2.2 Flashcards
(112 cards)
1
Q
Most common ion of Zinc
A
Zn 2+
2
Q
Most common ion of silver
A
Ag +
3
Q
What does the suffix ‘ate’ tell us
A
The compound usually includes oxygen
4
Q
Nitrate
A
NO3 -
5
Q
Carbonate
A
CO3 2-
6
Q
Sulfate
A
SO4 2-
7
Q
Hydroxide
A
OH-
8
Q
Ammonium
A
NH4 +
9
Q
Why does oxygen tend to make 2 bonds
A
It’s in group 6
10
Q
What does the suffix ‘ium’ tell us
A
The compound is very likely to have a positive charge
11
Q
Bromide
A
Br -
12
Q
Chloride
A
Cl -
13
Q
Sulfide
A
S 2-
14
Q
Nitride
A
N 3-
15
Q
Phosphate
A
PO4 3-
16
Q
Oxide
A
O 2-
17
Q
Phosphide
A
P 3-
18
Q
What does the suffix ‘ide’ usually tell us
A
The ion has no other elements in it
19
Q
How many electrons does each shell hold
A
Up to 2n^2 where n = shell number or quantum number
20
Q
What are the 4 sub shells in electrons
A
s
p
d
f
21
Q
Which sub shell is in the first electron shell
A
s
22
Q
Which sub shells are in the second electron shell
A
s and p
23
Q
Which sub shells are in the third electron shell
A
s, p and d
24
Q
Which sub shells are in the fourth electron shell
A
s, p, d and f
25
How many electrons does the ‘s’ sub shell hold
2
26
How many electrons does the ‘p’ sub shell hold
6
27
How many electrons does the ‘d’ sub shell hold
10
28
How many electrons does the ‘f’ sub shell hold
14
29
How many orbitals are in the ‘s’ sub shell
1
30
How many orbitals are in the ‘p’ sub shell
3
31
How many orbitals does the ‘d’ sub shell have
5
32
How many orbitals does the ‘f’ sub shell have
7
33
Atomic orbitals
A region within an atom, that can hold up to 2 electrons, which have opposite spins
34
Order of electron arrangement
Smallest to the largest
Electrons -> orbital -> sub shell-> shell
35
Shape of ‘s’ orbital
Spherical
36
Hund’s rule
Electrons prefer to occupy orbitals on their own, and only pair up when no empty orbitals of the same energy are available
37
Why does 4s fill up before 3d
4s has a lower energy than 3d
38
How can we explain the electron configuration for chromium and copper
Sub-shells like to be full or half full. As 4s and 3d are so close in energy, electrons can move between them easily, and rearrange to make full and half full sub shells
39
How are elements classified by blocks
The block refers to the the highest type of sub shell occupied by electrons e.g. Groups 1 and 2 are in the ‘s’ block because all of those elements highest sub shell is s
40
Which electrons are lost when a cation is formed
The highest energy electrons but the Sc - Zn elements are an exception
41
Why are the Sc - Zn elements an exception to the ion rule
4s and 3d energy levels are very close but once 4s is full, it’s energy is above 3d so empties first as well
42
How to write short structure electron configuration
Put noble gas closest to it then continue writing structure
| E.g. manganese = [Ar] 4s^2 3d^5
43
Electronegativity
Ability to attract a bonding pair of electrons in a covalent bond
44
The greater the electronegativity of an atom ...
... the more it attracts electrons towards it
45
Factors affecting electronegativity
Atomic charge
Distance from the nucleus
Electron shielding
46
The greater the difference in electronegativity between the bonded atoms ...
