Atoms have a tendency to obtain an electronic configuration like noble gases.
For this, they either donate or share electrons.
The force of attraction by which the atoms are bonded in a molecule is called a chemical bond.
Covalent Bonding:
It is the bond formed due to the sharing of an electron pair between two non-metallic elements.
The shared electron pair is attracted by two nuclei.
One covalent bond between two molecules is formed due to the sharing of two electrons.
Bond pair is the number of paired electrons in a bond or simply the number of covalent bonds.
Lone pair is the number of paired electrons left in the outermost shell after forming a covalent bond.
Dot Cross Diagram
Examples:
Some exceptional compounds:
Iodine Chloride & Iodine Fluoride
Iodine Flouride can form but Iodine Chloride cannot.
Chlorine has a larger atomic radius that Fluorine.
Seven Fluorine atoms can be arranged around Iodine but not seven Chlorine atoms.
Phosphorus (V) Chloride & Nitrogen (V) Chloride
Phosphorus (V) Chloride can form but not Nitrogen (V) Chloride.
The second shell of Nitrogen can only contain a maximum of 8 electrons.
But the third shell of phosphorus can contain a maximum of 18 electrons.
Electron deficient compounds
Compounds where the central atom has an incomplete octet.
Examples include Boron trifluoride and Aluminium Chloride.
These compounds can form a dative bond.
Dative/Co-ordinate bond
It is a bond formed when shared pairs of electrons are provided from the same species.
Examples:
Ionic Bonds
The bond formed due to the strong electrostatic force of attraction between oppositely charged ions.
During formation, metals donate electrons and non-metals gain them.
Metals form positively charged cations and non-metals form negatively charged anions.
Strength
The strength of an ionic bond depends on the charge of the ion and the ionic radius.
With the increase of charge the strength of the ionic bond increases.
With the increase of ionic radius strength of the ionic bond decreases.
Physical properties
High melting and boiling points
Ionic compounds have giant ionic lattices.
They have strong ionic bonds.
To melt the strong electrostatic force is needed to be overcome.
A higher amount of energy is needed to break the ionic bonds.
MgO vs NaCl
The charge of ions in MgO is higher than in NaCl.
Mg2+ ion has a smaller ionic radius than Na+ ion.
Oxide ion has a smaller ionic radius than Chloride ion.
The electrostatic attraction between ions is stronger in MgO than in NaCl
KCl vs NaCl
NaCl has a higher melting point than KCl.
Na ion has a smaller ionic radius than K ion.
Ionic compounds are non-conductor of electricity as solid but conductors as a liquid.
At room temperature, the ions are closely packed so cannot move to conduct electricity.
Soluble ionic compounds produce free ions in water which move to conduct electricity.
Solubility
Most of the ionic compound is soluble in water or polar solvent.
Ionic compounds are insoluble in non-polar solvents.
Process of solubility
Water molecules have a partial positive charge(𝛿+) on hydrogen and a partial negative(𝛿-) charge on oxygen.
The positively charged ions are attracted to the partial charge of oxygen(𝛿-).
The negatively charged ions are attracted to the partial charge of hydrogen(𝛿+).
Coordination number:
Six Sodium ions are surrounded around each Chloride ion.
Six Chloride ions are surrounded by each sodium ion.
Metallic bond
Metals have positively charged ions and delocalised electrons.
The bond formed between positively charged ions and delocalised electrons is called a metallic bond.
Strength of metallic bond: It depends on the following
Charge on the ions - Strength increases with increasing charge.
The number of delocalised electrons - Strength increases with increasing number.
Ionic radius- Strength decreases with increasing number.
Across the period strength of the metallic bond increases.
Charge of the metal ions increases.
Ionic radius decreases.
The strength of the electrostatic force of attraction between delocalised electrons and metal ions increases.
Physical properties of metal:
Metals have high melting and boiling points:
They have giant metallic lattices.
They have a strong electrostatic force of attraction between delocalised electrons and metal ions.
To overcome the force high amount of energy is required.
Metals are good conductors of heat and electricity:
They have delocalised electrons to transfer heat and electricity.
Electronegativity:
The measurement of power to attract the shared electron pair between two bonded atoms.
Across the period(left to right) the electronegativity increases.
Down the ground (top to bottom) electronegativity decreases.
