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C2 > Bonding and structure > Flashcards

Bonding and structure Flashcards

1
Q

define ionic bonding

A

the strong electrostatic attraction between positively and negatively charged ions IN A LATTICE

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2
Q

define giant ionic lattice

A

oppositely charged ions strongly attracted in all directions in a repeating pattern

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3
Q

why are ionic compounds soluble in water

A
  • dissolves in polar solvents
  • cations attracted to O, anions attracted to H
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4
Q

why are ionic compounds electrical conductors when aq or molten but electrical insulators when solid

A
  • when aq or molten, ions are mobile and can carry a charge
  • when solid, the ions are fixed in place in a lattice structure
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5
Q

why are ionic compounds solid at rtp

A

strong attractions between ions, lots of energy needed to overcome attractions and break the bonds

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6
Q

define covalent bonding

A

strong electrostatic force of attraction between a shared pair of electrons and the nuclei of the bonded atoms

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7
Q

define dative covalent bonding

A

both electrons in a covalent bond are donated by one atom

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8
Q

examples of giant covalent lattices

A
  • graphite
  • diamond
  • silicon dioxide
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9
Q

examples of giant covalent polymers

A
  • polymers
  • fullerenes
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10
Q

examples of simple covalent molecules

A
  • methane
  • oxygen
  • hydrogen chloride
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11
Q

examples of exceptions to the octet

A
  • BF3 ( hypovalent - too little e- )
  • SF6 (hypervalent - too many e-)
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12
Q

properties of simple covalent molecules

A
  • low melting and boiling points
  • insoluble in water
  • don’t conduct electricity as they have no freely moving charged particles
  • may react with water
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13
Q

define average bond enthalpy

A
  • average amount of energy needed to break a specific type of bond homolytically in a gaseous molecule
  • measurement of covalent bond strength
  • always +ve as bond breaking is endothermic
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14
Q

define metallic bonding

A

strong electrostatic attraction between cations and delocalised e-

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15
Q

structure of metals

A

metal atoms tightly packed together in lattice structure

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16
Q

are metals soluble

A
  • no
  • more likely to lead to reactions
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17
Q

why do alloys have lower melting point

A
  • disrupted lattices
  • weaker attraction between cations and delocalised sea of e-
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18
Q

how does charge density of metal affect bond strength

A
  • small, highly charged cations have greater attraction to sea of delocalised electrons
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19
Q

why are electron pairs as far apart as possible in molecules

A

to minimise repulsion

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20
Q

properties of metals

A
  • malleable - layers can slide over each other without breaking the bond
  • ductile - layers can be pulled into a wire one atom thick
  • high melting + boiling point - lot of energy needed to overcome the electrostatic attraction between cations and sea of delocalised electrons
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21
Q

the shape of a compound or ion is determined by

A
  • no of e- pairs around central atom
  • the nature of these pairs: bonding or lone pairs
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22
Q

what are bonding pairs

A
  • pairs of electrons that are involved in bonding
  • repel each other equally
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23
Q

what are lone pairs

A
  • pairs of e- that aren’t involved in bonding
  • repel more than bonding pairs
  • more electron dense
24
Q

