Shapes of Molecules and Intermolecular Forces (F6examonly)

Question 1

A lone pair decreases a bond angle by?

Answer 1

2.5°

Question 2

Shape and bond angle? Draw it out separately:
2 bonding pairs
0 lone pairs

Answer 2

linear
180°

Question 3

Shape and bond angle? Draw it out separately:
3 bonding pairs
0 lone pairs

Answer 3

trigonal planar
120°

Question 4

Shape and bond angle? Draw it out separately:
2 bonding pairs
1 lone pair

Answer 4

bent/non-linear
117.5°

Question 5

Shape and bond angle? Draw it out separately:
4 bonding pairs
0 lone pairs

Answer 5

tetrahedral
109.5°

Question 6

Shape and bond angle? Draw it out separately:
3 bonding pairs
1 lone pair

Answer 6

trigonal pyramidal
107°

Question 7

Shape and bond angle? Draw it out separately:
2 bonding pairs
2 lone pairs

Answer 7

bent/non-linear
104.5°

Question 8

Bond angle? Draw it out separately:
5 bonding pairs
0 lone pairs

Answer 8

trigonal bipyramidal
90°, 120°

Question 9

Bond angle? Draw it out separately:
6 bonding pairs
0 lone pairs

Answer 9

octahedral
90°

Question 10

What is the electron pair repulsion theory?

Answer 10

• e- pairs repel each other
• when around a central atom, e- pairs arrange themselves as far apart as possible to minimise repulsion
• this determines the molecules shape

Question 11

Order of relative repulsive strengths of bonded pairs and lone pairs of electrons:

Answer 11

increasing repulsion →
bond pair - bond pair < bond pair - lone pair < lone pair - lone pair

Question 12

Using electron pair repulsion theory, explain the shape and bond angles of CO2:

Answer 12

• the double bonds between C and O are treated as a bonding region
• therefore there are 2 bonding regions in CO2
• because these bonding regions repel each other, the bonding regions around the central atom (C) arrange themselves as far apart as possible to minimise repulsion
• this determines the linear shape of CO2, as the bond regions are as far apart as possible
• the bond angle is therefore 180°

Question 13

Using electron pair repulsion theory, explain the shape and bond angles of BF3:

Answer 13

• BF3 has 3 bonding pairs and 0 lone pairs
• because electron pairs repel, the electron pairs in around the central atom (B) will arrange themselves as far apart as possible to minimise repulsion
• this determines the shape of BF3: trigonal planar
• the bond angle is therefore 120°

Question 14

Using electron pair repulsion theory, explain the shape and bond angles of NH4-:

Answer 14

• NH4- has 4 bonding pairs and 0 lone pairs
• because electron pairs repel, the e- pairs around the central atom (N) arrange themselves as far apart as possible to minimise repulsion
• this determines the shape of NH4-: tetrahedral
• the bond angle is therefore 109.5°

Question 15

Using electron pair repulsion theory, explain the shape and bond angles of NH3:

Answer 15

• NH4- has 3 bonding pairs and 1 lone pair
• because electron pairs repel, the e- pairs around the central atom (N) arrange themselves as far apart as possible to minimise repulsion
• this determines the shape of NH3: trigonal pyramidal
• because lone pairs repel more than bonding pairs, the bond angle decreases by 2.5 for every lone pair (as they bonding pairs are pushed closer together)
• therefore the bond angle is 107°

Question 16

Using electron pair repulsion theory, explain the shape and bond angles of H2O:

Answer 16

• H2O has 2 bonding pairs and 2 lone pairs
• because electron pairs repel each other, the e- pairs around the central atom (O) arrange themselves to be as far apart as possible to minimise repulsion
• this determines the shape of H2O: non-linear
• because lone pairs repel more than bonding pairs, the bond angle decreases by 2.5° for each lone pair (as the bonding pairs are pushed closer together)
• since there are 2 lone pairs, the bond angle is 104.5°

Question 17

Using electron pair repulsion theory, explain the shape and bond angles of SF6:

Answer 17

• SF6 has 6 bonding pairs and 0 lone pairs
• because electron pairs repel each other, the e- pairs around the central atom (S) arrange themselves to be as far apart as possible to minimise repulsion
• this determines the shape of SF6: octahedral
• therefore the bond angle is 90°

Question 18

What is electronegativity?

Answer 18

an atom’s ability to attract the electron pair (bonding electrons) in a covalent bond

Question 19

Which element is the most electronegative? Which elements are also very electronegative?

Answer 19

• fluorine is the most electronegative
• oxygen, nitrogen, and chlorine are also very electronegative

Question 20

How is electronegativity measured?

Answer 20

using the Pauling scale - the higher the Pauline value, the higher the electronegativity

Question 21

How does electronegativity change down a group? Why?

