Test for Cations using aqueous NaOH and aqueous NH3

When testing for cations, the common reagents used are aqueous NaOH and aqueous NH3. 

It is important to be careful when using these reagents, because different products are got / observe when we put these reagents in drops from  when we apply these reagents in excess.

For example when NaOH reacts with zinc ions it forms insoluble zinc hydroxide which is usually observed as a white gelatinous precipitate and on further reaction with excess NaOH a soluble complex compound of zinc is formed which causes the white gelatinous precipitate to dissolve

The following video shows the action of NaOH and NH3 solutions on Al3+, Zn2+, and Pb2+ ions

OXIDES OF CARBON at a glance

 

CARBON (IV) OXIDE

Carbon (iv) oxide occurs in the  atmospheric. About 0.03%.

  

Laboratory preparation

Carbon (iv) oxide is prepared in the laboratory by the action of dilute hydrochloric acid on calcium trioxocarbonate (iv) (marble chips or limestone).

 

 

 

 

CaCO_3(s)  +  2HCl_(aq) →CaCl_2(aq) + H₂O_(l) + CO_2(g)

2. It is also prepared by heating metallic trioxocarbonates (iv) (except those of Na and K), or the hydrogen trioxocarbonate (iv) of Na or K.

CuCO_3(s) → CuO_(s) +  CO_2(g)

 Dry CO2 is obtained by  passing the gas  through potassium hydrogen trioxocarbonate (IV) solution (to  remove any acid fumes, and then through fused Calcium chloride in a U-tube to remove the water vapour.)

 

 

 

 

The dry gas is then collected by downward delivery as it is heavier than air.

 The reaction  can also be prepared in Kipp’s apparatus

 

INDUSTRIAL PREPARATION

CO₂ is prepared industrially as a by product of fermentation or when limestone is heated strongly make quicklime.

 

PHYSICAL PROPERTIES

i.   CO₂ is a colourless gas 

ii.  It is an  odourless gas with a sharp refreshing taste.

iii.  It is about 1.5 times denser than air.

iv.  It is soluble in water.

v.   It turns damp blue litmus paper pink.

vi.  It solidifies on cooling (-78⁰C) to form a white solid known as dry ice.

 

CHEMICAL PROPERTIES

1. Reaction with water: Carbon (iv) oxide dissolves in water to form trioxocarbonate (iv) acid (Soda water), a weak, dibasic acid which ionizes slightly.

(a)  CO_2(g) + H₂O_(l) →H₂CO_3(aq) 

2. Reaction with alkalis: It reacts with alkalis to yield trioxocarbonate (iv) salts.

      CO_2(g) + 2NaOH_(aq) →Na₂CO_3(aq) + H₂O_(l)

Limited

with excess CO₂ reacts with alkalis to produce Hydrogen trioxocarbonate (iv) salt.

    CO_2(g)+ NaOH_(aq)→NaHCO_3(aq)

 

4.  When passed over  red hot coke. CO₂ is reduced to CO.

          CO_2(g)+ C_(s)→2CO_(g)

 

Test for CO₂: 

When CO2 is bubbled through  lime water (Calcium hydroxide), it will turn lime water turn milky. ( because of the formation of insoluble calcium trioxocarbonate)

     Ca(OH)_2(aq)+ CO_2(g) →CaCO_3(s)+ H₂O_(l).

If the gas is bubbled in excess, the milkiness disappears and turns to a clear solution due to the formation of soluble calcium hydrogen trioxocarbonate (iv).

CaCO_3(s) + H₂O_(l) +CO_2(g) →Ca(HCO₃)_(aq)

 

 

Uses of carbon (IV) oxide

i.     It is used in making  carbonated (aerated) drinks. It is responsible their refreshing taste. 

 ii.   It is used in fire extinguishers because it does not support combustion.

iii..  It is used in the Solvay Process for the manufacture of Na₂CO₃ (washing soda)  

iv..    It is used as a leavening agent in the baking of bread. 

v.      Solid CO₂ (i.e dry ice) is used as a refrigerant for perishable goods e.g ice cream.

vi.     Gaseous CO₂ is used to preserve fruits.

vii.    CO₂ is also used as a coolant in nuclear reactors.

 

CARBON (II) OXIDE

LABORATORY PREPARATION

Carbon (II) oxide is prepared by the dehydration of methanoic (formic) acid or ethanedioic (oxalic) acid, using concentrated tetraoxosulphate (vi) acid.

 

 

 

 

 

 

HCOOH_(l)^{Conc. H2SO4}→CO_(g) + H₂O
Methanoic acid

 

  COOH
   |        ^{Conc. H2SO4}→  CO2 + CO
  COOH
ethanedioic 

 The CO2 is removed by passing the gaseous mixture through concentrated NaOH 

                                                                                                 

Physical Properties Of Carbon (ii) Oxide

i. Carbon (ii) oxide is a colourless, odourless and tastless

ii. It is a poisonous gas

(2) It is insoluble in water, but dissolves in a solution of ammoniacal copper (i) chloride.

(3) It is neither lighter nor heavier than air.

(4) It is neutral to litmus.

Chemical Properties of Carbon (ii) oxide

(1) As a reducing agent:-Most metallic oxides are reduced to the metals on reaction with CO  oxidizing it to CO₂.

CuO_(s) + CO_(g) →Cu_(s)+ CO_2(g)

Fe₂O_3(s) + 3CO_(g) →2Fe_(s) + 3CO_2(g

2.  Combination reaction

i.  With oxygen: CO burns in air with a faint pale blue flame to form CO₂ .

       2CO_(g)+ O_2(g) →2CO_2(g)

ii.  With haemoglobin: CO has equal affinity for the red blood cells as oxygen, and when exposed to as little as 0.005% of the gas  it combines irreversibly with haemoglobin in the red blood cells to form carboxy-haemoglobin. These prevents oxygen from reaching the blood and this can cause death by suffocation.

