: ACIDS, BASES AND SALTS

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Unit Checklist:

  1. Acids;
  • Meaning
  • Strong and weak acids
  • Concentrated and weak acids
  • Comparing strength of acids
  • Using evolution of hydrogen and carbon (IV) oxides
  • Using electrical conductivity
  • Using PH
  • Role of solvents in acidic properties of solvents
  • Hydrogen chloride in water
  • Hydrogen chloride in methylbenzene

 

  1. Bases:
  • Meaning
  • Strong and weak bases (alkalis)
  • Measuring strength of alkalis
  • Using electrical conductivity
  • Using PH values
  • Effect of solvent type on properties of ammonia solution
  • Ammonia n water
  • Ammonia in methylbenzene
  1. Uses of acids and bases

 

  1. Oxides and hydroxides
  • Basic oxides
  • Acidic oxides
  • Neutral oxides
  • Amphoteric oxides
  • Meaning of amphoteric oxides
  • To verify amphoteric oxides
  • Reactions of amphoteric oxides

 

  1. Salts
  • Meaning
  • Preparation methods (summary)
  • Solubility of salts
  • Qualitative tests for cations using NaOH and NH4OH;
  • Effect of heat on metal oxides and hydroxides.
  • Effect of sodium carbonate on salt solutions
  • Properties of cations with sodium chloride, sodium sulphate and sodium sulphite

 

  1. Solubility and solubility curves
  • Definition of solubility
  • Factors affecting solubility
  • Solubility curves and related calculations.
  • Fractional crystallization

 

  1. Water
  • Hardness of water
  • Temporary hardness (Meaning; causes; and removal)
  • Permanent hardness (meaning; cause; and removal)
  1. Acids

– Are substances whose molecules yield hydrogen ions in water; or

– Are substances, which contain replaceable hydrogen, which can be wholly or partially replaced by a metal.

 

HCl (aq)                      H+(aq)  + Cl(aq)

 

OR: – Acids are proton donors i.e. a substance which provides protons or hydrogen ions.

 

Strength of Acids

– Acids can be categorized as either strong or weak acids;

 

  • Strong acids

– Are those which dissociate or ionize completely to a large extent in water, to yield many hydrogen ions.

– They yield to the solution as many protons as they possibly can.

 

  • Examples

Hydrochloric acid; HCl(aq)                             H+ (aq)  + Cl(aq)

Sulphuric acid; H2SO4(aq)                              2H+ (aq)  + SO4 2-(aq)

Nitric acid; HNO3(aq)                                    H+ (aq)    NO3 aq)

 

  • Weak acids.

– Are acids, which undergo partial dissociation to yield fewer hydrogen ions.

– They do not ionize in water completely or to a large extent i.e. some of their molecules remained unionized in solution.

 

Examples:

Carbonic acid:

H2CO3(aq)              water           H+(aq) +  HCO3(aq)

 

Ethanoic acid

CH3COOH (aq)            water                  H+(aq) + CH3COO(aq)

 

Note: – concentrated acids and dilute acids

Concentrated acids

– Is an acid with a high number of acid molecules per given volume.

 

Dilute acids:

Are acids with a low number of acid molecules per given volume.

 

– Thus there are concentrated strong acids or dilute strong acids; as well as concentrated weak acids and dilute weak acids.

 

 

Comparing the strength of acids

(i). Using rate of evolution of hydrogen

Apparatus:

– Boiling tubes; 1M HCl/ H2SO4/ HNO3; Methanoic acid/ tartaric acid; magnesium ribbon.

 

Procedure:

– One boiling tube is half filled with 1M HCl; while another is half filled with 1M Ethanoic acid.

– 2 pieces of magnesium ribbons are cleaned to remove a layer of oxide on the surface.

– One of the two pieces is put in each tube of the acid.

 

Observations:

– Hydrochloric acid evolves hydrogen much more quickly than Ethanoic acid yet they were of equal concentration.

 

Conclusion

– Hydrochloric acid is a strong acid;

– Ethanoic acid is a weak acid.

 

Note:

– The same experiment can be repeated with marble chips (CaCO3) in acids of same concentration.

-The marble chips dissolve more quickly in HCl, which is a strong acid.

 

(ii). Using electrical conductivity

Procedure:

– 50cm3 of 2M-hydrochloric acid solution is placed into a beaker and set up apparatus as shown below.

– The switch is closed and the brightness of the bulb noted.

 

Diagram: Electrolytic circuit

 

 

 

 

 

 

 

 

 

 

 

Observations

– Strong acids like HCl, HNO3 and sulphuric acid gave a brighter bulb light than weak acids like ethanoic, carbonic acids e.t.c

 

 

 

 

Explanations

– Strong acids are completely dissociated and have more H+ in solution and hence have got a higher electrical conductivity; than solutions of weak acids which are only partially ionized thus have fewer hydrogen ions in solution

 

 

(iii). Using PH

Procedure

– 2cm3 solutions of different acids of equal concentrations are paired into different test tubes.

– To each test tube 2 drops of universal indicator are added.

– Acids tested: HCl, H2SO4, HNO3; ethanoic acid, carbonic acid, and tartaric acid.

– All acids are of 2M solutions

– The indicator colour and hence the PH number of each is noted; by comparing against the indicator chart.

 

Observations

Substance (1M)Colour of universal indicatorPH
Sulphuric acid

Hydrochloric acid

Nitric acid

Ethanoic

Carbonic acid

Tartaric acid

Red

Red

Red

Orange

Yellow

orange

3

3

3

5

6

5

 

Explanations

– Solutions of strong acids contain a higher concentration of hydrogen ions than those of weak acids

– Strong acids have low PH usually less than 3.

– Weak acids have higher PH values usually between 5 – 6.

 

Role of solvents on acidic properties of a solute.

Experiment: – to find out if solutions of HCl in different solvents display acidic properties

Apparatus and reagents

– Hydrogen chloride gas, water and methylbenzene

– Beakers and a funnel

– Blue and red litmus papers.

 

Procedure

– Solutions of hydrogen chloride gas are made by bubbling the dry gas from a generator into water and into methylbenzene contained in separate beakers.

– The hydrochloric gas is passed into the solution using an inverted funnel to prevent sucking back.

