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chem > topic 2: structure and bonding > Flashcards

topic 2: structure and bonding Flashcards

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1
Q

ionic bonding

A

strong electrostatic attraction between oppositely charged ions

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2
Q

ions formation

A
  • metals lose electrons from their valence shell, forming positive cations
  • these electrons are transferred to non-metal atoms, which gain these to form negative anions
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3
Q

ions electron configuration

A

SAME AS NOBLE GAS

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4
Q

Ionic bonds high melting points

A

lots of energy is required to overcome the very strong electrostatic attraction between oppositely charged ions

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5
Q

ionic compound structure + charge

A
  • lattice
  • evenly distributed crystalline structure
  • ions are in a regularly repeating pattern so positive cancels out negatice
  • so the lattice is NEUTRAL
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6
Q

isoelectronic radius increases with…

A

INCREASING NEGATIVE CHARGE

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7
Q

ionic radius decreases with…

A

INCREASING POSITIVE CHARGE

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8
Q

Why do negative ions with more negative charge have a larger radius?

A
  • outermost e- further away from the positive nucleus
  • weak attraction
  • more shielding
  • larger radius
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9
Q

Why do positive ions with higher charge have a smaller radius?

A
  • same nuclear charge
  • less sheielding
  • fewer electrons
  • stronger electrostatic force of attraction to nucleus
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10
Q

isoelectronic ions

A

ions with the same electronic configuration

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11
Q

ionic radii of isoelectronic ions

A
  • higher number of protons means electrons are pulled in closer to nucleus
  • so largest negative charge is largest
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12
Q

why are ionic compounds solid at room temperature

A
  • high melting point
  • lots of energy required to over come strong electrostatic forces of attraction between oppositely charged ions in the lattice
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13
Q

charge of ions effect on melting point

A
  • higher the charge the higher the melting point
  • stronger efa between oppositely charged ions
  • more energy required to overcome
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14
Q

ionic compounds electrical conductivity

A
  • when solid, ions are in fixed positions in lattice and are unable to move
  • can only conduct when molten or in solution has ions can move
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15
Q

ionic compounds dissolve in …

A

POLAR solvents eg water

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16
Q

how do ionic compounds dissolve

A

ions are hydrated

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17
Q

general rule of solubility of ionic compounds

A

the higher the ionic charge the less soluble

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18
Q

electrolysis (evidence of ionic)

A
  • positive ions in solution attracted to negative electrode
  • negative ions in solution attracted to positive electrode
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19
Q

cu2+ ions are

A

blue

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20
Q

CrO4 2- ions are

A

yellow

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21
Q

covalent bond is

A

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

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22
Q

dative bonding

A
  • lone pair of electrons can be donated to form a bond with an ELECTRON DEFICIENT ATOM
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23
Q

bond energy

A

is the energy required to break one mole of a particular covalent bond in the gaseous states

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24
Q

bond energy relationship to covalent bond

A

STRONGER bond = LARGER bond energy

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25
Q

bond length

A

INTERNUCLEAR DISTANCE of 2 covalently bonded atoms

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26
Q

triple bonds bond length and strength

A
  • large electron density between 2 atoms
  • increases attraction between nuclei and electrons between
  • so atoms are pulled closer
  • stronger covalent bond
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27
Q

how does VSEPR work (eg NH3)

A
  • nitrogen (central atom) has 4 total e- pairs
  • 3 bp, 1 lp
  • lp-bp repulsion > bp-bp repulsion
  • bond angle reduced from 109.5 to 107
  • achieves max separation and minimum repulsion
  • trigonal pyramidal
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28
Q

electronegativity

A

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

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29
Q

most electronegative element is

A

Fluorine

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30
Q

nuclear charge effect on electronegativity

A
  • more protons = increased attraction for outer shell e-
  • INCREASED electronegativity
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31
Q

atomic radius effect on electronegativity

A
  • larger = electrons further, less strongly attracted to nucleus
  • decreased electronegativity
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32
Q

shielding effect on electronegativity

A
  • more = increased inner shells = less attractive force from nucleus = decreased electronegativity
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33
Q

electronegativity DOWN a group

A

DECREASES
- number of protons increase
- each element has an extra filled electron shell, so increased shielding
- increased atom radii
- decreased attraction between nucleus and bonding e-el

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34
Q

electronegativity across a period

A

INCREASES
- nuclear charge increases
- constant shielding as electrons added to same shells
- attraction increases between nucleus and bonding e-
- smaller atomic radii

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35
Q

non polar

A

both atoms have same electronegativity

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36
Q

electron density drawing map polar

A

delta + side is smaller, delta negative side is fatter

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37
Q

what is stronger between INTRAmolecular and INTERmolecular forces

A

INTRA - eg covalent > hydrogen

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38
Q

strength of imfs

A

hydrogen
permanent dipole- permanent dipole
London

39
Q

how are London forces formed

A
  • electron charge cloud constantly randomly moving
  • INSTANTANEOUS, temporarily more on one side, causing a temporary INSTANTANEOUS DIPOLE
  • induces a dipole on neighbouring molecules
  • delta positive + delta negative attract
40
Q

what affects strength of London forces

A
  • number of electrons
  • if its linear
41
Q

permanent dipole-dipole forces cause

A
  • between permanent dipoles on polar molecules
  • delta + and negative attract
42
Q

hydrogen bonding

A

TYPE of permanent dipole-dipole bonding
- needs O N F with H
- h is positive, forms a bond with LONE PAIR on ONF

