Carbon IV oxide at a glance

 CARBON (IV) OXIDE: - About 0.03% of atmospheric air is Carbon (IV) oxide by volume while dissolved air contains about 0.50% by volume. This percentages are usually maintained by processes which use up and releases CO2 into the atmosphere, such processes include burning of fossil fuels and organic materials, respiration, deforestation and Photosynthesis 

 

Laboratory preparation

Carbon (IV) oxide is prepared in the laboratory by the action of dilute hydrochloric acid on CaCO₃ which can be in the form of   marble chips or limestone. 

Reaction between CaCO₃ and HCl can be carried out in a Kipp’s apparatus.

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

Note: The dry gas 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 dry the gas. The dry gas is then collected by downward delivery because it is heavier than air.

INDUSTRIAL PREPARATION

CO₂ is obtained industrially as a byproduct in fermentation processes and when limestone is heated to make quicklime.

 

PHYSICAL PROPERTIES

(1) CO₂ is a colourless.

(2) It is an odourless gas with a sharp refreshing taste.

(3) It is about 1.5 times denser than air.

(4) It is soluble in water.

(5). CO₂ dissolves in water to yield trioxocarbonate (IV) acid.

(6) It readily liquefies and solidifies at -78⁰C on cooling to form a white solid known as dry ice.

 

CHEMICAL PROPERTIES

1). It turns damp blue litmus paper pink because

1. Reaction with water: Carbon (IV) oxide dissolves in water to form trioxocarbonate (IV) acid (Soda water). It is a weak dibasic acid ( i.e it ionizes slightly)

(a)  CO_2(g) + H₂O_(l)→H₂CO_3(aq)

On heating rioxocarbonate (IV) acid it decomposes to form H₂O_(l) and CO_2(g).

 

2. Reaction with alkalis: It reacts to form  trioxocarbonate (IV)

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

Limited

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

      CO_2(g) + NaOH_(aq)→NaHCO_3(aq)

     Excess.

3.  Reaction with burning Na, K or Mg:  when passed over burning Na, K andd Mg CO₂ is reduced to carbon.

      CO_2(g) + 2Mg_(s)  →C_(s) + 2MgO_(s)

Note: CO₂ does not support combustion.

 

4.  Reaction with red hot carbon: CO₂ is reduced to CO when passed over red-hot coke.

          CO_2(g) + C_(s) →2CO_(g)

The reaction is important in the blast furnace and in the manufacture of gaseous fuels.

 

Test for CO₂: Bubble the unknown gas through a solution of lime water (Calcium hydroxide) if the lime water turns milky due to the formation of insoluble calcium trioxocarbonate (IV) then the unknown gas is CO₂

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

If the CO_2(g) is bubbled in excess, the milkiness will disappear and turn to a clear solution. This is due to the formation of soluble calcium hydrogen trioxocarbonate (IV).

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

Finally, if the clear solution is heated, the milkiness reappears due to the decomposition of soluble Ca(HCO₃)₂ to form insoluble CaCO₃

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

 

Uses of carbon (iv) oxide

1.  It is used as fire extinguishers since it does not support combustion.

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

3.  It is used in the manufacture of Na₂CO₃ (washing soda) by the Solvay process.

4.  It is used as a leavening agent in the baking of bread. Yeast and baking powder produces CO₂ which make the dough of bread to rise.

5.  It is used in the manufacture of fertilizer (such as urea and (NH₄)₂SO₄.

6.  Solid CO₂ (i.e dry ice) is used as a refrigerant for perishable goods e.g ice cream. (It sublimes on warming and provides a lower temperature).

7.  Gaseous CO₂ is used to preserve fruits.

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

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.  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

Theory Questions 

1) State the property of CO₂ that makes it to be used in

 (i) carbonated drinks (ii) fire extinguishers

(b). State what is observed when 

(i) excess CO₂ is bubbled through lime water. (ii) the solution in b(i) above is heated.