... the greater the permanent dipole
47
Electronegativity across a period
Atomic radius decreases
Greater nuclear charge
Stronger attraction between bonding pair of electrons and nucleus
Increased electronegativity
48
Electronegativity down a group
Decreases due to increased shielding and distance from nucleus
Less attraction between nucleus and bonding pair of electrons
49
Electronegativity up a group
Increases
50
Basic rules of VSEPR theory
All single-bonded and lone pairs arrange themselves as far apart in space as possible
Lone pairs repel more strongly than bonding pairs/regions
51
Shape of molecule if there are 2 bonding pairs and 0 lone pairs
Linear
52
Shape of molecule if there are 2 bonding pairs and 1 lone pair
Non linear
53
Angle of linear shape
180
54
Angle of non linear shape
117.5
55
Shape of molecule with 3 bonding pairs and 0 lone pairs
Trigonal planar
56
Angle of trigonal planar shape
120
57
Ionic bonding
Electrostatic attraction between positive and negative ions
58
Covalent bonding
The strong electrostatic bond between a shared pair of electrons and the nuclei of the bonded atoms
59
Shape of 4 bonding pairs and 0 lone pairs
Tetrahedral
60
Angle of tetrahedral
109.5
61
Shape of 3 bonding pairs and 1 lone pair
Trigonal pyramid
62
Angle of trigonal pyramid
107
63
Shape of 2 bonding pairs and 2 lone pairs
Bent
64
Angle of bent shape (2 bonding, 2 lone)
104.5
65
Shape of 6 bonding pairs and 0 lone pairs
Octahedral
66
Angle of octahedral
90
67
Shape of 5 bonding pairs and 1 lone pairs
Distorted square pyramid
68
Angle of distorted square pyramid
89
69
Shape of 4 bonding pairs and 2 lone pairs
Square planar
70
Angle of square planar
90
71
When does covalent bonding occur
Between atoms of the same element (N2, O2)
Between atoms of diff. elements on the RHS of the table
When one of the elements is in the middle of the table
With head-of-the-group elements with high ionisation energies
72
Permanent dipole
A small charge difference across a bond due to a difference in the electronegativities of the bonded atoms
73
Polar covalent bond
A covalent bond with a permanent dipole
74
Polar molecule
A molecule with an overall dipole, when you take into account any dipoles across the bonds
Not symmetrical
75
Sizes of atoms increases
Down a group
| Decrease across period
76
How does nuclear charge affect electronegativity
More protons
| Stronger attraction between nucleus and bonding pairs of electrons
77
How does atomic radius affect electronegativity
The smaller the radius
Closer electrons are to nucleus
Stronger attraction between nucleus and bonding pair of electrons
78
How does shielding affect electronegativity
Less shells
Less shielding
Stronger the attraction between the nucleus and bonding pair of electrons
79
When does a molecule have London Forces
When it is not polar
80
When does a molecule have hydrogen bonds
When it is polar and has N-H/ O-H/ F-H bonds
81
When does a molecule have permanent dipole-dipole forces
When it is polar and doesn’t have any N-H/ F-H/ O-H bonds
82
Hydrogen bonds
Strong dipole-dipole attraction between electron-deficient H atoms and a lone pair of electrons in a highly electronegative atom on a different molecule
83
Relative strength of forces
Ionic/ covalent - 1000
Hydrogen bonds - 50
Dipole-dipole - 10
London forces - 1
84
How do london forces arise
Electrons in shells are continually moving and fluctuating
Creates uneven distribution of electrons
Sets up instantaneous dipole
Induces dipole in neighbouring atoms
Creates weak forces of attraction
85
Anomalous properties of water
Ice is less dense than water - molecules in ice held apart by hydrogen bonds
Ice has a relatively high mp- hydrogen bonds are stronger than other intermolecular forces
86
Bond angle in CH4
109.5
87
Bond angle in NH3
107
88
Bond angle in H2O
104.5
89
How does lone pairs affect the mp
The more the lone pairs, the higher the melting point
90
What colour is anhydrous copper sulfate
White
| Blue when hydrated
91
Dative covalent bond
Both electrons have been donated by one atom
92
Why does iodine have a higher bp than fluorine
More electrons
Greater LF’s
More energy needed to overcome
93
All lattices are ...
Giant
94
Graphite
Giant covalent lattice w/ layers
Has delocalised electrons
Soft - layers can slide over each other and has LF
95
Why does pure water not conduct
No free charge carriers
96
Bonding in NH4 +
4 bonding pairs (one dative)
| Tetrahedral
97
Why is the 1st IE lower for Al
e- released from 3p which has a higher energy than 3s
98
Why is the 1st IE lower for S
Pair of e- removed from 3p
| Repulsion between paired e-
99
Graphene
One layer of graphite
100
How are atoms vaporized before going into a mass spectrometer
Dissolved in a volatile solvent
101
How does a mass spec work
High voltage supply
Vaporized atoms accelerate through capillary tube
Removes electrons to form ions
The ones with a lowest m/z are identified firs
Bigger current = more abundance
102
Shape of 5 bonded pairs
Bipyramidal
103
Angle in trigonal bipyramidal
90 and 120
103
Drawing trig. bipyramidal
Three on the plane
One doing back into the plane
One coming out of the plane
103
Polar bonds
C-O
C-N
N-O
S-F
106
What happens to the bond angle when you add a lp
Decreases by 2.5
107
s block elements
Gp 1 and 2
108
p block elements
Gp 13 to 18
109
d block elements
Transition metals
110
f bock elements
Lanthanoids and actanoids
111
Why is there a drop in IE between Be and B
2p has a higher energy than 2s so outermost e- is easier to remove
112
Why is there a drop in IE between N and O
Both outermost e- in 2p
| In O paired e- in 2p repel each other so its easier to remove the outermost e-
