Comprehensive Study Notes on Metals and Non-metals

Physical Properties of Metals

Metallic Lustre: Metals in their pure state have a surface that shines. This specific property is defined as metallic lustre.
Hardness: Metals are generally hard in nature, though this hardness varies significantly from one metal to another.
- Activity 3.2 Observation: While metals like iron ( $Fe$ ), copper ( $Cu$ ), and magnesium ( $Mg$ ) are hard, sodium ( $Na$ ) is soft enough to be cut with a knife.
- Alkali Metals: Lithium ( $Li$ ), sodium ( $Na$ ), and potassium ( $K$ ) are so soft they can be sliced with a knife. They also possess low densities and low melting points.
Malleability: This is the property that allows some metals to be beaten into thin sheets.
- Gold and Silver: These are recognized as the most malleable metals.
Ductility: The ability of metals to be drawn into thin wires is known as ductility.
- Gold Specifics: Gold is the most ductile metal. A wire of approximately $2\,km$ in length can be drawn from just $1\,g$ of gold.
Thermal Conductivity: Metals are generally good conductors of heat and possess high melting points.
- Best Conductors: Silver ( $Ag$ ) and copper ( $Cu$ ) are the superior conductors of heat.
- Poor Conductors: Lead ( $Pb$ ) and mercury ( $Hg$ ) are comparatively poor conductors of heat.
- Low Melting Point Exceptions: Gallium ( $Ga$ ) and caesium ( $Cs$ ) have exceptionally low melting points and will melt if held on a human palm.
Electrical Conductivity: Metals are good conductors of electricity.
- Electric wires are typically coated with polyvinylchloride (PVC) or a rubber-like material for insulation and safety.
Sonority: Metals that produce a ringing sound upon striking a hard surface are described as sonorous. This property explains why school bells are manufactured from metal.

Physical Properties of Non-metals and Exceptions

General State: Non-metals are fewer in number compared to metals. Examples include carbon ( $C$ ), sulphur ( $S$ ), iodine ( $I$ ), oxygen ( $O$ ), and hydrogen ( $H$ ). They are mostly solids or gases.
Bromine ( $Br$ ): This is the only non-metal that exists as a liquid at room temperature.
Physical Characteristics: Non-metals generally do not share the malleable, ductile, or sonorous properties of metals.
Exceptions to Classification Rules:
- Mercury ( $Hg$ ): The only metal that is liquid at room temperature.
- Iodine ( $I$ ): A non-metal that possesses a lustrous (shiny) surface.
- Allotropes of Carbon: Carbon exists in different forms called allotropes.
  - Diamond: An allotrope of carbon that is the hardest known natural substance. It has an extremely high melting and boiling point.
  - Graphite: An allotrope of carbon that is a conductor of electricity.

Chemical Properties of Metals: Reaction with Oxygen

General Reaction: Almost all metals combine with oxygen when heated to form metal oxides. These are generally basic in nature.
- $\text{Metal} + \text{Oxygen} \rightarrow \text{Metal oxide}$
Specific Examples:
- Copper: When copper is heated in air, it forms copper(II) oxide (a black oxide).
  - $2Cu + O_2 \rightarrow 2CuO$
- Aluminium: Forms aluminium oxide.
  - $4Al + 3O_2 \rightarrow 2Al_2O_3$
Amphoteric Oxides: Some metal oxides react with both acids and bases to produce salts and water. These are called amphoteric oxides.
- Aluminium Oxide with Acid: $Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O$
- Aluminium Oxide with Base: $Al_2O_3 + 2NaOH \rightarrow 2NaAlO_2 + H_2O$ (Product: Sodium aluminate).
Solubility and Alkalis: Most metal oxides are insoluble in water, but some dissolve to form alkalis.
- Sodium Oxide: $Na_2O(s) + H_2O(l) \rightarrow 2NaOH(aq)$
- Potassium Oxide: $K_2O(s) + H_2O(l) \rightarrow 2KOH(aq)$
Reactivity Variations:
- Sodium ( $Na$ ) and Potassium ( $K$ ): React so vigorously they catch fire in the open. They are stored immersed in kerosene oil to prevent accidental fires.
- Magnesium ( $Mg$ ), Aluminium ( $Al$ ), Zinc ( $Zn$ ), Lead ( $Pb$ ): At ordinary temperatures, these are covered with a thin protective layer of oxide that prevents further oxidation.
- Iron ( $Fe$ ): Does not burn on heating, but iron filings burn vigorously when sprinkled into a flame.
- Copper ( $Cu$ ): Does not burn but gets coated with black copper(II) oxide.
- Silver ( $Ag$ ) and Gold ( $Au$ ): Do not react with oxygen even at high temperatures.