 

 

CARBON AND ITS ALLOTROPES at a glance

   Carbon is found in group IV period II in the periodic table.  It has an electronic configuration of 1s22s22p2.

OCCURRENCE

 It exists naturally as a free element in both crystalline and non-crystalline forms (allotropes).  0v 

ALLOTROPES OF CARBON

Allotropy is the phenomenon whereby an element can exist in two or more different forms but in the same physical state. 

The different forms of the element are known as allotropes. Hence

Allotropes are different forms of an element but in the same physical state

  Allotropes have the same chemical properties but different physical properties.

Carbon exists in several allotropic forms:

(1). Crystalline Allotropes e.g Diamond, Graphite and Fullerene

(2). Non-crystalline Allotropes/Amorphous allotropes  e.g coal, charcoal, coke, lampblack and carbon black (soot)

 

Crystalline Allotropes of carbon

1. Diamond: Diamond is the purest form of carbon.  It is a giant molecule in which the carbon atoms are tetrahedrally bonded (i.e, the carbon atoms in diamond uses all four valence electrons for bonding), closely packed and held together by strong covalent bonds giving diamond an octahedral shape

 

 

Basic tetrahedral arrangement of C-atoms in Diamond Crystals

 

PROPERTIES OF DIAMOND

(1)   Diamond is the hardest substance know

(2)   It has a high melting and boiling point because of strong covalent bond.

(3)   It has a high density

(4)    It is resistant to chemical attack 

(5)    It does not conductor electricity because there are no free valence electrons in the crystal

(6)   It is transparent and has high refractive index (ability to scatter light.)

 

USES

(1) It is used industrially for making drilling machines

(2) It is used to sharpen very hard tools.

(3) It is used for cutting glass and metals.

(4) It is also used as pivot supports in precision instruments and as dies for drawing wires

(5) It is used as jewelry

Artificial diamond is made by subjecting graphite to a very high temperature and pressure for several hours in the presence of nickel or rhodium catalyst.

 

GRAPHITE:  

-Graphite is a dark and opaque allotrope.  

-The carbon atoms in graphite use only 3 out of the 4 valence electrons for bonding (hence graphite contain free mobile electrons) forming flat hexagonal layers.

- Each hexagonal layer is arranged one above the other held by week van der wall forces to form a crystal lattice, 

-These week forces of attraction cause each layer to easily slide over the other which make graphite to also flakes easily

 

 

PROPERTIES OF GRAPHITE

(1)    Graphite is soft and slippery because of weak forces holding its layers. Each layer can slide over one another. Hence, graphite acts as a lubricant.

(2)     It is less dense than diamond

(3)     It is not affected by chemical attack (due to its open structures in layers).

(4).    It is a good conductor of electricity (because of the presence of free delocalized electrons (mobile electron) in the crystal lattice.)

(5)     It has high melting and boiling point.

 

USES

(1) It is usually used on bicycle chains and for the bearings of some motor cars.

(2) It is used as a dry lubricant.

(3) It is used as electrodes in electroplating and in dry cells.

(4) It is used to line crucibles for making high-grade steel and other alloys (since it can withstand high temperature).

(6) It is used in making lead pencils i.e. combining it with clay makes lead in pencils.

(7) It is used as a black pigment in paints.

(8) It is used as a neutron moderator in atomic piles.

 

INDUSTRIAL PREPARATION OF GRAPHITE

Graphite is produced industrially by heating coke in an electric furnace to a very high temperature for about 20 to 30 hours in the absence of air and under sand. This process is called the Acheson process. .

 

 

 

DIFFERENCES IN PROPERTIES BETWEEN GRAPHITE AND DIAMOND

Graphite	Diamond
1. It has a density of 2.3gcm-3	1. It has a density of 3.5gcm-3
2. It is a black, opaque solid	2. It is a colourless, transparent solid
3. It is very soft, marks paper	3. It is the hardest known substance.
4. It is a good conductor of electricity	4. It is a non-conductor of electricity
5. Attacked by potassium trioxochlorate (v) and trioxonitrate (v) acid together.	5. Not attacked by these reagents.

