Bonding: Important Terms and Hybridization

Before we get started on this day's lesson, it would greatly benefit you if you re-read these posts:

Types of Bonds
Atomic Structure (electronic configurations in particular)

Done that? Good.

Let's go over some important terms:

Now if you recall, covalent bonds are formed due to the overlapping of atomic orbitals containing unpaired electrons with opposite spins. Such an orbitals are referred to as bond pairs, i.e they can participate in bonding.

An orbital which cannot participate in bonding, i.e one with paired electrons is called a lone pair.

Now, it is only logical that the strength of a bond directly correlates to the extent of overlapping. Greater the overlapping, stronger is the bond.

Orbitals can overlap in 2 ways:

(i) axial overlapping (along the axis): this results in the formation of a sigma(σ) bond.

(ii) sideways overlapping: results in the formation of a pi (π) bond. A pi bond is representative of a diffused electron cloud above and below the sigma bond that is shared by the two bonded atoms.

Note: That only one sigma bond can be formed between two atoms. The first bond formed is always a sigma bond.

Bond order: the no. of bonds present between atoms in a molecule.

a) Single Bond (eg. C-C): 1σ bond
b) Double Bond (eg. C=C ): 1σ + 1π bond
c) Triple Bond (eg. C≡C): 1σ + 2π bond

Bond Length: Is the average distance between the nuclei of two bonded atoms.

Study the energy diagram given below to better understand this phenomenon.

H2 Bond length

Quick Tip: Bear in mind that when two atoms approach one another, there are attractive (between electrons and nucleus) and repulsive (inter-electronic and inter-nuclear) forces at play. Initially the atoms move closer under the influence of attractive forces and as they move closer their potential energy decreases (generally, energy is released during bond formation). They reach a certain minimum distance at which attractive and repulsive forces are evenly matched. If the atoms move any closer, repulsive forces will dominate and cause the energy diagram to spike (as can be seen on the left). Since lower energy states are preferred by all systems, atoms are thus held a certain distance apart, where the the attractive and repulsive forces are evenly balanced and energy level is minimum. This minimum distance is called the bond length.

Bond Angle In water molecule

Remember: Bond Strength ∝ Bond Order ∝ 1/Bond length

Bond Angle: Angle Subtended by the bonded molecules.

Bond Disassociation Energy: Energy required to break a certain bond and take the bonded atoms to their gaseous states.

Okay, let us consider the water molecule. If we were to remove the Hydrogen on the right, we would need to expend a certain amount of energy, says, BE-1. Now, it is only natural that the oxygen atom bereft of a hydrogen atom would hold on more dearly to the second hydrogen on the left. So now if we were to remove the second hydrogen we would have to supply a slightly greater amount of energy (BE-2)

BE-2 > BE-1

So for the sake of convenience, we take the average of the two Bond Disassociation Energies and call it Bond Energy.

Bond Energy (O-H) = (BE-1+BE-2)/2

That pretty much covers all you need to know to follow whatever comes next.

Hybridization <This Shit IS IMPORTANT>

Central Atom: atom forming the maximum no. of bonds.

Orbitals of the central undergo hybridization ( in layperson language, it refers to the mixing of atomic orbitals) to form hybrid orbitals which overlap with "normal" orbitals of neighboring atoms to form bonds.

Conditions: i) Orbitals should have the same or almost the same energy.
ii) Hybrid orbitals formed should also have similar energy levels.
iii) Number of orbitals remains the same.
iv) Both bond pairs and lone pairs can take part in hybridization.

Types of Hybridization 

sp hybridization: forms 2 sp hybrids, each with 50% s and 50% p character.

sp² hybridization: forms 3 hybrid orbitals, each with 33.33% s and 66.66% p character.

sp³ hybridization: forms 4 orbitals each with 25% s character and 75% p character.

sp³ d hybridization: 5 orbitals with, 20% s, 60%p and 20% d character.

Sp³d² hybridization: 6 orbitals with 16.66% s, 50% p and 33.33% d character.

(in case you haven't already noticed, the letters (s,p,d etc.) indicate the orbitals participating in hybridization.)

Formation of Hybrid Orbitals

SP hybridization 

in BeCl₂ Beryllium is sp hybridized. 

Be in  ground state: 2s² 

One of the 's' electrons is excited and moves to the p orbital

Be excited state: 2s¹ 2p¹

Now, the s and the p orbitals undergo hybridization to form two sp-orbitals, with an unpaired electron in each.

Consider Chlorine. The chlorine atom is in the ground state and has the electronic configuration: 3s² 3p⁵

It is evident that the chlorine atom has an unpaired electron in a p-orbital. 

Now, the sp-hybridized berylium atom overlaps with two chlorine atoms forming the compound BeCl₂

Sp² hybridization

in BCl₃  Boron is sp²  hybridized.

B in the ground state: 2s² 2p¹

One of the 's' electrons is excited and moves to the p orbital. 

B in the excited state is now: 2s¹ 2p²

Now, the s and the 2 p orbitals undergo hybridization to form three sp²-orbitals, with an unpaired electron in each. Each of these hybrid orbitals bonds with the chlorine p-orbital with a bond pair (unpaired electron).

      

      

Sp³ hybridization

Carbon in CCl₄  is sp³  hybridized.

The process is pretty much the same.

C in ground state: 2s² 2p²

C in the excited state: 2s¹ 2p³

Four hybrid orbitals are thus formed, with an unpaired electron in each. These then overlap with the Chlorine p-orbital with a bond pair.