 

Apparatus

 

 

 

 

 

 

 

 

 

The resultant solutions are each separately subjected to various tests as shown below and observations recorded

 

Tests and observations

 

Test.Aqueous HCl solutionSolution of HCl in methylbenzene
1.       A piece of dry blue litmus paper is dropped into solution

2.       Dry universal indicator paper

3.       Add magnesium ribbon

4.       Add small marble chips

5.       Electrical conductivity

Blue litmus turns red

 

– Turns red (strong acid)

– Evolution of hydrogen

– CO2 evolved

– Good conductor

No effect on litmus

 

– Turns green (neutral)

– No reaction

– No reaction

– Does not conduct

 

Explanation

– The results show that the aqueous solution of hydrogen chloride behaves as an acid; but the solution in methylbenzene lacks acidic properties

– When HCl gas dissolves in water it changes from molecules to ions;

 

Equation:

HCl (aq)        water                 H+(aq) + Cl(aq)

 

– It is the hydrogen ions which give the acidic properties and these can only be formed in the presence of water

– HCl in water conducts electric current due to presence of free ions in solution

– HCl gas in methylbenzene does not conduct electric current because the HCl exists as molecules hence lack free ions

 

Note: – hydrogen chloride gas dissolves in water because both HCl and water polar molecules;

– This causes mutual attraction of both ends of HCl molecule by different water molecules causing the dissociation of HCl molecules into ions.

 

Illustration:

 

Hence:

HCl (g) + water                       HCl (aq)

HCl (aq)                      H+(aq)  + Cl(aq)

– The presence of hydrogen ions in aqueous solution of hydrogen chloride explains the electrical conductivity and acidic properties of hydrogen chloride

  • Acidic properties: –
  • turns blue litmus paper red;
  • evolves hydrogen gas when reacted with magnesium;
  • evolves carbon dioxide on reaction with CaCO3;

 

2H+ (aq)  +  CaCO3 (s)                           Ca2+(aq)  +  CO2(g) + H2O(l)

 

2H+(aq)  +  Mg(s)                         Mg 2+(aq)   + H2  (g)

 

– Methylbenzene has a weak attraction for hydrogen chloride and hence hydrogen chloride remains as molecules in methylbenzene

 

  1. Bases

– Are substances which accept the protons donated by acids and are hence proton acceptors

NH3 (aq)   +   H+ (aq)                          NH4+(aq)

 

CuO(s) +  2H+ (aq)                         Cu 2+(aq)   +   H2O(l)

 

Alkalis

– An alkali is a soluble base i.e. a base that is soluble in water.

– They are compounds, which produce hydroxyl ions in aqueous solutions.

 

NaOH(aq)                           Na+(aq)   +  OH (aq)

 

Note: –

When an acid proton reacts with a base (hydroxyl ions) in aqueous solution, a neutralization reaction occurs.

 

Strength of an Alkali

– Alkalis can be grouped as either strong or weak alkalis.

 

(a). Strong alkalis

– Are alkalis that undergo complete dissociation in aqueous solution; yielding a large number of hydroxyl (OH-) ions

 

Examples:

– Sodium hydroxide.

– Potassium hydroxide.

 

 

(b). Weak alkalis

– Are alkalis that undergo only partial dissociation in aqueous solution (water) yielding fewer numbers of hydroxyl ions.

 

Examples

– Calcium hydroxide

– Ammonium hydroxide

Measuring the strength of alkalis

(i). Using electrical conductivity

Procedure

50 cm3 of 2 M sodium hydroxide solution is put into a beaker and the apparatus set as shown below

Apparatus

 

 

 

 

 

 

 

 

 

Procedure

– The same procedure is repeated using other alkalis like NH4OH; Ca (OH)2 e.t.c.

 

Observation

– The bulb lights brightly with KOH and NaOH as electrolyte than with NH4OH and Ca (OH)2

 

Explanation

– NaOH and KOH are strong alkalis and are completely dissociated and have more ions in solution and hence have got a higher electrical conductivity than the weak alkalis of NH4OH and Ca(OH)2(aq)

 

(ii). Using PH values

Procedure

– 2 cm3 of NaOH and 2 cm3 of NH4OH are each poured into 2 different test tubes separately

– Into each test tube 2 drops of universal indicator are added.

– The colour change is noted and the corresponding PH scale recorded

 

 

 

 

 

Observations

 

AlkaliColour of universal indicatorPH
– Ammonium hydroxide (1M)

– Calcium hydroxide (1M)

– Sodium hydroxide (0.1M)

– Sodium hydroxide (1M)

– Potassium hydroxide

Blue.

Blue.

Purple.

Purple.

Purple.

11

10

13

14

14

 

Note: the PH scale

– Is a scale which gives a measure of the acidity of alkalinity of a substance.

 

Illustration: a PH scale.

 

1     2     3     4     5     6     7     8     9     10     11     12     13     14

 

 

Strong acid               Weak acids       Neutral   Weak alkali                                    Strong alkali

 

 

Increasing acidity                                                                      Increasing alkalinity

(High H+ ion concentration)                                                   (Low H+ ion concentration)

 

 

Indicator colours:

 

PH1234567891011121314
ColourRedOrange/ redYellow/ GreenGreenGreen/ BlueBlue/ PurplePurple

 

Effects of type of solvent on the properties of ammonium solution

Procedure

– Ammonium solution is prepared by bubbling the gas from a generator into methylbenzene (toluene) and into water contained in separate beakers

– The solutions are each divided into 3 portions and tested with litmus paper; universal indicator and for electrical conductivity

 

Apparatus

 

 

 

 

 

 

Observations

TestSolution of NH3 in waterSolution of NH3 in methylbenzene (toluene)
Dry litmus paperRed litmus paper turns blue;No effect
Dry universal indicator paperColour turns purple (alkaline PH)Turns green (Neutral PH)
Electrical conductivityPoor conductorNon-conductor

Explanations

– When NH3(g) dissolves in water it changes from molecules to ions.

Equation:

NH3(g) + H2O(l)                            NH4+(aq) + OH(aq)

 

– It is the hydroxide ions that cause alkaline properties.

– Since ammonium hydroxide is a weak alkali, it dissociates partially releasing fewer hydroxide ions hence the poor electric conductivity.

– Ammonium gas in methylbenzene or trichloromethane exists as molecules without free ions hence no alkaline properties and the electrical conductivity.