Note: Diamond is transparent to x-rays while glass is almost opaque.

 

Fullerenes 

Fullerenes are a type of carbon molecule, consisting of 60 carbon atoms (C60) arranged in a unique spherical structure. They're also known as buckyballs.

*Properties and applications*

1. *Unique structure*: Fullerenes have a hollow, cage-like structure, making them suitable for applications like drug delivery and nanoencapsulation.

2. *Electronic properties*: Fullerenes exhibit interesting electronic properties, making them potential candidates for use in materials science and electronics.

3. *Superconductivity*: Some fullerene derivatives have shown superconducting properties.

*Potential uses*

1. *Medicine*: Fullerenes are being explored for potential medical applications, such as drug delivery, imaging, and therapy.

2. *Materials science*: They're being researched for use in creating new materials with unique properties.

3. *Energy storage*: Fullerenes might have applications in energy storage and conversion.

Would you like to know more about fullerenes or their potential applications?

AMORPHOUS CARBON

These non-crystalline structures which are not considered to be true allotropes include:

 

CHARCOAL: This is made by burning wood, bones, or sugar in a limited supply of  air. Charcoal is used to remove colour from substances. Wood charcoal is used in absorbing poisonous gases while animal charcoal is used in absorbing colours.

 

CARBON BLACK AND LAMP BLACK: Lamp black is obtained by burning vegetable  oil lamp  that it leaves a deposit of soot  while carbon black is obtained from burning coal gas, natural gas or petroleum. Carbon black and lamp black are used as an additive to rubber tyres. They are also used in making printer’s ink, carbon paper, black shoe polish, type writing.

COAL

Coal is an impure form of carbon. Coal is a complex mixture of compounds composed mainly of carbon, hydrogen and oxygen with small amounts of nitrogen, sulphur and phosphorus as impurities.

Carbonization of coal.

Coal was formed by the gradual decomposition of plant vegetation under pressure and in the absence of air under sand. A time  known as the carboniferous Era. Carbon (iv) oxide, methane, and steam were liberated, leaving behind a material that contained a very high percentage of carbon.

During this process of carbonization, the vegetable material was converted in stages into several stages of coal namely

 

Types of Coal

There are 4 different types of coal namely:

(1) Peat-like coal: It contains about 60% of carbon by mass.

(2) Lignite coal (brown coal): It contains about 67% of carbon by mass.

(3) Anthracite coal (or hard coal): It is tough and hard. It contains about 94% of carbon by mass. Impurities present may include nitrogen, sulphur and phosphorus. Anthracite is the last stage of coal.

(4) Bituminous (soft) coal: These are use every day at home. It contains about 88% by mass of carbon.

 

Destructive Distillation of Coal

This is when coal is heated to a very high temperature in the absence ofair.

Yielding the following products

Coal            Coal gas   + Coal tar   Ammoniacal liquor  + Coke

 

Uses of coke

(i) Coke is mainly used as a fuel.

(ii) It is a very important industrial reducing agent and is used in the extraction of metals, especially iron, from their ores.

(iii) It is also used in the production of gaseous fuels, like water gas and producer gas.

(iv) It is used for the manufacture of graphite, calcium carbide, silicon carbide and carbon (iv) sulphide.

2. 

(a) Ammoniacal liquor: is a solution of NH₃ in water. It is used to make Fertilizers

(b) Coal tar :- it is used for road construction and also to produce other chemicals like toluene, phenol, benzene, naphthalene and anthracene which are used in the synthesis of important commercial product like dyes, paints, insecticides, drugs, plastics and explosives

Distillates of Coal	Uses
1.Ammoniacal liquor	To produce (NH4)2SO4 for fertilizer.
2.Coal tar	To produce useful chemicals such as disinfectants and perfumes
3.Coal gas	Used as industrial fuel.

