CHEM C3: Intermolecular Forces

Here’s the topics to cover:

  • Intermolecular Forces
    • Hydrogen bonding
    • Electronegativity
    • Bond energy, length, polarity
    • Dipoles
    • van der waals forces
  • Bonding & Physical Properties

Here we go!

Intermolecular Forces
There are a few types of intermolecular force (between MOLECULES):

  • van der Waals’ forces
  • permanent dipole-dipole forces
  • hydrogen bonds

Before we jump into the types of forces, we need a bit of background on ELECTRONEGATIVITY & POLARITY.

What is Electronegativity?
The ability of an atom to attract electrons towards itself in a covalent bond.
The more ELECTRONEGATIVE an atom is, the higher the TENDENCY to receive electrons.

Thus, when in a covalent bond, the more electronegative atoms receive more electrons.
They have a more NEGATIVE net charge compared to the other atom.

Factors affecting Electronegativity

  • Number of Protons
  • Size of Atom
  • Shielding Effect
Number of Protons Higher number of protons
= Higher positive charge
= higher attraction for negative electrons
= HIGHER electronegativity
Size of Atom Larger atomic radius
= Weaker attraction between protons in nucleus to valence electrons
= LOWER electronegativity
Shielding Effect Higher number of shells
= Stronger shielding effect
= Weaker attraction between protons in nucleus to valence electrons
= LOWER electronegativity

Thus, it can be concluded that:
Electronegativity increases when there is

  • A high proton number
  • A smaller number of shells

Trend down a group:
Electronegativity DECREASES
Because number of shells increases.

Trend across a period:
Electronegativity INCREASES
Because proton number increases.

Here’s a periodic table showing the electronegativity values for each element:


Electronegativity can also be expressed on the PAULING SCALE, which looks like this:

Image result for electronegativity chart

Q: What unit is Electronegativity measured in?

What is a Polar Bond?
A bond between atoms of DIFFERENT electronegativities.
The distribution of electrons (& thus charges) is NOT equal (the charge distribution is ASYMMETRIC).
The more ELECTRONEGATIVE atom attracts the shared electrons CLOSER to it.
It has a higher NEGATIVE charge compared to the more electropositive atom.
It is PARTIALLY negatively-charged.

The small difference in charge is denoted by:
δ+ (for the partially positive pole)
δ- (for the partially negative)

For example:
In Hydrogen Chloride, the chlorine atom is MORE ELECTRONEGATIVE compared to hydrogen.
The shared electrons are thus CLOSER to the Cl than to H.
The Cl side is slightly more negative.
It is a polar molecule, & Cl is the negative pole, while H is the positive pole.
Image result for polar bond

Polar molecules can be identified by the number of poles.
A DIPOLE molecule has 2 poles: 1 positive & 1 negative.

What is a Dipole Moment (m)?
A measure of molecular polarity.
It is defined as the product of magnitude of charges & distance of separation between charges.

We don’t have to know this, but it’s good knowledge:
m = charge x distance
Measured in DEBYE (D).

The higher the Dipole Moment, the MORE polar a molecule is.

What is a Non-Polar Bond?
Covalent bond between same atoms (which have the SAME electronegativity).
Electrons are shared equally, & neither atom has a more negative charge.

For example, this H2:

Image result for polar bond

What is a Polar MOLECULE?
A molecule with an uneven distribution of charges across all atoms.
A polar molecule MUST contain polar bonds,
but NOT ALL polar bonds create polar molecules.
The SHAPE of the molecule must be taken into consideration.

For example:

Water contains 2 polar BONDS.

For each bond, the O is more electronegative compared to H.

You’d think that since there are 2 partially positive H’s on both sides of the O, they’d cancel out & the molecule would be non-polar.

However, water ISN’T shaped LINEARLY.

Due to O having 2 lone pairs, the 2 H atoms are repelled to one side of the molecule.

Now, the charge distribution is UNEVEN,
& water as a whole is POLAR.

Related image
Methane contains 4 polar BONDS

However, it is a NON-POLAR MOLECULE.

The 4 positive poles are distributed EVENLY, with the negative pole stuck in the middle.

Thus, it acts as a polar molecule since the distribution of charges of the molecule as a whole is EVEN.

Image result for methane polarity

Polar molecules always contain polar bonds,
but non-polar molecules may contain polar OR non-polar bonds.

Now that we understand Polarity, let’s tackle

van der Waals forces
Force of attraction caused by INDUCED (temporary) dipole.

Electrons in an atom move highly randomly.
For a majority of the time, the distribution of electrons are UNEVEN around the nucleus.
This creates a BRIEF & TEMPORARY dipole: the area with more electrons is more negative.

A temporary dipole in one atom induces temporary dipoles in other atoms by attracting/repelling the electron charge cloud.

This causes weak forces of attraction between atoms.

Factors Affecting VDW Forces

  • Size of molecule / Number of electrons
  • Shape of molecule

Bigger molecule = more electrons.
Therefore, bigger molecules = higher VDW forces.

