Chapter 6 Intermolecular Forces and Shapes

Question 1

What is the bond angle of methane?

Answer 1

109.5 degrees

Question 2

What is the effect of lone pairs on bond angle of 4 bonding regions?

Answer 2

Reduces bonding angle by 2.5 degrees

Question 3

What means towards you and away from you in displayed diagrams?

Answer 3

Solid Wedge= Towards us (too real too close)
Dashed Wedge= Away from us (running away, fading)

Question 4

What are electrons and what is the Electron Repulsion Theory?

Answer 4

Electrons are a cloud of negative charge.
The theory is that electrons pairs will repels each other, occupying the position which will minimise repulsion the most.

Question 5

What will be the shape for 4 electron pairs, 3 electron pairs and a lone pair, and 2 electron pairs and 2 lone pairs?

Answer 5

1= Tetrahedral
2= Pyramidal
3= Non Linear

Question 6

What is the shape and angle of a molecule with 3 bonding regions and no lone pairs? What is the shape and angle of a molecule with 6 bonding regions?

Answer 6

1=Trigonal Planar, 120 degrees
2=Octahedral, 90 degrees

Question 7

What are the general steps of determining the shape of a molecule?

Answer 7

1. Find the number of bonding regions (double bonds count as one entity) on the central atom
2. Find the number of lone pairs of electrons
3. The total number of both of these will determine the starting angle e.g PH3, 4, start with 109.5 degrees, minus 2.5*n.of.lone pairs
4. Figure out shape dependent on number of bonding regions and lone pairs leg PH3, 3 bonds, 1 lone, pyramidal

Question 8

Order the repulsion from least to most, with the different type of electron interactions

Answer 8

Bonded Pair/ Bonded Pair
Bonded Pair/ Lone Pair
Lone Pair/ Lone Pair

Question 9

Define a covalent bond

Answer 9

The strong electrostatic force of attraction between a shared pair of electrons and the nuclei of the bonded atoms

Question 10

What is electronegativity?

Answer 10

The measure of the tendency of an atom to attract a shared pair of electrons in a covalent bond

Question 11

When is a polar bond formed and why?

Answer 11

It is formed when there is a different in electronegativity between the atoms in a covalent bond. The one with higher electronegativity forms a slight negative change, pulling the electrons more, and the other a slight positive charge. It is caused by the unequal distribution of electron density in a covalent bond.

Question 12

What happens to electronegativity down a group and why?

Answer 12

Atomic radius increases as more shells are present
There is more electron shielding
Electronegativity decreases

Question 13

What happens to electronegativity across a period and why?

Answer 13

Same number of shells, so same amount of shielding
Nuclear charge increases as number of protons increases
Atomic radius increases slightly due to increased charge, pulled inwards
So electronegativity increases

Question 14

Define a dipole

Answer 14

The separation of opposite charges

Question 15

At what difference in the Pauling scale is a bond pure covalent, polar and ionic?

Answer 15

0-0.4= Pure Covalent
0.4-1.8=Polar
>1.8=Ionic

Question 16

What allows a molecule to be polar?

Answer 16

A difference in electronegativity
The molecules is not symmetrical, so the dipoles do not directly oppose and do not cancel out
(Dependent on shape)

Question 17

Define intermolecular forces

Answer 17

Weak interactions between dipoles of different molecules

Question 18

How do London forces occur?

Answer 18

Due to the random movement of electrons, there is an unequal distribution of negative charge. At any instant, an instantaneous dipole is produced.
An instantaneous dipole will induce a dipole on a neighbouring molecules
These induced dipoles will induce dipoles on further neighbouring molecules, which are then attracted to one another

Question 19

Factors affecting London Forces

Answer 19

The larger the number of electrons, the larger the instantaneous dipoles, producing greater attraction
More energy is needed to overcome the stronger London forces

Question 20

Why does H-Cl have a larger boiling point than F2?

Answer 20

F2 has only London forces to overcome.
H-Cl has both London forces and permanent-dipole-dipole interactions as it is a polar molecule.
The intermolecular forces is therefore stronger. This means more energy is needed to overcome the intermolecular forces and a higher melting point

Question 21

What is the arrangement of molecules in simple molecular substances? Why do they have low melting points?

Answer 21

In solid, simple molecular lattice
Molecules held together by weak intermolecular forces- London forces, maybe permanent dipole
These are broken at low temperatures, as little energy is needed to overcome them
Atoms in the molecule are covalently bonded, not overcome

Question 22

Solubility of non-polar simple molecular substances

Answer 22

Will dissolve in non-polar solvents
Interactions between non polar simple and non polar solvent, cause dissolving

Will not dissolve in polar solvents
Intermolecular forces within solvent too strong to overcome by interactions between the two

Question 23

Solubility of polar simple molecular substances

Answer 23

Depends on the strength of the dipoles, shape…
Some have hydrophobic and hydrophilic components

Question 24

Do simple molecular substances conduct electricity?

Answer 24

No. No mobile charge carriers

Question 25

What is a hydrogen bond?

Answer 25

A special type of permanent dipole-dipole interaction between hydrogen bonded to a more electronegative element and a lone pair of electrons on a very electronegative element (usually O, F or N).

Question 26

What are anomalous properties of water and why do they exist?

Answer 26

Solid less dense than liquid:
Hydrogen bonds position molecules further apart in the open lattice structure than in liquid. (2 lone pairs, 4 bonds, tetrahedral, 180 degree bond angle virtually)
Relatively high melting point
Both London forces and hydrogen bonds to overcome. Lots of energy required

Question 27

Why does ethanol have a higher boiling point than ethylamine?

Answer 27

The hydrogen bonds in ethanol are stronger than with ethylamine as oxygen is more electronegative than nitrogen

Question 28

How many hydrogen bonds can ethanol form?

Answer 28

2 donating, 2 electron pairs on the oxygen,
1 accepting of the hydrogen

Question 29

How many hydrogen bonds can water form?

Answer 29

2 donating (2 lone pairs)
2 accepting (2 H slight +)

Question 30

What is the shape of a complex with 2 lone pairs and 4 bonding regions?

Answer 30

Square planar
Lone pairs above and below