 

Uses of coal

1.  Coal is used mainly as fuel to generate power for steam engines, factories and electrical plants.

2.  It is also used

 

FUEL GASES/GASIFICATION OF COKE

There are 3 types of fuel gases.

1.     Producer gas: Producer gas is a mixture of nitrogen and carbon (ii) oxide. It is prepared by passing a stream of air through red hot coke.

2C_(s)h   +  O_2(g)   +  N_2(g)           2CO_(g)     +     N_2(g)   +    Heat

Producer gas

2. Water gas: Water gas is a mixture of hydrogen and carbon (ii) oxide gas. It is prepared by passing steam over white hot coke.

H₂O_(g)    +        C_(s)                     CO_(g)      +       H_2(g)

Steam         white hot coke               Water gas

2.      Hydrogen gas:-water gas is then mixed with excess steam, and the mixture passed over iron (iii) oxide catalyst at 450⁰C.The carbon (ii) oxide decomposes the steam and the product are hydrogen and carbon (iv) oxide.

CO_(g)   +   H_2(g)      +    H₂O_(g)            CO_2(g)   +    2H_2(g)

 

Caustic soda or water is used to absorbed carbon (iv) oxide from the mixture. Ammoniacal copper (i) chloride can be used to remove unreacted carbon (ii) oxide. The final product is hydrogen.

 

Differences between Producer Gas and Water Gas

(1) Producer gas has a lower heating ability than water gas. ( because water gas consists of equal volumes of hydrogen  and carbon (ii) oxide both of which are combustible whereas producer gas consists of 33% combustible CO and 67% non-combustible N₂.

Water gas is an important industrial fuel and is used in the manufacture of hydrogen and other organic compounds e.g. methanol and butanol.

3.  Synthetic gas: It is a mixture of hydrogen and carbon (ii) oxide gas. It is prepared by mixing steam with methane (obtained as natural gas) and passing them over Nickel catalyst at about 800⁰C.

CH_4(g)     +      H₂O_(g)             CO_(g)   +   3H_2(g)

Synthetic gas is not a major source of air pollution because sulphur is removed in the gasification process/it does not contain sulphur or sulphur compounds.

 

CHEMICAL PROPERTIES OF CARBON

(1) Combustion:

(a) All forms of carbon burn in excess oxygen to produce carbon (iv) oxide gas.

C_(s)       +     O_2(g)          CO_2(g)          ( Complete combustion)

(b) All forms of carbon also burn in a limited supply of air to produce carbon (ii) oxide.

C_(s)      +     O_2(g)           CO_(g)               ( Incomplete combustion)

(2) Combination reaction: Carbon combines directly with certain elements such as Sulphur, Hydrogen, Calcium and Aluminium at very high temperatures.

C_(s)    +    2S_(s)                     CS_2(l)

Carbon (iv) sulphide

C_(s)   +    2H_2(g)                   CH_4(g)

Methane

2C_(s)   +     Ca_(s)                  CaC_2(s)

Calcium carbide

3C_(s)   +     4Al_(s)                   Al₄C_3(s)

Aluminium carbide.

(3) As a reducing agent: Carbon is a strong reducing agent. It reduces the oxides of the less active metals to the metals, while carbon is itself oxidized to either carbon (iv) oxide or carbon (ii) oxide, depending on the reaction conditions.

Fe₂O_3(s)   +   3C_(s)           2Fe_(s)     +       3CO_(g)

2CuO_(s)      +   C_(s)          2Cu_(s)      +        CO_2(g)

 

(4) Reaction with strong oxidizing agents: When carbon is heated with conc. HNO₃ or conc. H₂SO₄, it is oxidized to Carbon (iv) oxide.

C_(s)     +     4HNO_3(aqp          2H₂O_(l)    +      4NO_2(g)     +      CO_2(g)

C_(s)     +     2H₂SO_4(aq)             2H₂O_(l)    +      2SO_2(g)       +  CO_2(g)

 

 

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