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