makes the kettle less efficient.
- Temporary hardness is caused by the
presence of the hydrogencarbonate ion, which decomposes to give the
magnesium carbonate or calcium carbonate deposit.
- Hard water caused e.g. by dissolved
magnesium sulfate or calcium sulfate is permanent hardness because the
sulfate salts don't decompose on heating, so, the hardness is NOT
removed on boiling the hard water.
- The scale formation (e.g. the furring up of kettles) reactions are the exact
opposite the acid carbonated water dissolving
limestone, chalk minerals etc. So, the chemical equations for the removal
of temporary hardness by boiling are ...
- calcium hydrogencarbonate ==> calcium
carbonate + water + carbon dioxide
- Ca(HCO3)2(aq) ==>
CaCO3(s) + H2O(l)
+ CO2(g)
- similarly to give the other possible
insoluble carbonate ...
- magnesium
hydrogencarbonate ==> magnesium carbonate + water + carbon dioxide
- Mg(HCO3)2(aq)
==> MgCO3(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.
- Concentrated hydrochloric acid is used
in some 'limescale removers'
- 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?).
- The equation for limescale dissolving in
hydrochloric acid is ...
- 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)
- more simply and the more correct ionic
equation ...
- CaCO3(aq)
+ 2H+(aq) ==> Ca2+(aq)
+ H2O(l) + CO2(aq)
- Apart from boiling water, which only removes
temporary hardness,
- How else can we soften all types of
hardness in water?
- Hard water can be made soft by removing the
dissolved calcium and magnesium ions.
- This is the essential science behind how to
soften hard water, whether it be temporary hardness or permanent
hardness.
- If the hardness in water is due to calcium
hydrogencarbonate or 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 then precipitates out the calcium or magnesium
ions as their insoluble carbonates(s)
- The equations to show the precipitate
formation i.e. the removal of calcium or magnesium ions, that is the
removal of the source of hardness in the water ..
- calcium sulphate + sodium carbonate
==> calcium carbonate + sodium sulphate
- CaSO4(aq)
+ Na2CO3(aq) ==> CaCO3(s) + Na2SO4(aq)
- ionic equation: Ca2+(aq)
+ CO32–(aq) ==> Ca2+CO32–(s)
or CaCO3(s)
- magnesium sulphate + sodium carbonate
==> magnesium carbonate + sodium sulphate
- MgSO4(aq)
+ Na2CO3(aq)
==> MgCO3(s) + Na2SO4(aq)
- ionic equation: Mg2+(aq)
+ CO32–(aq)
==> Mg2+CO32–(s)
or MgCO3(s)
- If the calcium or magnesium ions are no
longer in the water, which is now softened, so they cannot cause
scum formation with soap.
- Packs of ion
exchange resins can hold or release ions in an ion exchange process.
- Ion exchange resins are solid
polymer materials with 'immobile' positive or negative ionic groups
on the polymer chains which can hold onto 'mobile' oppositely
charged ions.
- Negatively charged 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 positive calcium or magnesium ions are held on the
oppositely charged, negatively charged
resin (effectively trapped on the resin).
- The freed hydrogen or sodium ions do not form a scum with soap.
- e.g. using simple ionic equations
to illustrate the ion exchange
- 2[resin]–H+(s)
+ Ca2+(aq) ==> [resin]–Ca2+[resin]–(s)
+ 2H+(aq)
- 2[resin]–Na+(s)
+ Mg2+(aq) ==> [resin]–Mg2+[resin]–(s)
+ 2Na+(aq)
- 2[resin]–H+(s)
+ Mg2+(aq) ==> [resin]–Mg2+[resin]–(s)
+ 2H+(aq)
- 2[resin]–Na+(s)
+ Ca2+(aq) ==> [resin]–Ca2+[resin]–(s)
+ 2Na+(aq) etc.
- What I'm trying to illustrate
(hopefully successfully!), is the simple ion exchange
mechanism by which hydrogen ions and sodium ions replace the calcium
ions and magnesium ions which cause the hardness in water.
- The negative charges on the ion
exchange resin are on immobile groups that form part of the
molecular structure of the resin, BUT the attached positive ions are
mobile and can be exchanged with other ions.
- Extra Note on water
purification and ion exchange resins: 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.
 The
Determination of hardness in a sample of water
How can you measure the hardness in
water?
The apparatus and chemicals needed: A
standard soap solution (NOT detergent), a conical flask & rubber
bung, burette and stand. A 50cm3 measuring cylinder
should be accurate enough for the soft/hard water titration with
the soap solution. Its not a very accurate titration.
Procedure for determining the
hardness in water
This method of determining hardness
involves a soap titration.
You fill the burette (10 cm3
or 50 cm3 capacity, depending on how big the
titrations are) and level of the reading to 0.0. The bottom of
the meniscus should be on the reading you take e.g. initially
zero.
Measure out 50 cm3 of an
unboiled water sample into a conical flask and a rubber
bung to provide a good seal. Keep the tip of the burette down
into the flask.
Add the soap solution in small
portions e.g. 0.5 cm3 at a time. After each addition,
put the bung on the conical flask and give the mixture a good
shake for a few seconds.
