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Physical Chemistry > Bonding > Flashcards

Bonding Flashcards

1
Q

Define ionic bonding

A

the strong electrostatic attraction between oppositely charged ions

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2
Q

Give five properties of ionic compounds

A
  • 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
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3
Q

Why do ionic compounds have high melting and boiling points?

A

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

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4
Q

Why do ionic compounds not conduct electricity when solid?

A

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

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5
Q

Why do ionic compounds conduct electricity when molten / aqueous?

A

ions are free to move and carry charge

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6
Q

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

A

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

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7
Q

Why are ionic compounds hard and crystalline substances?

A

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

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8
Q

Why are ionic compounds brittle?

A

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

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9
Q

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

A
  • high melting and boiling point
  • conduct electricity when molten and aqueous
  • brittle
  • tend to be soluble in water
  • migration of ions
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10
Q

Define ionic radius

A

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

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11
Q

Explain the trend in ionic radius down a group

A

ionic radius increases; number of shells increase

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12
Q

Explain the trend in ionic radius along a period

A

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)

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13
Q

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

A

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

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14
Q

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

A

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

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15
Q

Define lattice energy

A

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

more energy released indicates stronger ionic bonding

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16
Q

Define metallic bonding

A

strong electrostatic attraction between cations and delocalised electrons

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17
Q

Give 3 properties of metallic bonding

A
  • high melting and boiling points
  • good conductor of electricity
  • good conductor of heat
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18
Q

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

A

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

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19
Q

Define covalent bond

A

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

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20
Q

What is a lone pair?

A

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

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21
Q

What is a dative covalent / coordinate bond?

A

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

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22
Q

What is a sigma bond?

A

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

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23
Q

When do pi bonds occur?

A

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

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24
Q

Where does electron density lie in a pi bond?

A

above and below the plane of the molecule

this means pi bonds can’t rotate

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25
Q

Explain why a sigma bond is stronger than a pi bond

A

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

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26
Q

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

A

the shorter the bond length, the stronger the bond

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27
Q

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

A

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

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28
Q

Explain why the double bond is stronger than the single bond

A

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

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29
Q

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

A

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

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30
Q

Give three properties of simple molecular structures

A
  • low melting and boiling points
  • don’t conduct electricity as solids, liquids or gases
  • more soluble in non-polar solvents than water
31
Q

Explain why simple molecular structures have low melting and boiling points

A

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

32
Q

Explain why simple molecular structures don’t conduct electricity

A

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

33
Q

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

A

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

34
Q

How do electrons arrange themselves around the central atom?

A

in order to minimise repulsions and maximise separation

35
Q

Which of lone pairs and bonding pairs have stronger repulsions?

A

lone pairs repel more strongly than bonding pairs

36
Q

Describe how to answer an exam question about shapes of molecules

A

central atom has () bonding pairs and () lone pairs; these are arranged to maximise separation and minimise repulsion

shape:
bond angle(s):
draw 3D diagram

37
Q

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

A

linear
180*

38
Q

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

A

bent
118*

39
Q

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

A

bent
104.5*

40
Q

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

A

trigonal planar
120*

41
Q

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

A

trigonal pyramidal
107*

42
Q

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

A

T shape
89*

43
Q

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

A

tetrahedral
109.5*

44
Q

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

A

seesaw
89 and 119

45
Q

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

A

square planar
90*

46
Q

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

A

trigonal bipyramidal
120 and 90

47
Q

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

A

square pyramid
89*

48
Q

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

A

octahedral
90*

49
Q

Define electronegativity

A

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

50
Q

Give 3 factors affecting electronegativity

A
  • nuclear charge
  • atomic radius
  • shielding
51
Q

Explain the trend in electronegativity down a group

A

decreases; more shells so greater atomic radius

this decreases the attraction of the bonding pair to the nucleus

52
Q

Explain the trend in electronegativity across a period

A

increases; greater nuclear charge and decreased atomic radius (shielding remains the same)

this increases the attraction of the bonding pair to the nucleus

53
Q

Describe what a pure covalent bond is in terms of electronegativity

A

0 - 0.5

54
Q

Describe what a polar covalent bond is in terms of electronegativity

A

0.5 - 1.7

55
Q

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

A

1.7 - 2.0

56
Q

Describe what an ionic bond is in terms of electronegativity

A

2.0 - 4.0

57
Q

What is a polar molecule?

A

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

58
Q

Explain how London forces arise

A

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

59
Q

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

A

more electrons = stronger London forces

60
Q

Explain how permanent dipole interactions occur

A

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

61
Q

Explain how hydrogen bonds arise

A

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

62
Q

What must you show on a drawing of hydrogen bonds?

A

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

63
Q

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

A

the alkanes with the longest chain and least branching have the highest boiling point

molecules can lie closer together so more points of contact and stronger London forces

64
Q

What intermolecular forces are present in water?

A

london forces; permanent - dipole interactions; hydrogen bonding

65
Q

What intermolecular forces are present in hexane?

A

london forces

66
Q

How soluble are ionic compounds in water?

A

mostly

67
Q

How soluble are ionic compounds in hexane?

A

insoluble

68
Q

How soluble are compounds with hydrogen bonding in water?

A

usually soluble (short chains at least)

69
Q

How soluble are compounds with hydrogen bonding in hexane?

A

usually insoluble

70
Q

How soluble are non-polar substances in water?

A

insoluble

71
Q

How soluble are non-polar substances in hexane?

A

soluble

72
Q

Explain why ionic compounds are mostly soluble in water

A

ion-dipole interactions are formed between oppositely charged ions in the ionic lattice and dipoles which releases energy

if enough energy is released then the lattice breaks down and ions become hydrated

73
Q

Why are compounds with hydrogen bonding usually soluble in water?

A

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

74
Q

Why are non-polar substances insoluble in water?

A

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

Physical Chemistry (4 decks)

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