 

 

Test for Carbon (ii) oxide

Inserted a lighted splinter into a test tube containing the unknown gas, if it burns with a pale blue flame and turns and some lime water after burning, the the gas is carbon (ii) oxide 

 

Uses of Carbon (ii) oxide

(1) CO is used for extraction of metals from their ores.

(2) It is an important constituent of gaseous fuels like producer gas and water gas.

(3) CO gas is used in the manufacture of organic compounds like methyl alcohol, synthetic petrol. 

 

OBJECTIVE QUESTIONS 

1. Kipp’s apparatus is important in the laboratory because it

 (a) allows intermittent supply of gases. 

(b) is used for preparing poisonous gases. 

(c) is used to prepare light gas. 

(d) is used to prepare sensitive gas

2. Gas prepared by the reaction between methanoic acid and concentrated tetraoxosulphate (vi) acid is 

(a) SO₂           

(b) CO             

 (c) CO₂           

(d) H₂S.

3. Gas which dissolves in ammoniacal copper (i) chloride but insoluble in water is

(a) NH₃ 

(b) CO 

(c) N₂O 

(d) CO₂.

4. Where else is CO₂ found in free state apart from the atmosphere?

(a) In carbonated drinks. 

(b) Dissolved form in water. 

(c) In corals. 

(d) In limestone region

5. It is dangerous to stay in a badly ventilated room which has a charcoal fire because of the presence of 

(a) carbon (ii) oxide 

(b) carbon (iv) oxide 

(c) hydrogen sulphide 

(d) producer gas.

 

THEORY QUESTIONS 

1(a)i  Describe the laboratory preparation of dry Carbon (iv) oxide.

   ii. write the equation for the preparation of CO2

   iii. mention two properties of CO2

1b. State what is observed when 

(i) excess CO₂ is bubbled through lime water. 

(ii) the solution in b(i) above is heated.

2(a)i. What  property of CO₂ makes it to be used in 

(I) carbonated drinks (II ) fire extinguishers

3(a)Draw  the laboratory preparation of carbon (ii) oxide done in a fume chamber?

2ii  Explain why Carbon (ii) oxide cannot be collected by any method of delivery

3.  Write two equations to show the chemical properties of Carbon (ii) oxide

 

 

 

 

 

 

OXYGEN AT A GLANCE

 

OXYGEN AND ITS COMPOUNDS

Oxygen is the 8^th element on the periodic table. It has an atomic number of 8 and a mass number of 16  (₁₆⁸O). it has an electronic configuration of 1s²2s²2p⁴. It exhibits oxidation states of -2, -1,0, and It exists in isotopic mixtures it also has two allotropic forms that is, molecular O₂ and ozone O₃. It was discovered separately by Carl. W Scheele in 1972 and Priestley 1974 but it was named by Antoine Lavoisier in 1777.

OCCURRENCE

It occurs  freely as molecular oxygen in the atmosphere (O₂) about 21% of the atmosphere is Oxygen. It also occurs in the combined state as Trioxosilicate (IV) (Al₂(SiO₃)₃), trioxocarbonates (IV) e.g  (CaCO₃), it is present most oxides that make up rocks and clays as well as in water (H₂O).

Laboratory Preparation

Oxygen is prepared in the laboratory by several methods the commonest being

1.  The decomposition of Potassium trioxochlorate V

2KClO₃  →^heat 2KCl + 3O₂

2.  Oxidation of Hydrogen peroxide using MnO₂ or acidified potassium tetraoxomanganate VII (KMnO₄) as the oxidizing agent.

      2H₂O₂  →^MnO₂     2H₂ + O₂            

     MnO₂ here is acting as a catalyst

5H₂O_2(aq) + 2KMnO_4(aq) + 3H₂SO_4(aq)  →K₂SO_4(aq) + MnSO_4(aq) + H₂O_(l) + O_2(g)

Industrial preparation of Oxygen

Oxygen is prepared industrially by 

1. Electrolysis of acidified water and 

2. Fractional distillation of liquid air

By electrolysis: hydrogen is discharged at the cathode while oxygen is discharged at the anode.

FRACTIONAL DISTILLATION OF LIQUEFIED AIR

1. Air is first passed through caustic soda to remove CO₂. The air is then subjected to a series of conditions which includes high pressures, low temperature and expansions which causes the air to liquefy (that is become a liquid).

2. The liquid air is then passed into the fractionating column and heated, Nitrogen with a lower boiling point of -196⁰C distills first followed by Oxygen with a boiling point of -183⁰C. Oxygen distils over as a gas and is collected, dried and stored in steel cylinders.

    Physical properties

1.  Oxygen  is a colourless gas 

2. it is an odourless 

3. it is a tasteless gas

4. It is slightly soluble in water

5. It is neutral to litmus 

Chemical properties 

Oxygen combines readily with almost all substances  as discussed below

Reaction with

1.  metals: - most metals burn in Oxygen to yield basic oxides

            2Mg_(s) + O_2(g) →2MgO_(s)

2. Non-metals: - non-metals burn in oxygen to yield acidic oxides

          S(s)  + O2(g) → SO2(g)

3. Reaction with organic compounds: - most organic compounds burn in oxygen to yield CO2, H2O and the oxide of any other element except oxygen present in the compound. E.g 

i.  C2H6 +O2  → CO2 + H2O

ii.  

   

  TEST FOR OXYGEN

When oxygen gas is brought close to a dying flame it rekindles the flame 

USES OF OXYGEN

1.      It is used for breathing by divers and mountain climbers

2.      It is combined with ethyne by welders to produce very hot flames

3.      Liquid oxygen and fuels are used as propellants for space rockets.

                             

                        OXIDES

Oxides are binary compounds formed when elements burn in oxygen.