Answer 21

• it decreases down a group
• although the nuclear charge increases down a group
• the increase in atomic radius and shielding outweighs this
• therefore the pull/attraction between the nucleus and the shared pair of e- decreases

Question 22

How does electronegativity change across a period? Why?

Answer 22

• it increases across a period
• the nuclear charge increases across a period
• the atomic radius decreases slightly across a period
• the shielding stays the same across a period
• so the pull/attraction between the nucleus and the shared pair of e- increases

Question 23

Explain a polar bond in terms of electronegativity:
Give an example of a polar bond:

Answer 23

• a polar bond occurs when the two atoms in a covalent bond have a different electronegativity
• the more electronegative atom has a greater share of the two electrons, as it has a stronger ability to attract the electron pair in the covalent bond
• therefore, because the two electrons are not shared equally, the more electronegative atom is slightly negative (𝛿-) and the less electronegative atom is slightly positive (𝛿+)
• e.g. H-Cl (H is 𝛿+ and Cl is 𝛿-)

Question 24

Explain why a permanent dipole occurs in a polar bond:

Answer 24

because the more electronegative atom has a greater share of the two electrons, there is a shift in electron density, causing a slight difference in charge between the two atoms - this is called a permanent dipole

Question 25

The greater the difference in electronegativity, the more polar a bond. True or false?

Answer 25

true

Question 26

Explain a non-polar bond in terms of electronegativity: Give an example of a non-polar bond:

Answer 26

- a non-polar bond occurs when the two atoms in a covalent bond have the same electronegativity - therefore both atoms have the same ability to attract the electron pair in the covalent bond - so the two electrons are shared equally - e.g. Cl-Cl

Question 27

Are C-H bonds considered polar even though there is a small difference in electronegativity between C and H?

Answer 27

no

Question 28

Explain a polar molecule in terms of the arrangement of polar bonds: Give an example of a polar molecule:

Answer 28

- if polar bonds in a polar molecule are arranged in a way that they don't cancel each other out, the charge is spread unevenly across the molecule, and there will be an overall dipole, making the molecule polar - e.g. H2O - the charge is unevenly distributed across the molecule, with oxygen side of the molecule being slightly more negative than the hydrogen side

Question 29

Why is a CF4 non-polar, even though it contains polar C-F bonds?

Answer 29

- although CF4 contains polar C-F bonds, where the fluorine is slightly more negative than the carbon - the dipoles are arranged in a way where they cancel each other out - so the molecule has no overall dipole and so is non-polar

Question 30

What are the three intermolecular forces you need to know?

Answer 30

- induced dipole-dipole forces (London forces) - permanent dipole-dipole interactions - hydrogen bonding

Question 31

What are induced dipole-dipole forces (London forces)? Where are they present?

Answer 31

- electrons are constantly moving around and there will be an uneven electron distribution at any given moment of time - this causes a temporary dipole within a molecule - this temporary dipole induces a temporary dipole in a neighbouring molecule - so there is an attraction between these molecules - theses forces are present in all molecular substances

Question 32

What affects the strength of London forces?

Answer 32

- the size of the molecule (the number of e-): the bigger the molecule, the more electrons it has, and so the greater the London forces - the shape of the molecule: straight chain molecules have stronger London forces than branched molecules because the straight chain molecules have more points of contact between them

Question 33

What are permanent dipole-dipole interactions? Where are they present?

Answer 33

- the 𝛿+ and 𝛿- charges on polar molecules cause weak electrostatic forces of attraction between molecules - they are present between polar molecules - they are not present between non-polar molecules, even if they have polar bonds, e.g. CO2

Question 34

What is hydrogen bonding?

Answer 34

the attraction between the hydrogen bonded to an N, O, or F on one molecule and the lone pair of a N, O, or F atom on another molecule

Question 35

What are the relative strengths of intermolecular forces?

Answer 35

H-bonding > permanent dipole-dipole > London forces

Question 36

Are intermolecular forces stronger than covalent bonds?

Answer 36

no, they are relatively very weak

Question 37

Why does water have relatively high melting and boiling points?

Answer 37

- there is hydrogen bonding between water molecules - this occurs between the H on one water molecule and the lone pair of e- on the O of another water molecule - because hydrogen bonds are relatively strong compared to other intermolecular bonds, more energy is required to break the bonds, resulting in high melting and boiling points

Question 38

Why is ice less dense than liquid water (in terms of hydrogen bonding)?

Answer 38

- in ice, water molecules are held together in a lattice by hydrogen bonds - when ice melts into liquid water, hydrogen bonds break, so there are less hydrogen bonds in liquid water compared to ice - because hydrogen bonds are relatively long, the water molecules are held farther apart in ice than in liquid water, as there are more hydrogen bonds in ice - this means that ice is less dense that liquid water