Anodising

Definition: Anodising is the process of forming a thick oxide layer of aluminium.
Process: A clean aluminium article is made the anode and electrolysed with dilute sulphuric acid ( $H_2SO_4$ ). The oxygen gas evolved at the anode reacts with the aluminium to create a thicker, protective oxide layer.
Uses: This layer makes the metal resistant to further corrosion and can be dyed for attractive finishes.

Chemical Properties of Metals: Reaction with Water

General Reaction:
- $\text{Metal} + \text{Water} \rightarrow \text{Metal oxide} + \text{Hydrogen}$
- $\text{Metal oxide} + \text{Water} \rightarrow \text{Metal hydroxide}$
Specific Metal Behaviors:
- Potassium and Sodium: React violently with cold water. The reaction is highly exothermic, and the hydrogen ( $H_2$ ) catches fire immediately.
  - $2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g) + \text{heat energy}$
- Calcium ( $Ca$ ): Reaction is less violent. The metal floats because bubbles of hydrogen stick to its surface.
  - $Ca(s) + 2H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g)$
- Magnesium ( $Mg$ ): Does not react with cold water; reacts only with hot water to form magnesium hydroxide. It also floats due to hydrogen bubbles.
- Aluminium ( $Al$ ), Iron ( $Fe$ ), and Zinc ( $Zn$ ): Do not react with cold or hot water. They react only with steam to form the metal oxide.
  - $2Al(s) + 3H_2O(g) \rightarrow Al_2O_3(s) + 3H_2(g)$
  - $3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$
- Lead ( $Pb$ ), Copper ( $Cu$ ), Silver ( $Ag$ ), Gold ( $Au$ ): Do not react with water at all.

Chemical Properties of Metals: Reaction with Acids

General Reaction: $\text{Metal} + \text{Dilute acid} \rightarrow \text{Salt} + \text{Hydrogen}$
Reactivity Order: Based on the rate of bubble formation and temperature change, the reactivity order is Mg > Al > Zn > Fe.
Reaction with Nitric Acid ( $HNO_3$ ): Hydrogen gas is generally not evolved. $HNO_3$ is a strong oxidising agent that oxidises the produced $H_2$ to water and is itself reduced to nitrogen oxides ( $N_2O$ , $NO$ , or $NO_2$ ).
- Exception: Magnesium ( $Mg$ ) and Manganese ( $Mn$ ) react with very dilute $HNO_3$ to evolve $H_2$ gas.
Aqua Regia: A freshly prepared mixture of concentrated hydrochloric acid ( $HCl$ ) and concentrated nitric acid ( $HNO_3$ ) in a ratio of $3:1$ . Known as "royal water," it is highly corrosive and can dissolve gold and platinum.

The Reactivity Series and Displacement Reactions

Displacement Reaction: A more reactive metal can displace a less reactive metal from its salt solution.
- $\text{Metal A} + \text{Salt solution of B} \rightarrow \text{Salt solution of A} + \text{Metal B}$
Reactivity Series Table (Decreasing Order):
1. Potassium ( $K$ ) - Most reactive
2. Sodium ( $Na$ )
3. Calcium ( $Ca$ )
4. Magnesium ( $Mg$ )
5. Aluminium ( $Al$ )
6. Zinc ( $Zn$ )
7. Iron ( $Fe$ )
8. Lead ( $Pb$ )
9. [Hydrogen ( $H$ )]
10. Copper ( $Cu$ )
11. Mercury ( $Hg$ )
12. Silver ( $Ag$ )
13. Gold ( $Au$ ) - Least reactive

How Metals and Non-metals React: Ionic Compounds

Noble Gas Stability: Noble gases like Helium ( $He$ , config: $2$ ), Neon ( $Ne$ , config: $2, 8$ ), and Argon ( $Ar$ , config: $2, 8, 8$ ) have completely filled valence shells and are chemically inactive.
Reactivity Logic: Elements react to attain a completely filled valence shell (a stable octet).
Ionic Bonding: This involves the transfer of electrons from a metal (which becomes a cation) to a non-metal (which becomes an anion).
- Formation of Sodium Chloride ( $NaCl$ ):
  - Sodium ( $Na, 2, 8, 1$ ) loses one electron to become $Na^+ (2, 8)$ .
  - Chlorine ( $Cl, 2, 8, 7$ ) gains one electron to become $Cl^- (2, 8, 8)$ .
  - Oppositely charged ions are held by strong electrostatic forces.
- Formation of Magnesium Chloride ( $MgCl_2$ ):
  - Magnesium ( $Mg, 2, 8, 2$ ) loses two electrons to become $Mg^{2+} (2, 8)$ .
  - Two Chlorine atoms each gain one electron.
Definition: Compounds formed by the transfer of electrons from a metal to a non-metal are ionic or electrovalent compounds.

