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HL Chem > 4 chemical bonding > Flashcards

4 chemical bonding Flashcards

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1
Q

chemical bond definition

A

an electrostatic force which holds two or more atoms or ions together

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2
Q

bond breaking and forming – exothermic or endothermic?

A

breaking bonds = absorbing energy = endothermic
formation = releases energy = exothermic

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3
Q

ionic/electrostatic bond definition

A

the electrostatic attraction experienced between the electric charges of a cation and an anion

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4
Q

dot and cross diagrams

A

diff symbols for the electrons of diff atoms just draw without the circle
- in order to find the charges, cross the numbers of each
eg MgCl2 – Mg2+ and Cl-

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5
Q

nature of ionic bonds 3

A
  • electrostatic
  • usually strong (ionic = high melting and boiling)
  • non-directional (equally strong in all directions)
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6
Q

what does lattice energy indicate

A

strength of electrostatic attraction in ionic bonding

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7
Q

lattice energy of ionic componud definition

A

the enthalpy change when one mole of a solid ionic compound is seperated into gaseous ions under standard conditions

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8
Q

lattice energy magnitude depends on

A
  • proportional to charges of ions
  • inversely proportional to the distance
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9
Q

exception to lattice energy

A

melting point of Al2O3 is lower than that of MgO

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10
Q

structure of ionic compounds

A

giant crystal lattice structure
- with strong ionic bonds

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11
Q

typical properties of ionic compounds 4

A
  • high melting and boilign points
  • hard but brittle
  • good conductors of electricity (molten and aqueous states)
  • soluble in polar solvents, insoluble in non-polar solvents
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12
Q

why do ionic compounds have high melting and boiling points

A

large amt of thermal energy is required to overcome the strong electrostatic attractive forces between oppositely-charged ions

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13
Q

why are ionic compounds hard and not easily scratched

A

the ions are bound strongly to the lattice and are not easily displaced

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14
Q

why are ionic compounds brittle (easy to shatter/fracture under stress)

A

distortion causes ions of like charges to come close together and repel sharply

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15
Q

why are ionic compoudns good/bad electrical conductors in respective states

A

solid:
strong electrostatic attractive forces prevent the movement of charged ions – ions unable to move, cannot carry electrical charge

molten/aq:
free mobile ions present to carry the electrical charges

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16
Q

why are ionic compounds soluble in water and polar solvents

A

formation of ion-dipole attraction = release of energy = detachment of ions from crystal lattice for solvation

water is polar molecule – attacks ionic lattice and pulls it apart. ion surrounded by water molc – strongly attracted bc of electrostatic ion-dipole attraction

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17
Q

why are ionic compounds not soluble in non-polar solvents

A

non-polar sol eg hexane and benzene
- form weak forces of attraction with ions = do not pull away from the lattice = does not dissolve

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18
Q

definition of covalent bonding

A

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

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19
Q

overlapping of atomic orbitals

A

in order to form covalent bond = atomic orbitals with unpaired electrons partially overlap

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20
Q

2 ways of electron sharing (covalent)

A
  1. normal covalent bonds: both atoms give 1
  2. coordinate (dative) covalent bonds: both electrons originiate from the same atom
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21
Q

what are bond pairs and lone pairs

A

bond pairs: shared pairs of electrons between 2 atoms

lone pairs: pairs of electrons not shared

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22
Q

how to draw lewis structure

A

LINES for each pair!!! visulaise it

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23
Q

octet rule definition

A

atoms acheive stable noble gas electronic configuration with 8 valence electrons – gain, loss, sharing of electrosn

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24
Q

exceptions to octet rule 2

A
  1. incomplete octet structure -
    - molecules w central atoms w less than 8 e after bonding
    - eg NO, NO2 (N = 5e)
  2. expanded octet structure
    - molecules w atoms w more than 8 e after bonding
    - eg SF6 (usually compounds of PERIOD 3 onwards)
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25
Q

definition of coordinate covalent bond

A

the shared pair of electrons is contributed by only one of the bonded atoms

(drawn as an arrow is lewis structure)

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26
Q

example of dative bond

A

NH4+

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27
Q

how are pure ionic compounds formed

A

through the complete transder of electrons from metallic atoms to non-metallic atoms – forms cations and anions

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28
Q

how are pure covalent compounds formed

A

by EQUAL sharing of electrons between covalently bonded atoms

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29
Q

what do the electron clouds of an ionic bond with covalent cahracter look like

A

large cloud oval towards small cloud
- partial transfer of electrons

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30
Q

what do the electron clouds of a covalent bond with ionic character look like

A

one big cloud bigger on the bigger side
- unequal sharing of electrons

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31
Q

how do intermediate types of bonds 2

A
  • polarisation of anion by cation (due to incomplete electron transfer)
  • partial charges on covalent bond (due to diff in electroneg)
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32
Q

oversimplification of ionic bonding – how? 3

A
  • electron charge assumed to be COMPLETELY transferred from one atom to another
  • CHARGE of ions is assumed to be EVENLY distributed in the ions
  • cations and anions are assumed to be point charges
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33
Q

polarity definition

A

the distribution of electric charge between 2 bonded atoms

34
Q

polarisation definition

A

distortion of the anions electron cloud by a cation

35
Q

polarising power definition

A

the ability of the cation to distort/polarise the anions electron cloud

36
Q

polarisability definition

A

the ease with which the anion’s electron cloud can be distorted (polarised)

37
Q

ionic bonds between NaCl and Al2O3

A

NaCl – ideal ionic bond
Al2O3 – cation is small, highly charged (high charge density) = strong attraction, distorts the cloud
- ionic bond now has covalent character bc of POLARISATION of the anion
- electron density moves to a region between the cation and anion