They are classified as

1.      Basic oxides: - these are oxides formed when metals burn in oxygen. 

Examples of basic oxides are Na2O, CaO, MgO, K2O

2Mg_(s) + O_2(g) →2MgO_(s)

Properties of basic oxides

a.      They are mainly solids

b.      Soluble oxides dissolve in water to form Alkalis

      Na₂O_(s) + H₂O_(l) → 2NaOH_(aq) + H₂O_(l)

c.       They react with acids to form salt and water

      K₂O_(s) + HCl_(aq) →KCl_(aq) + H₂O_(l)  

2.      Acidic Oxides: - these are oxides of non-metals; they are formed when non-metals burn in oxygen. Examples of acidic oxides are SO2, SO3, CO2, NO2, P2O5

            S_(s) + O_2(g) → SO_2(g)

Properties of acidic oxides

a.      They dissolve in water to form corresponding acids ( also called acid anhydride)

SO_2(g) +H₂O_(l) → H₂SO_3(aq)

b.      They react with bases to form salt and water.

         2NaOH_(aq) + SO_2(g) → Na₂SO₃ + H₂O_(l)

3.      Amphoteric Oxide: - these are oxide of metals that behave both as acidic and basic oxides. They are oxides of Al, Sn Pb and Zn. Examples of Amphoteric oxides are PbO₂, ZnO, Al₂O₃, SnO₂

1. With acids they form salt and water only

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

2. With alkalis they form complex salts 

ZnO + 2KOH + H2O → K2Zn(OH)4

4.      Neutral Oxides: - these are non-metallic oxides that are neutral to litmus.

 E.g carbon II oxide (CO), dinitrogen (I) oxide (N₂O) and water (H₂O) which is the only neutral oxide that is liquid at room temperature.

5.      Peroxides: - these are oxides that contain a higher proportion than the usual oxides. E.g sodium peroxide Na₂O₂, H₂O₂,

Objective Questions 

1. Amphoteric oxides are oxides which 

a) react with water to form acids 

b) react with water to form alkali

c) show neither acid nor basic properties 

d) react with both acids and alkalis 

2.  The component of air that is removed when air is bubbled through alkaline pyrogallol solution is 

a) Carbon (IV) oxide 

b) oxygen

 c) water vapour 

d) nitrogen

3.  When the trioxonitrate (V) salt of an alkali metal Y is heated, the formula of the residue is 

 a) Y₂O 

b) YNO₂ 

c) Y₂O₃ 

d) Y(NO₂)₂

3.   Which of the following oxides is amphoteric 

a) Na₂O 

b) Fe₂O₃ 

c)Al₂O₃

d) CuO.

5.      The following oxides react with both acids and bases to form salts except

 a) zinc oxide 

b) lead (II) oxide

 c) aluminum oxide

 d) tin (IV) oxide.

6.      The following oxides reacts with water except 

a) Na₂O

 b) SO₃ 

c) NO₂ 

d) CuO

7.      If X is a group III element, its oxide would be represented as 

a) X₃O₂ 

b) X₂O

 c) X₂O₃ 

d) XO₃

8.      Which of the following elements is diatomic? 

a) Iron 

b) Neon 

c) Oxygen 

d) Sodium.

9.       Which of the following substances is mainly responsible for the depletion of the ozone layer?

 a)  Chlorofluorocarbon

 b) Carbon (IV) oxide

 c) Nitrogen

 d) Oxygen

10.    Which of the following oxides is ionic

a) P₄O₁₀ 

b) MgO

 c) Al₂O₃

d)SO₂

11. What term is used to describe an oxide whose aqueous solution turns red litmus blue

a. Strong electrolyte

b. Acid anhydride

c. Amphoteric oxide

d. Basic oxide

 THEORY

1.  (a)(i) what are acidic oxides? 

(ii). give one example of each of the following oxides I. acidic oxide II. Basic oxide III. Amphoteric oxide IV. Neutral oxide

b (i) Explain what is meant by acid anhydride 

(ii) give one example of the oxide mentioned in b(ii) above

2.(a).Draw a well labelled diagram for the laboratory preparation of oxygen.

(b). Write the formulae of three different oxides of period 3 elements that react with water.

3a(I). Classify each of the following oxides as acidic, basic, neutral or amphoteric.         

(a)(i). I. ZnO  (II) CO (III) NO₂

(b).  Consider the following oxides: CaO, SiO₂, CO, NO₂ and ZnO. Which of the oxide(s)

 (i). is an acidic oxide that is insoluble in water?

 (ii). Reacts with water to give alkaline solution

 (iii). Is amphoteric? 

 (iv). Is neutral 

(v) is/are gaseous at room temperature.

4.   ZnO is an amphoteric oxide. Write equations to illustrate this statement.

ii. Explain why NaNO3 is preferred to AgNO3 in the preparation of oxygen by thermal decomposition of trioxonitrate (V) salts? 

 

Redox Reactions at a glance

Oxidation And Reduction reaction (Redox)

 Oxidation-Reduction(redox) reactions are two opposite and complementary reactions which occur simultaneously

Redox reactions have been defined in several ways before attaining a more general and simplified definition. These definitions are as follows

1.     In term of addition of oxygen:  Oxidation is defined as the addition of oxygen to a substance while reduction is defined as the removal of oxygen from a substance. E.g.

_Reduction              

CuO + C_(s) → Cu_(s) +  CO_(g)   

 ^{O.A        R.A}                       
   the example above, carbon (C) is oxidized to carbon (II) oxide (CO) while Cupper (II) oxide (CuO) is reduced to metallic Copper (Cu).

   Carbon is removing oxygen from CuO and so is the reducing agent because it causes CuO to become reduced to Cu. CuO  supplies the oxygen atom that causes carbon to become oxidized to CO and so CuO  is the oxidizing agent.