Hydrogen Bond:
A type of attractive intermolecular force that exists between two partial electric charges of opposite polarity. This happens due to a difference in electronegativity.
Hydrogen Bond in water
Conditions for Hydrogen Bond:
The hydrogen atom must be bonded between Oxygen/Nitrogen/ Fluorine.
Due to the difference in electronegativity of atoms dipoles ( 𝛿) form.
The oppositely charged diples attract each other forming hydrogen bonds.
It forms between water, ammonia, Hydrogen Flouride, alcohol, hydrazine etc.
Types of intermolecular force:
Induced dipole attraction(van der Waals)
Permanent dipole attraction
Hydrogen Bond
Polarity:
Polarity in a molecule is caused due to the difference in electronegativity.
A more electronegative atom will "pull" the electron towards itself.
This results in the formation of dipoles.
Permanent dipoles & Polar molecules :
A molecule becomes polar when the dipoles do not cancel out and have a resultant dipole.
This resultant dipole is permanent in the molecule.
Temporary dipoles:
Electrons in a molecule can sometimes get unevenly distributed.
This results in the formation of short temporary dipoles.
This phenomenon occurs in simple molecular substances.
Simple molecular substances
Simple molecular substances exist as solid, liquid or gas mainly liquid and gas.
They have a low melting and boiling point as their intermolecular force (van der Waals) is very weak.
The melting and boiling points of simple molecular substances depend on the number of electrons and the presence of permanent dipoles.
The more electrons the stronger the strength of temporary dipoles.
Down the group seven the melting and boiling point increases and becomes darker in colour.
From fluorine to Astatine number of electrons in the molecule increases.
Strength of van der Waals's force of attraction increases.
Greater energy is needed to overcome this force of attraction.
Molecules are closer together.
Colour becomes darker.
Carbon monoxide vs Nitrogen
Carbon monoxide has a higher boiling point than nitrogen.
Carbon monoxide is a polar molecule so has permanent dipoles.
Nitrogen is non-polar, so only have temporary dipoles.
Permanent dipoles are stronger than temporary dipoles.
Hydrogen vs water
Hydrogen molecules only have van der Waals's force of attraction.
Water has van der Waals force of attraction and Hydrogen bond.
Hydrogen bonds are stronger than van der Waals force
The attraction between water molecules is stronger.
Volatility: It describes how easily a substance vaporizes.
Volatile liquids have weak intermolecular forces and are mainly simple molecular.
Evaporation: Changing of liquid into a gas at any range of temperature.
It occurs from the liquid surface
It is a slow process.
Enthalpy change of vaporisation can be used to measure the strength of van der Waals force of attraction.
Shapes and angles of molecules
1. Tetrahedral Number of bond pairs - 4 Number of lone pair - 0 Bond angle - 109.5° Examples: Ammonium, Methane. 2. Linear Number of bond pairs - 2 Number of lone pair - 0 Bond angle - 180° Examples: Carbon dioxide, Beryllium hydride 3. Trigonal Planar Number of bond pairs - 3 Number of lone pair - 0 Bond angle - 120° Examples: Boron trifluoride, Boron hydride 4. Trigonal Bipyramidal Number of bond pairs - 5 Number of lone pair - 0 Bond angle - 120°, 90° Examples: Phosphorus pentafluoride, Sulfur tetrafluoride. 5. Octahedral Number of bond pairs - 6 Number of lone pair - 0 Bond angle - 90° Examples: Sulfur hexafluoride. 6. Pyramidal Number of bond pairs - 3 Number of lone pairs - 1 Bond angle - 107° Examples: Ammonia 7. V-shaped Number of bond pairs - 2 Number of lone pairs - 2 Bond angle - 104.5° Examples: Water
Hybridization:
A process by which different orbitals are mixed with each other to make a new hybrid orbital.
Sp³ Hybridization:
A process by which one S orbital and three p-orbitals are mixed with each other to make four hybridized sp³ orbitals.
Sigma Bond (σ): Bond formed due to the direct overlap between two orbitals. (A single bond is a sigma bond)
Sp² Hybridization:
A process by which one S orbital and two p-orbitals are mixed with each other to make four hybridized sp² orbitals.
During the formation of one p-orbital is not involved.
This orbital makes the pi (π) bond.
Pi Bond (π): Bond formed due to the sideways overlapping between two p-orbitals.
Every double bond has a sigma bond and a pi(π) bond
then finally to
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