38
Q

how to determine whats more polar

A
  1. charge density of cation + size
  2. compare anion size – bigger = more polarisable
39
Q

how does charge and size of cation affect polarity

A

higher charge = higher polarising power
SMALLER size (ionic radius) = greater polarising power

40
Q

how does charge and size of anion affect polarity

A

higher charge = greater polarisability
larger ionic radius/size = greater polarisabiliy u

41
Q

ionic compound with high covalent character (very polar) will likely have (cations and anions)

A
  • large anion (high polarisability)
  • small cation (high polarising power)
    either one highly charged
42
Q

factors affecting bond strength

A
  • lattice energy
  • covalent character
43
Q

lattice energy definition

A

enthalpy change when one mole of a solid ionic compound is seperated into gaseous ions under stnadard conditions

44
Q

how to calculate lattice energy

A

charge x charge / dist + dist ???

45
Q

effect of covalent character on strength of ionic bond

A

incr strength
eg. AgCl experimental lattice e is higher than theoretical bc theoretical considers it as purely ionic BUT its actually ionic with covalent character

46
Q

properties of ionic compounds with covalent compounds 2

A
  • lower melting point
  • may be soluble in organic solvents
47
Q

process of stating reasons for determining which is more covalent

A

state which is more covalent and which is more ionic

  1. size of cation
  2. polarising power of cation
  3. extent of distortion of anion electron cloud
  4. extent of covalent character eg BeCl2 > MgCl2
48
Q

define electronegativity

A

the relative attraction an atoms has for the shared pair of electrons in a covalent bond

49
Q

across period and down group electronegativity

A

across - incr
down - decr

50
Q

where are the most electronegateive elements

A

top right hand corner

51
Q

why does electronegativity incr across the period

A

atoms smaller = decr dist between elec and nuclei

charge incr but shielding effect constant

electrostatic attraction incr

52
Q

why does electronegativity decr down a group

A

larger atoms = incr dist = decr electrostatic attraction

53
Q

assumtions in the model of coavlent bonding

A

assumes that the electron pair is shared equally between the 2 atoms

54
Q

polarity of identical atoms

A

NON POLAR – eg H2, CL2

55
Q

how to explain a polar covalent bond 3

A
  • electrons shared unequally – diff in electronegative values
  • valence electron density distorts towards higher electroneg = polarisation
  • polar bond – partial electrical charges due to uneven dist of e
56
Q

what is a dipole

A

2 equal and opposite charges seperated over a distance

57
Q

why is there dipole moment

A

the more electroneg atom with higher charge density – will have a partial negative charge

58
Q

2 things needed for a covalent molecule to be polar

A
  • at leas tone polar bond
  • dipole moment must not cancel out (needs net dipole moment) (the arow with the cross)
59
Q

factors affecting covalent bond strenght 3

A
  • bond length (incr length, decr strenth)
  • bond multiplicity (incr multiple bonds = incr strength)
  • bond polarity (incr polarity = incr ionic character = incr strentgh)
60
Q

how does incr bond length affect a covalent bond strength

A

longer length = nuclei further from shared pair of electrons
exception: F2 (weak bc realtively large electron + internuclear repulsions + small overlap of bondign orbitals)

61
Q

how are shapes of covalent molecules predicted

A

valence shell electron pair repulsion theory
VSEPR
- any atom bonded to more than one other atom = central atom
- each bond/lone pair = electron domain

62
Q

how to calculate number of lone pairs

A

(no. of e in valence shell - no. of e in bondign pairs) / 2

63
Q

how do electron pairs arrange themselves (in terms of repulsion)

A

to MINIMISE repulsion

64
Q

order of strength of repulsion

A

lone-lone > lone-bond > bond - bond

65
Q

explain streght of repulsion

A

bonded pair held more tightly by TWO nuclei, lone pair held only by ONE atom = electron cloud more diffuse

66
Q

steps to determine shapes of molecules and ions (VSEPR) 5

A
  1. draw dot and cross
  2. count no of electron pairs around the central atom (single/double/triple/dative all counted as 1 pair)
  3. determine basic shape (by total no of e pairs aroudn central)
  4. determine actual shape (exact no of lone and bond pairs)
  5. draw structure
67
Q

0 lone pairs, 2 electron domains

A

linear, 180

68
Q

0 lone pairs, 3 electron domains

A

trigonal planar, 120

69
Q

1 lone pair, 3 electron domains

A

v shaped / bent, <120 (trigonal planar)

70
Q

0 lone pair, 4 electron domains

A

tetrahedral, 109.5

71
Q

1 lone pair, 4 electron domains

A

trigonal pyramidal, <109.5 (tetrahedral)

72
Q

2 lone pairs, 4 electron domaisn

A

v shaped/ bent, «109.5 (tetrahedral)

73
Q

0 lone pairs, 5 electron domains

A

trigonal bipyramid, 90, 120

74
Q

1 lone pair, 5 electron domains

A

see-saw, trigonal bipyramid (lone pairs replace equatorial electron pairs)

75
Q

2 lone pairs, 5 electron domains

A

T-shaped, trigonal bipyramid (lone pairs replace equatorial electron pairs)

76
Q

what happends when there are lone pairs in a trigonal bipyramid structure

A

they will replace the equatorial electron pairs first to minimise repulsion

77
Q

0 lone pair, 6 electron domains

A

octahedral, 90

78
Q

1 lone pair, 6 electron domaisn

A

square pyramidal, octahedral

79
Q

2 lone pairs, 6 electron domains

A

square planar, octrahedral

80
Q

behaviour of lone pairs in an octahedral geomertry

A

1st lone pair: any position as they are all equivalent
2nd lone pair: position opposite

81
Q
A

HL Chem (28 decks)

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