Uses of acids and bases

  1. Acids

– Refer to the various acids for uses of sulphuric, nitric and hydrochloric acids.

 

  1. Bases/ alkalis

– Some weak bases e.g milk of magnesia, are used to relieve stomach disorders.

 

Amphoteric oxides and hydroxides.

  • Oxides

– An oxide is a binary compound of oxygen and another element.

– Are of four categories:

  • Basic oxides
  • Acidic oxides
  • Neutral oxides
  • Amphoteric oxides

 

(i). Basic oxides

– Are usually oxides of metals (electronegative elements)

– They react with acids to form salt and water only.

Examples

CaO, MgO, CuO etc.

 

(ii). Acidic oxides

– Are usually oxides of non metals (electronegative elements).

– Many of them react with water to form (give) acids and are known as acid anhydrides

Examples

CO2; SO2; SO3; P2O5 and NO2

 

(iii). Amphoteric oxides

– Are oxides, which behave as both bases and acids.

– Are mainly oxides of certain metals in the middle group of the periodic table.

Examples

Oxides of Zn, Al, Pb

 

Experiment: – To verify amphoteric oxides

Procedure

– A small sample of aluminium oxide is placed in a test tube and 5 cm3 of 2M nitric acid added to it and the mixture shaken.

– The procedure is repeated in different test tubes with ZnO, PbO, CuO and CaO.

– The experiments are repeated using excess 2M sodium hydroxide in place of nitric acid

 

Observations

 

Name of solidObservations when
acid is addedhydroxide is added
Aluminium oxide

Zinc oxide

Lead II oxide

Zinc hydroxide

Lead hydroxide

Aluminium hydroxide

Oxide dissolves

Oxide dissolves

‘’

Hydroxide dissolves

‘’

Oxide dissolves

‘’

‘’

Hydroxide dissolves

‘’

Explanations

– These oxides are soluble in acids as well as in the alkalis (NaOH)

  • Reaction with acids

– Oxides react with acids to form a salt and water only in a reaction called neutralization reaction.

 

Equations

  • Oxides

(i).PbO(s) + 2H+(aq)                   Pb2+(aq) + H2O(l)

 

(ii). Al2O3(s) + 6H+(aq)                  2Al3+(aq) + 3H2O(l)

 

(iii). ZnO(s) + 2HCl(aq)                 ZnCl2(aq) + H2O(l)

 

  • Hydroxides:

(i). Pb(OH)2(s) + 2HCl(aq)                        PbCl2(aq) + 2H2O(l)

 

(ii). Zn(OH)2(s) + 2HCl(aq)                    ZnCl2(aq) + 2H2O(l)

 

(iii). Al(OH)3(s) + 3HCl(aq)                      AlCl3(aq) + H2O(l)

 

Note: in these reactions the metal oxides are reacting as bases

 

  • Reaction with alkalis

– These oxides and hydroxides also react with alkalis e.g sodium hydroxide in which case they are reacting as acids.

– Their reactions with alkalis involve the formation of complex ions; M(OH)2-4

 

Equations

  • Oxides

(i). PbO(s) + 2NaOH(aq) + H2O(l)                   Na2Pb(OH)4(aq)

 

Ionically: PbO(s) + 2OH(aq) + H2O(l)                 [Pb(OH)4]2-(aq)

 

(ii). Al2O3(s) + 2OH(aq) +  3H2O (l)                   2[Al(OH)4](aq) + 3H2O(l)

 

Ionically: Al2O3(s) + 2OH(aq) + 3H2O(l)                       2[Al(OH)4](aq)

 

(iii). ZnO(s) + 2NaOH(aq) + H2O(l)                     Na2Zn(OH)4(aq)

 

Ionically: ZnO(s) + 2OH(aq) + H2O(l)                [Zn(OH)4]2-(aq)

  • Hydroxides:

(i). Al(OH)3(s) + NaOH(aq)                             NaAl(OH)4(aq)

 

Ionically: Al(OH)3(s) + OH(aq)                     [Al(OH)4](aq)

 

(ii). Zn(OH)2(s) + 2NaOH(aq)                            Na2Zn(OH)4(aq)

 

Ionically: Zn(OH)2(s) + 2OH(aq)                      [Zn(OH)4]2-(aq)

 

(iii). Pb(OH)2(s) + 2NaOH(aq)                            Na2Pb(OH)4(aq)

 

Ionically: Pb(OH)2(s) + 2OH(aq)                      [Pb(OH)4]2-(aq)

 

 

Salts

– Is a compound formed when cations derived from a base combine with anions derived from an acid.

– Salts are usually formed when an acid reacts with a base i.e. when the hydrogen ions in an acid re wholly or partially by a metal ion or ammonium (NH4+) radical.

 

Laboratory preparations of salts

– Salts are prepared in the laboratory using various depending on property of the salt especially solubility

Examples

(a). Preparations by direct synthesis

Equation:

Fe(s) + Cl2(g)                        2FeCl3(s)

 

(b). Reactions of acids with metals, metal oxides, metal hydroxides and metal carbonate

Equations:

Zn(s) + H2SO4(aq)                         ZnSO4(aq) + H2(g)

 

CuO(s) + H2SO4(aq)                      CuSO4(aq) + H2O(l)

 

NaOH(aq) + HCl(aq)                      NaCl(aq) + H2O(l)

 

PbCO3(s) + 2NHO3(aq)                 Pb(NO3)2(aq) + CO2(g) + H2O(l)

 

Note:

– Acid + metal method will not be suitable if:

  • The metal is too reactive e.g. sodium or potassium.
  • The salt formed is insoluble; as it will form an insoluble layer on the metal surface preventing further reaction.
  • The metal is below hydrogen in the reactivity series.

 

(c). Double decomposition/ precipitation

– Mainly for preparations of insoluble salts

– Involves formation (precipitation) of insoluble salts by the reaction between two solutions of soluble salts.

 

Equations:

Pb(NO3)2(aq) + 2NaCl(aq)                      PbCl2(s) + 2NaNO3(aq)

 

AgNO3(aq) + HCl(aq)                   AgCl(s) + HNO3(aq)

 

Types of salts:

– Are categorized into three main categories:

  • Normal salts
  • Acid salts
  • Double salts

Solubility of salts: – a summary

– All common salts of sodium, potassium and ammonium are soluble.

– All common nitrates are soluble.

