M2 - Chemical Bonding (2)

Question 1

Explain what is meant by the term electronegativity [2].

Answer 1

The ability of an atom to attract electrons ✔
(Electron pair) in a (covalent) bond ✔

Question 2

How is electronegativity measured?

Answer 2

On the Pauling scale

Question 3

What does a higher Pauling value mean?

Answer 3

A higher electronegativity and thus a greater attraction for an electron pair in a covalent bond

Question 4

What makes a bond polar? Provide an example of a polar bond

Answer 4

Two atoms with differing electronegativities, bonding electrons are pulled towards the more electronegative atom, making it polar.

For example, H-Cl which has a polar bond.

Question 5

What force does a polar bond form?

Answer 5

A permanent dipole-dipole force

Question 6

How are the differences in electronegatives displayed?

Answer 6

With the use of a dipole:
𝛿− is more electronegative
𝛿+ is less electronegative

Question 7

Are diatomic gases (e.g. H₂, Br₂, N₂, I₂ etc) polar or non-polar, why?

Answer 7

Non-polar
Because the atoms have equal electronegativities and so the electrons are equally attracted to both nuclei

Question 8

What happens to the polarity of the molecule if polar bonds are arranged symmetrically, for example in CCl₄

Answer 8

• The CCl₄ molecule is symmetrical
• The dipoles cancel each other out
• Molecule has no overall dipole and is non-polar

Question 9

What happens to the polarity of the molecule if polar bonds are arranged unsymmetrically, for example in CHCl₃

Answer 9

• The CHCl₃ molecule is unsymmetrical
• The dipoles do not cancel each other out
• Molecule has an overall dipole and is polar

NOTE: Lone pairs in a compound automatically make them a Polar molecule

Question 10

What is the general strength of intermolecular forces (IMFs)?

Answer 10

Very weak intermolecular forces between molecules

Question 11

What is the order of the intermolecular forces from strongest to weakest?

Answer 11

1) Hydrogen bonding
2) Permanent dipole-dipole forces
3) Induced dipole-dipole forces

Question 12

How are induced dipole-dipole forces formed?

Answer 12

• At any moment there may be an uneven distribution of electrons in a molecule
• This causes a temporary dipole
• The instantaneous dipole in one molecule causes an induced dipole in a neighbouring molecule
• The δ+ of a dipole in one molecule attracts the
  δ- of a dipole in a neighbouring molecule to produce an induced dipole-dipole force

Question 14

When can hydrogen bonding occur?

Answer 14

When hydrogen is covalently bonded to fluorine, nitrogen or oxygen.

Question 15

Why are hydrogen bonds brought about?

Answer 15

Hydrogen has a high charge densiy and F, N and O are very electronegative. The bond is so polarised that a weak bond forms between the hydrogen of one molecule and a lone pair on a neighbouring molecule’s F, N or O.

Question 16

What effect does hydrogen bonding have on a molecule?

Answer 16

Soluble in water.
Higher boiling and freezing points than molecules of a similar size that don’t form hydrogen bonds.

Question 17

What is an interesting property of ice caused by hydrogen bonding?

Answer 17

In ice, water molecules are held together in a lattice. When ice melts, hydrogen bonds are broken, so ice has more of these than water. As hydrogen bonds are long, this causes ice to be less dense than water.

Question 18

How do intermolecular forces explain simple covalent compounds having low melting and boiling points?

Answer 18

Weak intermolecular forces to overcome.

Question 19

How do intermolecular forces explain simple covalent compounds sometimes being soluble in water?

Answer 19

Water is also a polar molecule, hydrogen bonded molecules can form these with water molecules so are soluble.

Question 20

How do intermolecular forces explain simple covalent compounds not conducting electricity?

Answer 20

Overall covalent molecules are uncharged, permanent dipoles are not strong enough.

Question 21

what is the shape of molecules table?

Answer 21

Question 22

Describe and explain two anomalous properties of water from hydrogen bonding (4)

Answer 22

Liquid H2O is more dense than solid
Ice has open lattice
H2O has a relatively high boiling point
Hydrogen bonds need to be broken

Question 23

Magnesium melting point: 650°C
Chlorine melting point: -101°C
Describe the structure and bonding of these elements and explain the difference in melting points (6)

Answer 23

Mg has a giant structure
Mg has metallic bonding
Electrostatic attraction between positive ions and electrons
Cl has a simple covalent lattice
Cl has induced dipole-dipole
Less energy is needed to overcome induced dipole-dipole than metallic bonds

Question 24

Why does H2S have a much lower boiling point than H2O? (2)

Answer 24

No hydrogen bonding

Weaker intermolecular forces

Question 25

What is meant by hydrogen bonding? (1)

Answer 25

Interaction between lone pair of F, N or O and a H in a neighbouring molecule

Question 26

Explain electronegativity (2)

Answer 26

Attraction of an atom for electrons in a covalent bond

Question 27

Why do O-H and N-H bonds have dipoles? (1)

Answer 27

Oxygen and nitrogen are more electronegative than hydrogen

Question 28

What is a multiple covalent bond? Provide some examples

Answer 28

Two atoms share more than one pair of electrons:

Question 29

What is a dative covalent bond? Provide an example of a compound that has a dative covalent bond

Answer 29

The shared pair of electrons has been supplied by one of the bonding atoms only.

The N atom on the ammonia molecule (NH₃) donates its lone pair of electrons to an H⁺ ion forming an ammonium ion, NH₄⁺.

Question 30

Why is water a polar molecule, but CO2 is not?

Answer 30

CO2 is symmetrical

In CO2, dipoles cancel