Bonding

Question 1

Define ionic bonding

Answer 1

the strong electrostatic attraction between oppositely charged ions

Question 2

Give five properties of ionic compounds

Answer 2

• high melting and boiling points
• don’t conduct electricity when solid
• conduct electricity when molten / aqueous
• soluble in water (and other polar solvents)
• hard brittle crystalline substances

Question 3

Why do ionic compounds have high melting and boiling points?

Answer 3

strong electrostatic attraction between oppositely charged ions requires a lot of energy to break

Question 4

Why do ionic compounds not conduct electricity when solid?

Answer 4

ions are in fixed positions so they are not free to move and can’t carry a charge

Question 5

Why do ionic compounds conduct electricity when molten / aqueous?

Answer 5

ions are free to move and carry charge

Question 6

Why are ionic compounds soluble in water and other polar solvents?

Answer 6

the delta positive Hydrogen atoms of water are attracted to negative ions; the delta negative Oxygen atoms are attracted to positive ions

this forms ion-dipole interactions and releases energy, which is enough to break the ionic lattice apart

Question 7

Why are ionic compounds hard and crystalline substances?

Answer 7

they have a very regular arrangement of ions in a giant lattice

Question 8

Why are ionic compounds brittle?

Answer 8

applying a force causes layers to slide; like charged ions line up and repel

Question 9

What properties of ionic compounds provide evidence for the ionic model?

Answer 9

• high melting and boiling point
• conduct electricity when molten and aqueous
• brittle
• tend to be soluble in water
• migration of ions

Question 10

Define ionic radius

Answer 10

a measure of the space occupied by an ion in a crystal lattice

Question 11

Explain the trend in ionic radius down a group

Answer 11

ionic radius increases; number of shells increase

Question 12

Explain the trend in ionic radius along a period

Answer 12

Group 1, 2, and 3 ions are isoelectronic; atomic radius decreases as nuclear charge increases so electron cloud is drawn in

there is a big jump in ionic radius between group 3 and 5 as an extra shell is added; Group 5, 6, and 7 are also isoelectronic; radius decreases again due to increasing nuclear charge (extra shell means that decrease is less marked as more shielding)

Question 13

Why is the ionic radii of positive ions smaller than their corresponding atomic radii?

Answer 13

they have lost their outer shell electrons so have fewer quantum shells

Question 14

Why is the ionic radii of negative ions greater than their corresponding atomic radii?

Answer 14

although they have the same number of quantum shells as their corresponding atoms, they have more electrons so there are increased repulsions and more electrons being attracted to the same nuclear charge, increasing the ionic radii

Question 15

Define lattice energy

Answer 15

the energy change when one mole of an ionic solid is formed from its gaseous ions

more energy released indicates stronger ionic bonding

Question 16

Define metallic bonding

Answer 16

strong electrostatic attraction between cations and delocalised electrons

Question 17

Give 3 properties of metallic bonding

Answer 17

• high melting and boiling points
• good conductor of electricity
• good conductor of heat

Question 18

Explain why magnesium (2) has a higher melting point than sodium (1)

Answer 18

magnesium has 2 delocalised electrons per atom whereas sodium has 1 in the sea of electrons; magnesium also has a smaller ionic radius, leading to greater charge density

this means that magnesium has stronger metallic bonds than sodium

Question 19

Define covalent bond

Answer 19

strong electrostatic attraction between two nuclei and the shared pair of electrons between them

Question 20

What is a lone pair?

Answer 20

a pair of electrons in the valence shell which are not involved in bonding

Question 21

What is a dative covalent / coordinate bond?

Answer 21

a bond in which two atoms share a pair of electrons, both of the electrons donated by one atom

Question 22

What is a sigma bond?

Answer 22

overlap of 2x s orbitals, 1x s orbital and 1x p orbital, or 2x p orbitals
can rotate

Question 23

When do pi bonds occur?

Answer 23

occur in double or triple bonds (first bond is sigma bond)

Question 24

Where does electron density lie in a pi bond?

Answer 24

above and below the plane of the molecule

this means pi bonds can’t rotate

Question 25

Explain why a sigma bond is stronger than a pi bond

Answer 25

sigma bonds are stronger because electron density is between the two nuclei whereas in pi bonds, the electron density lies above and below the plane of the molecule

Question 26

What is the general relationship between bond length and bond strength?

Answer 26

the shorter the bond length, the stronger the bond

Question 27

Explain why the Cl-Cl (period 3) bond is stronger than the Br-Br (period 4) bond

Answer 27

Cl atoms have a smaller atomic radius than Br atoms, so there is less shielding and the shared pair of electrons in the bond are closer to the two nuclei, meaning the bond is shorter

Question 28

Explain why the double bond is stronger than the single bond

Answer 28

two pairs of electrons are shared in the double bond so electron density between the two nuclei increases so there is stronger attraction of the nuclei to the electrons and the atoms are pulled in tighter

Question 29

Why is the F-F bond shorter than the Cl-Cl bond, but much weaker?

Answer 29

atomic radius of F is so much smaller than Cl; lone pairs around the F atoms repel and this weakens the F-F bond

Question 30

Give three properties of simple molecular structures

Answer 30

- low melting and boiling points - don't conduct electricity as solids, liquids or gases - more soluble in non-polar solvents than water

Question 31

Explain why simple molecular structures have low melting and boiling points

Answer 31

weak intermolecular forces between molecules; doesn't require much energy to break

Question 32

Explain why simple molecular structures don't conduct electricity

Answer 32

molecules are not charged; no ions or delocalised electrons to carry charge

Question 33

Explain why simple molecular structures are more soluble in non-polar solvents than water

Answer 33

the strength of the intermolecular forces are equal to or greater than the strength of the intermolecular forces within the solvent

Question 34

How do electrons arrange themselves around the central atom?

