Nature of Oxides-Oxides of s-Block Elements, Acidic and Basic Nature, Reactions, Practice Problems, and FAQ

You must have seen and admired a stunning superhero in a movie theatre at some point in your life. 

What if I asked you to create your own fictional superhero? 

You probably envision someone who is extremely versatile and reacts to an attacker on the spur of the moment. You might as well endow him with the qualities of a good Samaritan who is also a life-saving figure. 

You're probably wondering what all of this has to do with here, where we're all set to study oxides!

Actually, oxygen can be personified as one of the planet's superheroes. It is a life-giving element that also spontaneously reacts with almost all other elements (metals, metalloids, and nonmetals) to form oxides with diverse properties. Not only that, these oxides are superheroes in and of themselves. They have distinct characteristics and qualities, and they can be used for a variety of purposes that benefit humanity.

For now, we shall study a part of the oxides formed by the elements of the s-block in our periodic table and try to understand their nature. That might as well help you out to later innovate new utilities pertaining to these oxides! Stay Tuned!

TABLE OF CONTENTS

• Oxides of Group 1
• Nature of Oxides of Group 1
• Reactions of Alkali Metal Oxides with Water
• Oxides of Group 2
• Nature of Oxides of Group 2 and their Reaction with Water
• Acidic and Basic Nature of Oxides
• Order of Basic Nature Across a Period
• Order of Basic Nature Down a Group
• Practice Problems
• Frequently Asked Questions - FAQ

Oxides of Group 1

• The alkali metals tarnish in dry air due to the formation of their oxides. Lithium forms monoxide, sodium forms oxide as well as peroxide, and the other metals form oxide, peroxide as well as superoxides.

4Li (s)+ O₂(g) ⟶ 2Li₂O (s) [Oxide] 

4Na (s)+ O₂(g) ⟶ 2Na₂O (s) [Oxide] 

2Na (s)+ O₂(g) ⟶ Na₂O₂ (s) [Peroxide]

M(s) + O₂(g) ⟶ MO₂(s) [Superoxide]

[Where, M = K, Rb, Cs]

Sodium Oxide (Na2O)

• It is a white amorphous substance. It dissolves violently in water, yielding caustic soda (NaOH) with the evolution of a large amount of heat.

Na₂O + H₂O ₂NaOH

• Sodium oxide (Na₂O) can be prepared by the following reactions.

Na₂O₂(s)+ 2Na (s) 2Na₂O(s)

2NaNO₃ + 10Na 6Na₂O + N₂

2NaNO₂+ 6Na 4Na₂O + N₂

3NaN₃ + NaNO₂ 2Na₂O + 5N₂

Sodium Peroxide (Na2O2)

• Sodium peroxide is formed in a two-stage reaction in the presence of excess air.

2Na + O₂ Na₂O

Na₂O + O₂Na₂O₂

• Some important reactions undergone by sodium peroxide (Na2O2) are as follows: 
  

Potassium Superoxide (KO₂)

• It is prepared by burning potassium in excess of oxygen, free from moisture.
  

• Some important reactions undergone by potassium superoxide (KO₂) are as follows: 
  

Potassium Sesquioxide (K₂O₃).

• It is obtained when oxygen is passed through liquid ammonia containing potassium.
  

• K₂O₃ exists as (K⁺)₄^2- (O₂)2 (O₂).

Nature of Oxides of Group 1 (Alkali Metals)

• They are normally kept in kerosene oil because of their high reactivity towards air and water.

• As the size of the metal ions increases, the stability of superoxide or peroxides increases due to the stabilisation of large anions by larger cations through lattice energy effects.

• Li forms stable oxide (Li₂O), Na forms peroxide (Na₂O₂) and the rest of the metals form superoxides besides oxides and peroxide.

• Oxides of alkali metals are basic in nature and their basicity increases from Li to Cs as the ionic character increases.

• Peroxides and superoxides behave as strong oxidising agents. Superoxides on treatment with dilute acids form H₂O₂, O₂ and hydroxide.

Reactions of Alkali Metal Oxides with Water

• The oxides of the alkali metals react with moisture to form hydroxides and hydrogen gas. 

• Alkali metal peroxides react with water to form metal hydroxides and hydrogen peroxide.

• Superoxides of alkali metals react with water to form hydroxides, and hydrogen peroxide and liberate oxygen gas.

• Hydroxides of alkali metals are alkaline in nature and the basicity of hydroxides increases down the group. 

LiOH<NaOH<KOH<RbOH<CsOH

• The internuclear distance between the metal cation and the oxygen of the hydroxide group grows as the size of the metal cation grows. The hydroxide ion can ionise more easily as a result.

Example:

Li₂O (s) + 2H₂O (l)⟶ 2LiOH (aq)+ H₂(g) [Oxide + Water]

Na₂O₂(s)+ 2H₂O (l) ⟶ 2NaOH (aq)+ H₂O₂ (aq) [Peroxide + Water]2KO₂ (s)+ 2H₂O (l) ⟶ 2KOH (aq)+ H₂O₂ (aq)+O₂(g) [Superoxide + Water]

2RbO₂ (s)+ 2H₂O (l) ⟶ 2RbOH (aq)+ H₂O₂ (aq)+O₂(g) [Superoxide + Water]

2CsO₂ (s)+ 2H₂O (l) ⟶ 2CsOH (aq)+ H₂O₂ (aq)+O₂(g) [Superoxide + Water]

Oxides of Group 2

• Alkaline earth metals are less reactive than alkali metals. The reactivity of these elements increases down the group as the size increases.