2       In terms of removal of hydrogen:  Oxidation is defined as the removal of hydrogen from a substance while reduction is defined as the addition of hydrogen to a substance                    

                   _{——————
                     ↓Reduction       ↓}
H₂S   +   Cl₂ → S_(s) + HCl_(aq)
 ^{R A            O.A}    
 ↑   ^{Oxidation          ↑
      ——————}

                     

          Similarly in this reaction, H₂S is oxidized to atomic S_(s) due to the removal of hydrogen as chlorine is reduced by gaining or addition of hydrogen. H₂S is action as the reducing agent while Cl₂ is the oxidizing agent.

3        In terms of change in the oxidation number of an element: Oxidation is defined as the increase in oxidation number of an element while reduction is the decrease in oxidation number of an element. 

4.  Definition in terms of electronegative elements: - Oxidation is the addition of an electronegative element to a substance or the removal of electropositive element from a substance while reduction is the removal of electronegative element from a substance or the addition of electropositive element to a substance 

4        In terms of electron transfer: Oxidation is defined as the loss of electrons while reduction is defined as the gain of electrons. 

  When an element loses an electron to become an ion; the O.N increase to a higher number while a gain of electrons by an element will lead to a decrease in O.N of an element. For example,

        ₂₀Ca   → ₂₀Ca²⁺+ 2e^–
        20 protons      20 protons
   20 electrons         18 electrons
    O.N 0 (zero)        O.N (+2)

        

        ₁₇Cl    +   e^–   → 17Cl^–
     17 protons            17 protons
     17 electrons         18 electrons
     O.N 0 (zero)         O.N( –1)

          In other words, as noted from the above examples, loss of electrons means a higher O.N while gain of electrons means a decrease in O.N of an element.

                            _Oxidation
               Mg + Cl   →MgCl₂       
               ^{R.A   O.A     reduction}

Example of redox reactions that occurs generally around us include

1. Photosynthesis

2. Rusting of iron

  Fe_(s) + nH₂O → Fe₂O₃.nH₂O_(s)

3.   Combustion

           _{–——————}
        ↓ _Oxidation         ↓
      C₄H₁₀ + O₂ → CO₂ + H₂O
                      ↑ ^reduction       ↑
                         ^{——————}

Redox reactions always involve the movement of electrons ( i.e loss and gain of electrons) For example 

    Pb° →Pb²⁺ + 2e^–  

   Pb²⁺→Pb⁴⁺ + 2e^–. 

O.N increased from 0 to +2 and then to +4 in Pb. i.e., oxidation involves loss of electrons which will lead to increase in O.N.

In contrast reduction involves the gains of electrons which will lead to a decrease in the O.N of the element for example                   S° + e → S^– + e^– → S^2–.

Some examples of redox recitations are

1        Fe_(s)+ S_(s)→ FeS

          ^{R.A        O.A}

          Fe losses electrons in the above reaction to become iron (II) ions (Fe²⁺),  its O.N from 0 to +2, it is oxidized, and so it is the Reducing Agent. Sulphur on the other hand, gains electrons from the iron, its O.N decreases from 0 to (–2) and so it is reduced and so is the Oxidizing Agent

2. Pb⁺²O^-2 + C⁺²O^-2 → Pb⁰+ C⁴⁺O^2-₂

     In the above example, the O.N of Pb decreased from +2 to 0, so PbO is reduced to Pb and so PbO is the oxidizing agent while  the O.N of C increased from +2 to +4, so CO is oxidized and so CO is the reducing agent.

3        H_2(g)+ O_2(g) → H₂O_(l) 
          ^{R.A           O.A}

          O.N of H increased from 0 to (+1) i.e. H is oxidized.

          The above reaction is a combustion reaction, and at this point it  is important to note that all combustion reactions are redox reactions with oxygen as the oxidizing agent.

   i.     Pb(NO3)2 + 2NaCl → PbCl2 + 2NaNO3

          The above reaction is a double decomposition reaction; it is not a redox reaction as there is no change in the O.N number of all the element s involved. Another non-redox reaction is a neutralization reaction, there is no change in the O.N of element involved in the reaction.

  ii  KOH_(aq) +HCl_(aq) → NaOH_(aq) + H₂O_(l)    (neutralization reaction)

          

Oxidizing and reducing agents (in summary)

OXIDIZING AGENT 1.      Supplies Oxygen 2.      Removes Hydrogen 3.      Decreases in oxidation number 4.      Gains electrons

REDUCING AGENT Supplies hydrogen Removes Oxygen Increases in oxidation number Loss electrons

Test for oxidizing agents

To common test or reactions that are used to test for an oxidizing agent involves the action on iron (II) chloride and hydrogen sulphide.

a)       Reaction with FeCl₂

          When an oxidizing agent is added to green iron (II) chloride; the green iron (II) ions become oxidized to yellow or brown Fe³⁺.

          Fe²⁺ →     Fe³⁺ + e^–

          green         yellow/brown

b)      Reaction with hydrogen sulphide

          When hydrogen sulphide is bubbled through a solution of an oxidizing agent, the sulphide ions S^2– becomes oxidized to elemental sulphur; and this is seen or observed as yellow deposits sulphur,   i.e. 

  S^2– → S(s) + 2e^–.