Answer 34

in order to minimise repulsions and maximise separation

Question 35

Which of lone pairs and bonding pairs have stronger repulsions?

Answer 35

lone pairs repel more strongly than bonding pairs

Question 36

Describe how to answer an exam question about shapes of molecules

Answer 36

central atom has () bonding pairs and () lone pairs; these are arranged to maximise separation and minimise repulsion ## Footnote shape: bond angle(s): draw 3D diagram

Question 37

Give the shape and bond angle of a molecule with 2 bonding pairs and 0 lone pairs

Answer 37

linear 180*

Question 38

Give the shape and bond angle of a molecule with 2 bonding pairs and 1 lone pair

Answer 38

bent 118*

Question 39

Give the shape and bond angle of a molecule with 2 bonding pairs and 2 lone pairs

Answer 39

bent 104.5*

Question 40

Give the shape and bond angle of a molecule with 3 bonding pairs and 0 lone pairs

Answer 40

trigonal planar 120*

Question 41

Give the shape and bond angle of a molecule with 3 bonding pairs and 1 lone pair

Answer 41

trigonal pyramidal 107*

Question 42

Give the shape and bond angle of a molecule with 3 bonding pairs and 2 lone pairs

Answer 42

T shape 89*

Question 43

Give the shape and bond angle of a molecule with 4 bonding pairs and 0 lone pairs

Answer 43

tetrahedral 109.5*

Question 44

Give the shape and bond angle of a molecule with 4 bonding pairs and 1 lone pair

Answer 44

seesaw 89 and 119

Question 45

Give the shape and bond angle of a molecule with 4 bonding pairs and 2 lone pairs

Answer 45

square planar 90*

Question 46

Give the shape and bond angle of a molecule with 5 bonding pairs and 0 lone pairs

Answer 46

trigonal bipyramidal 120 and 90

Question 47

Give the shape and bond angle of a molecule with 5 bonding pairs and 1 lone pair

Answer 47

square pyramid 89*

Question 48

Give the shape and bond angle of a molecule with 6 bonding pairs

Answer 48

octahedral 90*

Question 49

Define electronegativity

Answer 49

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

Question 50

Give 3 factors affecting electronegativity

Answer 50

- nuclear charge - atomic radius - shielding

Question 51

Explain the trend in electronegativity down a group

Answer 51

decreases; more shells so greater atomic radius ## Footnote this decreases the attraction of the bonding pair to the nucleus

Question 52

Explain the trend in electronegativity across a period

Answer 52

increases; greater nuclear charge and decreased atomic radius (shielding remains the same) ## Footnote this increases the attraction of the bonding pair to the nucleus

Question 53

Describe what a pure covalent bond is in terms of electronegativity

Answer 53

0 - 0.5

Question 54

Describe what a polar covalent bond is in terms of electronegativity

Answer 54

0.5 - 1.7

Question 55

Describe what an ionic with covalent character bond is in terms of electronegativity

Answer 55

1.7 - 2.0

Question 56

Describe what an ionic bond is in terms of electronegativity

Answer 56

2.0 - 4.0

Question 57

What is a polar molecule?

Answer 57

a molecule with polar bonds (0.5 - 1.7) with polarities that don't cancel out

Question 58

Explain how London forces arise

Answer 58

fluctuations in the electron clouds of molecules set up temporary dipole; this induced temporary dipoles in a second molecule

Question 59

What is the relationship between number of electrons and the strength of London forces?

Answer 59

more electrons = stronger London forces

Question 60

Explain how permanent dipole interactions occur

Answer 60

set up between polar molecules; molecules with polar bonds that are not symmetrical so dipoles don't cancel out

Question 61

Explain how hydrogen bonds arise

Answer 61

a hydrogen atom bonded to N, O, F; lone pairs on an N, O, F on a second molecule

Question 62

What must you show on a drawing of hydrogen bonds?

Answer 62

lone pairs on N, O, F; dipoles; Hydrogen bond as dashed line; angle of 180* labelled

Question 63

Explain the trend in boiling points for alkanes with the same molecular formula with different amounts of branching

Answer 63

the alkanes with the longest chain and least branching have the highest boiling point ## Footnote molecules can lie closer together so more points of contact and stronger London forces

Question 64

What intermolecular forces are present in water?

Answer 64

london forces; permanent - dipole interactions; hydrogen bonding

Question 65

What intermolecular forces are present in hexane?

Answer 65

london forces

Question 66

How soluble are ionic compounds in water?

Answer 66

mostly

Question 67

How soluble are ionic compounds in hexane?

Answer 67

insoluble

Question 68

How soluble are compounds with hydrogen bonding in water?

Answer 68

usually soluble (short chains at least)

Question 69

How soluble are compounds with hydrogen bonding in hexane?

Answer 69

usually insoluble

Question 70

How soluble are non-polar substances in water?

Answer 70

insoluble

Question 71

How soluble are non-polar substances in hexane?

Answer 71

soluble

Question 72

Explain why ionic compounds are mostly soluble in water

Answer 72

ion-dipole interactions are formed between oppositely charged ions in the ionic lattice and dipoles which releases energy ## Footnote if enough energy is released then the lattice breaks down and ions become hydrated

Question 73

Why are compounds with hydrogen bonding usually soluble in water?

Answer 73

compounds can form hydrogen bonds with water; the energy released from forming these hydrogen bonds is enough to break the hydrogen bonds within water

Question 74

Why are non-polar substances insoluble in water?

Answer 74

e.g. hexane; hexane won't form hydrogen bonds with water; only forms London forces so there isn't enough energy released to break the hydrogen bonds within water