Selected Compounds of Metals (Na, Mg, Ca, Al, Fe, Zn, Cu, and Pb) Metal Oxides
What Are Metal Oxides?
An oxide is a binary compound made up of oxygen and another element. Examples include MgO, PbO, Al₂O₃, H₂O₂, NO₂, and SO₂.
Metal oxides are compounds formed between oxygen and a metal, such as PbO, FeO, MgO, etc.
Note: Compounds like OF₂ are not called oxides of fluorine. Instead, they are called oxygen fluorides because fluorine is more electronegative than oxygen.
General Methods for Preparing Metal Oxides
There are two main ways to prepare metal oxides:
1. Direct Method
In this method, the metal reacts directly with oxygen, air, or another reagent to form the oxide.
- Burning the metal in air or oxygen: Many metals form oxides when heated strongly in air or pure oxygen.
- Reacting red-hot metal with steam: Some metals produce oxides when steam is passed over them while red hot.
- Reacting the metal with an oxidizing agent: For example, using concentrated nitric acid (HNO₃).
2. Indirect Method
Here, the oxide is obtained by heating other compounds of the metal, such as carbonates, hydroxides, or nitrates.
Example: Metal carbonate → Metal oxide + CO₂ (upon heating)
Types of Metal Oxides
Metal oxides can be classified into several types based on their chemical behavior:
- Basic Oxides
- Acidic Oxides
- Amphoteric Oxides
- Peroxides
- Superoxides
- Mixed Oxides
1. Basic Oxides
These oxides react with acids to form salt and water only. They can also combine with acidic oxides to form salts. Basic oxides are usually formed by metals in their lower oxidation states and may be ionic or covalent in nature.
2. Acidic Oxides
These are typically formed by metals in higher oxidation states. They are generally covalent and dissolve in water to form oxy-acids (they act as acid anhydrides).
3. Amphoteric Oxides
These oxides show both basic and acidic properties. They react with both acids and bases. Common examples: ZnO, Al₂O₃, BeO, PbO, SnO₂, Cr₂O₃.
- As bases: React with acids to form salt + water.
- As acids: React with bases to form salt + water.
Examples of reactions:
- With base → Sodium zincate or sodium aluminate, etc.
- With acid → Normal salt.
4. Peroxides
Peroxides contain the peroxide ion (O₂²⁻). Examples include Na₂O₂. They are mainly formed by alkali metals and alkaline earth metals.
Preparation:
- Heating the metal in excess oxygen/air.
- Heating the normal oxide (monoxide) in oxygen.
- Action of oxygen on the metal dissolved in liquid ammonia (for K, Rb, Cs).
- Reaction of H₂O₂ with a metal salt solution in alkaline medium.
Properties of Peroxides:
- Stability increases with the electropositive character of the metal.
- More stable when dry than in solution.
- Many form hydrates due to hydrogen bonding (e.g., Na₂O₂·8H₂O, CaO₂·8H₂O).
- React with water (ice-cold) to give alkaline solution + hydrogen peroxide.
- With dilute acids → H₂O₂.
- On heating → release O₂ and act as oxidizing agents.
Note: PbO₂ is not a peroxide because it does not liberate H₂O₂ with acids.
5. Superoxides
Superoxides contain the superoxide ion (O₂⁻). The common ones are KO₂, RbO₂, and CsO₂.
Preparation: Burning the metal (K, Rb, Cs) in excess oxygen or air.
Properties:
- Yellow-colored solids.
- Stability order: KO₂ < RbO₂ < CsO₂.
- Strong oxidizing agents.
- React with water to form H₂O₂ and O₂.
- Paramagnetic due to the presence of one unpaired electron in the O₂⁻ ion.
6. Mixed Oxides
These oxides consist of two different simple oxides combined together. The metals may be in different oxidation states.
Examples:
- Red lead (Pb₃O₄) = 2PbO + PbO₂
- Magnetite (Fe₃O₄) = FeO + Fe₂O₃
Other examples include magnesium ferrite (MgFe₂O₄) and zinc ferrite (ZnFe₂O₄).
Mixed Oxides (Continued)
Red Lead (Pb₃O₄) – Also known as Tri-lead tetroxide
- It is a brilliant scarlet (bright red) solid that is insoluble in water.
Chemical Behavior: Red lead behaves as if it is a loose combination of lead(II) oxide (PbO) and lead(IV) oxide (PbO₂), i.e., 2PbO + PbO₂.
