Chapter 6 - Shapes of Molecules + Intermolecular Forces

LINEAR shape

bond angle = 180, has 2 bonding pairs

example = carbon dioxide

TRIGONAL PLANAR shape

bond angle = 120, has 3 bonding pairs

example = boron trifluoride

TETRAHEDRAL shape

bond angle = 109.5, has 4 bonding pairs

example = methane

OCTAHEDRAL shape

bond angle = 90, has 6 bonding pairs

example = sulfur hexafluoride

how many degrees do lone pairs decrease the bond angle?

2.5

PYRAMIDAL shape

bond angle = 107, has 3 bonding pairs, has 1 lone pair

example = ammonia

NON-LINEAR shape

bond angle = 104.5, has 2 bonding pairs, has 2 lone pairs

example = water

electronegativity

the ability of an atom to attract bonding electrons in a covalent bond

TRENDS in electronegativity -period

electronegativity increases

-nuclear charge increases
-atomic radius decreases
-electron shielding remains the same
-nuclear attraction experienced by bonding electrons increases

TRENDS in electronegativity -group

electronegativity decreases

-atomic radius increases
-increase in nuclear charge is outweighed by increase in electron shielding
-nuclear attraction decreases

permanent dipole

there is a small charge difference across a bond resulting in a large difference in electronegativity (>1.4) of the bonded atoms

polar covalent bond

-when a covalent bond contains a permanent dipole

-the bonded electron pair is shared unequally as the bonded atoms are different and have different electronegativity values so one element is more electronegative creating the dipoles

covalent bond with no permanent dipole

-the bonded electron pair is shared equally as the bonded atoms are the same and have similar/the same electronegativity values

polar molecule

-has an overall dipole

-the permanent dipoles do not cancel each other out because the bonds are arranged asymmetrically as they act in different directions

non-polar molecule

-no overall dipole

-the bonds are arranged symmetrically so the dipoles act in equal and opposite directions so cancel each other out as they oppose each other

TYPES of intermolecular forces

1) London forces (induced dipole-dipole)

2) permanent dipole-dipole interactions

3) hydrogen bonding

-also in increasing strength order

London forces

-weakest type of intermolecular force

-exists between all molecules (polar or non-polar)

STRENGTH of London forces

the greater the number of electrons in the molecule, the larger the induced dipoles so the stronger the attractive forces between the molecules so the melting and boiling point increase as more energy is needed to overcome

permanent dipole-dipole interactions

-exist in polar molecules that contain permanent dipoles

-are stronger and never change unlike London forces but not as common

-S+ region of one molecule is attracted to S- region of another molecule

hydrogen bonding

-need very large difference in electronegativity to exist → fluorine, nitrogen and oxygen are the most electronegative

-strong intermolecular forces formed between molecules containing N-H, F-H AND O-H

effects of HYDROGEN BONDING on WATER 1)

ice is less dense than liquid water

-in solid ice, the water molecules are held apart by hydrogen bonds so causes it to have an open lattice structure

-when heated, the hydrogen bonds break and the water molecules can move closer together so liquid water is denser

effects of HYDROGEN BONDING on WATER 2)

water has a relatively high melting/boiling point

-has hydrogen bonding which is stronger than other intermolecular forces so more heat energy required to overcome