CHEM1001 Lecture 21 Worksheet Polar Bonds, Polar Molecules & Intermolecular Forces

You already know about the forces within molecules – bonds. Intermolecular Forces, the forces between molecules, depend on two things: (i) the shape of the molecule and (ii) how the charge (electron cloud) is distributed.

1. Molecular Shape (VSEPR Revision – you can start this before lecture) As you already know, we can use Lewis structure to describe the bonding and then VSEPR theory to predict molecular shape. Let’s compare the shapes of the first row hydrides CH4, NH3, H2O and HF using VSEPR. The first thing we need to do is determine how many (valence) electron pairs there are around the central atom of each compound so we can work out the Lewis structure.

(a) Why is H never the central atom in a compound?

(b) Calculate the total number of valence electrons around the central atom in each compound, and then the total number of electron pairs. (Leave the last column empty for now.

Valence electrons electron pairs bonds

CH4

NH3

H2O

HF

(c) They’re all the same, right? Now draw out each of their Lewis Structures below.

(d) Circle the arrangement or shape that these electron pairs make around the central atom.

Linear trigonal planar tetrahedral trigonal bipyramidal octahedral

When we describe a molecule (or an ion) we are describing the arrangements of the atoms, not the electron pairs. The electrons are basically invisible, so we have to visualize how the remaining atoms (hydrogens in this case) are arranged around the central atom.

(e) Connect each of these four hydrides to the correct molecular shape

CH4 NH3 H2O HF

linear tetrahedral bent trigonal pyramid

2. Bond Polarity and Molecular Polarity

The sharing of electrons between unlike atoms in a bond can lead to a slight excess of negative charge (-) around the more electronegative end of a bond, leaving a slight shortfall (+) at the less electronegative end. This creates a bond dipole.

(a) Sketch HF and, using what you know about electronegativity, label the H and F atoms to show which has excess positive and negative charges.

(b) We can also label a dipole using and arrow running from the positive to the negative end of the bond thus: . We cross the arrow as shown to remind us which end is “+ve.” Draw an arrow to label the dipole in the HF bond.

If the (only) bond in a diatomic molecule is polar, then the molecule is polar (of course).

(c) Decide which of the following molecules will be polar and draw the dipole arrow.

N2 H2 HBr CO Br2 BrCl TiO

(d) In polyatomic molecules the bond dipoles may cancel out making the molecule non-polar, or they may not quite cancel out, giving the molecule a net or overall dipole, and making it polar. This is easiest to see in molecules with six electron pairs around the central atom, based on the octahedron. Look at the VSEPR structures of the molecules below, and decide which bond dipoles do and do not cancel, and hence label the dipoles of the polar molecule(s).

Br

F

FF

FF