6 Shapes of molecules and intermolecular forces

Question 1

What is the electron-pair repulsion theory?

Answer 1

An electron has a negative charge, so electron pairs repel one another

Question 2

What do the electron pairs impact?

Answer 2

• Electron pairs surrounding a central atom determine the shape of the molecule or ion
• Repel one another so that they are arranged as far apart as possible
• Arrangement of EP minimises repulsion and thus holds the bonded atoms in a definite shape
• Different numbers of electron pairs result in different shapes

Question 3

What is the difference between bonded pairs vs lone pairs repulsions?

Answer 3

A lone pair of electrons is slightly closer to the central atom and occupies more space than a bonded pair
The lone pair repels more strongly than a bonding pair

Question 4

How much is the bond angle reduced by for each lone pair?

Answer 4

2.5 degrees
This is because lone pairs repel bonded pairs slightly closer together, decreasing the bond angle

Question 5

What is a bonding region?

Answer 5

A pair or pairs of electrons occupying one area in space

Question 6

Define electronegativity

Answer 6

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

Question 7

When will a covalent bond possibly experience more attraction from one of the bonded atoms than the other?

Answer 7

The nuclear charges are different
The atoms may be different sizes
The shared pair of electrons may be closer to one nucleus than the other

Question 8

What are the trends across a period?

Answer 8

Electronegativity increases
- Atomic radius decreases
- More nuclear charge (more protons)
Therefore stronger attraction between nucleus and bonding pair of electrons

Question 9

What are the trends down a group?

Answer 9

• Atomic radius increases –> more electron shells
• More shielding –> electrons between the electron of interest and positive nucleus therefore less attraction between nucleus and bonding pair of electrons

Question 10

What are the factors affecting electronegativity?

Answer 10

Nuclear charge
Atomic radius
Shielding

Question 11

What is the electronegativity difference numbers?

Answer 11

0-0.4 Nonpolar Covalent Bond
0.4-1.8 Polar Covalent Bond
1.8+ Ionic Bond

Question 12

Describe the characteristics of non-polar covalent bonds

Answer 12

Atoms have equal electronegativities so electrons equally attracted to both nuclei

Question 13

Describe the characteristics of polar covalent bonds

Answer 13

The difference in electronegativity between the two atoms causes a dipole

Question 14

What is a dipole

Answer 14

The difference in charge between two atoms caused by a shift in electron density in the bond

Question 15

What happens when dipoles act in the same direction?

Answer 15

Overall dipoles cancel and non-polar molecule

Question 16

Why are ionic compounds soluble in polar solvents e.g. water?

Answer 16

E.g. Sodium chloride dissolved
Water molecules attract Na+ (attracted to oxygen) and Cl- ions (attracted to hydrogen)
The ionic lattice breaks down as it dissolves
Water molecules surround the Na+ and Cl- ions

Question 17

What are intermolecular forces?

Answer 17

The weak interactions between dipoles of different molecules
Exist between all simple covalent molecules

Question 18

What are the three types of intermolecular forces?

Answer 18

London forces (induced dipole-dipole interactions)
Permanent dipole-dipole interactions
Hydrogen bonding

Question 19

What do intermolecular forces determine?

Answer 19

Physical properties such as melting and boiling points

Question 20

What are London Forces?

Answer 20

Weak intermolecular forces that exist between all molecules, whether polar or non-polar

Question 21

Explain induced dipole-dipole interactions (London Forces)

Answer 21

1. As electrons orbit there are fluctuations in the electron density forming a temporary, changing dipole –> an instantaneous dipole will exist, but its position is constantly shifting
2. Instantaneous dipole induces a dipole on a neighbouring molecule
3. The induced dipole induces further dipoles on neighbouring molecules which then attract one another

Question 22

What happens to the strength of London forces the more electrons in the molecule?

Answer 22

The larger the instantaneous and induced dipoles
The greater the induced dipole-dipole interactions
The stronger the attractive force between the molecules

Question 23

What are the physical properties of simple molecular substances and why?

Answer 23

Low melting and boiling points
- Atoms within each molecules are bonded together strongly by covalent bonds but the molecules themselves are held in place by only weak intermolecular forces
- Only the weak intermolecular forces break
- The weak intermolecular forces can be broken by energy at low temps

Question 24

What are permanent dipole-dipole interactions?

Answer 24

Permanent dipoles in polar molecules (also have induced dipole interactions as well)

Question 25

What are the patterns of solubility of non-polar simple substances? Why is this?

Answer 25

Non polar simple molecular substances are soluble in non-polar solvents. - Intermolecular forces form between the molecules and the solvent -Interactions weaken the intermolecular forces in the simple molecular lattice - the intermolecular forces break and the compound dissolves

Question 26

Why are non-polar simple molecular substances insoluble in polar solvents?

Answer 26

Insoluble in polar solvents (like water) The intermolecular forces within the polar solvent are too strong to be broken There is little interaction between the molecules in the lattice and the polar solvent molecules

Question 27

What are polar simple molecular substances solubility?

Answer 27

Often polar covalent substances dissolve in polar solvents They do not dissolve in non-polar solvents as the intermolecular forces are too strong and are not overcome by interactions (London forces) between the solvent and the molecule

Question 28

What does solubility depend on?

Answer 28

The strength of the dipole and how many polar groups are present, it can be difficult to predict

Question 29

What do some biological molecules have? (solubility)

Answer 29

Hydrophobic (non polar, comprised of a carbon chain) and hydrophilic (polar, contains electronegative atoms that interact with water) parts of the molecule

Question 30

What is a hydrogen bond?

Answer 30

The force of attraction between the slightly positive hydrogen atom in one molecule and the lone electron pair on the nitrogen, oxygen or fluorine atom in another molecule

Question 31

What do molecules contain to have a hydrogen bond?

Answer 31

- An electronegative atom with a lone pair of electrons e.g. oxygen, nitrogen or fluorine - A hydrogen atom attached to an electronegative atom, for example H-O, H-N or H-F

Question 32

How strong are hydrogen bonds relative to the other IMF bonding types?

Answer 32

Stronger than London or permanent dipole-dipole interactions, but the latter two will also be present in addition to any hydrogen bonds between molecules

Question 33

What are the 2 anomalous properties that water has?

Answer 33

Ice is less dense than the liquid Water has a relatively high melting point and boiling point

Question 34

Why is ice less dense than water?

Answer 34

- Hydrogen bonds hold water molecules apart in an open lattice structure - two lone pairs on the oxygen atom and two hydrogen atoms, each water molecule can form four hydrogen bonds - Water molecules in ice are further apart than in water - Solid ice is less dense than liquid water and floats --> the holes in the open lattice structure decrease the density of water on freezing. Melting the ice lattice collapses and the molecules move closer together

Question 35

Why does water have a relatively high melting and boiling point?

Answer 35

- Hydrogen bonds are extra forces, over and above the London forces - Good quantity of energy is needed to break the hydrogen bonds, so water has much higher melting and boiling points than would be expected from just London Forces - When the ice lattice breaks, the rigid arrangement of hydrogen bonds in ice is broken. When water boils, the hydrogen bonds break completely