Covalent Bonding

The Octet Rule:

Elements will try to achieve a stable electron configuration by having 8 electrons in their valence shell, similar to noble gases, making the element stable and less likely to react with other atoms
Violations of the rule:

Hydrogen and Helium: these elements achieve stability with just two electrons in their outer shell
Beryllium and Boron: beryllium is stable with four electrons in its outer shell, and boron is stable with six electrons
Elements in the Third Period or Below: these elements can use their d-orbitals to be stable with more than eight electrons in their outer shell

Single Covalent Bond:

The electrostatic attraction between two positive nuclei and a shared pair of electrons
Occurs between non-metal atoms, where two electrons are shared between both atoms
Elements will try to fulfil the octet rule using covalent bonds

This also applies to the violations stated above (e.g. hydrogen will tend to seek a total of 2 valence electrons)

Naming Covalent Molecules

This was discussed in detail here
Summary: take the names of both elements and add suitable prefixes (mono, di, tri, etc.) to them based on their subscripts

Multiple Covalent Bonds

Single Bonds: one pair of electrons are being shared (2 electrons in total)
Double Bonds: two pairs of electrons are being shared (4 electrons in total)
Triple Bonds: three pairs of electrons are being shared (6 electrons in total)

Coordinate Covalent Bonds:

A covalent bond where both of the electrons being shared in the bond have been donated by one atom
This does not affect the properties of the covalent bond; coordinate bonds have the same properties as regular covalent bonds
To form a coordinate covalent bond, one atom/species (the donor) must have a lone pair of electrons available, while the other atom/species (the receiver) must have space to accept the electrons
To show a coordinate bond, use an arrow pointing from the donor to the receiver to demonstrate the origin of the electrons

Lewis Formulas

What are Lewis Formulas?

Lewis formulas show all valence electrons in a molecule, whether these are pairs of electrons involved in covalent bonds or non-bonding pairs of electrons
Uses lines between atoms to represent covalent bonds, and dots around atoms to represent non-bonding electrons
Also known as electron dot structures or Lewis structures

Drawing Lewis Formulas for Binary Molecules

Calculate the total number of valence electrons from both atoms in the molecule
Draw the skeletal structure of the molecule to show how the atoms are linked to each other with a single bond
Add remaining pairs of electrons as non-bonding pairs to atoms that do not have complete valence shells
Check if the final structure contains the correct number of valence electrons and all atoms have a filled valence shell (octet rule or exceptions)

Remember, each bond represents two shared electrons
If there are not enough electrons to fill the octets of all atoms, create double or triple bonds as necessary

Drawing Lewis Formulas for Molecules With More Than Two Atoms

Calculate the total number of valence electrons from all atoms in the molecule
Identify the central atom, which is usually the least electronegative element (excluding hydrogen)
Arrange the atoms around the central atom and connect all surrounding atoms to it using single bonds
Distribute the remaining valence electrons as lone pairs to the outer atoms first, then to the central atom if needed, to help each atom achieve a full valence shell (octet rule or exceptions)
If any atoms still lack a complete valence shell, convert lone pairs from surrounding atoms into double or triple bonds with the central atom as necessary

Always check that the total number of electrons used matches the total number of valence electrons calculated at the start

Covalent Bond Polarity

Electronegativity:

A measure of how much an atomic nucleus attracts the shared electrons that are involved in a covalent bond
Atoms with a higher electronegativity will have greater possession of electron pairs, creating unequal sharing
On the periodic table, elements in the bottom-left (metals) have low electronegativities, while elements in the top-right (non-metals) have high electronegativities

Ionic Bonds:

Exist between atoms with low electronegativities and high electronegativities (generally metals and non-metals)
The large electronegativity difference results in electrons transferring to the more electronegative element, instead of being shared between both elements
According to IB, ionic bonds occur when the electronegativity difference between two elements is greater than 1.7

Polar Covalent Bonds:

Exist between different atoms that have a moderate electronegativity difference
Result from electron pairs being unequally shared between atoms
According to IB, polar covalent bonds occur when the electronegativity difference between two elements is between 0.4 and 1.7

Non-Polar Covalent Bonds:

Exist between identical nonmetal atoms, or nonmetal atoms with similar electronegativities
Result from electron pairs being equally shared between atoms
According to IB, non-polar covalent bonds occur when the electronegativity difference between two elements is less than 0.4

Valence Shell Electron Pair Repulsion (VSEPR) Theory

What is VSEPR?

