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Y12 Organic Chemistry > Bonding > Flashcards

Bonding Flashcards

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1
Q

Ionic Bond

A

Electrostatic attraction between oppositely charged ions.

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2
Q

Melting point for ionic compounds

A

High - strong electrostatic attraction between oppositely charged ions

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3
Q

Ionic bonding is stronger when…

A

Ions are smaller and/or have higher charges. (A higher charge density)

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4
Q

Covalent Bond

A

Shared pair of electrons

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5
Q

Dative covalent bond

A

When the shared pair of electrons both come from one of the bonding atoms. Also called co-ordinate bonding.

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6
Q

Metallic bonding

A

electrostatic force of attraction between cations and a sea of delocalised electrons

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7
Q

Factors that affect strength of metallic bond

A

Number of protons (more = stronger), number of delocalised electrons (more = stronger), size of ion (smaller = stronger)

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8
Q

4 crystal structures are…

A

ionic, simple molecular, macromolecular and metallic

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9
Q

conductivity of ionic compounds

A

poor when solid - ions can’t move. Higher when molten or dissolved - ions free to move

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10
Q

boiling points of molecular compounds

A

low - weak intermolecular forces between molecules (Van der Waals)

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11
Q

melting point of macromolecular compounds

A

high - many strong covalent bonds to break

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12
Q

melting point of metallic substances

A

high - strong electrostatic forces of attraction between cations and delocalised electrons

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13
Q

Conductivity of molecular substances

A

Poor - no charged particles

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14
Q

Conductivity of macromolecular

A

Diamond and silica - poor. Not electrons or ions.

Graphite - good as delocalised electrons able to move between layers

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15
Q

Conductivity of metals

A

Good - delocalised electrons able to move

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16
Q

Shape of molecule with 2 bp, 0 lp

A

linear 180 angles

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17
Q

Shape of molecule with 3 bp, 0 lp

A

trigonal planar, 120 angles

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18
Q

Shape of molecule with 4 bp, 0 lp

A

tetrahedral, 109.5 angles

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19
Q

Shape of molecule with 3 bp, 1 lp

A

trigonal pyramidal, 107 angles

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20
Q

Shape of molecule with 2 bp, 2 lp

A

V shape or bent. 104.5 angles

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21
Q

Shape of molecule with 5 bp, 0 lp

A

trigonal bypyramidal, 120 and 90 angles

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22
Q

Shape of molecule with 5 bp, 1 lp

A

See-saw

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23
Q

Shape of molecule with 6 bp, 0 lp

A

Octahedral 90 angles

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24
Q

How to explain shape

A
  1. State number of bonding pairs and lone pairs.
  2. State that electron pairs repel to get as far apart as possible
    3a. If no lone pairs state that electron pairs repel equally
    3b. If there are lone pairs state that lone pairs repel more than bonding pairs
  3. State shape and bond angle
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25
Q

Shape of molecule with 4 bp, 2 lp

A

Square planar 90 angles

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26
Q

Shape of molecule with 5 bp, 1 lp

A

Square based pyramid

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27
Q

How to calculate number of electron pairs

A

Group of central atom + number of electrons from bonding atoms +/- electrons from charge = total number of electrons. Divide by 2 = number of pairs. Work out whether they are bonding or lone pairs