For example, when warmed with dilute nitric acid (HNO₃), the PbO part reacts to form lead nitrate and water, while PbO₂ remains unreacted as a brown solid.
Note: This reaction is a redox reaction, showing that red lead can act as an oxidizing agent.
Uses:
- Used as a pigment in oil paints (especially for protective coatings on iron and steel).
Ferrous-Ferric Oxide (Fe₃O₄) – Also written as FeO·Fe₂O₃
- It occurs naturally as Magnetite.
- Can be prepared by heating iron with oxygen or steam.
Properties:
- Black in color.
- Strongly ferromagnetic (strongly attracted to magnets).
- Chemically quite unreactive (inactive).
- Reacts with acids as a mixed oxide, producing a mixture of ferrous (Fe²⁺) and ferric (Fe³⁺) salts in solution.
Hydroxides of Metals
Metal hydroxides are compounds containing the hydroxide ion (OH⁻) as the only negative ion. Examples: NaOH, Mg(OH)₂, Zn(OH)₂, Fe(OH)₃, etc.
Preparation of Metal Hydroxides
There are two main methods:
1. Direct Method
This works well for strongly electropositive metals (Group IA and some Group IIA). Examples: LiOH, NaOH, Ca(OH)₂.
2. Indirect Method
- Action of water on metal oxides: Metal oxide + Water → Metal hydroxide (This gives soluble hydroxides).
- Using milk of lime (Ca(OH)₂) on metal carbonates: Used in the preparation of NaOH and KOH. The mixture is allowed to settle, and the alkali solution is decanted.
- Precipitation method: Adding sodium hydroxide or ammonia solution to a solution of the metal salt to precipitate the hydroxide.
- Electrolysis: Electrolysis of brine (NaCl solution) is the industrial method for preparing NaOH (caustic soda).
Properties of Metal Hydroxides
A. Alkali Metal Hydroxides (Group IA – MOH)
Physical Properties:
- White crystalline solids.
- Melt at moderate temperatures without decomposition (except LiOH).
- Deliquescent (absorb moisture from air and become sticky).
- Highly soluble in water, forming strongly alkaline solutions.
Chemical Properties:
- Basic strength increases down the group (NaOH < KOH < RbOH < CsOH).
- React with chlorine differently depending on conditions:
- Cold, dilute alkali → Hypochlorite (bleach) + chloride.
- Hot, concentrated alkali → Chlorate + chloride. (Chlorine undergoes disproportionation in both cases).
- Absorb CO₂ from the air to form carbonates.
- React with acids in neutralization reactions to form salt + water only.
Uses of NaOH and KOH:
- Absorb acidic gases (e.g., CO₂).
- Used in neutralization and precipitation reactions.
- Caustic soda (NaOH): Manufacture of soap, paper, artificial silk, etc.
- Caustic potash (KOH): Manufacture of soft soaps.
B. Alkaline Earth Metal Hydroxides (Group IIA – M(OH)₂)
Physical Properties:
- White crystalline solids.
- Solubility in water increases down the group: Be(OH)₂ (almost insoluble) → Mg(OH)₂ (slightly soluble) → Ca(OH)₂ → Sr(OH)₂ → Ba(OH)₂ (more soluble).
- Solubility of Ca(OH)₂ decreases with increase in temperature.
- Be(OH)₂ is covalent due to the small size and high polarizing power of Be²⁺.
Chemical Properties:
- NaOH and KOH can precipitate many metals as their hydroxides from salt solutions.
- They liberate ammonia gas from ammonium salts.
- React with amphoteric metals (Zn, Al, Pb, Sn) to form complex ions: – Zincate, aluminate, plumbate, stannate.
Reaction with CO₂:
- Forms carbonate first.
- With excess CO₂ → soluble hydrogen carbonate.
Limewater Test: When CO₂ is bubbled through limewater [Ca(OH)₂], it turns milky due to CaCO₃. Excess CO₂ makes it clear again (formation of soluble Ca(HCO₃)₂).
Action with Heat:
- Thermal stability increases down the group. They decompose to oxide + water at higher temperatures.
Action with SO₂:
- Turns limewater milky (CaSO₃), excess SO₂ clears it (Ca(HSO₃)₂).
Uses:
- Limewater [Ca(OH)₂] – Test for CO₂.
- Milk of Magnesia [Mg(OH)₂] – Antacid.
- Slaked lime [Ca(OH)₂] – Making mortar, bleaching powder, whitewash, soil neutralization, water softening, and sugar refining.
-
Related posts:
Leave a Reply