Test for reducing agents

Two commonest reagents that are used to test for a reducing agent are

1 Acidified potassium tetraoxomanganate(VI) (KMnO₄) and acidified potassium heptaoxodichromate(I) (K₂Cr₂O₇).

a)   Action of potassium hyptaoxodichromate (VI) (K₂Cr₂O₇)

  When acidified potassium heptaoxodichromate (VI) (K₂Cr₂O₇) is added to a sample of a reducing agent, its colour changes from orange to green, due to the reduction of the dichromate (VI) ion  (Cr⁶⁺) (orange) to chromium (III) (Cr³⁺) ion green

     Cr⁶⁺  +  3e^– → Cr^{3+
    Orange                green}

 b) Test using acidified potassium tetraoxomangane(VI) (KMnO₄)

  When acidified potassium tetraoxomanganate (VII) to a sample of reducing agent, the purple colour changes to colourless: due to the reduction of the manganate ion from (+7) which is purple to (+2) which is colourless and a more stable oxidation state.

MnO₄^- + 8H⁺ + 5e^– →Mn²⁺ + 4H₂O
 ^{purple              colourless}         

Mn⁷⁺ + 5e^– → Mn^{2
purple              colourless}

   This reaction is reversible as the purple colour is restored when an oxidizing agent is reintroduced into the mixture.

         Mn²⁺ + 5e →Mn^{7+   
       colourless          purple}

OBJECTIVE QUESTIONS.

1.How many electrons are removed from Cr^2- when it is oxidized to CrO₄^2- ?

a) 0

b) 2

c) 4

d) 8

2. Rusting of iron is an example of 

a) deliquescence

b)  decomposition 

c) displacement reaction

d) redox reaction

3.

THEORY QUESTIONS 

1(a) State what you will see 

i)  on bubbling SO₂ into acidified KMnO4 solution   [neco 2025]

ii). when hydrogen sulphide is bubbled into a solution of acidified potassium heptaoxodichromate VI

(b)(i). Write the ionic equation for the reaction between zinc powder and silver trioxonitrate (V) solution 

(ii). Which substance in bi above is I. Oxidized  II. Reduced 

WATER at a glance

 

WATER

CONTENT

·     Types, Uses and Structure of Water.

·      Laboratory Preparation of    Water.

·      Test for Water

·      Causes/ Removal of Hardness of Water.

·      Purification of Water for Municipal Supply.

 

WATER

Water is said to be a universal solvent , because it can dissolve almost all other substances.

SOURCES OF WATER

The following are the sources of water: the sources of water may be grouped into two

1.             Natural water: Rainwater, Well water, Spring water and Sea water, rivers and lakes

2.             Treated water: Distilled water, Pipe – borne water, deionized water and chlorinated water

STRUCTURE OF WATER

Because of the repulsion between the two lone pairs of electrons in the oxygen atom the two bonding pair are pushed towards each other resulting to a V-shape or angular shape or bent shape for water.

 

                   O
                 ∕    \
               H    H

 

LABORATORY PREPARATION OF WATER 

When dry hydrogen gas is lighted in air. It burns with a faint blue flame to give steam, which condenses when it comes in contact with any cold surface to form water.

 

PHYSICAL PROPERTIES OF WATER

1.  Water has a boiling point of  100^oC and freezes at 0^oC

2.  It has a maximum density of 1gcm^-3 at 4^oC

3.  It is neutral to litmus.

 4. It is a liquid at room temperature

 

CHEMICAL PROPERTIES

1.     Water reacts with electropositive metals like K, Na and Ca to form alkali and liberate hydrogen gas. E.g                   Na_(s)+ H₂O_(aq) →NaOH_(aq) +  H_2(g)

  Mg & Zn react with steam to form an alkaline solution while Cu, Hg, Ag, Au, do not react with water

2.       Non-metal especially the halogens chlorine reacts with water to form acid solution

   i.   H₂O_(aq)  +  Cl_2(g) →HCl_(aq)  +  HOCl_(aq)

 ii.   H₂O_(aq)  +Br_2(g) →HBr_(aq)  +  HOBr_(aq)

    

TEST FOR WATER

When few drops of water are added to

1.    White anhydrous copper (II) tetraoxosulphate (VI), it turns blue.

2.     Blue cobalt (II) chloride, it turns pink.

NOTE: These two tests are not specific for water. They only indicate the presence of water. Any aqueous solution or substance containing water will give a positive test for water

 HARDNESS OF WATER

Hard water is any water that does not form lather (foam) readily with soap. 

There are two types of hardness / hard water 

I. Temporary hardness or temporary hard water

II. Permanent hardness/ permanent hard water.

I. Temporary hardness is caused by the presence of Ca(HCO3) or Mg(HCO3) in any water sample and these can be removed by boiling the water.

II. permanent hardness is caused by the presence of CaSO4 or MgSO4 or CaCl2 or MgCl2 it can not be removed by boiling

 

REMOVAL OF TEMPORARY HARDNESS

1.  Physical method: By boiling

                             _heat
    Ca(HCO₃)_2(aq) →     CaCO_3(s)   +   H₂O_(l)   +  CO_2(g)

2.      Chemical method: By using  calculated amount of slaked lime (calcium hydroxide solution)

              Ca(HCO₃)_2(aq) + Ca(OH)_2(aq)  →2CaCO_3(s)+ 2H₂O_(l)

3         Addition of washing soda :-

               Ca(HCO₃)_2(aq) + Na₂CO_3(aq) → CaCO_3(s)+ NaHCO_{3 (aq)}

EFFECTS OF TEMPORARY HARDNESS:  

Hard water causes

1.         Furring of kettles and boilers.

2.        Stalagmite and stalactites in caves.

 

Removal of permanent hardness: 

1. by physical method : Distillation 

2.By chemical method only

i.           Addition of washing soda

    Na₂CO_3(aq) + CaSO_4(aq) →CaCO_3(s) + Na₂SO_4(aq)

ii.   Addition of caustic soda

  2NaOH_(aq) +CaSO_4(aq) → Ca(OH)_2(s)+ Na₂SO_4(aq)

iii.    Ion exchange resin

       CaSO_4(aq) + Sodium zeolite →Calcium zeolite + NaSO_4(aq)
                                                                          (insoluble)

ADVANTAGES OF HARD WATER

i.  Hard water taste better than soft water because of the presence of ions

ii.  Calcium salts in it helps to build strong teeth and bones.

iii.  It provides CaCO₃, that crab and snail use to build their shells.

iv. It does not dissolve Lead, hence it can be supplied in lead pipes.

