Electron redistribution in chemical bonds

Metal and Non-metals of Class 8

ELECTRON REDISTRIBUTION IN CHEMICAL BONDS

Chemical Bonds can be classified in a very broad way based on whether electrons are transferred or shared into

Ionic Bond and
Covalent Bonds.

IONIC BONDS OR ELECTROVALENT BOND:

An Ionic Bond is a chemical bond formed between two atoms (usually a metal and a non-metal) by transfer of electrons from valence shell or one atom to valence shell of another atom.

The transfer of electron takes place in such a way that both the atoms attain inert gas configurations.

The actual bond is formed due to electrostatic force of attraction between the resulting two oppositely charged ions.

e.g. Ionic bond formation in NaCl

Consider Na and Cl atoms

Na has a configuration (2, 8, 1), it has a tendency to loose one electron (oxidation)to form Na+ ion. On loosing the electron the Na+ ion that is formed has the stable configuration of Neon-an inert gas. Hence the ion of sodium is more stable than the atom of sodium. This is also the reason why sodium does not occur in the free state as sodium but occurs as compounds of sodium only.

Cl has a configuration (2, 8, 7). It has a tendency to accept one electron (reduction) to form Cl– ion. On gaining one electron the Cl- ion that is formed has the stable configuration of Argon an inert gas. Hence the ion of Chlorine is more stable than the atom of chlorine. This is also the reason why chlorine will not occur in the free state but occurs as compounds of chlorine only.

The actual bond is formed by attraction of Na+ ion and Cl- ion – opposite ions attract each other.

Examples:

		Cation formation configuration changes	Anion formation configuration changes	Ionic bond formation
1	Ion formation	K – 1e– → K+ (2, 8, 8, 1) → (2, 8, 8, - Argon)	Cl + 1e– → Cl– (2, 8, 7) → (2, 8, 8 – Argon)	K+ + Cl– → KCl
2	Ion formation	Mg – 2e– → Mg2+ (2, 8, 2) → (2, 8 – Neon)	O + 2e– → O2– (2, 6 → (2, 8 – Neon)	Mg2+ + O2– → MgO
3	Ion formation	2Al – 6e– → 2Al3+ (2, 8, 3) → (2, 8 – Neon)	3O + 6e– → 3O2– (2, 6) → (2, 8 – Neon)	2Al3+ + 3O2– → Al2O3
4	Ion formation	Ca – 2e– → Ca2+ (2, 8, 8, 2) → (2, 8, 8 – Argon)	2Cl + 2e– → 2Cl– (2, 8, 7) → (2, 8, 8 – Argon)	Ca2+ + 2Cl– → CaCl2

COVALENT BOND, 1916, LEWIS (AMERICA):

Sharing of electrons such that each atom donates an electron to the shared pair forms Covalent Bonds. On sharing both atoms get either a duplet or octet configuration. They are also known as Atomic Bond or Electron Pair Bond.

Lewis Dot diagram

To denote the valence shell electrons, dots or crosses around the symbol of the atom are used. This is called Lewis diagram. It is used to represent the bond formation between atoms.

eg. H2, Cl2, HCl, N2, O2, F2, H2O, NH3, CO2, CH4

H2 molecule	:	or H – H	or H2
Cl2 molecule	:	or Cl – Cl	or Cl2
HCl molecule	:	or H – Cl	or HCl
N2 molecule	:	or N ≡ N	or N2
H2O molecule	:		or H2O
CO2 molecule	:		or CO2
NH3 molecule	:		or NH3

Based on covalency

Covalency is the number of shared Pairs of electrons between atoms in a covalent bond. It is never greater than 3

Single Covalent Bond molecules have one shared pair of electrons. Their Covalency is 1.

eg. H2, Cl2 etc.

Double Covalent Bond molecules have two shared pairs of electrons. Their covalency is 2.

eg. O2

Triple Covalent Bond molecules have three shared pairs of electrons. Their covalency is 3.

eg. N2

DIFFERENCES IN CHEMICAL PROPERTIES OF METALS AND NON-METALS:

PROPERTY	METAL	NON-METAL
Number of electrons in the valence shell:	Have 1 to 3 electrons in their valence shells.	Have 4 to 7 electrons in their valence shells. Exceptions: Hydrogen has one electron in its valence shell because the first shell is its valence shell.
Formation of ions:	Lose electrons from their valence shells to attain stable structures and form Cations	Accept electrons to attain stable structures and form Anions.
Reaction with oxygen	Metals on heating in air or oxygen react to form their respective oxides. Metal + Oxygen Metal Oxide The oxides are either basic or amphoteric in nature. Exceptions: Metals like Au and Pt do not form oxides. Examples 4Na + O2 -------- 2Na2O 4K + O2 ------- 2K2O 2Ca + O2 ------ 2CaO	Non-metals on heating in air or oxygen form their respective oxides. Non-Metal + O2 Non-Metal Oxide The oxides are either acidic or neutral in nature. Examples C + O2 CO2 S + O2 SO2 P4 + 5O2 2P2O5
Reaction with Acid	Metals, which are more reactive than hydrogen, replace hydrogen from the acid. Metal + Dil. Acid Salt + H2	The majority of non-metals do not react with acids.
Reaction with Water	Metals, which are more reactive than hydrogen, reacts with water and removes hydrogen and forms metal hydroxide or Metals Oxide. Metal + H2O Metal Oxide + H2 Metal+H2O Metal hydroxide H2	Non-metals react with water in various ways.
Reaction With Cl2	Metals react with chlorine to form metal chloride salts. Example: 2Na + Cl2 2NaCl Mg + Cl2 MgCl2	Non-metals react with chlorine to form compounds which are either volatile liquids or gases. Example: H2 + Cl2 2HCl P4 + 10Cl2 4PCl5
Electrolysis	Metals are generally liberated at cathode during electrolysis.	Non-metals are generally liberated at anode during electrolysis.
Oxidation/Reduction	Metals are good reducing agents as they readily lose electrons.	Non-metals are good oxidising agents as they readily gain electrons.