ACIDS, BASES AND SALTS - Chemistry Notes Form 4

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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; 
    1. 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; water
        HCl(aq)                   →              H+(aq) + Cl-(aq)
        Sulphuric acid;      water
        H2SO4(aq)                 →              2H+(aq) + SO42-(aq)
        Nitric acid;            water
        HNO3(aq)               →               H+(aq) + NO3-(aq)
    2. 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:   water
        H2CO3(aq)             →         H+(aq) + HCO3-(aq)
        Ethanoic acid     water
        CH3COOH (aq)     →       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

  1. 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.
  2. 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
      electrolyic 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
  3. 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 Indicator  pH

     Sulphuric acid 

     Hydrochloric acid 

     Nitric acid 

     Ethanoic acid 

     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
    acidic properties water
    acidic properties methyl
  • The resultant solutions are each separately subjected to various tests as shown below and observations recorded

Tests and observations

  Test  Aqueous HCl solution  Solution of HCl in methylbezene
 1. A piece of dry litmus paper is dropped  Blue litmus turns red  No effect on litmus
 2. Dry universal indicator paper  Turns red (strong acid)  Turns green (neutral)
 3. Add magnesium ribbon  Evolution of hydrogen  No reaction
 4. Add small marble chips  COevolved  No reaction
 5. Electrical conductivity  Good conductor  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:  water
    HCl(aq)         →     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 are polar molecules;
  • This causes mutual attraction of both ends of HCl molecule by different water molecules causing the dissociation of HCl molecules into ions.
  • 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)    → Mg2+(aq) + H2(g)
  • Methylbenzene has a weak attraction for hydrogen chloride and hence hydrogen chloride remains as molecules in methylbenzene


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) → Cu2+(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.
    1. Strong alkalis
      • Are alkalis that undergo complete dissociation in aqueous solution; yielding a large number of hydroxyl (OH-) ions
        Examples:
        - Sodium hydroxide.
        - Potassium hydroxide.
    2. 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

  1. Using electrical conductivity
    Procedure
    • 50 cm3 of 2 M sodium hydroxide solution is put into a beaker and the apparatus set as shown below
      sodium hydroxide conductivity1

    • 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)
  2. 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
     Alkali  Colour of Universal Indicator  pH

     Ammonium hydroxide (1M)

     Calcium hydroxide (1M)

     Sodium hydroxide (0.1M)

     Sodium hydroxide (1M)

     Potassium hydroxide (1M)

     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.
      a pH scale

      Indicator colours:
       pH  2  3  4  5  6  7  8  9  10  11  12  13  14
       Color  Red       Orange/red Yellow/green    Green   Green/Blue  Blue/purple  Purple

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
    basic properties water
    basic properties methy

Observations

 Test  Solution of NH3 in water  Solution of NH3 in methylbenzene (Toluene)
 Dry litmus paper  Red litmus paper turns blue  No effect
 Dry universal indicator paper  Colour turns purple (alkaline pH)  Turns green (Neutral pH)
 Electrical conductivity  Poor conductor  Non - 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.
  2. 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 solid     Observations when
 acid added  hydroxide added

 Aluminium oxide

 Zinc oxide

 Lead II oxide

 Zinc hydroxide

 Lead hydroxide

 Aluminium hyrdroxide

 Oxide dissolves

 Oxide dissolves

 Oxide dissolves

 Hyrdoxide dissolves

 Hydroxide dissolves

 Hydroxide dissolves

 Oxide dissolves

 Oxide dissolves

 Oxide dissolves

 Hydroxide dissolves

 Hydroxide dissolves

 Hydroxide dissolves

Explanations

  • These oxides and hydroxides are soluble in acids as well as in the alkalis (NaOH)

Reaction with Acids

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

    Oxides

    PbO(s) + 2H+(aq) → Pb2+(aq) + H2O(l)

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

    ZnO(s) + 2HCl(aq) → ZnCl2(aq) + H2O(l)

