TOPIC: OXYGEN

Oxygen is the most abundant element on Earth. It belongs to Group VI of the Periodic Table. The Other elements in the Group are sulphur, selenium, tellurium, and polonium. There is an increase in Metallic properties from oxygen to polonium; hence, the elements in the oxygen family vary widely in their Chemical properties. The following trends are seen in the Group;

– All the elements have an outer electronic configuration s2,p4

– The atomic radius increases down the Group

– The melting and boiling points increase from oxygen to tellurium

– Electronegativity decreases down the Group

General properties of Group VI elements

1. They are chemically reactive
2. They have high electron affinity and tend to form covalent compounds with other non-metals. Only polonium, which is a metal, is an exception to this property.
3. The most common oxidation state is -2. However, +2, +4, and +6 oxidation states are possible for Sulphur, selenium, tellurium, and polonium.

The bonding capacity of oxygen

The electronic configuration of oxygen is 1s2, 2s2, 2p4 (2, 6). It requires two more electrons to completely Fill its valence shell. This is achieved by gaining or sharing electrons. When oxygen gains two electrons from metals, it forms ionic compounds e.g. Na2O, CaO, MgO etc. When it shares electrons with other

Non-metal, covalent compounds are formed e.g., O2, SO2, CO2, etc. Oxygen can also form a coordinate Bonding in molecules e.g in CO

Laboratory preparation of oxygen

1. By the reaction of hydrogen peroxide with acidified potassium tetraoxomanganate (VII) in the Cold. It is a redox reaction in which KMnO4 is the oxidizing agent and H2O2 is the reducing agent.

5H2O2(aq) + 2KMnO4(aq) + 3H2SO4(aq)      →       K2SO4(aq) + 2MnSO4(aq) + 8H2O(l) + 5O2(g)

2. By catalytic decomposition of potassium trioxochlorate (V). KClO3 is heated with the catalyst MnO2. The KClO3 decomposes to oxygen and potassium chloride; 2KClO3(s)heat   →      2KCl(s) + 3O2(g) In both cases, oxygen is collected over water.

3. By catalytic decomposition of hydrogen peroxide: add 10 volume or 3% hydrogen peroxide solution very slowly to a catalyst such as manganese (IV) oxide (MnO2). Oxygen gas will be liberated;

 “H2O2(aq)         →               O2(g) + 2H2O(l)

Industrial preparation of oxygen

Oxygen is prepared industrially by the fractional distillation of air. The following are the steps involved’

– Cold air is compressed to about 150atm

– The compressed air is cooled by refrigeration to separate CO2 and water vapour. The two gases become solids and are removed by filtration.

– Cold air is allowed to expand through a nozzle. The expansion causes air to cool rapidly.

– The compression and expansion cycles are repeated until the temperature reaches about -200°C, when liquid air is produced.

– The liquid air is warmed up and then fractionally distilled.

Physical properties of oxygen

1. it is slightly denser than air
2. It is a colourless and odourless gas
3. It does not affect moist litmus paper
4. It is sparingly soluble in water
5. It does not burn but support combustion
6. It is non-poisonous

Chemical properties of oxygen

1. Reaction with metals: The reaction of oxygen with metals depends on the reactivity of the metals.

– Alkali metals react readily with oxygen; hence, they are usually kept under oil e.g. sodium

4Na(s) + O2(g)      →        2Na2O(s)

– Alkali earth metals like magnesium burn vigorously in oxygen to form oxides

2Mg(s) + O2(g)    →          2MgO(s)

– Metals such as iron and copper react with oxygen, only when heated strongly, to form oxides, e.g. 2Fe(s) + 2O2 (g)  heat    Fe3O4(s)

2. Reaction with non-metals:

– Hydrogen burns in oxygen with pale blue flame to form water

2H2 (g) + O2(g)      →        2H2O(l)

– Carbon burns in a limited supply of oxygen to produce carbon (II) oxide gas;

2C(s) + O2(g)       →       2CO(g)

– In excess oxygen, carbon burns to produce carbon (IV) oxide;

C(s) + O2(g)      →        CO2(g)

– Sulphur burns with a bluish flame to produce sulphur (IV) oxide;

S(s) + O2(g)               SO2(g)

The reactions of oxygen with metals and non-metals are all oxidation reactions. Oxidation reactions include burning of fuels, rusting of iron, and respiration.

