Lewis Structure Of Oxygen Essay

1081 words - 5 pages

Oxygen
Lewis structure of oxygen

Preparation of Oxygen
a. Thermal decomposition of potassium chlorate(V)

2 KClO3 ( s )  2 KCl ( s ) + 3O2 ( g ) →
Drying agent: CaCl2, CaO, concentrated H2SO4

heat ,300ο C MnO2

b. Catalytic decomposition of hydrogen peroxide solution

2 H 2O2 (aq)  2 H 2O(l ) + O2 ( g ) →
c. Reaction between sodium peroxide and water 2Na2O2(s) + 2 H2O (l) → 4NaOH (aq) + O2 (aq)

MnO2

Other substances that produce O2
2 F2 ( g ) + 2 H 2O(l )  4 HF (aq) + 2O2 ( g ) → 2 HClO(aq) → 2 HCl (aq) + O2 ( g )
heat 2 KNO3 ( s ) → 2 KNO2 ( s ) + O2 ( g )  sunlight

2 AgNO3 ( s ) → 2 Ag ( s ) + 2 NO2 ( g ) + O2 ( g ) 

heat

Industrial ...view middle of the document...

superoxides
• Very reactive alkali metals -K, Rb and Cs can also form superoxides (O2− ion) • Potassium burns in excess oxygen with a lilac flame to form a higher oxide, KO2 . • K(s) + O2 (g) → KO2 (s) (potassium superoxide) • Oxidation state of oxygen in superoxides is

1 − 2

nonReaction with non-metals
• Nonmetals burn in oxygen to form acidic oxides. • Example: 1 S (s) + O2 (g) → SO2 (g) 2 P4 (s) + 3 O2 (g) → P4O6 (s)

Reaction of oxygen with
a. hydrogen sulphide b. Methane c. ammonia
ignited 2 H 2 S ( g ) + 3O2 ( g )  2SO2 ( g ) + 2 H 2O(l ) →

CH 4 ( g ) + 2O2 ( g )  CO2 ( g ) + 2 H 2O(l ) → 4 NH 3 ( g ) + 5O2 ( g )  4 NO( g ) + 6 H 2O(l ) →
pt − Rh ,850ο C

Test for oxygen
• Rekindle a glowing splinter. • Dinitrogen monoxide also does this. • it can be distinguished from oxygen by reactions with NO NO + O2 (g) → NO2 Produces reddish-brown fumes of NO2 • N2O + O2 (g) → No reaction

Classification of Oxides
Basic oxides Metal oxides oxides Nonmetal oxides Acidic oxides Neutral oxides Amphoteric oxides

Basic Oxides
• most metal oxides are ionic compounds. • Example: CaO , BaO, K2O , Na2O …. • dissolve in water to form hydroxides. K2O(s) + H2O(l) → 2KOH (aq) • metallic oxide reacts with acid to produce a salt and water only. MgO(s) + 2HCl(aq) → MgCl2 (aq) + H2O(l)

Acidic oxides
• non-metal oxides are covalent compounds Example: CO2, SO3 , SO2 ,P4O10…. • dissolved in water to give acid • also known as acid anhydrides. • SO2 (g) + H2O(l) → H2SO3 (aq) • SO3 (g) + H2O(l) → H2SO4 (aq) • P4O6 (g) + 6 H2O(l) → 4 H3PO3 (aq) • P4O10 (g) + 6 H2O(l) → 4 H3PO4 (aq)

Amphoteric oxides
• a metallic oxide which can show both basic and acidic properties • can react with both acid and alkali to produce a salt and water only • Example: ZnO , PbO, SnO, Al2O3 , …. • For example, aluminium oxide reacts with hydrochloric acid to form aluminium chloride Al2O3 + 6 HCl → 2 AICl3(aq) + 3 H2O (aq)

Amphoteric oxides with alkali
• aluminium oxide reacts with sodium hydroxide to form sodium aluminate: Al2O3 + 2KOH + 3 H2O → 2 KAI(OH)4 (aq) Or Al2O3 + 2KOH → 2 KAIO2 (aq) + H2O • ZnO + 2NaOH → Na2ZnO2 (aq) + H2O or ZnO + 2 NaOH + H2O → Na2Zn(OH)4 (aq)

Neutral...

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