• Be and Mg are inert to oxygen due to the formation of an oxide film on their surface. Be is purely inert in air as its surface is passivated by the formation of a thin layer of BeO.

• Mg and Ca also tarnish in air with the formation of an oxide layer, but will burn completely to their oxides and nitrides when heated. 

• To form oxides and nitrides, calcium, strontium, and barium are readily attacked by air. As magnesium is more electropositive, it burns intensely in air to give MgO and Mg₃N₂.

2Mg(s)+ O₂(g) 2MgO(s)

3Mg(s)+ N₂(g) Mg₃N₂(s)

• Magnesium burns so bright because the reaction releases a lot of heat. As a result of this exothermic reaction, magnesium gives two electrons to oxygen, forming powdery magnesium oxide (MgO).

• Powdered Be burns brilliantly to form beryllium oxide which is amphoteric in nature.
  

• All the group 2 elements form normal oxides with oxygen except Ba, which forms a peroxide due to its larger size and lesser polarising power.

Ba(s)+ O₂(g) BaO₂(s)

Nature of Oxides of Group 2 and their Reaction with Water

• Oxides of alkaline earth metals are basic in nature (except BeO). They react with water to form hydroxides.

MO(s) + H₂O (l) M(OH)₂ (aq) where M= Ca, Sr, Ba

• Since BeO and MgO are not soluble in water, they do not react with water, but CaO reacts with water (hydration energy overcomes lattice energy) exothermically to form calcium hydroxide.

CaO (s)+ H₂O (l) Ca(OH)₂(aq)

• Similarly, SrO and BaO react with water to form Sr(OH)₂ and Ba(OH)₂, respectively. The thermal stability of alkaline earth metal oxides increased down the group from Be to Ba.

• The basicity of alkaline metal hydroxides increases down the group.

Be(OH)₂ < Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂< Ba(OH)₂

• Alkaline earth metal hydroxides are less basic and less stable than alkali metal hydroxides due to more charge on alkaline earth metal cations.

• Beryllium hydroxide is amphoteric in nature (reacts with both acid and base), whereas the hydroxides of all other alkaline earth metals are basic in nature.

Be(OH)₂(aq)+ 2OH^- [Be(OH)₄]^2- (Berrylate Ion)

Be(OH)₂(aq)+ H₂SO₄ (l) BeSO₄ + 2H₂O(l)

• When BeO is heated with C at 1900−2000 °C, a brick red coloured carbide of formula Be₂C is formed. Ca, Ba and Sr also give this reaction to form their respective carbides.

MO (s) + 3C (s) MC₂ (s) + CO (g)

• These carbides, on being treated with water, produce hydroxide and hydrocarbons.

Be₂C (s) + 4H₂O (l) 2Be(OH)₂(s) + CH₄(g)

CaC₂(s) + 2H₂O (l) Ca(OH)₂(s) + C₂H₂(g)

Acidic and Basic Nature of Oxides

The oxides of metals are basic in nature (for example, oxides of magnesium and calcium) and they react with acids. The oxides of non-metals are acidic in nature (for example, oxides of sulphur) and they react with bases.

• Amphoteric oxides (ZnO, Cr₂O₃) can react with both acids and bases.
• Neutral oxides (CO, N₂O) do not react with acids and bases.

Order of Basic Nature Across a Period

As we move along the period, the non-metallic character increases due to increase in the ionisation energy and the metallic character decreases. Hence, across the period from left to right, the basic character of oxides decreases and the acidic character increases.

Order of Basic Nature Down a Group

As we move down a group, the metallic character increases and the non-metallic character decreases, and hence the basic character of oxides increases and the acidic character decreases.

Practice Problems

Q 1. What is observed when Mg burns in air?

a. Red fumes
b. White gas
c. Dazzling white flame 
d. No reaction

Answer: Magnesium burns so bright because the reaction releases a lot of heat. As a result of this exothermic reaction, magnesium gives two electrons to oxygen, forming powdery magnesium oxide (MgO). So, option C) is the correct answer.

Q 2. Which alkaline earth metal alone forms peroxide?

a. Barium
b. Magnesium
c. Beryllium
d. Calcium

Answer: All the group 2 elements form normal oxides with oxygen except Ba, which forms a peroxide due to its larger size and lesser polarising power. 

Ba(s)+ O₂(g) BaO₂(s)

So, option A) is the correct answer.

Q 3. What is the oxidation state of potassium in K₂O?

a. +1
b. -2
c. -1
d. +2

Answer: Potassium is in +1 in K₂O and hence the anion is O^2-. 

So, option B) is the correct answer.

Q 4. The formation of oxides is an exothermic reaction. True or False?

Answer: Alkali metals react with water to form basic hydroxides and give off hydrogen. The reaction of the metal is exothermic due to the liberation of the intense heat of hydration. So, the statement is true.

Frequently Asked Questions - FAQ

Q 1. Why the oxides of s-block elements are basic?
Answer: S-block elements show a valency of +1 and +2 only. They form strong ionic bonds with oxides and on dissolution with water, the oxides form metal hydroxide. These hydroxides completely dissociate to produce hydroxide ions. Thus, they are alkaline in nature. 

Q 2. What are amphoteric oxides?
Answer: Amphoteric oxides are metal oxides that react with both acids and bases to produce salts and water. Example: BeO

Q 3. Peroxides of only barium are formed among all the alkaline earth metals?
Answer: Barium peroxide can form because the Ba²⁺ ion can stabilise the large peroxide ion and hence forms a stable compound.

Q 4. What are Oxides?
Answer: The reaction of oxygen with other elements produces oxides, which are binary compounds. In nature, oxygen is extremely reactive. They form oxides when they react with metals and non-metals.

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