    Hydroxides

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

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

    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

     PbO(s) + 2NaOH(aq) + H2O(l) → Na2Pb(OH)4(aq)
    Ionically: PbO(s) + 2OH-(aq) + H2O(l) → [Pb(OH)4]2-(aq)

    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)

     ZnO(s) + 2NaOH(aq) + H2O(l) → Na2Zn(OH)4(aq)
    Ionically: ZnO(s) + 2OH-(aq) + H2O(l) → [Zn(OH)4]2-(aq)

    Hydroxides:

    Al(OH)3(s) + NaOH(aq) → NaAl(OH)4(aq)
    Ionically: Al(OH)3(s) + OH-(aq) → [Al(OH)4]-(aq)

    Zn(OH)2(s) + 2NaOH(aq) → Na2Zn(OH)4(aq)
    Ionically : Zn(OH)2(s) + 2OH-(aq) → [Zn(OH)4]2-(aq)

    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 are replaced  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
    1. Preparations by direct synthesis
      Equation:
      Fe(s) + Cl2(g) → 2FeCl3(s)
    2. 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.
    3.  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 Aqueous Sodium Hydroxide and Aqueous Ammonium Hydroxides and Solubilities of Some Salts in Water

 Cation  Soluble compounds in water  Insoluble compounds in water  Reaction with NaOH(aq)  Reaction with NH4OH
 K+  all  None  No reaction  No reaction
 Na+  all  None  No reaction  No reaction
 Ca2+  Cl-; NO3-  CO32-; O2-; OH-  White precipitate insoluble in excess  No precipitate
 Al3+  Cl-; NO3-; SO42-  CO32-; OH-  White precipitate soluble in excess  White precipitate insoluble in excess
 Pb2+ NO3-, ethanoate  All others  White precipitate soluble in excess  White precipitate insoluble in excess
 Zn2+  Cl-; SO42-; NO3-  CO32-; OH-  White precipitate soluble in excess  White precipitate insoluble in excess
 Mg2+  Cl-; SO42-; NO3-  CO32-; OH-  White precipitate insoluble in excess  No precipitate
 Fe2+  Cl-; SO42-; NO3-  CO32-; O2-; OH-  (dark) green precipitate soluble in excess  Green precipiate insoluble in water
 Fe3+  Cl-; SO42-; NO3-  CO32-; O2-; OH-  (red)brown precipitate insoluble in excess  Brown precipiatate insoluble in excess
 Cu2+  Cl-; SO42-; NO3-  CO32-; O2-; OH-  Pale blue precipitate insoluble in excess  Pale blue precipitate soluble in excess forming a deep blue solution
 NH4+  all   none  Ammonium gas on warming  Not 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:
    insoluble hydroxides
  • The hydroxides of aluminum, zinc and lead dissolves in excess sodium hydroxide solution because of complexes are formed
    Equations:
    Tetra hydroxyl aluminium III ion
    complex salts
  • Note: - in these reactions KOH(aq) may be used instead of sodium hydroxide

With Ammonia Solution

  • Insoluble metals hydroxides are similarly formed.
    insoluble hydroxides with ammonia
  • However hydroxides of copper and zinc dissolve in excess ammonia solution due to formation of complex ions/ salts

    Equations:
    complex salts with ammonia

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
    effect of heat on copper hydroxide
    effect of heat on hydroxides
  • Note: Both iron (II) and iron (III) hydroxides give iron (III) oxide when heated.
    Equations:
    effect of heat on iron hydroxides
  • 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 containing  Observations 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 pecipitate and a colourless gas that forms a white ppt in lime water
     Al3+  A white pecipitate 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

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.