 

DISADVANTAGES OF HARD WATER

1.             It causes furring of kettles and boilers.

2.             It wastes soap.

3.             It cannot be used in dying and tanning.

4.             Effects is seen in stalactites and stalagmite

 

TREATMENT OF WATER FOR MUNICIPAL/ TOWN SUPPLY

The following are the processes involved in the treatment of river water for town supply

1.    Coagulation: Chemicals like potash alum, KAl(SO₄)₂, or sodium aluminate III, NaAlO₂ is added to water in a large settling tank.

 

2.    Sedimentation: The coagulated solid particles or flocs are allowed to settle in the settling tank to form sediments at the bottom of the tank.

 

3.     Filtration: The water above the sediment still contains some suspended particles.  The water is passed through a filter bed to remove the remaining fine dirt particles.

 

4.     Chlorination (Disinfection): Chemicals like chlorine is then added to the water to kill germs. 

5. Calculated amount of iodine to prevent goiter and fluorine to prevent tooth decay are added as food supplements to prevent goiter and tooth decay respectively. 

 The treated water is then stored in a reservoir and distributed to the town.

OBJECTIVE QUESTIONS

1.  Treated town water undergoes the following steps except

 (A). coagulation    

 (B). precipitation

 (C). sedimentation

 (D). chlorination

2.   Water is temporarily hard because it contains

 A. CaSO₄              

B MgSO₄    

 C. Chlorine

 D. Ca(HCO₃)₂

3.    Temporary hardness of water is removed by the use of one of the following

A. boiling

 B. use of use of Ca(OH)₂

C. use of Na₂CO₃

D. use of alum

4. A substance that turns white anhydrous CuSO₄ blue is

A. water

B. liquid ammonia

C. hydrochloric acid

D. molten Sulphur

5.   Distilled water is different from deionized water because

 A. distilled water is a product of condensed steam while deionized water is filtered laboratory water

B. distilled water is always pure and sold in packs while deionized is not packaged for consumption

C. distilled water is condensed steam but deionized water is produced using ion-exchange resins which absorbs undesired ions.

D. distilled water is man-made while deionized water is both natural and artificial

6.

 

 

THEORY QUESTIONS

1.a i    Mention two compounds that causes permanent hardness in water

  ii.    State two ways of removing permanent hardness in water

  iii.     List two advantages of hard water

  b.(ai )    State the steps involved in the treatment of river water for town supply.

       ii. Write two equations to show the removal of permanent hardness of water.

      iii. Name two cations that causes hardness of water?

c.i. Give two methods of removal removal 

 

 

SALTS at a glace

SALTS

A salt is a compound formed when all or part of the ionizable or replaceable hydrogen ion in an acid is replaced by a metallic or ammonium ion e.g.

i. HCl_(aq) + NaOH_(aq) →      NaCl_(aq) + H₂O_(l)

ii.  H₂SO_4(aq) + KOH_(aq) → KHSO_4(aq) + H₂O_(l)

TYPES OF SALTS

There are five main types of salts namely:

1. Normal salt.

2. Acid salts

3. Basic salts

4. Double salts.

5. Complex salts.

1. Normal salts: are the salts formed when all the replaceable hydrogen ion in the acid has been completely replaced by a metal or an ammonium ion e.g. NaCl, K₂SO₄, Na₃PO₄, NaNO₃ etc. 

Normal salts are neutral to litmus and does not contain any replaceable hydrogen ion (H+) 

i. HCl_(aq)+NaOH_(aq) → NaCl_(aq) + H₂O_(l)

ii. H₂SO_4(aq) + KOH_(aq) → K₂SO_4(aq) +   H₂O_(aq)

2. Acid salts: Acid salts are salts formed when the replaceable hydrogen atoms of an acid are only partially replaced by a metal or an ammonium ion. e.g. NaHSO₄, Na₂HPO₄, NaH₂PO₄, NaHCO₃. 

They are usually formed from acids which contain more than one replaceable hydrogen ion. Acids with two replaceable hydrogen ions can form only one acid salt while acids with three replaceable hydrogen ions can form two different acid salts.

H₂SO_4(aq) + NaOH_(aq) → NaHSO_4(aq)+ H₂O_(l)
                                         sodiumhydrogentetraoxosulphate (VI)

2H₃PO_4(aq) + NaOH_(aq) → NaH₂PO_4(aq) 
                                            monosodiumhydrogentetraoxophosphate (V)

NaH₂PO_4(aq)  + 2NaOH_(aq) → Na₂HPO_4(aq) +H₂O_(l)

disodiumhydrogentetraoxophosphate (V)

Na₂HPO_4(aq + NaOH_(aq) → Na₃PO_4(aq) +H₂O_(l) 
                                       sodiumhydrogentetraoxophosphate (VI)                                        

Properties of Acid salts 

i. Acid salts turn blue litmus red. 

ii. Acid salts react with bases to form salts 

 KHSO_4(aq) + KOH_(aq) → K₂SO_4(aq) + H₂O_(l)

3. Basic salts: Basic salts that still contain replaceable hydroxide ions.

Thay are are formed when only part of the hydroxide ions of a base are replaced by the negative ions from an acid.

e.g Zn(OH)Cl,   Mg(OH)Cl, Mg(OH)NO₃, Bi(OH)₂NO₃ e.t.c.

i. Zn(OH)_2(aq) + HCl_(aq)  → Zn(OH)Cl_(aq) +  H₂O_(l)

ii. Ca(OH)(aq) + HNO3 →    Ca(OH)NO3(aq) + H2O(l)

Properties of basic salts 

i. basic salts turn red litmus blue. 

ii. basic salts react with more acid to form a normal salt and water only.

