METAL NON METAL

Physical Properties of Metals

• Metallic Lustre: Metals have a shiny surface in their pure state, known as metallic lustre (e.g., gold, silver, platinum).

• Physical State: Metals are generally solids at room temperature.
  • Exception: Mercury is the only metal that is liquid at room temperature.

• Hardness: Most metals are hard, like iron and copper.
  • Exception: Metals like lithium, sodium, and potassium are soft and can be cut with a knife.

• Ductility: This refers to the ability of a metal to be drawn into thin wires. Metals are generally ductile, with gold being the most ductile.

• Malleability: This is the ability of a metal to be beaten into thin sheets. Gold and silver are the most malleable metals.

• Electrical Conductivity: This is the ability to conduct electricity. Most metals are good conductors. Silver is the best conductor, followed by copper, gold, aluminium, and tungsten.

• Electrical Conductivity: Metals conduct electricity due to free electrons.
  • Exception: Mercury is a poor conductor, while lead is almost non-conducting.
• Thermal Conductivity: Metals are good heat conductors (e.g., copper, silver).
  • Aluminum and copper are used in cookware for this reason.
  • Exception: Lead and mercury are poor heat conductors.
• Melting and Boiling Points: Metals usually have high melting and boiling points.
  • Exception: Gallium and caesium have low melting points and can melt in the hand.
• Sonority: Metals produce sound when struck and are called sonorous. This property makes them suitable for uses like school bells.

Chemical Properties of Metals

1. Reaction with Oxygen:

• Metals react with oxygen to form metal oxides.
• Example:

4Al (s)+3O₂ (g)→2Al₂O₃ (s)

This equation shows aluminium reacting with oxygen to form aluminium oxide.

Metal Oxides

1. Basic Nature:

• Metal oxides are generally basic.

2. Amphoteric Oxides:

• • Some metal oxides like aluminum oxide and zinc oxide can act as both acids and bases, making them amphoteric.
  • Example Reactions:

Al₂O₃ (s)+6HCl (aq)→2AlCl₃ (aq)+3H₂O (l)

Al₂O₃ (s)+2NaOH (aq)→2NaAlO₂ (aq)+H₂O (l)Al₂​O₃​ (s)

3. Solubility and Alkalis:

• • Metallic oxides are generally insoluble in water, but some form hydroxides (alkalis) when they dissolve.
  • Examples:

Na₂O (s)+H₂O (l)→2NaOH (aq)

K₂O (s)+H₂O (l)→2KOH (aq)K_2​O (s)+H₂​O (l)→2KOH (aq)

Note

• Alkalis are water-soluble bases that turn red litmus blue.

Order of Reactivity with Oxygen

• Highly Reactive Metals:
  • Sodium (Na) and Potassium (K): React vigorously with oxygen, often catching fire. Stored in kerosene to prevent accidents.
• Moderately Reactive Metals:
  • Magnesium (Mg) and Aluminium (Al): Require heating to burn in air; form a protective oxide layer.
  • Zinc (Zn): Burns only with strong heating.
  • Iron (Fe): Rods/blocks don’t burn easily, but iron filings burn vigorously in flame.

• Less Reactive Metals:
  • Copper (Cu): Does not burn but forms black copper(II) oxide on heating.
  • Silver (Ag) and Gold (Au): Do not react with oxygen, even at high temperatures.

Reactivity Order:

Na>Mg>Zn>Fe>Cu>Ag

Anodising

• Purpose: To form a thick oxide layer on aluminum for protection against further oxidation.
• Process:
  • Aluminum naturally forms a thin oxide layer when exposed to air.
  • Anodising enhances this layer by using an electric current.
  • The aluminum article is used as the anode, with dilute H₂SO₄ (sulfuric acid) as the electrolyte.
  • Oxygen gas is liberated at the anode, reacting with aluminum to thicken the oxide layer.

• Benefits:
  • Provides a protective, durable finish.
  • The oxide layer can be dyed for an attractive appearance.

Reaction of Metals with Water

1. General Reaction:
   • Metal + Water → Metal Oxide + Hydrogen Gas

• Metal Oxide + Water → Metal Hydroxide

2. Sodium (Na) and Potassium (K):

• • React violently with cold water.The reactions are very exothermic.
  • Potassium Reaction:

2K (s)+2H₂O (l)→2KOH (aq)+H₂ (g)

3. Sodium Reaction:

2Na (s)+2H₂O (l)→2NaOH (aq)+H₂ (g)

The heat produced can ignite the hydrogen gas, so these metals are stored in kerosene to prevent contact with air and water.

Reaction of Calcium with Water

1. Reaction:

Ca (s)+2H₂O (l)→Ca(OH)₂ (aq)+H₂ (g)

2. Characteristics:

• The reaction is less violent compared to sodium and potassium.
• The heat produced is not enough to ignite the hydrogen gas.
• Calcium floats on water because hydrogen gas bubbles stick to its surface.