 

 

Observations

 Solution containing   Sodium sulphate  Sodium chloride  Sodium sulphate
 Zn2+  Colourless solution  Colourless solution  Colourless solution
 Mg2+  Colourless solution  Colourless solution  Colourless solution
 Cu2+  Blue solution  Blue solution  Blue solution
 Fe2+  Greenish solution  Green solution  -
 Fe3+  Yellow solution  Yellow/dark brown solution  -
 Pb2+  White precipitate  White precipitate which dissolves on warming  White precipitate
 Ba2+  White precipitate  White precipitate  White 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)
    Pb 2+(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   Conclusion
 Solid  Aqueous solution (if soluble)
 1. White  Colourless  Compound of K+; Na+; Ca2+; Mg2+; Al3+; Zn2+; Pb2+; NH4+
 2. Yellow  Insoluble  Zinc oxide, ZnO (turns white on cooling); Lead oxide, PbO (remains yellow on cooling, red when hot)
 Yellow   Potassium or sodium chromate
 3. Blue  Blue  Copper (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. Brown

 Brown(sometimes yellow)

 Insoluble

 Iron (III) compounds, Fe2+

 Lead (IV) oxide, PbO2

 6. Pink

 Pink(almost colourless)

 Insoluble

 Manganese(II) compounds, Mn2+

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

 7. Orange  Insoluble  Red lead, Pb3O4 (could also be mercury (II) oxide, HgO)
 8. Black  Purple
 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

 CATION  SOLUBLE COMPOUNDS (IN WATER)  INSOLUBLE COMPOUNDS (IN WATER)  REACTION WITH NaOH(aq)  REACTION WITH NH4OH
 K+  all  None  No reaction  No reaction
 Na+  all  None  No reaction  No reaction
 Ca2+  Cl-; NO3-  CO32-; O2-; OH-  White precipitate insoluble in excess  No precipitate
 Al3+  Cl-; NO3-; SO42-  CO32-; OH-  White precipitate soluble in excess  White precipitate insoluble in excess
 Pb2+ NO3-, ethanoate  All others  White precipitate soluble in excess  White precipitate insoluble in excess
 Zn2+  Cl-; SO42-; NO3-  CO32-; OH-  White precipitate soluble in excess  White precipitate insoluble in excess
 Mg2+  Cl-; SO42-; NO3-  CO32-; OH-  White precipitate insoluble in excess  No precipitate
 Fe2+  Cl-; SO42-; NO3-  CO32-; O2-; OH-  (dark) green precipitate soluble in excess  Green precipiate insoluble in water
 Fe3+  Cl-; SO42-; NO3-  CO32-; O2-; OH-  (red)brown precipitate insoluble in excess  Brown precipiatate insoluble in excess
 Cu2+  Cl-; SO42-; NO3-  CO32-; O2-; OH-  Pale blue precipitate insoluble in excess  Pale blue precipitate soluble in excess forming a deep blue solution
 NH4+  all   none  Ammonium gas on warming  Not applicable

Qualitative analysis for common anions.

   SO42-(aq)   Cl-(aq)  NO3-(aq)  CO32-(aq)
 TEST

Add Ba 2+ (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)and warm
 Add FeSO4(aq);
 Tilt the tube and carefully add 1-2 cmof concentrated H2SO4(aq)
 Add dilute HNO3(aq);
bubble gas through lime water;
 OBSERVATION The 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- ions

 

Evolution of a colourless gas that forms a white precipitate with lime water, turns moist blue litmus paper red; and extinguishes a glowing splint shows presence of CO32- ions

 EXPLANATION  Only 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 AgCOcan 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 H2SOforms nitrogen (II)oxide with NO3-(aq) and this forms brown ring complex (FeSO4 .NO) with FeSO4 ;

All CO32- or HCO3will 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. Sublimation  White solids on cool, parts of a test tube indicates NH4+ compounds;
 Purple vapour condesing to black solid indicates iodine  crystals
 2. Water vapour (condensed)  Colourless droplets on cool parts of the parts of the test tube indicate water of crystallization or HCO3-
 3. Carbon (IV) oxide  CO32- of Zn2+, Pb2+, Fe2+, Fe3+, Cu2+
 4. Carbon (IV) oxide and water vapour  HCO3-
 5. Nitrogen (IV) oxide  NO3-; of Cu2+; Al3+; Zn2+; Pb2+; Fe2+; Fe3+
 6. Oxygen  NO3-; or BaO; MnO2; PbO2

Reduction-oxidation (Redox Reactions)

  1. Displacement reactions.
    1. 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.
    2. 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.
  2. 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 colourless Mn2+(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;
  3. 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.
  4. Oxidation of Fe 2+ (aq) to Fe 3+ (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.