Sp³d hybridization

P in the ground state:  3s² 3p³ 3d⁰
P in the excited state: 3s¹ 3p³ 3d¹

Five hybrid orbitals are formed, with an unpaired electrons in each. Each of these then overlap with the Chlorine p-orbital with a bond pair (i.e unpaired electron). Note: that the axial chlorine atoms project out of the plane of the molecule.

Sp³d² hybridization

 S in the ground state:  3s² 3p⁴ 3d⁰

 S in the excited state: 3s¹ 3p³ 3d²

6 hybrid orbitals are formed with unpaired electrons in each which in turn bond with orbital of fluorine atom with a bond pair. 4 fluorine atoms arrange themselves along the vertices of  a square in the plane of the molecule and two project outwards

Sp³d³ hybridization

I in the ground state: 5s² 5p⁵ 5d⁰
I in the excited state: 5s¹ 5p³ 5d³

7 hybrid orbitals are formed with unpaired electrons in each which in turn bond with orbital of fluorine atom with a bond pair. 5 Fluorine atoms arrange themselves along the vertices of a pentagon in the plane of the molecule and two axial atoms project outwards.

Okay that does it for today. In the next post I shall deal with VSEPR theory and arrangement of molecules in 3-D space.

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That's all Folks!

The Passive Observer Out!

Types of Bonds

A bond, as its names suggests, holds atoms together to form molecules.

In this article we study a few elementary types of bonds.


Ionic Bond

An ionic bond is formed as consequence of the electrostatic forces of attraction between oppositely charged ions. 

If the difference in electronegativity of the atoms is over 1.7 the bond is likely to be ionic.

Covalent Bond

A covalent bond is formed due to the sharing of valence electrons or in terms of the orbital concept, a covalent bond is formed due to the overlapping of atomic orbitals. 

If the difference in electronegativity of the atoms is less than 1.7 the bond is likely to be covalent. 

In practice most bonds are not 100% covalent or ionic.

If the difference in electronegativity between the 2 bonded atoms is zero, then the bond is purely covalent. If the difference in electronegativity between the 2 bonded atoms ranges from zero to 1.7 the bond is a polar covalent bond and exhibits some ionic character. If the difference in electronegativity between the 2 bonded atoms is 1.7 then the bond is 50% covalent and 50% ionic and if the difference in electronegativity is over 1.7 then the bond is said to be ionic.

Ready-Reference Chart (ΔE is the difference in electronegativity)

ΔE Value             Bond Character 

ΔE = 0;                 100% Covalent

ΔE =  0-1.7;          polar covalent

ΔE = 1.7;              50% covalent 50% ionic
ΔE > 1.7;              ionic

Dipole Moment: Accounts for the ionic character of covalent bonds.

Homo-nuclear, diatomic molecules (eg. H_{2  ,}  Cl₂ etc.) are non-polar.

In case of hetero-nuclear molecules, for a molecule to be non polar, the central atom mustn't have a lone pair of electrons and the central atom must be surrounded by the same atoms i.e the molecule must be symmetrical (explained below). (both conditions must be met).

However, if these conditions aren't met, then the more electronegative element will pull the electrons towards itself and thus acquire a slightly negative charge (given as  δ- ) and the other atom shall acquire a slight positive charge (given as  δ+ ).

Quick Tip: Think of the bond as a match of tug of war, the stronger (in this case more electronegative element) player will pull the rope (in this case electrons) towards himself.

Such a bond is termed a polar covalent bond and its polarity is given in terms of dipole moment. 

Dipole moment (m) = electric charge x distance of charge seperation (bond length)

_{m = q x d}

_{Dipole moment is measured in ‘Debye’ unit (D)}

Dipole moment is a vector quantity, therefore if a molecule with polar bonds is symmetrical then the opposite pulls will cancel each other resulting in zero net dipole. eg. Carbon dioxide, pictured below.

 

Fajans' Rules: Account for the covalent nature of ionic bonds. According to these small cations (or high positive charge) have greater polarising power and large anions have greater polarisibility, thus bonds formed between such species involves overlap of their electron clouds resulting in covalent character.

Quick Tip: Picture a large bunch of teenage girls (electrons) on a field trip with their teacher (nucleus) and let's call this group our anion. Now a small cation, hmm Justin Beiber for instance, walks in, and some of the girls move away from the group towards him. Thus the small cation, Justin Bieber in this case has pulled the electron cloud of the large anion towards himself and this mingling of electron clouds accounts for the covalent character of the bond.

If the group were smaller (i.e the anion was smaller) the teacher would've been able to check the movement of the girls. And  if the celebrity was somebody more important (i.e a larger cation), like Kazimierz Fajans(yes, he is the guy behind fajans rule) for instance, again fewer girls (electrons) would've strayed away from the group.


Ready-Reference Chart

Ionic                               Covalent

Low +ve charge              High +ve charge
small anion                      large anion
large cation                     small cation 



Metallic Bonding

Since Ionic and Covalent bonds cannot account for many physical properties of metals, such as strength, malleability, ductility, thermal and electrical conductivity, opacity, and luster, the concept of metallic bonding has been devised.

The positive charge in a metal aggregates to form a "kernel" whereas the electron separate out forming an electron cloud.

Metallic bonding constitutes the electrostatic attractive forces between these delocalized electrons, gathered in an electron cloud, and the kernel of positively charged metal ions.

Shitty illustration depicting kernel of positive charge and cloud of electrons.

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That's all folks!

The Passive Observer