– All chlorides are soluble except silver, mercury and lead chlorides.

– All sulphates are soluble except calcium, barium, lead and stomium sulphates.

– All carbonates are insoluble except sodium, potassium and ammonium carbonates.

– All hydroxides are insoluble except sodium, potassium ammonium and calcium hydroxides is        sparingly soluble.

Note:

– Lead(II) chloride is soluble in hot water.

– Calcium hydroxide is sparingly soluble in water.

 

Reactions of some cations with NaOH(aq)  and NH4OH(aq) and solubilities of some salts in water

CationSoluble compounds in waterInsoluble compounds in waterReaction with NaOH(aq)Reaction with NH4OH solution
K+allNoneNo reactionNo reaction
Na+allNoneNo reactionNo reaction
Ca2+Cl; NO3CO32-; O2-; SO42-; OH;White precipitate insoluble in excessNo precipitate
Al3+Cl; NO3; SO42-;CO32-; O2-; OH;White precipitate

soluble in excess

White precipitate insoluble in excess
Pb2+NO3; ethanoate;All othersWhite precipitate soluble in excessWhite precipitate insoluble in excess
Zn2+Cl; SO42-; NO3;CO32-; OH;White precipitate soluble in excessWhite precipitate soluble in excess
Mg2+Cl; SO42-; NO3;CO32-; OH;White precipitate insoluble in excessNo precipitate
Fe2+Cl; SO42-; NO3;CO32-; O2-; OH;(dark) green precipitate insoluble in excessGreen precipitate insoluble in water
Fe3+Cl; SO42-; NO3;CO32-; O2-; OH;(red) brown precipitate insoluble in excessBrown precipitate insoluble in excess
Cu2+Cl; SO42-; NO3;CO32-; O2-; OH;Pale blue precipitate insoluble in excessPale blue precipitate soluble in excess forming a deep blue solution
NH4+allnoneAmmonium gas on warmingNot applicable

Explanations

– In these experiments NaOH forms insoluble hydroxides with ions of Zn2+, Al3+, Cu2+, Fe2+, Ca2+, Mg2+, Fe3+, and Pb2+.

– These hydroxides have a characteristic appearance, which form the basis of their identification

Examples

Equations:

Zn2+(aq) + 2OH(aq)                         Zn(OH)2(s)

(White ppt).

 

Cu2+(aq) + 2OH(aq)                         Cu(OH)2(s)

(Pale blue ppt).

 

Fe2+(aq) + 2OH(aq)                          Fe(OH)2(s)

(Dirty green ppt).

 

Fe3+(aq) + 2OH(aq)                          Fe(OH)3(s)

(Red-brown ppt).

 

Pb2+(aq) + 2OH(aq)                          Pb(OH)2(s)

(White ppt).

 

– The hydroxides of aluminum, zinc and lead dissolves in excess sodium hydroxide solution because of complexes are formed

Equations:

Al(OH)3(s) + OH(aq)                      [Al(OH)4](aq)

(Tetra-hydroxyl-aluminium (III) ion)

 

Pb(OH)2(s) + 2OH(aq)                    [Pb(OH)4]2-(aq)

(Tetra-hydroxyl-lead (II) ion)

 

Zn(OH)2(s) + 2OH(aq)                        [Zn(OH)4]2-(aq)

(Tetra-hydroxyl-zinc (II) ion)

 

Note: – in these reactions KOH(aq) may be used instead of sodium hydroxide

 

With ammonia solution

– Insoluble metals hydroxides are similarly formed.

Zn2+(aq) + 2OH(aq)                         Zn(OH)2(s)

(White ppt).

 

Cu2+(aq) + 2OH(aq)                         Cu(OH)2(s)

(Pale blue ppt).

 

Fe2+(aq) + 2OH(aq)                          Fe(OH)2(s)

(Dirty green ppt).

 

Fe3+(aq) + 2OH(aq)                          Fe(OH)3(s)

(Red-brown ppt).

 

Pb2+(aq) + 2OH(aq)                          Pb(OH)2(s)

(White ppt).

 

– However hydroxides of copper and zinc dissolve in excess ammonia solution due to formation of complex ions/ salts

Equations:

Zn(OH)2 (s) + 4NH3(aq)                      [Zn(NH3)4]2+(aq)  + 2OH(aq)

(White ppt)                                            (Tetra-amine zinc (II) ion; colourless solution)

 

Cu(OH)2(s) + 4NH3(aq)                       [Cu(NH3)4]2+(aq)  + 2OH(aq)

(Pale blue ppt)                                          (Tetra-amine copper (II) ion; deep blue solution)

Effects of heat on metal hydroxides

Procedure

– Hydroxides of Zn, Ca, Pb, Cu e.t.c are strongly heated in a test tube each separately

 

Observation

– Most metal hydroxides are decomposed by heat to form metal oxides and water

– Sodium and potassium hydroxides only decompose at very high temperatures.

– Hydroxides of metals lower in the reactivity series are readily decomposed by heat than those metals higher in the series.

 

Examples

Cu(OH)2(s)            Heat                CuO(s) + H2O(l)

(Blue)                                                      (Black)

 

Pb(OH)2(s)             Heat                PbO(s) + H2O(l)

(White)                                                    (Red brown when hot; yellow when cold)

 

Zn(OH)2(s)            Heat                ZnO(s) + H2O(l)

(White)                                                    (Yellow when hot; white when cold)

 

Ca(OH)2(s)            Heat                CuO(s) + H2O(l)

(White)                                                    (White)

 

Note:

– Both iron (II) and iron (III) hydroxides give iron (III) oxide when heated.

Equations:

2Fe(OH)2(s)  + ½ O2(g)          Heat                Fe2O3(s) + 2H2O(l)

(Green)                                                                             (Red-brown)

 

Fe(OH)3(s)             Heat                Fe2O3(s) + 3H2O(l)

(Brown)                                                   (Red-brown)

 

These oxides do not decompose on further heating

 

Effects of sodium carbonate on various salt solutions

Procedure

– 3 drops of NaOH(aq) are added to 2cm3 of 1M solution containing magnesium ions in a test tube

the procedure is repeated with salt solutions containing

 

Solution containingObservations after adding sodium carbonate
Mg2+A white precipitate is formed
Ca2+A white precipitate
Zn2+A white precipitate
Cu2+A green precipitate
Pb2+A white precipitate
Fe2+A green precipitate
Fe3+A brown precipitate and a colourless gas that forms a white ppt. in lime water;
Al3+A white precipitate and a colourless gas that forms a white ppt. in lime water;

 

Explanations:

– Sodium carbonate, potassium and ammonium carbonate are soluble in water; all other metal carbonates are insoluble

– Hence their solutions may be used to precipitate the insoluble metal carbonates.