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28
Q

Shape CO2

A

linear 180 angles

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29
Q

Shape CS2

A

linear 180 angles

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30
Q

Shape HCN

A

linear 180 angles

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31
Q

Shape BeF2

A

linear 180 angles

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32
Q

Shape BF3

A

Trigonal planar, 120 angles

33
Q

AlCl3

A

Trigonal planar, 120 angles

34
Q

SO3

A

Trigonal planar, 120 angles

35
Q

NO3 -

A

Trigonal planar, 120 angles

36
Q

CO3 2-

A

Trigonal planar, 120 angles

37
Q

SiCl4

A

Tetrahedral, 109.5 angles

38
Q

CH4

A

Tetrahedral, 109.5 angles

39
Q

NH4+

A

Tetrahedral, 109.5 angles

40
Q

PF3

A

Trigonal planar, 107 angles

41
Q

NH3

A

Trigonal planar, 107 angles

42
Q

H30 +

A

Trigonal planar, 107 angles

43
Q

H2O

A

Bent/ V-shape 104.5

44
Q

OCl2

A

Bent/ V-shape 104.5

45
Q

H2S

A

Bent/ V-shape 104.5

46
Q

OF2

A

Bent/ V-shape 104.5

47
Q

SCl2

A

Bent/ V-shape 104.5

48
Q

PF5

A

Trigonal bypyramidal, 120 & 90

49
Q

PCl5

A

Trigonal bypyramidal, 120 & 90

50
Q

SF6

A

octahedral, 90

51
Q

XeF4

A

Square planar, 90

52
Q

BrF5

A

Square pyramid, 89

53
Q

I3 -

A

Linear, 180

54
Q

ClF3

A

T shape, 89

55
Q

SF4

A

see saw, 119 & 89

56
Q

IF4 +

A

see saw, 119 & 89

57
Q

Electronegativity

A

Ability of an atom to attract a pair of electrons in a covalent bond

58
Q

How electronegativity changes across a period

A

increases - more protons, same number of shells and shielding so easier to attract incoming pair of electrons

59
Q

How electronegativity changes down a group

A

decreases - more protons but greater distance to outer shell so less easy to attract an incoming pair of electrons

60
Q

Effect of large electronegativity difference between elements

A

perfect ionic compound

61
Q

Effect of small difference in electronegativity between elements

A

purely covalent compound

62
Q

Effect of some difference in electronegativity between elements

A

Polar bond or ionic bond with covalent character

63
Q

Polar bond

A

covalent bond with an unequal distribution of electrons in the bond. This causes a dipole with partially charged atoms

64
Q

Non-polar molecule

A

A symmetric molecule with all identical bonds and no lone pairs. Individual dipoles cancel out.

65
Q

Polar molecules

A

Asymmetric molecules with at least one polar bond

66
Q

3 types of intermolecular force in order of strength

A

weakest: van der Waals’ forces,
permanent dipole-dipole,
strongest: hydrogen bond

67
Q

How Van der Waals’ forces form

A

Electrons move constantly and randomly in any molecule. A temporary, instantaneous dipole can occur. This induces a dipole in a neighbouring molecule forming an attraction

68
Q

Main factors affecting the size of Van der Waals.

A

More electrons and larger surface area of a molecule increases the VdWs so boiling points are higher

69
Q

Trend in bp down group 7

A

Larger molecules with more electrons so greater temporary and induced dipoles therefore stronger attraction between molecules so higher bp

70
Q

Trend in bp in alkanes

A

Longer alkanes have more electrons so have greater VdWs between molecules. Less branched alkanes have a greater surface area of contact between molecules so also have stronger VdWs between molecules and so have higher bps

71
Q

Cause of permanent dipole-dipole forces

A

Occur between polar molecules with a dipole. Stronger than VdWs so compounds have a higher bp

72
Q

Explanation of hydrogen bonding

A

Occurs in compounds with a H atom attached to very electronegative F, O or N. The large difference of electronegativity creates a partial charge difference with d+ H and d- FON. An attraction is formed between the H and a lone pair on the FON

73
Q

3 marks for drawing a hydrogen bond

A

Partial charges shown,
All lone pairs shown
Attraction shown between lone pair on FON and H attached to a FON

74
Q

Requirements for hydrogen bonding

A

Hydrogen atom attached to a F,O or N. and a FON on another molecule

75
Q

Trend in bp of group 4 hydrides - CH4, SiH4, GeH4, SnH4

A

Increases down group as molecules become larger so have more electrons and stronger VdWs between molecules

76
Q

Trend in bp of group 5,6 and 7 hydrides

A

Generally increases down the group as molecules become larger and have stronger VdWs. HF, H2O and NH3 have unexpectedly high bp because they can form H bonds as well as VdWs

77
Q

What is unusual about ice?

A

It is less dense than liquid water. The H bonding between molecules holds them further apart from each other in a regular arrangement

78
Q

Diagram of iodine

A

herringbone structure of I2 molecules

Y12 Organic Chemistry (3 decks)

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