Binary compounds of oxygen

An oxide is a compound of oxygen and another element. They are grouped into five;

1. Acidic oxide
2. Basic oxide
3. Amphoteric oxide
4. Neutral oxide
5. Higher oxide

Acidic oxides: These are oxides of non-metals. They dissolve in water to form acids, e.g., SO2, CO2, SO3, NO2, P4O10.

SO2(g) + H2O(l)  →       H2SO3(aq)

Acidic oxides do not react with acids but react with alkalis to form a salt and water.

SO2(g) + 2NaOH(aq)              Na2SO3(aq) + H2O(l)

Silicon (IV) oxide, SiO2 is a solid at room temperature. It is an acidic oxide. It does not dissolve in water but reacts with sodium hydroxide to form sodium trioxosilicate (IV) and water.

SiO2(s) + 2NaOH(aq)     →         Na2SiO3(aq) + H2O(l)

Properties of acidic oxides

– They are usually gases at room temperature (except P4O10 and SiO2)

– If soluble, they dissolve in water to form an acidic solution

– They react with alkalis or bases to form a salt and water only.

Basic oxide: they are oxides of metals. Most of them are insoluble in water. A few of them, such as Na2O and K2O dissolve readily in water. These soluble oxides are called alkalis. Calcium oxide (quick lime)dissolves moderately in water (a vigorous reaction) to form lime water.

Basic oxides are solids at room temperature. They react with acids to form salts and water only, e.g.

CaO(s) + 2HNO3(aq)               Ca(NO3)2(aq) + H2O(l)

Amphoteric oxide: these are metallic oxides that react with both acids and bases to form salts and Water only e.g. zinc oxide-ZnO, aluminium-Al2O3 and lead (II) oxide – PbO.

Neutral oxides: Some non-metals form oxides that show neither basic nor acidic properties. They are called neutral oxides. They are insoluble in water (except water itself). Examples are water, carbon (II) Oxide and nitrogen (II) oxide.

Higher oxides: Higher oxides are oxides containing a higher proportion of oxygen than the ordinary oxides. Example Perioxides (H2O2, K2O2, CaO2, BaO2), Dioxides (PbO2, MnO2) and Mixed oxides (Pb3O4, Fe3O4)

Uses of oxygen

Oxygen is used:

1. In steel making industry to burn off or oxidize impurities such as carbon, sulphur, and phosphorus.
2. In the hospital to help patients with breathing difficulties
3. By mountain climbers, pilots of high-flying fighter planes and deep sea divers to help them breathe at high altitude and under water.
4. In the oxy-acetylene flame (about 35000C) to cut metals or weld them together.
5. In the manufacture of important chemical compounds like tetraoxosulphate (VI) acids, trioxonitrate (V) acid and ethanoic acid
6. Liquid oxygen and fuels are used as propellants for space rockets.

Test for oxygen

When a glowing splinter is inserted into the jar containing the unknown gas, if the gas is oxygen, the splinter will be rekindled.

Assignment

Dinitrogen (I) oxide behaves like oxygen when brought close to glowing splinter. Highlight the

Differences between oxygen and dinitrogen (I) oxide.

Tutorial questions

1. State the steps involved in the industrial preparation of oxygen
2. Give any three uses of oxygen
3. With two examples each, briefly describe the types of oxides
4. Give four physical properties of oxygen
5. With the aid of a balanced equation, show the reaction of oxygen with a metal and a non-metal
6. Describe one way by which oxygen can be produced in the laboratory