VSEPR is a model used to predict the 3D shape of molecules based on the repulsions between electron pairs in the valence shell of the central atom
Bonding and non-bonding electron pairs repel each other, and will arrange themselves as far apart as possible to minimize repulsion
This arrangement determines the molecular geometry and bond angles, which affects properties like polarity and intermolecular forces

Electron Domains:

The regions in which bonding and non-bonding pairs of electrons are most likely to be found
Non-bonding pairs, single bonds, double bonds, and triple bonds each count as one electron domain
The number of electron domains around the central atom determines the electron domain geometry and molecular geometry of the molecule

Electron Domain Geometry:

The shape of the arrangement of electron domains surrounding a central atom in a molecule
Non-bonding pairs of electrons are more repulsive than bonding pairs, as they have a greater degree of freedom to move around in the electron cloud

This greater repulsion caused by non-bonding pairs reduces the bond angles within molecules

Electron domain geometry does not indicate the arrangement of the atoms that are bonded to the central atom, and only relates to electron domains/pairs

Molecular Geometry:

The three-dimensional arrangement of atoms in a molecule, determined using VSEPR theory
The molecular geometry will only be different from the electron domain geometry for central atoms that have non-bonding pairs of electrons
To draw the 3D structure of a molecule, use wedge-and-dash notation:

Solid wedges (▲) represent bonds coming out of the plane of the paper toward the viewer
Dashed wedges (︽) represent bonds going behind the plane of the paper away from the viewer
Lines (—) represent bonds that are in the plane of the paper

Common Electron Domain and Molecular Geometries

Electron Domains	Electron Domain Geometry	Number of Lone Pairs	Molecular Geometry	Bond Angles	Example
2	Linear	0	Linear	180°	CO₂
3	Trigonal Planar	0	Trigonal Planar	120°	BF₃
3	Trigonal Planar	1	Bent	<120°	SO₂
4	Tetrahedral	0	Tetrahedral	109.5°	CH₄
		1	Trigonal Pyramidal	<109.5°	NH₃
		2	Bent	<109.5°	H₂O
5	Trigonal Bipyramidal	0	Trigonal Bipyramidal	90°, 120°, 180°	PCl₅
		1	See-Saw	<90°, <120°, <180°	SF₄
		2	T-Shaped	90°, 180°	ClF₃
		3	Linear	180°	XeF₂
6	Octahedral	0	Octahedral	90°, 180°	SF₆
		1	Square Pyramidal	<90°, <180°	BrF₅
		2	Square Planar	90°, 180°	XeF₄

Dipoles

Bond Dipoles

A bond dipole occurs when two atoms in a covalent bond have different electronegativities (polar bond), causing the shared electrons to be pulled closer to the more electronegative atom
This creates a partial negative charge (δ^–) on the more electronegative atom and a partial positive charge (δ⁺) on the less electronegative atom
The direction of the bond dipole is represented by an arrow pointing toward the more electronegative atom

Molecular Dipoles

A molecular dipole refers to the overall polarity of a molecule, which depends on both the individual bond dipoles and the molecular geometry
If the bond dipoles are symmetrically arranged (e.g. CO₂ or CCl₄), they cancel out, and there is no molecular dipole (nonpolar)
If the bond dipoles are asymmetrical (e.g. H₂O or NH₃), they do not cancel, and there is a molecular dipole (polar)
Thus, to have a molecular dipole, there must be both polar covalent bonds as well as an asymmetrical molecular geometry

Hybridization

Sigma Bonds (σ):

Bonds which are formed by the direct, head-on overlap of atomic orbitals (s, p, etc.) from two atoms
Result in a region of high electron density occurring between atoms
The strongest type of covalent bond
Single, double, and triple covalent bonds all contain one sigma bond

Pi Bonds (π):

Bonds which are formed by the side-to-side overlap of p-orbitals from two atoms
Weaker than sigma bonds, but a sigma + pi bond is stronger than just a sigma bond
Result in two regions of high electron density occurring above and below atoms
Single covalent bonds have no pi bonds, while double and triple bonds have one and two pi bonds respectively

Hybrid Orbitals:

Orbitals formed from the mixing of other atomic orbitals, which have intermediate energies
Used to explain the symmetry in molecular geometry resulting from VSEPR
Usually, only s and p orbitals will hybridize to form sp, sp², sp³, etc. orbitals, where the type depends on the number of electron domains of a bonding atom

Determining the Hybridization of an Atom

The number of electron domains (found from the Lewis formula), the electron domain geometry (VSEPR), and the hybridization of an atom are interchangeable
You can use the following table to determine an atom's hybridization:

Number of Electron Domains	Electron Domain Geometry	Hybridization
2	Linear	sp
3	Trigonal Planar	sp²
4	Tetrahedral	sp³

Resonance

Formal Charge:

The theoretical charge of an individual atom within a molecule, based on the electron distribution in its Lewis structure
For compounds that have multiple possible Lewis structures, the formal charge is used to check which one is the most likely
Formal Charge = V - (B + L), where V is the number of valence electrons in the unbonded atom, B is the number of bonds, and L is the number of non-bonding electrons on the atom
The structure that is most likely to exist will have a formal charge closest to zero for as many atoms in the structure as possible, and the most negative formal charge should be given to the most electronegative atom in the compound

Resonance Structures:

Used to show the structure of a compound when multiple Lewis structures are possible
In order to show the different possible Lewis structures, connect the Lewis diagrams with a double headed arrow
For example, the resonance structures for ozone can be communicated as shown below:

Resonance Hybrid Structures

Experimental evidence shows that the true structure of a resonance compound is actually a hybrid of all the possible Lewis structures

As such, none of the Lewis structures show the true structure, and the hybrid is an intermediate structure