Reaction of Magnesium with Water

1. With Hot Water:
   • Magnesium reacts with hot water to form magnesium hydroxide and hydrogen gas.

Mg (s)+2H₂O (l)→Mg(OH)₂ (aq)+H₂ (g)

2. With Steam:

• Magnesium reacts with steam to produce magnesium oxide and hydrogen gas.

Mg (s)+H₂O (g)→MgO (s)+H₂ (g)

3. Characteristics:

• Magnesium does not react with cold water.
• It floats on water due to hydrogen gas bubbles adhering to its surface.

Reaction with Steam

1. Aluminum (Al), Iron (Fe), and Zinc (Zn):

• Do not react with cold or hot water.
• React with steam to form metal oxides and hydrogen gas.

2. Aluminum Reaction:

2Al (s)+3H₂O (g)→Al₂O₃ (s)+3H₂ (g)

3. Iron Reaction:

3Fe (s)+4H₂O (g)→Fe₃O₄ (s)+4H₂ (g)

No Reaction with Water

• Lead (Pb), Copper (Cu), Silver (Ag), and Gold (Au):
  • Do not react with water at all.

Reactivity Order Towards Water

K>Na>Ca>Mg>Al>Fe>Pb>Cu>Ag>Au

• Reaction of Metals with Dilute HCl

• General Reaction:
  • Metal + Dilute Acid → Salt + Hydrogen Gas
• Examples:
  • Zinc (Zn):

Zn (s)+2HCl (aq)→ZnCl₂ (aq)+H₂ (g)

• Iron (Fe):

Fe (s)+2HCl (aq)→FeCl₂ (aq)+H₂ (g)

Reactivity Order:

• The rate of hydrogen gas formation decreases in the order:

Mg>Al>Zn>FeMg>Al>Zn>Fe

• This order indicates the decreasing reactivity of these metals with dilute HCl.
• Note:
  • Copper (Cu) does not react with dilute HCl.

Aqua-regia (Latin for “Royal Water”)

• Composition:
  • A mixture of concentrated hydrochloric acid and concentrated nitric acid in a ratio of 3:1.
• Properties:
  • Can dissolve gold and platinum, which neither acid can do alone.
  • Highly corrosive and fuming.
• Uses:
  • One of the few reagents capable of dissolving noble metals like gold and platinum.

Reaction of Metals with Solutions of Other Metal Salts

• Principle:
  • More reactive metals can displace less reactive metals from their compounds in solution or molten form.
• Application:
  • This principle is used in displacement reactions to extract or purify metals.

Displacement Reaction

1. Concept:

• If metal A is more reactive than metal B, it can displace B from its salt solution.

2. General Reaction:

Metal A + Salt Solution of B → Salt Solution of A + Metal B

3. Example:

• Copper (Cu) displacing silver (Ag) from silver nitrate:

Cu (s)+2AgNO₃ (aq)→Cu(NO₃)₂ (aq)+2Ag (s)

Type of Reaction:

• This is called a displacement reaction.

Reactivity Series of Metals

1. Definition:

• A list of metals arranged in order of decreasing reactivity.
• Based on their tendency to lose electrons and form cations.

2. Characteristics:

• More reactive metals have a greater tendency to lose electrons.

3. Classification:

• Most Reactive Metals:
  • Placed above hydrogen.
  • Examples: Potassium (K), Sodium (Na), Calcium (Ca).
• Least Reactive Metals:
  • Placed below hydrogen.
  • Examples: Noble metals like gold (Au) and platinum (Pt).

Hydrogen in the Reactivity Series

• Properties:
  • Hydrogen has non-metallic properties.
• Reason for Inclusion:
  • Due to its electropositive nature, hydrogen is placed in the reactivity series.

NON – METALS

Physical Properties of Non-Metals

• Lustre:
  • Non-metals generally do not have lustre.
  • Exception: Diamond, graphite (allotropic forms of carbon), and iodine have lustre.
• Physical State:
  • Non-metals are either solids or gases.
  • Exception: Bromine is liquid at room temperature.
• Softness:
  • Most non-metals are soft when solid.
  • Exception: Diamond is the hardest known substance.
• Malleability and Ductility:
  • Non-metals are neither malleable nor ductile.
• Brittleness:
  • Non-metals are brittle. They break into pieces when hammered.
• Electrical and Thermal Conductivity:
  • Non-metals are poor conductors of heat and electricity.
  • Exception: Graphite is a good conductor of electricity.
• Melting and Boiling Points:
  • Generally, non-metals have low melting and boiling points.
  • Solid non-metals have higher boiling points compared to gases. Diamond has a very high melting and boiling point.

Chemical Properties of Non-Metals

• Reactivity with Water and Acids:
  • Non-metals do not react with water, steam, or dilute acids to produce hydrogen gas.

