CBSE10th Science: CH03 Metals and Non-Metals || Best Revision Notes || By Sanjay Sir

Metals and Non-Metals

Metals:

  • They are good conductors of heat and electricity.
  • They have a shiny or metallic luster.
  • They are generally solid at room temperature (except for mercury).
  • They are malleable, which means they can be hammered into different shapes.
  • They are ductile, which means they can be drawn into wires.

Non-Metals:

  • They are generally poor conductors of heat and electricity.
  • They lack the shiny luster that metals have.
  • They can be in different states at room temperature (e.g., hydrogen and oxygen are gases, sulfur is a solid, and nitrogen is a gas).
  • They are typically brittle and cannot be hammered into thin sheets or drawn into wires.

Properties of Metals:

  1. Good Conductors: Metals are excellent conductors of heat and electricity.
  2. Luster: They have a shiny, metallic luster.
  3. Malleability: Metals can be hammered into thin sheets without breaking.
  4. Ductility: They can be drawn into thin wires.
  5. Solid State: Most metals are solid at room temperature (except mercury).
  6. High Melting and Boiling Points: Metals generally have high melting and boiling points.
  7. Sonorous: They produce a ringing sound when struck.
  8. Tendency to Lose Electrons: Metals tend to lose electrons and form positively charged ions (cations) in chemical reactions.

Properties of Non-Metals:

  1. Poor Conductors: Non-metals are poor conductors of heat and electricity.
  2. Dull Appearance: They lack the shiny luster of metals and may appear dull.
  3. Various States: Non-metals can exist in different states at room temperature (e.g., hydrogen and oxygen are gases, sulfur is a solid, and nitrogen is a gas).
  4. Brittle: They are typically brittle and cannot be drawn into thin wires or hammered into sheets.
  5. Low Melting and Boiling Points: Non-metals generally have low melting and boiling points.
  6. Not Sonorous: They do not produce a ringing sound when struck.
  7. Tendency to Gain Electrons: Non-metals tend to gain electrons and form negatively charged ions (anions) in chemical reactions.

1. Reaction with Oxygen:

  • Most metals react with oxygen to form metal oxides. This process is known as oxidation.
  • Example: 2Hg(l) + O2(g) → 2HgO(s) (Formation of mercury oxide)

Amphoteric Oxides:

Amphoteric oxides are certain chemical compounds that exhibit dual properties, acting as both acids and bases in different reactions.
Example (Acidic Reaction):
  • Aluminum oxide (Al2O3) reacts with hydrochloric acid (HCl) to form aluminum chloride (AlCl3) and water (H2O):
  • Al2O3 + 6HCl → 2AlCl3 + 3H2O
Example (Basic Reaction):
  • Zinc oxide (ZnO) reacts with sodium hydroxide (NaOH) to produce sodium zincate (Na2ZnO2) and water (H2O):
  • ZnO + 2NaOH → Na2ZnO2 + H2O

2. Reaction with Water:

  • Some metals react with water to produce metal hydroxides and hydrogen gas.
  • Example: 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) (Formation of sodium hydroxide and hydrogen gas)

3. Reaction with Dilute Acids:

  • Many metals react with dilute acids, such as hydrochloric acid, to form metal salts and release hydrogen gas.
  • Example: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) (Formation of zinc chloride and hydrogen gas)

Aqua Regia

Aqua Regia, derived from the Latin term "Royal Water," is a potent and highly corrosive mixture consisting of concentrated nitric acid (HNO3) and concentrated hydrochloric acid (HCl) in a 1:3 ratio.

4. Reaction with Bases:

  • Metals can react with bases to form metal salts and hydrogen gas.
  • Example: 2Na(s) + 2NaOH(aq) → 2Na2O(s) + H2(g) (Formation of sodium oxide and hydrogen gas)

5. Reaction with Salt Solutions:

  • Metals can displace less reactive metals from their salt solutions in a process known as the displacement reaction.
  • Example: Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s) (Copper displacing silver from silver nitrate)

Que. Why does Metal Reacts with Non-Metal?

Metals react with non-metals to form ionic bonds. This happens because metals have a tendency to lose electrons, becoming positively charged ions (cations), while non-metals tend to gain electrons, resulting in negatively charged ions (anions).

The strong attraction between these oppositely charged ions leads to the creation of ionic compounds.


Formation of Sodium Chloride (NaCl)

Formation of Sodium Chloride (NaCl):

Sodium chloride (table salt) is formed through the combination of sodium (Na), a metal, and chlorine (Cl), a non-metal. The reaction involves the transfer of electrons, resulting in the formation of an ionic compound.

Equation:

2Na(s) + Cl2(g) → 2NaCl(s)

In this reaction, sodium loses an electron to become a positively charged ion (Na+), and chlorine gains an electron to become a negatively charged ion (Cl-). These oppositely charged ions are attracted to each other, forming the ionic compound sodium chloride (NaCl).


Formation of Magnesium Chloride (MgCl2)

Formation of Magnesium Chloride (MgCl2):

Magnesium chloride is formed similarly, with magnesium (Mg) and chlorine (Cl) reacting to create an ionic compound.

Equation:

Mg(s) + Cl2(g) → MgCl2(s)

In this reaction, magnesium loses two electrons to become a doubly positively charged ion (Mg2+), while chlorine gains two electrons to become doubly negatively charged ions (Cl2-). The electrostatic attraction between these oppositely charged ions results in the formation of magnesium chloride (MgCl2).


Reactivity Series of Common Metals

  • Potassium (K)
  • Sodium (Na)
  • Calcium (Ca)
  • Magnesium (Mg)
  • Aluminum (Al)
  • Zinc (Zn)
  • Iron (Fe)
  • Lead (Pb)
  • Hydrogen (H)
  • Copper (Cu)
  • Silver (Ag)
  • Gold (Au)


Properties of Ionic Compounds

  1. Physical Nature:
    • Ionic compounds exist in the solid state at room temperature and have a characteristic crystal lattice structure.
    • The arrangement of positively charged cations and negatively charged anions in a repeating pattern forms this crystalline structure.
  2. Melting and Boiling Points:
    • Ionic compounds typically have high melting and boiling points.
    • This is because the electrostatic forces of attraction between oppositely charged ions in the crystal lattice are very strong.
    • It takes a significant amount of energy to break these ionic bonds and transition the compound from a solid to a liquid or gas state.
  3. Solubility:
    • Many ionic compounds are soluble in polar solvents, such as water.
    • When an ionic compound dissolves in water, the water molecules surround and separate the individual ions, allowing them to disperse evenly throughout the solution.
    • The solubility of ionic compounds can vary, with some being highly soluble, and others only sparingly so.
  4. Conduction of Electricity:
    • In the solid state, ionic compounds do not conduct electricity because the ions are fixed in place within the crystal lattice and cannot move.
    • However, when dissolved in water (or molten), ionic compounds can conduct electricity.
    • This is because the ions are now free to move in the solution, carrying electric charge with them.