Ionic equations:

Ca2+(aq) + CO32-(aq)                          CaCO3(s)

 

Note:

Iron (III) and Aluminium salts hydrolyse in water giving acidic solutions which react with carbonates to liberate carbon dioxide gas; hence effervescence.

 

Reaction of metal ions in salt solutions with sodium chloride, sodium sulphate and sodium sulphate

(i). Procedure

– 2cm3 of a 0.1M solution containing lead ions is placed in a test tube.

– 2-3 drops of 2M sodium chloride solution are added and the mixture warmed;

– The procedure is repeated using salt solutions containing Ba2+; Mg2+; Ca2+; Zn2+; Cu2+; Fe2+ and Fe3+

– Each experiment (for each salt) is repeated using Na2SO4 and Na2SO3 respectively, in place of sodium chloride.

 

(ii). Observations

 

Solution containingSodium sulphateSodium chlorideSodium sulphate
Zn2+– Colourless solution– Colourless solutionColourless solution
Mg2+– Colourless solution– Colourless solutionColourless solution
Cu2+– Blue solution– Blue solutionBlue solution
Fe2+– Greenish solution– Green solution
Fe3+– Yellow solution– Yellow/ dark brown solution
Pb2+– White precipitate– White precipitate which dissolve on warmingWhite precipitate
Ba2+– White precipitate– White precipitateWhite precipitate

 

Explanations

– All the listed cations soluble salts except Ba2+ and Pb2+

– Lead sulphate and barium sulphate are insoluble in water;

– Lead chloride and barium sulphite are insoluble; however PbCl2(s) dissolves on warming

 

Equations:

Pb2+(aq) + 2Cl(aq)                       PbCl2(s)

 

Pb2+(aq) + SO42-(aq)                     PbSO4(s)

 

Ba2+(aq) + SO42-(aq)                     BaSO4(s)

 

Ba2+(aq) + SO32-(aq)                     BaSO3(s)

 

 

 

Note:

– To distinguish the precipitate of barium sulphate from barium sulphite; dilute HNO3 (aq) or HCl(aq) is added to both;

– BaSO3(s) will dissolve in the dilute acid but barium sulphate will not.

 

Uses (importance) of precipitation reactions.

– Precipitation of metal carbonate from aqueous solutions is useful in softening hard water; usually by removing calcium and magnesium ions from water as insoluble carbonate

 

Useful information on salts (qualitative analysis)

Colours of substances in solids and solutions in water.

 

COLOUR 
SOLIDAQUESOUS SOLUTION

(IF SOLUBLE)

1. WhiteColourlessCompound of K+; Na+, Ca2+; Mg2+; Al3+; Zn2+; Pb2+; NH4+
2. YellowInsolubleZinc oxide, ZnO (turns white on cooling); Lead oxide, PbO (remains yellow on cooling, red when hot)
YellowPotassium or sodium chromate;
3. BlueBlueCopper (II) compound, Cu2+
4. Pale green

 

Green

Pale green (almost colourless)

Green

Iron (II) compounds,Fe2+

 

Nickel (II) compound, Ni2+; Chromium (II) compounds, Cr3+; (Sometimes copper (II) compound, Cu2+)

5. BrownBrown (sometimes yellow)

 

Insoluble

Iron (III) compounds, Fe3+;

 

Lead (IV) oxide, PbO2

6. PinkPink (almost colourless)

Insoluble

Manganese (II) compounds, Mn2+;

Copper metal as element (sometimes brown but will turn black on heating in air)

7. OrangeInsolubleRed lead, Pb3O4 (could also be mercury (II) oxide, HgO)
8. BlackPurple

Brown

Insoluble

Manganate (VII) ions (MnO) as in KMnO4;

Iodine (element)-purple vapour

Manganese (IV) oxide, MnO2

Copper (II) oxide, CuO

Carbon powder (element)

Various metal powders (elements)

 

Reactions of cations with common laboratory reagents and solubilities of some salts in water

 

CATIONSOLUBLE COMPOUNDS (IN WATER)INSUOLUBLE COMPOUNDS (IN WATER)REACTION WITH AQUEOUS SODIUM HYDROXIDEREACTION WITH AQUEOUS AMMONIA SOLUTION
Na+AllNoneNo reactionNo reaction
K+AllNoneNo reactionNo reaction
Ca2+Cl; NO3;CO32-; O2-; SO42-; OH;White precipitate insoluble in excessWhite precipitate insoluble in excess, on standing;
Al3+Cl; NO3; SO42-O2-; OH;White precipitate soluble in excessWhite precipitate insoluble in excess
Pb2+NO3; ethanoate;All others;White precipitate soluble in excessWhite precipitate insoluble in excess
Zn2+Cl; NO3; SO42-CO32-; O2-; SO42-; OH;White precipitate soluble in excessWhite precipitate soluble in excess
Fe2+Cl; NO3; SO42-CO32-; O2-; OH;(Dark) pale green precipitate insoluble in excess(Dark) pale green precipitate insoluble in excess
Fe3+Cl; NO3; SO42-CO32-; O2-; OH;(Red) brown precipitate insoluble in excess(Red) brown precipitate insoluble in excess
Cu2+Cl; NO3; SO42-CO32-; O2-; OH;Pale blue precipitate insoluble in excessPale blue precipitate soluble in excess forming a deep blue solution
NH4+AllNone;Ammonias gas on warmingNot applicable.

 

 

 

Qualitative analysis for common anions.

 

 SO42-(aq)Cl(aq)NO3(aq)CO32-(aq)
TESTAdd Ba2+(aq) ions from Ba(NO3)2(aq); acidify with dilute HNO3(aq)Add Ag+(aq) from AgNO3(aq).