Materials

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

 

Apparatus
solubility of potassium nitrate

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

 Temperature  20.0oC
 Mass of evaporating dish + solution  100.7 g
 Mass of evaporating dish  65.3g
 Mass of solution  35.4g
 Mass of evaporating dish + dry salt  73.8g

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)/26.9 = 31.6g of salt;
Therefore the solubility of KNO3 at 20o C = 31.6g per 100g of water

 

 

Factors Affecting Solubility

  1. 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.

      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.
       Experiment number  I  II  III  IV  V
       Volume of water added  5  7  9  11  13
       Temperature at which crystals appear oC          
       Solubility of K in g/100g of water          
      Questions:
      1. Complete the table and calculate the solubility of solid X in g/100g of water at different temperatures. (2 marks)
      2. Using the table above, plot a graph of solubility of solid X in g/100g of water against temperature. (5 marks)
      3. From the graph:
        1. calculate the mass of K that would be obtained if the saturated solution is cooled from 60oC to 40oC. (2 marks)
        2. determine the solubility at 70oC. (1mark)
        3. at what temperature would solubility of K be 100g/100g of water? (1mark)
  2. 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.
    solubilty curves

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)

Sample Question

  1. An experiment was carried out to determine the solubility of potassium nitrate and the following results were obtained.
     Temperature  10  15  30  40  50  60
     Mass of KNO3 per 100g of water  20  25  45  63  85  106
    1. What is meant by solubility? (1 mark)
    2. Plot a graph of mass of potassium nitrate against temperature. (3 marks)
    3. 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)
    4. Explain what would happen if 100g of KNO3 was put in cold water and heated to 50oC. (2 marks)
  2. The table below shows the solubility of sulphur (IV) oxide at various temperatures.
     Temperature (oC)  0  5  10  15  20  25  35  40  45  50  55  60
     Mass of SO2 per 100g of water  22  18.4  15.4  13.0  10.8  9.05  7.80  6.80  5.57  4.80  4.20  3.60

    1. On the grid provided plot a graph of solubility against temperature. (3 marks)
    2. From the graph determine:
      1. The lowest temperature at which 100cm3 of water would contain 11.6g of sulphur dioxide.(1 mark)
      2. 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)
    3.  
      1. Sulphur (IV) oxide reacts with sodium hydroxide solution to form sodium sulphite and water. (1 mark)
      2. Write the equation for this reaction. (1 mark)
      3. 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 and 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+
    2C17H35COO- Na+(aq) + Ca2+(aq) → (C17H35COO-)2Ca(s) + 2Na+(aq)
    Sodium stearate                                   Calcium stearate

    With Mg 2+
    2C17H35COO-Na+(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 Hardness

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:

  1. Boiling:
    • Boiling decomposes and an insoluble chalk of CaCO 3 and MgCO 3 respectively is deposited in the sides of the vessel.
    • This forms an encrustation commonly known as furr the process being furring.
      Equations:
      removal of hardness by boiling
  2. 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.
  3. 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).
  4. 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)
  5. 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.
      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

  1. 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
  2. By permutit softener (ion exchange); explanations as before
  3. Distillation.

Advantages of Hard Water

  1. It is good for drinking purposes as calcium ions contained in it helps to form strong bones and teeth.
  2. 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.
  3. It is good for brewing and the tanning industries; it improves wine or beer flavour in brewing industries.

Disadvantages of Hard Water

  1. 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, Ariel, Sunlight fab e.t.c.
  2. 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
  3. Formation of kettle fur which makes electrical appliances inefficient hence increasing running costs.
  4. Formation of scum on clothing reduces their durability and aesthetic appearance

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