     Mg(OH)NO_3(aq) +HNO_3(aq) → Mg(NO₃)_2(aq) + H₂O_(l)

4. Double salts: Double salts are salt which ionize to produce three different types of ions in solution. Usually, two of these are positively charged (metallic or NH4+ ion) while the other is negatively charged e.g. (NH₄)2Fe(SO₄)₂.6H₂O, KAl(SO₄)₂.12H₂O, KCr(SO₄)₂.12H₂O.

(NH₄)2Fe(SO₄)₂.6H₂O:  Ammonium iron (II) tetraoxosulphate (VI) hexahydrate.

KAl (SO₄)₂.12H₂O:  Aluminium Potassium tetraoxosulphate (V) dodecahydrate (Potash alum).

KCr(SO₄)₂.12H₂O:   Chromium (III) Potassium tetraoxosulphate (VI) dodecahydrate (Chrome alum).

5. Complex salts: Complex salts contain complex ion i.e ion consisting of a charged group of atoms e.g. Na₂Zn(OH)₄, K₄Fe(CN)₆, NaAl(OH)₄.

Na₂Zn(OH)₄:     Sodium tetrahydroxozincate (II)

K₄Fe (CN)₆:      Potassium hexacyanoferrate (II)

NaAl(OH)₄:       Sodium tetrahydroxoaluminate (III)

Na₂Zn(OH)₄  → 2Na⁺  + [Zn(OH)₄]^2-

K₄Fe(CN)₆  → 4K⁺ + [Fe(CN)₆]^4-.

Properties of complex salts

i. they are soluble in water

HYDROLYSIS OF SALT

Some salts when dissolved in water, undergoes hydrolysis to give an acidic or alkaline solution.

 e.g. Na₂CO₃, NaHCO₃, AlCl₃, Na₂S, NH₄Cl, CH₃COONa e.t.c. It is like the reverse of neutralization. A salt dissolves in water to give the initial acid and alkali or hydroxide from which it was formed. for example 

1. Na₂CO_3(s) +H₂O_(l) → Na⁺ + CO3^2-
 From water   OH^-     

    H⁺
                       strong.    weak

                        base.      acid                                              pH < 7

2. AlCl_3(s) + H₂O_(l) → Al³⁺ + Cl^-
 From water             3(OH)^-    H⁺
                             weak.    strong

                               base      acid                                           pH < 7

3. (NH₄)₂CO_3(s) +H₂O_(l) ⇌ NH₄⁺ + CO₃^2-

     from water       OH^-        H⁺
                            Weak      weak

                             base.       acid                                         pH = 7

Hydrolysis of salt occurs when a salt reacts with water e.g, salt of strong acid and weak base to give an acidic solution. The change in pH of solution is due to hydrolysis.

USES OF SALTS

	SALT	USES
1.	NH4Cl	is used as an electrolyte in dry cell (Leclanché cell)
2.	CaCO3	is used as medicine to neutralize acidity in the stomach.
3.	CaCl2	i. is used as antifreeze while fused CaCl2 is used as a drying agent and also in desiccators. ii.  is used in dyeing and calico printing.
4.	CaSO4	is used for making plaster of Paris.
5.	MgSO4	is used as a laxative.
6.	KNO3	is used for making gunpowder, matches and soil fertilizer.
7.	NaCl	is used for preserving food and in glazing pottery.
8.	ZnCl2	is used in petrol

METHODS OF PREPARATION OF SALTS

The method of preparing  a salts in general depends on its:

i. Solubility in water

ii. Stability to heat.

It is important for us to know the simple rules of solubility indicated above. If we know  the solubility of a salt, it will enables us to determine which method will be used for its preparation

SOLUBLE SALT

Soluble salts can be prepared by any one of the following method:

1. Neutralization of an acid by an alkali

2. Action of dilute acid on a metal.

3. Action of dilute acid on an insoluble base.

4. Action of dilute acid on trioxocarbonate (IV).

OBTAINING SOLUBLE SALTS FROM SOLUTION

This can be done by:

1. Heating to dryness (Evaporation): This is used to recover soluble salts which do decomposed or destroyed by heat e.g. most chlorides such as NaCl, ZnCl2, FeCl2 and FeCl3 are recovered by heating.

2. Crystallization: This method is used to prepare salts which are easily decomposed or destroyed by heating to dryness. All trioxonitrate (V) salts and tetraoxosulphate (VI) are recovered by crystallization.

INSOLUBLE SALTS

Insoluble salts can be prepared by the following method:

1. Double decomposition or precipitation.

i. Pb (NO₃)_2(aq) + 2NaCl_(aq) → 2NaNO_3(aq) + PbCl_2(s)

ii. AgNO_3(aq) + NH₄Cl_(aq) →NH₄NO_3(aq) + AgCl_(s)

2. Direct combination of 2 elements.

i. Fe_(s) + S_(s) → FeS_(s)

ii. 2Fe_(s) + 3Cl_2(g) →2FeCl_3(s)

ANHYDROUS AND HYDRATED SALT

Anhydrous salts: These are salts which do not contain water of crystallization. They cannot be crystallized out from aqueous solution. 

Hydrated salts are salts which contain water of crystallization, when heated, such salt loses their water of crystallization.

Water Of Crystallization: This is a specific amount of water molecules that is embedded in crystals of salts as they form during crystallization.

Cu(NO₃)₂.3H₂O:  Copper (II) trioxonitrate (V) trihydrate.

MgSO₄.7H₂O:      Magnesium tetraoxosulphate (VI) heptahydrate.

FeSO₄.7H₂O:       Iron (ii) tetraoxosulphate (VI) heptahydrate.