43
Q

Water high mp/bp

A
  • hydrogen bonding
  • strong imf
  • lots of energy required to separate molecules
44
Q

density of ice

A
  • ice LESS dense than water
  • water molecules in an open lattice
  • more space between molecules
  • HBs longer than covalent bonds
45
Q

alkanes imfs

A

only london

46
Q

boiling temp of alcohol vs alkane

A
  • alkanes have only London
  • alcohol have hydrogen
  • more energy required to separate alcohol compared to alkane
47
Q

hydrogen halides boiling point

A
  • increases from HCl to HI
  • but HF high due to hydrogen bonds
48
Q

general solvent rule

A

‘like dissolves like’
- non polar dissolves non polar
- polar dissolves polar

49
Q

water dissolving alcohols

A

GOOD SOLVENT
- hydrogen bonds
- larger the alcohol, the LESS soluble as the polar part is smaller

50
Q

metallic bonding

A
  • metal atoms in tight lattice
  • delocalised electrons
  • strong efa between positive metal atoms and delocalised electrons
51
Q

giant lattices are present in …

A

ionic
giant covalent (diamond, graphite silicon IV dioxide)
giant metallic

52
Q

giant covalent lattice mp

A
  • very high
  • lots of many strong covalent bonds require lots of energy to break
53
Q

graphite is soft?

A
  • metal cations arranged in layers
  • weak imfs so can slide over each other
54
Q

diamond and silicon IV dioxide hard

A

each carbon atom bonded to 4 other carbon atoms
- many strong covalent bonds difficult to break

55
Q

covalent solubility

A

INSOLOUBLE

56
Q

graphite electricity conduction

A
  • each carbon atoms boned to 3 others
  • one free electron per carbon atom
  • electrons can move therefore conducts electricity
57
Q

diamond imfs

A

NO IMFS

58
Q

GRAPHENE

A
  • single layer of graphite
  • 1 layer (atom) thick
59
Q

metals solubility

A

NOT soluble in water

60
Q

graphene ring shape

A

hexagonal

61
Q

graphene structure

A

2 dimensional

62
Q

dative bonding explanation (2)

A
  • donation of lone pair from x
  • to y which is electron deficient
63
Q

bond angles in graphene

A

120

64
Q

explain why x is polar but y isn’t

A
  • x is linear so DIPOLES CANCEL
  • y is v shaped so dipoles dont cancel
65
Q

bonding in caco3

A

ionic AND covalent

66
Q

non polar covalent

A

0-0.4

67
Q

polar covalent

A

0.41-1.69

68
Q

ionic

A

1.7

69
Q

WHY does h2o have more hydrogen bonds than HF

A
  • 2x the amount
  • as 2 atoms of hydrogen for every molecule
70
Q

why has a branched molecule isomer got a lower boiling point

A
  • branching means fewer London forces
  • due to less surface area/ points of contact
71
Q

angle of the hydrogen bond

A

180

72
Q

why is x more dense?

A
  • larger atoms
  • packed more closely together
73
Q

why can x make that molecule but y cant

octet

A
  • x can expand its octet
  • y can only accommodate 8 outer shell e-
74
Q

who can expand the octet

A

if they have 3p subshell, cos they can move into 3d

75
Q

why is the hydrogen bond 180

A
  • 2 pairs of bonding e-
    • lp
  • max separation
76
Q

graph of ionic radius

A

linear decrease
group 5,6,7,1,2,3

77
Q

ionic radius down group

A

INCREASES
- more shells of e-
- more shielding, further distance etc

78
Q

why are ions 5,6,7 > radius than their atoms

A
  • same no. protons
  • more electrons
  • attraction per electron is less
79
Q

ionic bonding electron density map

A
  • discrete (e- density is 0 between ions)
  • negative ion is larger
80
Q

CuCrO4 2- electrolysis

A
  • Cu2+ ions (BLUE) migrate to cathode
  • CrO4 2- ions (YELLOW) migrate to anode
  • EVIDENCE FOR EXISTENCE OF IONS
81
Q

COVALENT ELECTRON DENSITRY MAP

A

ellipse around both nuclei

82
Q

pauling electronegativity range

A

0-4

83
Q

how to know if a molecule is symmetrical

A
  • identical bonds
  • 0 lps
84
Q

London forces aren’t present in

A

IONIC substances

85
Q

how many HB’s can water form per molecule?

A

2
- oxygen has 2 lone pairs of e-

86
Q

ionic dissolving in water

A
  • lattice bonds broken
  • negative ions attracted to delta +, positive attracted to delta -
  • the higher the charge the less soluble as a stronger attraction
87
Q

factors affecting strength of metallic bonding

A
  1. more protons, the stronger
  2. more delocalised electrons, the stronger
  3. smaller the ion the stronger
88
Q

ionic compounds brittle

A
  • layers slide; REGULAR LATTICE DISRUPTED
  • positive ions aligned with positive
  • like charges repel
  • structure breaks apart
89
Q

description of solubility

A

IMFs the same strength in both

90
Q

ionic bonding solubility terminology

A

HYDRATION of na+ and cl-

91
Q

halogenoalkanes in water

A

INSOLUBLE
- hydrogen bonds between water stronger than permanent dipole-dipole

92
Q

do giant covalent form liquid

A

SUBLIMES rather than melting

93
Q

longer hydrocarbons alcohol solubility

A

LES SSOLUBLE

chem (16 decks)

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