Properties of Ionic Compounds

Physical Nature: They are solids and generally hard due to strong inter-ionic attraction. They are brittle and break into pieces under pressure.
Melting and Boiling Points: They have very high melting and boiling points because significant energy is needed to break the strong inter-ionic bonds.
- Example: $NaCl$ has a melting point of $1074\,K$ and boiling point of $1686\,K$ .
Solubility: Generally soluble in water but insoluble in organic solvents like kerosene or petrol.
Electrical Conduction:
- Solid State: Do not conduct electricity because the rigid structure prevents ion movement.
- Molten/Solution State: Conduct electricity as the ions are free to move toward opposite electrodes.

Occurrence and Extraction of Metals

Minerals: Elements or compounds occurring naturally in the earth's crust.
Ores: Minerals from which a metal can be extracted profitably.
Gangue: Large amounts of impurities like soil and sand found in mined ores.
Extraction Based on Reactivity:
- Low Reactivity ( $Hg, Cu$ ): Oxides can be reduced by heating alone.
  - Cinnabar ( $HgS$ ): $2HgS + 3O_2 \xrightarrow{\text{Heat}} 2HgO + 2SO_2 \rightarrow 2HgO \xrightarrow{\text{Heat}} 2Hg + O_2$
- Medium Reactivity ( $Zn, Fe, Pb, Cu$ ): Found as sulphides or carbonates. Must be converted to oxides first.
  - Roasting: Heating sulphide ores strongly in excess air.
    - $2ZnS + 3O_2 \xrightarrow{\text{Heat}} 2ZnO + 2SO_2$
  - Calcination: Heating carbonate ores strongly in limited air.
    - $ZnCO_3 \xrightarrow{\text{Heat}} ZnO + CO_2$
  - Reduction: Metal oxides are reduced using carbon (coke) or reactive metals (like $Al$ ).
  - Thermit Reaction: The reaction of iron(III) oxide ( $Fe_2O_3$ ) with aluminium powder produce molten iron; used to join railway tracks.
    - $Fe_2O_3(s) + 2Al(s) \rightarrow 2Fe(l) + Al_2O_3(s) + \text{Heat}$
- High Reactivity ( $K, Na, Ca, Mg, Al$ ): Extracted via electrolytic reduction of molten chlorides (for $Na, Mg, Ca$ ) or oxides (for $Al$ ).
  - Sodium is deposited at the cathode (negative electrode), Chlorine is liberated at the anode (positive electrode).

Refining of Metals

Electrolytic Refining: The most common method to obtain pure metal.
- Anode: A thick block of impure metal.
- Cathode: A thin strip of pure metal.
- Electrolyte: A solution of the metal's salt.
- Mechanism: On passing current, pure metal from the anode dissolves into the electrolyte and is deposited on the cathode. Insoluble impurities settle at the bottom as anode mud.

Corrosion and Prevention

Corrosion Examples:
- Silver: Reacts with sulphur in air to form black silver sulphide.
- Copper: Reacts with moist carbon dioxide to form a green coat of basic copper carbonate.
- Iron: Reacts with moist air to form brown flaky rust.
Prevention Methods: Painting, oiling, greasing, galvanising, chrome plating, anodising, or alloying.
Galvanisation: Coating steel or iron with a thin layer of zinc ( $Zn$ ). The zinc protects even if the coating is scratched.
Alloying: Mixing a metal with other metals or non-metals to change properties.
- Steel: Iron mixed with carbon ( $0.05\,\%$ ) becomes hard.
- Stainless Steel: Iron mixed with nickel and chromium; it is hard and does not rust.
- Brass: Alloy of Copper ( $Cu$ ) and Zinc ( $Zn$ ).
- Bronze: Alloy of Copper ( $Cu$ ) and Tin ( $Sn$ ).
- Solder: Alloy of Lead ( $Pb$ ) and Tin ( $Sn$ ); has a low melting point for welding wires.
- Amalgam: An alloy where one of the metals is mercury ( $Hg$ ).
- Gold Purity: Pure gold (24 carat) is too soft; typically alloyed with copper or silver to 22 carat (22 parts gold, 2 parts alloying metal) for jewellery.
Ancient Metallurgy: The iron pillar near Qutub Minar in Delhi, over $1600$ years old, weighs $6\,tonnes$ and is $8\,m$ high; it remains rust-free due to advanced ancient processing techniques.