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School Chemistry Notes: Explaining hard and soft water

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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.

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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.

volumetric apparatus

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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1. Water cycle, treatment, pollution  *  2. Colloids – sols, foam and emulsions 

3. Hard & soft water: causes & treatment (this page)

4. Gas and salt solubility in water and solubility curves  *  5. Calculation of water of crystallisation

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