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
bond length
INTERNUCLEAR DISTANCE of 2 covalently bonded atoms
26
triple bonds bond length and strength
- large electron density between 2 atoms - increases attraction between nuclei and electrons between - so atoms are pulled closer - stronger covalent bond
27
how does VSEPR work (eg NH3)
- 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
28
electronegativity
ability of an atom to attract the bonding electrons in a covalent bond
29
most electronegative element is
Fluorine
30
nuclear charge effect on electronegativity
- more protons = increased attraction for outer shell e- - INCREASED electronegativity
31
atomic radius effect on electronegativity
- larger = electrons further, less strongly attracted to nucleus - decreased electronegativity
32
shielding effect on electronegativity
- more = increased inner shells = less attractive force from nucleus = decreased electronegativity
33
electronegativity DOWN a group
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
34
electronegativity across a period
INCREASES - nuclear charge increases - constant shielding as electrons added to same shells - attraction increases between nucleus and bonding e- - smaller atomic radii
35
non polar
both atoms have same electronegativity
36
electron density drawing map polar
delta + side is smaller, delta negative side is fatter
37
what is stronger between INTRAmolecular and INTERmolecular forces
INTRA - eg covalent > hydrogen
38
strength of imfs
hydrogen permanent dipole- permanent dipole London
39
how are London forces formed
- 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
what affects strength of London forces
- number of electrons - if its linear
41
permanent dipole-dipole forces cause
- between permanent dipoles on polar molecules - delta + and negative attract
42
hydrogen bonding
TYPE of permanent dipole-dipole bonding - needs O N F with H - h is positive, forms a bond with LONE PAIR on ONF
43
Water high mp/bp
- hydrogen bonding - strong imf - lots of energy required to separate molecules
44
density of ice
- ice LESS dense than water - water molecules in an open lattice - more space between molecules - HBs longer than covalent bonds
45
alkanes imfs
only london
46
boiling temp of alcohol vs alkane
- alkanes have only London - alcohol have hydrogen - more energy required to separate alcohol compared to alkane
47
hydrogen halides boiling point
- increases from HCl to HI - but HF high due to hydrogen bonds
48
general solvent rule
'like dissolves like' - non polar dissolves non polar - polar dissolves polar
49
water dissolving alcohols
GOOD SOLVENT - hydrogen bonds - larger the alcohol, the LESS soluble as the polar proportion is smaller - h bonds between alcohol and water similar in strength to those within each
50
metallic bonding
- metal atoms in tight lattice - delocalised electrons - strong efa between positive metal atoms and delocalised electrons
51
giant lattices are present in ...
ionic giant covalent (diamond, graphite silicon IV dioxide) giant metallic
52
giant covalent lattice mp
- very high - lots of many strong covalent bonds require lots of energy to break
53
graphite is soft?
- metal cations arranged in layers - weak imfs so can slide over each other
54
diamond and silicon IV dioxide hard
each carbon atom bonded to 4 other carbon atoms - many strong covalent bonds difficult to break
55
covalent solubility
INSOLOUBLE
56
graphite electricity conduction
- each carbon atoms bonded to 3 others - one free electron per carbon atom - electrons can move therefore conducts electricity
57
diamond imfs
NO IMFS
58
GRAPHENE
- single layer of graphite - 1 layer (atom) thick
59
metals solubility
NOT soluble in water
60
graphene ring shape
hexagonal
61
graphene structure
2 dimensional
62
dative bonding explanation (2)
- donation of lone pair from x - to y which is electron deficient
63
bond angles in graphene
120
64
explain why x is polar but y isn't
- x is linear so DIPOLES CANCEL - y is v shaped so dipoles dont cancel
65
bonding in caco3
ionic AND covalent
66
non polar covalent
0-0.4
67
polar covalent
0.41-1.69
68
ionic
1.7
69
WHY does h2o have more hydrogen bonds than HF
- 2x the amount - as 2 atoms of hydrogen for every molecule
70
why has a branched molecule isomer got a lower boiling point
- branching means fewer London forces - due to less surface area/ points of contact
71
angle of the hydrogen bond
180
72
why is x more dense?
- larger atoms - packed more closely together
73
why can x make that molecule but y cant | octet
- x can expand its octet - y can only accommodate 8 outer shell e-
74
who can expand the octet
if they have 3p subshell, cos they can move into 3d
75
why is the hydrogen bond 180
- 2 pairs of bonding e- - - lp - max separation
76
graph of ionic radius
linear decrease group 5,6,7,1,2,3
77
ionic radius down group
INCREASES - more shells of e- - more shielding, further distance etc
78
why are ions 5,6,7 > radius than their atoms
- same no. protons - more electrons - attraction per electron is less
79
ionic bonding electron density map
- discrete (e- density is 0 between ions) - negative ion is larger
80
CuCrO4 2- electrolysis
- Cu2+ ions (BLUE) migrate to cathode - CrO4 2- ions (YELLOW) migrate to anode - EVIDENCE FOR EXISTENCE OF IONS
81
COVALENT ELECTRON DENSITRY MAP
ellipse around both nuclei
82
pauling electronegativity range
0-4
83
how to know if a molecule is symmetrical
- identical bonds - 0 lps
84
London forces aren't present in
IONIC substances
85
how many HB's can water form per molecule?
2 - oxygen has 2 lone pairs of e-
86
ionic dissolving in water
- lattice bonds broken - negative ions attracted to delta +, positive attracted to delta - - the higher the charge the less soluble as a stronger attraction
87
factors affecting strength of metallic bonding
1. more protons, the stronger 2. more delocalised electrons, the stronger 3. smaller the ion the stronger
88
ionic compounds brittle (4)
- layers slide; REGULAR LATTICE DISRUPTED - positive ions aligned with positive - like charges repel - structure breaks apart
89
description of solubility
IMFs the same strength in both
90
ionic bonding solubility terminology
HYDRATION of na+ and cl-
91
halogenoalkanes in water
INSOLUBLE - hydrogen bonds between water stronger than permanent dipole-dipole
92
do giant covalent form liquid
SUBLIMES rather than melting
93
longer hydrocarbons alcohol solubility
LES SSOLUBLE
94
bond angle in graphite and graphene
120
95
in what shape are the carbon atoms arranged in graphite and graphene
hexagonal
96
importance of a molecule w polar bonds unsymmetrical
dipoles DONT CANCEL
97
MORE difference in eneg MEANS MORE OF WHICH CHARACTER?
ionic
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