• Reaction with Concentrated Acids:
  • When heated, non-metals readily form oxides or salts with concentrated acids.

• Examples:
  • Sulfur reacts with concentrated sulfuric acid:

S(s)+2H₂SO₄ (conc.)→Heat 3SO₂ (g)+2H₂O(l)

Sulfur reacts with concentrated nitric acid:

S(s)+6HNO₃ (conc.)→HeatH₂SO₄ (aq)+6NO₂ (g)+2H₂O(aq)

NO2​ is released as a reddish-brown vapor.

Displacement Reaction of Non-Metals

• Example:
  • Chlorine displaces bromine in sodium bromide:

Cl_2 (g)+2NaBr (l)→2NaCl (l)+Br₂ (g)

Chlorine gas reacts with liquid sodium bromide to produce liquid sodium chloride and bromine gas.

• Note:
  • Most non-metals produce acidic oxides when dissolved in water.

Ionic Bond Formation

• Reactivity of Elements:
  • Elements aim for a completely filled valence shell (2 or 8 electrons).
• Tendencies:
  • Metals tend to lose electrons to form cations (positive ions).
  • Non-metals tend to gain electrons to form anions (negative ions).
• Ionic Bonds:
  • Formed by the complete transfer of electrons from metals to non-metals.
  • This results in ionic or electrovalent compounds.
• Example:
  • Formation of Sodium Chloride (NaCl):
    • Sodium (Na) is a metal with electronic configuration 2,8,1
    • To achieve a stable configuration, sodium loses one electron from its outer shell, forming a cation (Na⁺).
    • The sodium atom then has 11 protons and 10 electrons.

• Sodium Chloride (NaCl):
  • Chlorine:
    • Electronic configuration: 2,8,7.
    • Gains one electron to complete its octet, forming a chloride ion (Cl⁻).
  • Sodium:
    • Loses one electron to form a sodium ion (Na⁺).
  • Ionic Bond:
    • Na⁺ and Cl⁻ attract each other through electrostatic forces, forming NaCl.
    • NaCl does not exist as discrete molecules but as aggregates of ions.

• Magnesium Chloride (MgCl₂​):
  • Magnesium:
    • Electronic configuration: 2,8,2.
    • Loses two electrons to form a magnesium ion (Mg²⁺)
  • Chlorine:
    • Gains one electron each, forming two chloride ions (Cl⁻).
  • Ionic Bond:
    • Mg^{2 +} attracts two Cl⁻ ions, forming MgCl₂​.

Properties of Ionic Compounds

• Physical Nature:
  • Ionic compounds are hard crystalline solids.
  • They have strong forces of attraction between positive and negative ions.
  • Generally brittle and break into pieces when pressure is applied.
• Melting and Boiling Points:
  • Ionic compounds have high melting and boiling points.
  • A large amount of energy is required to break the strong inter-ionic attractions.

• Solubility:
  • Ionic (electrovalent) compounds are soluble in water.
  • Insoluble in organic solvents like kerosene, benzene, ether, and petrol.

Conduction of Electricity:

• • Good conductors of electricity in molten form or in aqueous solution.In solution or molten form, ions move freely and conduct electricity.
  • Do not conduct electricity in solid form due to the rigid structure preventing ion movement.

Occurrence of Metals

• Sources:
  • The earth’s crust is a major source of metals.
  • Sea water contains soluble salts like sodium chloride and magnesium chloride.
• Minerals and Ores:
  • Elements or compounds that occur naturally in the earth’s crust are called minerals.
  • Minerals from which metals can be profitably extracted are called ores.

Extraction of Metals: Metallurgy

• Importance:
  • Metals are crucial in daily life.
  • Found in the earth’s crust in free state or as compounds (often oxides).

• Extraction Process:
  • The process of obtaining pure metal from ore is called extraction of metals or metallurgy.
  • The method depends on the nature of the ore, impurities, and metal to be extracted.

Extraction of Pure Metals

1. Concentration or Enrichment of Ore:
   1. Initial step to remove impurities and concentrate the ore.

• Classification by Reactivity:
  • Metals of High Reactivity:
    • Extraction by electrolysis of molten ore to obtain pure metal.
  • Metals of Medium Reactivity:
    • Carbonate Ore: Undergoes calcination to form metal oxide.
    • Sulphide Ore: Undergoes roasting to form metal oxide.
  • Metals of Low Reactivity:
    • Sulphide Ore: Undergoes roasting to form metal.

• Reduction to Metal:
  • Metal oxides are reduced to obtain the metal.

• Purification of Metal:
  • Final refining to achieve pure metal.

Common Steps in Metallurgical Operations

Refining of Impure Metal: To achieve purity.

Enrichment or Concentration of Ore: To increase the metal content.

Extraction from Concentrated Ores: To obtain crude metal.