what does each lone pair reduce the bond angle by

A

2.5 degrees

25
Q

normal line represents

A

bond in plane of the paper

26
Q

dotted wedge represents

A

bond going into paper, away from you

27
Q

bold wedge represents

A

bond coming out of paper towards you

28
Q

linear molecule (2 pairs) :
1. bond pairs
2. lone pairs
3. angle
4. example

A

1) 2 bond pairs
2) 0 lone pairs
3) 180 degrees
4) CO2

29
Q

linear molecule (1 pair) :
1. bond pairs
2. lone pars
3. angle
4. example

A

1) 1 bond pair
2) 0 lone pairs
3) 180 degrees
4) H2

30
Q

trigonal planar molecule :
1. bond pairs
2. lone pars
3. angle
4. example

A

1) 3 bond pairs
2) 0 lone pairs
3) 120 degrees
4) BF3

31
Q

tetrahedral molecule :
1. bond pairs
2. lone pars
3. angle
4. example

A

1) 4 bond pairs
2) 0 lone pairs
3) 109.5 degrees
4) CH4

32
Q

pyramidal molecule :
1. bond pairs
2. lone pars
3. angle
4. example

A

1) 3 bond pairs
2) 1 lone pair
3) 107 degrees
4) NH3

33
Q

non- linear molecule :
1. bond pairs
2. lone pars
3. angle
4. example

A

1) 2 bond pairs
2) 2 lone pairs
3) 104.5 degrees
4) H2O

34
Q

octahedral molecule :
1. bond pairs
2. lone pars
3. angle
4. example

A

1) 6 bond pairs
2) 0 lone pairs
3) 90 degrees
4) SF6

35
Q

define electronegativity

A

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

36
Q

define polar bond

A

A permanent dipole within a molecule containing covalently bonded atoms with different electronegativities

37
Q

define polar molecule

A

A molecule containing polar bonds with dipoles that do not cancel due to their direction

38
Q

factors affecting electronegativity

A
  • atomic radius
  • number of shells/ shielding
  • nuclear charge
39
Q

why does electronegativity decrease down a group

A
  • more shells, so more shielding
  • atomic radius increases (outer shell further from nucleus)
  • less atraction between nucleus and outer e-
40
Q

elements in the same period have the same…

A

number of electron shells

41
Q

why does electronegativity increase across a period

A
  • nuclear charge increases
  • same shielding (no. of shells)
  • atomic radius decreases (outer shells closer to nucleus)
  • increased attraction between nuclei and outer e-
42
Q

what’s the most electronegative element

A

F

43
Q

why do symmetrical molecules like CCl4 not have a permanent dipole

A

the dipoles within the molecule cancel out

44
Q

why do unsymmetrical molecules like CH3Cl have permanent dipoles and are polar

A

the dipoles within the molecule don’t cancel out

45
Q

what are the 3 types of intermolecular forces

A
  • London forces (induced dipole-dipole)
  • permanent dipole-dipole interactions
  • Hydrogen bonding (strongest)
46
Q

what are london forces

A
  • weak interactions
  • exist between all molecules
  • when there’s an uneven distribution of e- causing one end of the molecule to become temporarily charged due chance and this induces a charge in a nearby molecule
47
Q

how does the size of an atom/molecule affect london forces

A
  • larger atoms/molecules have more e- so stronger london forces
  • stronger london forces = higher melting + boiling points
48
Q

what are permanent dipole-dipole interactions

A
  • stronger than london forces
  • exists when there’s a polar covalent bond
  • attraction between partially +ve and partially -ve ends of a molecule
  • can induce a non-polar molecule
49
Q

what’s hydrogen bonding

A
  • strongest type of intermolecular force
  • occurs when H is bonded to F,O,N
50
Q

what are the unusual properties of water

A
  • ice is less dense than water
  • water has a relatively high melting and boiling point
51
Q

why is ice less dense than water

A
  • due to H bonding, water molecules in ice are further apart than in water
52
Q

why does water have relatively high melting and boiling points

A
  • water has H bonding as well as london forces
  • each water molecule can form 4 H bonds so a large amount of energy is needed to break them
53
Q

what’s a simple molecular lattice

A

the solid structure of covalently bonded molecules attracted together by intermolecular forces

54
Q

example of simple molecular lattices

A
  • I2
  • ice
  • P
  • S
55
Q

why do simple molecular lattices have low melting + boiling points

A

weak London forces between molecules only require little energy to break

56
Q

why can’t simple molecular lattices conduct electricity

A

no charged particles to carry charge throughout the structure

57
Q

why are most simple covalent lattices not soluble

A

most are not polar enough to be soluble in water but anything with H bonding will dissolve

C2 (7 decks)

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