Chemical Properties Of Metals

Metals exhibit a wide range of chemical properties that distinguish them from non-metals. These properties are a result of their tendency to lose electrons and form positive ions. Below is a detailed explanation of the chemical properties of metals along with examples, applications, and practice questions.

Reaction of Metals with Oxygen

Metals react with oxygen to form metal oxides. The reactivity of metals with oxygen varies, with some metals reacting vigorously and others reacting slowly. The metal oxides formed are generally basic in nature, but some metal oxides can be amphoteric (exhibiting both acidic and basic properties).

Examples:

Magnesium: When magnesium burns in oxygen, it produces magnesium oxide, which is basic in nature.
$\boldsymbol{2\textbf{Mg} + \textbf{O}_2 \rightarrow 2\textbf{MgO}}$
Magnesium oxide dissolves in water to form magnesium hydroxide, a basic solution.

Aluminium: Aluminium reacts with oxygen to form aluminium oxide, which is amphoteric. It can react with both acids and bases to form salts and water.
$\boldsymbol{4\textbf{Al} + 3\textbf{O}_2 \rightarrow 2\textbf{Al}_2\textbf{O}_3}$

Amphoteric Oxides:

Aluminium oxide ( $\boldsymbol{\textbf{Al}_2\textbf{O}_3}}$ ) and zinc oxide ( $\boldsymbol{\textbf{ZnO}}$ ) are examples of amphoteric oxides. They react with acids to form salts and water.
$\boldsymbol{\textbf{Al}_2\textbf{O}_3 + 6\textbf{HCl} \rightarrow 2\textbf{AlCl}_3 + 3\textbf{H}_2\textbf{O}}$

Aluminium oxide also reacts with bases like sodium hydroxide to form sodium aluminate:
$\boldsymbol{\textbf{Al}_2\textbf{O}_3 + 2\textbf{NaOH} \rightarrow 2\textbf{NaAlO}_2 + \textbf{H}_2\textbf{O}}$

Real-life Application:

Anodizing: Aluminium is anodized to form a protective layer of aluminium oxide on its surface, making it resistant to corrosion. This is widely used in the manufacturing of aluminium cookware, window frames, and aircraft parts.

Reaction of Metals with Water

Metals react with water to form metal oxides or metal hydroxides, depending on their reactivity. More reactive metals react with cold water, while less reactive metals only react with steam.

Examples:

Sodium and Potassium: These metals react vigorously with cold water, producing metal hydroxides and hydrogen gas.
$\boldsymbol{2\textbf{Na} + 2\textbf{H}_2\textbf{O} \rightarrow 2\textbf{NaOH} + \textbf{H}_2}$
The reaction is highly exothermic, and the hydrogen gas released catches fire.

Magnesium: Magnesium reacts slowly with cold water but reacts more vigorously with hot water to form magnesium hydroxide and hydrogen gas.
$\boldsymbol{\textbf{Mg} + 2\textbf{H}_2\textbf{O} \rightarrow \textbf{Mg(OH)}_2 + \textbf{H}_2}$

Iron: Iron does not react with cold water but reacts with steam to form iron oxide and hydrogen gas.
$\boldsymbol{3\textbf{Fe} + 4\textbf{H}_2\textbf{O} \rightarrow \textbf{Fe}_3\textbf{O}_4 + 4\textbf{H}_2}$

Reaction of Metals with Acids

Metals react with dilute acids to produce salts and hydrogen gas. The reactivity of metals with acids varies according to their position in the reactivity series. Metals like copper and silver do not react with dilute acids, while highly reactive metals like zinc and magnesium react vigorously.

Examples:

Zinc: Zinc reacts with dilute hydrochloric acid to form zinc chloride and hydrogen gas.
$\boldsymbol{\textbf{Zn} + 2\textbf{HCl} \rightarrow \textbf{ZnCl}_2 + \textbf{H}_2}$

Magnesium: Magnesium reacts rapidly with dilute sulfuric acid to form magnesium sulfate and hydrogen gas.
$\boldsymbol{\textbf{Mg} + \textbf{H}_2\textbf{SO}_4 \rightarrow \textbf{MgSO}_4 + \textbf{H}_2}$

Real-life Application:

Acid Cleaning: Reactivity with acids is used in cleaning metal surfaces. For example, zinc or iron parts are often treated with dilute acids to remove rust and corrosion before further processing.

Reaction of Metals with Other Metal Salts (Displacement Reactions)

A more reactive metal can displace a less reactive metal from its salt solution. This is known as a displacement reaction. These reactions help in determining the relative reactivity of different metals.

Example:

Copper and Iron: When an iron nail is placed in a solution of copper sulfate, the more reactive iron displaces copper from the solution and forms iron sulfate, while copper is deposited on the nail.
$\boldsymbol{\textbf{Fe} + \textbf{CuSO}_4 \rightarrow \textbf{FeSO}_4 + \textbf{Cu}}$
This principle is used in extracting metals and refining them from ores.

Reactivity Series

The Reactivity Series is a list of metals arranged in decreasing order of their reactivity. Highly reactive metals like potassium and sodium are placed at the top, while less reactive metals like silver and gold are placed at the bottom. The reactivity series helps predict how metals will react with oxygen, water, acids, and other compounds.

Metal	Reactivity
Potassium (K)	Very Reactive
Sodium (Na)	Reacts vigorously
Calcium (Ca)	Moderately Reactive
Magnesium (Mg)	Reacts with acids
Zinc (Zn)	Reacts with acids
Iron (Fe)	Reacts with steam
Lead (Pb)	Reacts slowly
Copper (Cu)	Unreactive
Silver (Ag)	Very Unreactive
Gold (Au)	Least Reactive

Application:

The reactivity series helps in extracting metals from their ores. Metals higher up in the reactivity series are extracted through electrolysis, while those lower down are extracted through reduction with carbon.

Practice Questions with Answers

Q1: What happens when magnesium reacts with dilute hydrochloric acid?

Answer: Magnesium reacts with hydrochloric acid to form magnesium chloride and hydrogen gas.
- Equation: $\boldsymbol{\textbf{Mg} + 2\textbf{HCl} \rightarrow \textbf{MgCl}_2 + \textbf{H}_2}$

Q2: Why does sodium need to be stored under kerosene?

Answer: Sodium is highly reactive with moisture and oxygen in the air. It must be stored under kerosene to prevent accidental fires.

Q3: Write the balanced equation for the reaction between zinc and sulfuric acid.

Answer: $\boldsymbol{\textbf{Zn} + \textbf{H}_2\textbf{SO}_4 \rightarrow \textbf{ZnSO}_4 + \textbf{H}_2}$

Q4: What is the product of the reaction between iron and steam?

Answer: Iron reacts with steam to form iron(III) oxide and hydrogen gas.
- Equation: $\boldsymbol{\textbf{Fe} + 4\textbf{H}_2\textbf{O} \rightarrow \textbf{Fe}_3\textbf{O}_4 + 4\textbf{H}_2}$