Repeat this with further portions of
the soap solution.
The titration is complete when a
good soapy bubble lather is obtained that persists for at least
30 seconds.
The whole procedure is repeated with
50 cm3 of the same water sample, BUT this time from the water is
boiled for a few minutes before doing the soap titration.
By analysing boiled and unboiled samples of water,\you can
deduce the relative amounts of temporary and permanent hardness
in the water samples.
Record all your results in a
suitable table for analysis and drawing conclusions about
whether the water is hard or not, AND how much of the hardness
is temporary or permanent?
The procedure is illustrated in the
diagram above where three situations are described
 0.5 cm3
of soap solution added to 50 cm3 of distilled or
de-ionised water in the
conical flask. Put the bung on and shake vigorously. You should
get a good lather immediately. You can also consider this as a
'blank titration' and subtract it from the final titration
reading. You should not detect any hardness in pure water, but
it always takes a little soap to get a good lather of soap
bubbles. This 0.5 cm3 titration on relatively poor
water acts as what we call a 'blank titration' which would apply
to all samples. Therefore all your hard water titration results
should be corrected by subtracting 0.5.
 If
the water contains any hardness a grey-white precipitate will
form ('scum') instead of a lather, on further additions (e.g. 0.5
cm3 portions) of soap solution, shaking after each
addition.
 Eventually
enough soap is added to precipitate all the hardness and a good
lather will then form.
|
SAMPLE |
Volume of soap solution
(cm3)
to give a good lather of soap bubbles
Relative
units of hardness |
|
unboiled water |
boiled water |
|
|
burette reading |
corrected reading |
burette reading |
corrected reading |
|
A. |
0.5 |
- |
0.5 |
- |
|
B |
12.0 |
11.5 |
11.0 |
10.5 |
|
C |
8.5 |
8.0 |
1.0 |
0.5 |
|
D |
14.0 |
13.5 |
7.5 |
7.0 |
Some typical Results
Think of the
titration volumes as 'units of hardness', arbitrary,
but enables comparisons to be made between water samples as
long as it involves the same standard soap solution.
(i) Subtracting
0.5 'blank' value from the titration value of the unboiled
water gives you the total hardness units.
(ii) The
difference between the titration volumes of the boiled and
unboiled hard water, gives you the units of temporary
hardness.
(iii) The
corrected titration value for the boiled water gives you the
permanent hardness units. The boiled sample titration will
often be less because boiling destroys temporary hardness.
Water sample
A
Soft water
containing no magnesium or calcium ions: It should just take
a small amount of soap solution e.g. 0.5 cm3 to give a good
lather on shaking the 'stoppered' conical flask. So you can
consider the 'blank titration' as 0.5 cm3.
Samples B to D
all contain some hardness because the soap titration volumes
are >0.5 cm3.
Water sample
B
(i) Total
hardness = 11.5 units.
(ii) temporary
hardness = 11.5 - 10.5 = 1.0 units
(iii) permanent
hardness = 10.5 units
Conclusion:
Mainly permanent hardness
Water sample
C
(i) Total
hardness = 8.0 units
(ii) temporary
hardness = 8.0 - 0.5 = 7.5 units
(iii0 permanent
hardness = 0.5 units
Conclusion:
Mainly temporary hardness
Water sample
D
(i) Total
hardness = 13.5 units
(ii) temporary
hardness = 13.5 - 7.0 = 6.5 units
(iii) permanent
hardness = 7.0 units
Conclusion:
A roughly equal mixture of permanent hardness and
temporary hardness
Extra Notes:
This method does not distinguish
between hardness caused magnesium compounds and hardness
caused by calcium compounds in the water.
|
Any other apparatus you are
likely to need, should be on the diagram above!
Extra Aqueous Chemistry
Index:
1. Water cycle, treatment, pollution
2. Colloids – sols, foam and emulsions
3. Hard & soft water: causes &
treatment
4. Gas and salt
solubility in water and solubility curves
5.
Calculation of water of crystallisation
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chemistry equation keywords:
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+ +
2C17H35COO–
==> (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
CaCO3 +
2H+Cl– ==> Ca2+(Cl–)2 + H2O + CO2 CaCO3 + 2H+ ==> Ca2+ + H2O + CO2 CaSO4 +
Na2CO3 ==> CaCO3
+ Na2SO4 Ca2+ + CO32– ==> Ca2+CO32– CO2 + H2O <==> H+ + HCO3– what
is hard water? what causes hardness in water? what is permanently hard
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soft water, detailed revision notes on hard and soft water, why can
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hardness in water include calcium sulfate, calcium hydrogencarbonate,
magnesium sulfate and magnesium hydrogencarbonate, how can you estimate
hardness in water using a soap solution? what apparatus do you need to
measure hardness in water, hard water titrations using soap solution, what
is the scum when soap is added to hard water? why do kettles fur up in hard
water limestone areas? why do get hard walk in chalk landscapes? how can you
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hard & soft water causes treatment chemistry
1. Water cycle, treatment, pollution
* 2. Colloids – sols, foam and emulsions
* 3. Hard & soft water: causes &
treatment
4. Gas and salt
solubility in water and solubility curves * 5.
Calculation of water of crystallisation
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