This happens because the molecule delocalises pairs of electrons, which spread over multiple positions as they are no longer fixed to one bonding position
This electron delocalisation is represented using using dashed lines, as seen below for ozone:

Bond Order:

Refers to the number of covalent bonds between a pair of atoms, and relates to properties like bond strength and bond length
Single bonds have a bond order of 1, double bonds have a bond order of 2, and triple bonds have a bond order of 3
In resonance hybrid structures, the bond order is the average number of bonds between two atoms

For example, for the average of a single and double bond (like in ozone), the bond order is closer to 1.5

Ozone

Ozone contains 3 oxygen atoms, where one oxygen only has a single bond with one other oxygen, while the third oxygen has a double bond (as seen above)
The double bond can be attributed to either of the oxygens, creating two possible structures
However, the actual structure of ozone is a resonance hybrid structure, where all oxygens have intermediate bonds which are between a single and double bond

Benzene

Benzene is a hydrocarbon with the chemical formula C₆H₆
Each hydrogen is attached to one carbon, and each carbon has a single and double bond with other carbons, creating a ring structure as shown below:

As such, two possible resonance structures can form, however experimental evidence shows that the observed bond lengths between the carbons are all the same, and are of a length between that expected for a single and a double carbon-carbon bond
In order to show benzene's hybrid structure, chemists often draw a circle in the center of the Lewis structure to show the delocalisation of 6 electrons:

Intermolecular Forces (IMFs)

What are IMFs?

The forces that act between different molecules are called intermolecular forces
These are the forces that keep particles together to form solids and liquids
Intermolecular forces are weaker than intramolecular forces, which are the chemical bonds within an individual molecule

London Dispersion Forces (LDF):

The electrostatic attractive forces that exist between all molecules, whether polar or nonpolar
Caused by temporary (instantaneous) dipoles that form when the electrons in a molecule become unevenly distributed at a given moment, due to the random motion of electrons
These temporary dipoles can induce dipoles in neighbouring molecules, creating a momentary attraction between them
The strength of London dispersion forces increases with the number of electrons and the surface area of the molecules (larger, heavier atoms and molecules experience stronger dispersion forces)
LDFs are the weakest IMFs

Dipole-Dipole Forces:

The electrostatic attractive forces that exist between polar molecules that have a molecular dipole
Polar molecules have partially positive and negative ends, which lead to electrostatic forces
The attractive forces (opposite charge) are stronger than the repulsive forces (same charge), so there is an overall attraction between polar molecules
Dipole-dipole forces are moderately strong IMFs

Hydrogen Bonds

A special case of dipole-dipole forces
Seen among molecules where H is bonded to a highly electronegative atom, such as N, O, or F
Hydrogen bonds are the strongest IMFs

IMFs and Physical Properties

Stronger IMFs lead to higher boiling and melting points (more energy required to separate molecules)
Stronger IMFs lead to higher surface tension and viscosity (molecules stick together more)
Stronger IMFs lead to lower vapour pressure (fewer molecules escape into the gas phase)
Molecules with strong IMFs tend to be solids or liquids (e.g. H₂O), while those with weak IMFs tend to be gases (e.g. CO₂)

In these examples, H₂O is a liquid due to having hydrogen bonds, while CO₂ is a gas due to only having LDFs

Giant Covalent Structures

Allotropes:

Different structures of the same element
Carbon has 4 allotropes: diamond, graphite, graphene, and C₆₀ fullerene
Graphite, diamond, and graphene are examples of covalent network solids (giant covalent structures)

The other allotrope of carbon, C₆₀ fullerene, is a molecule

General Properties of Covalent Network Solids:

High melting points (>1000°C) due to many strong covalent bonds
Usually poor conductors
Usually insoluble in most substances
Usually very hard

Diamond

Each carbon is bonded to 4 other carbon atoms in a tetrahedral geometry
Forms a lattice structure
Very strong bonds: high melting & boiling point, one of hardest substance on earth
Not conductive (no delocalized electrons)
Insoluble in most solvents

Graphite

Each carbon is bonded to 3 other carbon atoms in a trigonal planar geometry
Carbon atoms form layers of hexagonal rings
Layers are connected by weak London dispersion forces
Conductive (delocalized electrons)

Graphene

A single layer of graphite
Atoms are arranged hexagonally and form a one-atom-thick layer
One of thinnest and strongest materials known
Excellent conductor
Flexible (can be rolled into nanotubes)

C₆₀ Fullerene

Each carbon is bonded to 3 other carbons
Has molecules containing 60 carbon atoms, which are connected by weak London dispersion forces
Insoluble in water, but soluble in non-polar solvents
Not conductive, as its delocalized electrons can't move between molecules

Pure Silicon

Forms a covalent network solid
Each silicon is bonded to 4 other silicon atoms in a tetrahedral geometry
Similar structure and properties to diamond

Silicon Dioxide (Quartz)

A covalent network solid containing silicon and oxygen atoms
Each silicon is bonded to 4 oxygens
Each oxygen is bonded to 2 silicon atoms
High melting and boiling point due to strong covalent bonds
Not conductive

Oakridge IB Guide

Chemistry

The Covalent Model