Calculations of water of crystallization

1.  14g of hydrated H₂C₂O₄.xH₂O was heated to give an anhydrous salt weighing 9.99g.

(a). Calculate the value of x.

(b). Give the formula of the hydrated salt.

(c). Calculate the % of water of crystallization.

Solution

(a).  Mass of hydrated salt      = Molar mass of hydrated salt
       Mass of water molecule       Molar mass of water molecule
               

                mass of water lost = (14-9.99) = 4.01
                     14 = (90+18x)
                     4.01        18x

        14(18x) = 4.01 (90 + 18x)

        252x = 360.9 + 72.18x

        252x – 72.18x = 360.9

        179.82x = 360.9

        x =    360.9
                 179.82
                   x = 2.007

        x = 2 to the nearest whole number.

(b). Formula of hydrated salt = H₂C₂O₄.2H₂O.

(c) To calculate the % of water of crystallization:

% of water of crystallization = Mass of water x 100%
                                                         Total mass
                       = 36 x 100
                          (90 + 36)
            
                =   36    x   100   = 28.57%
                          126

   

EFFLORESCENCE, DELIQUESCENCE AND HYGROSCOPIC

When certain compounds are exposed to the air, they either lose some or all of their water of crystallization or they absorb moisture from their surroundings to become either moist or form solutions. The term efflorescent, deliquescent and hygroscopic are used to describe such compound/ or phenomenon.

EFFLORESCENCE: This is a phenomenon whereby some salts/ compounds when exposed to the atmosphere loss all or part of their water of crystallization.

EFFLORESCENT SAALTS: are substances which on exposure to air, lose some or all of their water of crystallization. The phenomenon or process is efflorescence. There is loss of weight or mass of the substances.

e.g Na₂CO₃.10H₂O →   Na₂CO₃.H₂O + 9H₂O

Other examples are Na₂SO₄.10H₂O, MgSO₄.7H₂O and CuSO₄.5H₂O e.t.c

DELIQUESCENCE: This a phenomenon whereby some salts when exposed to air absorbs so much water from the air that they form a solution.

DELIQUESCENTS SALTS: are substances that absorb so much water from air and form a solution e.g. NaOH, CaCl₂, FeCl₃, MgCl₂, KOH and P₄O₁₀. There is a gain in weight.

HYGROSCOPIC SUBSTANCES: are substances which absorb moisture on exposure to the atmosphere without forming a solution but only become sticky or wet. If they are solids, no solution will be formed but if a liquid, they absorb water and become diluted  e.g Conc. H₂SO₄, NaNO₃, CuO, CaO and anhydrous Na₂CO₃.

DRYING AGENTS

These are substances which have high affinity for water or moisture. They are either deliquescent or hygroscopic substances. They remove water molecules attached to wet substances to effect physical change. Drying agents are different from dehydrating agents which removes elements of water i.e hydrogen and oxygen atoms or intra-molecular water.

Drying agents which react with gases are not used to dry the gas e.g conc. H₂SO₄ is not used to dry NH₃ and H₂S gas.

NH_3(g) + H₂SO_4(aq) → (NH₄)₂SO_4(aq)

H₂S_(g) + H₂SO_4(aq) → 2H₂O_(l) + SO_2(g) + S_(s)

Drying agent For Gases

Concentrated H₂SO₄ is used to dry All gases except NH₃ & H₂S

Fused CaCl₂ is used to dry All gases except NH₃

CaO (quicklime) is used to dry Ammonia 

P₂O₅ All gases except Ammonia

Silica gel All gases

Salts are usually placed inside desiccators to dry

             a desicator

OBJECTIVE QUESTIONS

1. A substance is said to be hygroscopic if it absorbs 

a. water from the atmosphere to form a solution

b. heat from the surrounding 

c. carbon (iv) oxide from the atmosphere

d. moisture from the atmosphere 

2. The gas given off when NH4Cl is heated with an alkali is 

a. H₂

b. Cl₂

c. N₂ 

d. NH₃

3. A major factor considered in selecting a suitable method for preparing a simple salt is its

a. crystalline form 

b. melting point 

c. reactivity with dilute acids 

d. solubility in water

4. Which of the following salts solutions will have a pH greater than 7

a. NaCl_(aq)

b. Na₂CO_3(aq)

c. Na₂SO_4(aq)

d.NaHSO_4(aq)

5.  Which of the following compound will leave a metal residue when heated 

a. Cu(NO₃)₂

b. AgNO₃

c. K₂CO₃

d.CaCO₃

6. 

THEORY QUESTIONS

1. Give one example of the following salts 

i. Hydrated salt  

ii. Acidic salt

iii. Basic salt 

2.(a)(i) State three methods of preparing salts, giving one example in each case of a salt so prepared.

    (ii). What type of salt is each of the following?   

 i. NaH₂PO₄; 

ii. (CH₃COO)₂Pb; 

iii. KAl(SO₄).12H₂O

b. Write an equation for the reaction between dilute HCl and a solution of AgNO3 

3.(a) Rock salt is an impure form of sodium chloride.

  (i). Outline a suitable procedure for preparing a pure sample of sodium chloride from rock salt.

 (b). Classify each of the following as normal salt/ acid salt/basic salt/double salt

(i). Sodium hydrogen trioxocarbonate (IV) 

(ii). Iron (III) chloride 

(iii). Sodium ethanoate 

4. When a sample of a crystalline salt X was exposed to air, there was a loss in mass.

i. What phenomenon was exhibited by X ?

ii. Suggest two substances which X could be. 

iii. On heating 5.00g of a fresh sample of X to constant mass, 1.80g was lost in the form of water vapour.  Calculate the number of moles of water of crystallization in one molecule of X [ H=1.00, O=16.00; anhydrous form of X=160g/mol]