Remember that the strength of intermolecular forces affects the melting/boiling points
(because the stronger the intermolecular forces, the more energy needed to pull these molecules apart during phase changes).

the LARGER the molecule, the HIGHER the melting/boiling points.

The shape of the molecule also affects this.If the molecules can be CLOSER to each other (due to shape), then there will be a STRONGER VDW force between them.
Thus, more energy will be needed.

Saturated vs Unsaturated Fats

Saturated Fats Has single C-C bond
Cannot open up to form branches
Forms straight chains
Molecules can be placed closer together
Higher VDW forces of attraction
Higher MP
Unsaturated Fats Has double C=C bond
Opens up to form branches
Forms twisted chains
Molecules have more space between them
Lower VDW forces of attraction
Lower MP


Permanent Dipole-Dipole Attraction
(As the name implies), a force of attraction between polar molecules with PERMANENT dipoles.

A polar molecule will have separate poles (2 poles being a DIPOLE).
For example: HCl

If there is another identical molecule near that HCl molecule, then the opposite poles of the other molecule will be attracted to the poles of the 1st HCl.

Hydrogen Bonding
Attraction between a HYDROGEN atom IN a covalent POLAR compound & the LONE PAIRS of an atom in another molecule.

Examples of compounds with lone pairs:
H2O, HF, NH3

Q: Why does N, O, F induce H-bonding but not the other elements in the same groups?
A: These elements are MORE electronegative than their other group members.
They pull the electrons in the H CLOSER than other elements would.

The proton of H is MORE EXPOSED to the lone pairs in other molecules.
Thus, there is a stronger force of attraction between the exposed proton (H) & the lone pairs.

Trend Down Groups
Here’s a graph of the boiling points of various hydrogen compounds, according to the group of the element bonded with the H.

Light Blue indicates group 6, (O, S, Se, Te)
Red indicates group 7 (F, Cl, Br, I)
Green indicates group 5 (N, P, As, Sb)
Dark Blue indicates group 4 (C, Si, Ge, Sn)

Image result for boiling point hydrogen bonding
Going down each group, the atomic mass of the other (non-H) atom increases.
Thus, the molecular mass increases.

It is expected that VDW forces (& thus BP) increases as the molecular mass/size increases,
but H2O, NH3 & HF are exceptions.
They have a HIGH BP due to the STRONG hydrogen bonds between molecules.
The other compounds do not have hydrogen bonds due to their lower electronegativities.
Only hydrogen compounds with N, O, F experience this effect.

Factors Affecting Strength of H-Bonds

  1. of H-Bonds per molecule
  2. of Lone Pairs per molecule
  3. Difference of Electronegativity
No. of H-Bonds per Molecule The more hydrogen atoms you have, the higher possibility of forming H-bonds.

NH3 vs HF

NH3 can form 3 H-bonds with neighbouring atoms:

HF can only form 1:

No. of Lone Pairs per Molecule The more lone pairs you have,
the higher the possibility of forming H-bonds.Ex:
NH3 vs H2ONH3 has 1 lone pair:H2O has 2:
Difference in Electronegativity The larger the difference in electronegativity,
the more exposed the H is.Ex:
NH3 vs HFF is the MOST electronegative element
There is a higher difference between H & F compared to H & N

However, ALL 3 of these factors must be taken into account before concluding which compound has a higher BP.
In the end, the compound which has a highest BP is H2O.

Let’s see why:

Compound No. of H No. of Lone Pairs Diff. in Electroneg.
NH3 3 1 Low
HF 1 3 Very High
H2O 2 2 High

Hydrogen bonds can form between different molecule types.
As long as there are exposed, covalently-bonded hydrogen atoms in one type of molecule & a lone pair in the other.
For example: Alcohol & Water
This is why, if you mix 50cm3 of water with 50cm3 of alcohol,
they are MISCIBLE,
& the resulting solution will be LESS than 100cm3.

Ice: A Strange Phenomenon
When most liquids freeze, they become more compact & more DENSE.
Water is the exception – it EXPANDS when it freezes.
This is because H2O can form hydrogen bonds when it freezes,
& the shape of the H2O molecule causes them to arrange in a crystal lattice structure:

This lattice takes up MORE volume than liquid water does.
Thus, the molecules are further apart in ice than in water.

Why Hydrogen Bonding is REALLY Important
Hydrogen bonding enables many organic compounds to link & form more complex chains.
Things like DNA & proteins rely on H-bonding to exist.

Living organisms depend on these bonds!

A Table for Comparison

Van Der Waals Dipole-Dipole Hydrogen
Weak Strong Strongest
In all molecules In polar molecules In hydrogen compounds with N, O or F with lone pairs

One thought on “CHEM C3: Intermolecular Forces

  1. Pingback: PHY C12: Thermal Energy & Thermal Capacity – ProDuckThieves

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