Acidify with dilute HNO3

Alternatively;

Add Pb2+ from Pb(NO3)2 and warm

Add FeSO4(aq);

Tilt the tube and carefully add 1-2 cm3 of concentrated H2SO4(aq)

Add dilute HNO3(aq); bubble gas through lime water;
OBSERVATIONThe formation of a white precipitate shows presence of SO42- ion;The formation of a white precipitate shows presence of Cl ion;

Formation of a white precipitate that dissolves on warming shown presence of Cl(aq) ions

The formation of a brown ring shows the presence of NO3 ionsEvolution of a colourless gas that forma a white precipitate with lime water, turns moist blue litmus paper red; and extinguishes a glowing splint shows presence of CO32- ions
EXPLANATIONOnly BaSO4 and BaCO3 can be formed as white precipitates.

BaCO3 is soluble in dilute acids and so BaSO4 will remain on adding dilute nitric acid

Only AgCl and AgCO3 can be formed as white precipitates.

AgCO3 is soluble in dilute acids but AgCl is not;

– PbCl2 is the only white precipitate that dissolves on warming

Concentrated H2SO4 forms nitrogen (II) oxide with NO3(aq) and this forms brown ring complex (FeSO4.NO) with FeSO4;All CO32- or HCO3 will liberate carbon (IV) oxide with dilute acids

 

Checklist:

  1. Why is it not possible to use dilute sulphuric acid in the test for SO42- ions;
  2. Why is it not possible to use dilute hydrochloric acid in the test for chloride ions?
  3. Why is it best to use dilute nitric acid instead of the other two mineral acids in the test for CO32- ions?
  4. How would you distinguish two white solids, Na2CO3 and NaHCO3?

 

What to look for when a substance is heated.

 

1. SublimationWhite solids on cool, parts of a test tube indicates NH4+ compounds;

Purple vapour condensing to black solid indicates iodine crystals;

2. Water vapour (condensed)Colourless droplets on cool parts of the test tube indicate water of crystallization or HCO3 (see below)
3. Carbon (IV) oxideCO32- of Zn2+; Pb2+; Fe2+; Fe3+; Cu2+;
4. Carbon (IV) oxide and water vapour (condensed)HCO3
5. Nitrogen (IV) oxideNO3of Cu2+; Al3+; Zn2+; Pb2+; Fe2+; Fe3+
6. OxygenNO3 or BaO2; MnO2; PbO2;

 

 

 

 

 

 

Reduction-oxidation (Redox reactions)

(a). Displacement reactions.

(i) More reactive halogens metals will displace less reactive metals from solutions of their salts in the series:

Zn                               Fe  Pb Cu

More reactive                                Less reactive

 

Example:

– Zinc powder placed in a solution of copper (II) sulphate, which contains Cu2+(aq) ions, will become Zn2+(aq) ions and brown copper solid (metal) will be deposited.

– The Cu2+(aq) is reduced to copper by addition of electrons and the zinc is oxidized to Zn2+(aq) by removal of electrons.

 

(ii). More reactive halogens will displace less reactive halogens from solutions of their salts in series:

Cl2                             Br2 I2

More reactive                              Less reactive.

 

Example:

– Chlorine bubbled into a solution of potassium iodide (colourless), which contains I(aq) ions will turn grey (black) as iodine is liberated.

– The chlorine is reduced to Cl(aq) ions by addition of electrons and the I(aq) ions are oxidized to iodine by removal of electrons.

 

(b). Decolourisation of purple potassium manganate (VII) ions.

When a few drops of purple KMnO4 solution are added to a compound and the purple colour disappears, then this shows that the MnO4(aq) ions have been reduced to almost colourlessMn2+ (aq) ions

The substance in the solution has been oxidized.

 

Example:

– KMnO4 will oxidize Fe2+(aq) ions to Fe3+(aq) ions; pale green solution turns red-brown.

– KMnO4 will oxidize Cl(aq) ions to Cl2(g); colourless solution results to a green gas with a bleaching action;

 

(c). Orange potassium chromate (VI) turning to a green solution.

– Orange solution of dichromate ions, Cr2O72-(aq), changes to green Cr3+(aq) ions when the dichromate is reduced.

– The substance causing this change is oxidized.

 

Example:

– K2Cr2O7 will oxidize Fe2+(aq) to Fe3+(aq)

– K2Cr2O7 will oxidize SO2(g) to SO42-(aq)

– Formation of sulphate ions in solution from sulphur (IV) oxide gas is often used in the test for sulphur (IV) oxide gas.

 

(d). Oxidation of Fe2+(aq) to Fe3+(aq) ions by concentrated nitric acid.

 

 

 

 

Solubility and solubility curves

  • Solubility

– Is the maximum number of grams of a solid which will dissolve in 100g of solvent (usually water) at a particular temperature

– A solution is made up of two parts: – a solute and a solvent.

 

Solute

– The solid part of a solution usually dispersed in the solvent e.g. a salt.

 

Solvent

– The liquid part of the solution into which the solute is dissolved.

 

Experiment: to determine the solubility of potassium nitrate at 20oC.

(i). Materials

– Beakers, evaporating dish, measuring cylinder, burner, scales, thermometer, distilled water and potassium nitrate

 

(ii). Apparatus

 

 

(iii). Procedure

– About 50cm3 of distilled water is placed in a beaker

– Potassium nitrate is added to it a little at a time stirring continuously.

– The nitrate is added until no more will dissolve and there is an excess undissolved salt present. This is the saturated solution of KNO3 at the temperature.

Note:

  • Saturated solution: solution that cannot dissolve any more of the solid/ solute at a particular temperature

– The solution is allowed to settle and it is temperature recorded.

– About 25 cm3 of clear solution is poured in a previously weight evaporating dish.

– The mass of the dish and solution is recorded.

– The dish is then heated in a water bath (to avoid spurting) till the solution is concentrated.

– The concentrated solution is allowed to cool and the dish weighted with its contents.

 

Results and calculations

 

Temperature20.0oC
Mass of evaporating dish + solution100.7g
Mass of evaporating dish65.3g
Mass of solution35.4g
Mass of evaporating dish + dry salt73.8g

Calculating:

Mass of salt dissolved = (73.8 – 65.3)g = 8.5g;

Mass of water (solvent) = (35.4 – 8.5)g = 26.9g

Thus:

If 26.9g of water dissolves 8.5g of KNO3 at 20oC;

Then 100g of water will have ? = 100 x 8.5 = 31.6g of salt;

26.9

Therefore the solubility of KNO3 at 20oC = 31.6g per 100g of water

 

Factors affecting solubility.

(i). Temperature

– For most salts solubility increases with rise/ increase in temperature.

Reason

– Increased temperature increases the kinetic energy, and hence the momentum and velocity of the solvent molecules so that they can disintergrate the solute molecules more effectively.

– However solubilities of certain salts remain almost constant with temperature change

– Solubility of gases however decreases with increase in temperature;

Reason:

Increase in temperature causes the gas molecules to expand and hence escape from the solvent.

 

Experiment: To investigate the effect the effect of temperature on solubility.

Requirements: potassium nitrate, distilled water, test tube, thermometer, stirrer, bunsen burner, 250 cm3 glass beaker, 4.5g of potassium nitrate.

 

(ii). Procedure.

Using a 10ml measuring cylinder, measure 5 cm3 of distilled water and add it to the boiling tube containing solid potassium nitrate.  Insert a thermometer into the boiling tube and heat the mixture gently in a water bath or while shaking to avoid spillage. Continue heating until all the solid has dissolved.  Stop heating and allow the solution to cool while gently stirring with a thermometer. Record the temperature at which the crystals of potassium nitrate first appear. Note this in the table below.

Retain the boiling tube and its contents for further experiments.

Measure 2 cm3 of distilled water and add to the mixture in the boiling tube. Heat until the crystals dissolve, then cool while stirring with a thermometer. Record the temperature at which the crystals again first stat to reappear. Repeat this procedure, each time adding more 2 cm3 of distilled water, heating, cooling and recording the crystallization temperature until the table is completely filled.

 

Table 2:

 

Experiment numberIIIIIIIVV
Volume of water added5791113
Temperature at which crystals appear (oC)
Solubility of K in g/100g of water

 

 

 

 

Questions:

(a). Complete the table and calculate the solubility of solid X in g/100g of water at different temperatures.                                                                                                                                        (2 marks)

 

(b). Using the table above, plot a graph of solubility of solid X in g/100g of water against temperature.                                                                                                                                               (5 marks)

(c). From the graph:

(i). calculate the mass of K that would be obtained if the saturated solution is cooled from 60oC to 40oC.                                                                                                                                           (2 marks)

 

(ii). determine the solubility at 70oC.                                                                                              (1mark)

 

(iii). at what temperature would solubility of K be 100g/100g of water?                                      (1mark)

 

(ii). Stirring

– Stirring increases the solubility of a solid

Reason

– Stirring causes the molecules of solvent and solute to move faster causing the solute particles to disintergrate more effectively

 

Solubility curves

– Are curves showing the variation of solubility with temperature.

 

Uses / importance of solubility curves

– Can be used to determine the mass of crystals that would be obtained by cooling a volume of hot saturated solution from one known temperature to another.

– Solubility differences can be used to separate substances i.e. recrystallization or fractional crystallization (refer to separation of mixtures)

– Separation of salts from a mixture of salts with differing solubilities e.g. extraction of sodium carbonate from Trona (refer to carbon and its compounds)

– Manufacture of certain salts e.g. sodium carbonate by the Solvay process (refer to carbon and its compounds)

 

Worked examples

  1. An experiment was carried out to determine the solubility of potassium nitrate and the following results were obtained.

 

Temperature101530405060
Mass of KNO3 per 100g of water2025456385106

 

(a). What is meant by solubility?                                                                                                   (1 mark)

 

(b). Plot a graph of mass of potassium nitrate against temperature.                                            (3 marks)

 

(c). From the graph work out the mass of KNO3 that would crystallize if a solution containing 70g of KNO3 per 100g of water was cooled from 45oC to 25oC.                                                                 (2 marks)

 

(d). Explain what would happen if 100g of KNO3 was put in cold water and heated to 50oC.                                                                                                                                                              (2 marks)

  1. The table below shows the solubility of sulphur (IV) oxide at various temperatures.

 

Temperature (oC)0510152025354045505560
Mass of SO2 per 100g of water2218.415.413.010.89.057.806.805.574.804.203.60

 

(a). On the grid provided plot a graph of solubility against temperature.                                    (3 marks)

 

(b). From the graph determine:

(i). The lowest temperature at which 100cm3 of water would contain 11.6g of sulphur dioxide.(1 mark)

 

(ii). The maximum mass of sulphur (IV) oxide that would dissolve in 2 litres of solution at 10oC. (Assume that the density of the solution is 1gcm-3)                                                                          (3 marks)

 

(c) (i). Sulphur (IV) oxide reacts with sodium hydroxide solution to form sodium sulphite and water.                                                                                                                                                 (1 mark)

(ii). Write the equation for this reaction.                                                                                        (1 mark)

 

(iii). Using the information from the graph, determine the volume of the saturated sulphur (IV) oxide solution that can neutralize 153 cm3 of 2M sodium hydroxide solution at 25oC.                     (3 marks)

Water

– Can be pure or impure

Pure water

– Is a pure substance which is a compound of hydrogen and oxygen; that boils at 100oC; melts at 0oC and has a density of 1gcm-3 at sea level.

 

Impure water

– Are the natural waters constituted of dissolved solutes in pure water.

 

Hardness of water

– Water without dissolved substances (salts) hence lathers easily with soap is referred to as soft water while water with dissolved substances that does not lather easily with soap is termed as hard water

 

Experiment: effect of water containing dissolved salts on soap solution

Procedure

– 2 cm3 of distilled water is put in a conical flask.

– Soap solution from a burette is added into the water and shaken until formation of lather is noted.

– If the soap fails to lather more soap solution is added from the burette till it lathers and the volume of the soap required for lathering recorded.

– The procedure is repeated with each of the following: tap water, rain water, dilute solutions of MgCl2, NaCl, CaCl2, a(NO3)2, CaSO4, MgSO4, Mg(HCO3)2, Ca(HCO3)2, ZnSO4 , NaHCO3, and KNO3.

– The procedure is repeated with each of the solutions when boiled.

 

Observations

 

Explanations

– Distilled water requires very little soap to produce lather because it lacks dissolved salts and hence termed soft water.

– Solutions containing NaCl, ZnSO4, KNO3 and NaHCO3 do not require a lot of soap to form lather

Water containing Ca2+ and Mg2+ ions do not lather easily (readily) with soap

Reason:

– These ions react with soap (sodium stearate) to form an insoluble salt (metal stearate) called (Mg and Ca stearate respectively); which is generally termed scum.

 

Equations:

With Ca2+

2C17H35COONa+(aq) + Ca2+(aq)                                  (C17H35COO)2Ca(s) + 2Na+(aq)

Sodium stearate                                                                                             calcium stearate

 

With Mg2+

2C17H35COONa+(aq) + Mg2+(aq)                                 (C17H35COO)2Mg(s) + 2Na+(aq)

Sodium stearate                                                                                             Magnesium stearate

 

– Thus water with Mg and Ca is termed hard water and can only be made soft by removing these ions upon which the water will lather easily with water

– When Ca(HCO3)2(aq) and Mg(HCO3)2(aq) are boiled the amount of soap required for lathering decreases than before boiling

 

Reason

– Boiled decomposes the 2 salts into their respective carbonate s which precipitates from the solution leaving soft water which leathers easily with water

 

– The amount of soap solution used with  solutions containing sulphates and chlorides of calcium and magnesium did not change significantly even after boiling

Reason

– The soluble sulphates and chlorides of Mg and Ca do not decompose upon boiling hence can not be precipitated out.

 

Types of water hard ness

Temporary hardness

– Is hardness due to the presence of CaHCO3 or Mg(HCO3)2 in water; and can usually be removed by boiling.

 

Removal of temporary hardness in water:

(i). Boiling:

– Boiling decomposes and an insoluble chalk of CaCO3 and MgCO3 respectively is deposited in the sides of the vessel.

– This forms an encrustation commonly known as furr the process being furring.

 

Equations:

Ca(HCO3)2(s)           Heat        CaCO3(s) + 2CO2(g) + H2O(l)

 

Mg(HCO3)2(s)           Heat        MgCO3(s) + 2CO2(g) + H2O(l)

 

(ii). Distillation:

– Water containing dissolved salts is heated in a distillation apparatus;

– Pure water distils over first leaving dissolved salts in the distillation flask (refer to separation of mixtures)

– Is of less economic value as it is too expensive hence disadvantageous.

 

(iii). Addition of calcium hydroxide:

– Involves adding correct amount of lime water where CaCO3 is precipitated out.

– This method is cheap and can be used on large scale at water treatment plants.

– However if excess lime (Ca2+) ions is added this will make water hard again.

 

Equation:

Ca(HCO3)2(aq) + Ca(OH)2(aq)                            2CaCO3(s) + 2H2O(l).

 

(iv). Addition of ammonia solution:

– Addition of aqueous ammonia to water containing calcium and magnesium hydrogen carbonates (temporary hard) precipitates calcium and magnesium ions as corresponding carbonates.

 

Equations:

Ca(HCO3)2(aq) + 2NH4OH(aq)                          CaCO3(s) +2H2O(l) + (NH4)2CO3(aq)

 

 

(ii). By permutit softener (ion exchange).

– Uses a complex sodium salt (NaX), such as sodium aluminium silicate commonly known as sodium permutit.

– Permutit is a manufactured ion exchange resin.

 

  • Iron exchange resin: materials that will take ions of one element out of it’s compounds and replace it with ions another element

 

 

 

 

Working principle

– The permutit is contained in a metal cylinder

– The hard water is passed through the column of permutit in the cylinder and it emerges softened at the other end

– As hard water passes through the column ion exchange takes place.

– The Ca2+ and Mg2+ remain in the column while sodium ions from the permutit pass into water thus softening it.

 

Diagram: permutit water softener.

 

 

Equations:

NaX(aq) + Ca2+(aq)                        CaX(s) + 2Na+(aq)

 

NaX(aq) + Mg2+(aq)                       MgX(s) + 2Na+(aq)

 

– When all the Na+ ions in the permutit have been replaced by Ca2+ and Mg2+ ions the permutit can not go on softening water.

– It is then regenerated by washing the column with brine (a strong NaCl solution); during which calcium and magnesium chlorides are washed away.

 

Equation:

CaX(s) + 2NaCl(aq)                        CaCl2(aq) + Na2X(s)

 

MgX(s) + 2NaCl(aq)                       MgCl2(aq) + Na2X(s)

 

Permanent hardness

– Is that due to soluble sulphates and or chlorides of calcium and or magnesium and cannot be removed by boiling

 

Removal of permanent hardness

(i). By the addition of washing soda (sodium carbonate)

– Washing soda softens hard water by causing the formation of insoluble CaCO3 or MgCO3

– The soluble sodium salts left in water do not react with soap.

 

Equations:

Na2CO3(aq) + CaCl2(aq)                 2NaCl(s) + CaCO3(S)

 

Ionically:

Ca2+(aq) + CO32-(aq)                    CaCO3(s)

 

Na2CO3(aq) + MgSO4(aq)                          Na2SO4(s) + CaCO3(S)

 

Ionically:

Mg2+(aq) + CO32-(aq)                   MgCO3(s)

– This method is very convenient and economical on large scale. It softens both temporary and permanent hardness

 

(ii). By permutit softener (ion exchange); explanations as before

(iii). Distillation.

 

Advantages of hard water

(i). It is good for drinking purposes as calcium ions contained in it helps to form strong bones and teeth.

 

(ii). When soft water flows in lead pipes some lead is dissolved hence lead poisoning. However when lead dissolves in hard water insoluble PbCO3 are formed, coating the inside of the lead pipes preventing any further reaction; this reduces any chances of lad poisoning.

 

(iii). It is good for brewing and the tanning industries; it improves wine or beer flavour in brewing industries.

 

Disadvantages of hard water

(i). Soap forms insoluble salts with magnesium and calcium ions; scum (calcium or magnesium stearate) thereby wasting soap.

– For these reason soapless detergents are preferred to ordinary soaps because they do not form scum; but rather form soluble salts with Mg2+ and Ca2+

  • Examples of soapless detergents: omo, perfix, persil, fab e.t.c.

 

(ii). Deposition of insoluble magnesium and calcium carbonates and sulphates formed from hard water result into blockage of water pips due to the formation of boiler scales

(iii). Formation of kettle fur which makes electrical appliances inefficient hence increasing running costs.

 

(iv). Formation of scum on clothing reduces their durability and aesthetic appearance

 

F4 SALTS

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KEVIN OWINO

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