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OCR Chemistry A COPY > 2.2 > Flashcards

2.2 Flashcards

1
Q

Most common ion of Zinc

A

Zn 2+

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2
Q

Most common ion of silver

A

Ag +

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3
Q

What does the suffix ‘ate’ tell us

A

The compound usually includes oxygen

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4
Q

Nitrate

A

NO3 -

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5
Q

Carbonate

A

CO3 2-

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6
Q

Sulfate

A

SO4 2-

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7
Q

Hydroxide

A

OH-

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8
Q

Ammonium

A

NH4 +

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9
Q

Why does oxygen tend to make 2 bonds

A

It’s in group 6

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10
Q

What does the suffix ‘ium’ tell us

A

The compound is very likely to have a positive charge

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11
Q

Bromide

A

Br -

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12
Q

Chloride

A

Cl -

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13
Q

Sulfide

A

S 2-

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14
Q

Nitride

A

N 3-

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15
Q

Phosphate

A

PO4 3-

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16
Q

Oxide

A

O 2-

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17
Q

Phosphide

A

P 3-

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18
Q

What does the suffix ‘ide’ usually tell us

A

The ion has no other elements in it

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19
Q

How many electrons does each shell hold

A

Up to 2n^2 where n = shell number or quantum number

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20
Q

What are the 4 sub shells in electrons

A

s
p
d
f

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21
Q

Which sub shell is in the first electron shell

A

s

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22
Q

Which sub shells are in the second electron shell

A

s and p

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23
Q

Which sub shells are in the third electron shell

A

s, p and d

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24
Q

Which sub shells are in the fourth electron shell

A

s, p, d and f

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25
Q

How many electrons does the ‘s’ sub shell hold

A

2

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26
Q

How many electrons does the ‘p’ sub shell hold

A

6

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27
Q

How many electrons does the ‘d’ sub shell hold

A

10

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28
Q

How many electrons does the ‘f’ sub shell hold

A

14

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29
Q

How many orbitals are in the ‘s’ sub shell

A

1

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30
Q

How many orbitals are in the ‘p’ sub shell

A

3

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31
Q

How many orbitals does the ‘d’ sub shell have

A

5

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32
Q

How many orbitals does the ‘f’ sub shell have

A

7

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33
Q

Atomic orbitals

A

A region within an atom, that can hold up to 2 electrons, which have opposite spins

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34
Q

Order of electron arrangement

A

Smallest to the largest

Electrons -> orbital -> sub shell-> shell

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35
Q

Shape of ‘s’ orbital

A

Spherical

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36
Q

Hund’s rule

A

Electrons prefer to occupy orbitals on their own, and only pair up when no empty orbitals of the same energy are available

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37
Q

Why does 4s fill up before 3d

A

4s has a lower energy than 3d

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38
Q

How can we explain the electron configuration for chromium and copper

A

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

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39
Q

How are elements classified by blocks

A

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

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40
Q

Which electrons are lost when a cation is formed

A

The highest energy electrons but the Sc - Zn elements are an exception

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41
Q

Why are the Sc - Zn elements an exception to the ion rule

A

4s and 3d energy levels are very close but once 4s is full, it’s energy is above 3d so empties first as well

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42
Q

How to write short structure electron configuration

A

Put noble gas closest to it then continue writing structure

E.g. manganese = [Ar] 4s^2 3d^5

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43
Q

Electronegativity

A

Ability to attract a bonding pair of electrons in a covalent bond

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44
Q

The greater the electronegativity of an atom …

A

… the more it attracts electrons towards it

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45
Q

Factors affecting electronegativity

A

Atomic charge
Distance from the nucleus
Electron shielding

46
Q

The greater the difference in electronegativity between the bonded atoms …

A

… the greater the permanent dipole

47
Q

Electronegativity across a period

A

Atomic radius decreases
Greater nuclear charge
Stronger attraction between bonding pair of electrons and nucleus
Increased electronegativity

48
Q

Electronegativity down a group

A

Decreases due to increased shielding and distance from nucleus
Less attraction between nucleus and bonding pair of electrons

49
Q

Electronegativity up a group

A

Increases

50
Q

Basic rules of VSEPR theory

A

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
Q

Shape of molecule if there are 2 bonding pairs and 0 lone pairs

A

Linear

52
Q

Shape of molecule if there are 2 bonding pairs and 1 lone pair

A

Non linear

53
Q

Angle of linear shape

A

180

54
Q

Angle of non linear shape

A

117.5

55
Q

Shape of molecule with 3 bonding pairs and 0 lone pairs

A

Trigonal planar

56
Q

Angle of trigonal planar shape

A

120

57
Q

Ionic bonding

A

Electrostatic attraction between positive and negative ions

58
Q

Covalent bonding

A

The strong electrostatic bond between a shared pair of electrons and the nuclei of the bonded atoms

59
Q

Shape of 4 bonding pairs and 0 lone pairs

A

Tetrahedral

60
Q

Angle of tetrahedral

A

109.5

61
Q

Shape of 3 bonding pairs and 1 lone pair

A

Trigonal pyramid

62
Q

Angle of trigonal pyramid

A

107

63
Q

Shape of 2 bonding pairs and 2 lone pairs

A

Bent

64
Q

Angle of bent shape (2 bonding, 2 lone)

A

104.5

65
Q

Shape of 6 bonding pairs and 0 lone pairs

A

Octahedral

66
Q

Angle of octahedral

A

90

67
Q

Shape of 5 bonding pairs and 1 lone pairs

A

Distorted square pyramid

68
Q

Angle of distorted square pyramid

A

89

69
Q

Shape of 4 bonding pairs and 2 lone pairs

A

Square planar

70
Q

Angle of square planar

A

90

71
Q

When does covalent bonding occur

A

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
Q

Permanent dipole

A

A small charge difference across a bond due to a difference in the electronegativities of the bonded atoms

73
Q

Polar covalent bond

A

A covalent bond with a permanent dipole

74
Q

Polar molecule

A

A molecule with an overall dipole, when you take into account any dipoles across the bonds
Not symmetrical

75
Q

Sizes of atoms increases

A

Down a group

Decrease across period

76
Q

How does nuclear charge affect electronegativity

A

More protons

Stronger attraction between nucleus and bonding pairs of electrons

77
Q

How does atomic radius affect electronegativity

A

The smaller the radius
Closer electrons are to nucleus
Stronger attraction between nucleus and bonding pair of electrons

78
Q

How does shielding affect electronegativity

A

Less shells
Less shielding
Stronger the attraction between the nucleus and bonding pair of electrons

79
Q

When does a molecule have London Forces

A

When it is not polar

80
Q

When does a molecule have hydrogen bonds

A

When it is polar and has N-H/ O-H/ F-H bonds

81
Q

When does a molecule have permanent dipole-dipole forces

A

When it is polar and doesn’t have any N-H/ F-H/ O-H bonds

82
Q

Hydrogen bonds

A

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
Q

Relative strength of forces

A

Ionic/ covalent - 1000
Hydrogen bonds - 50
Dipole-dipole - 10
London forces - 1

84
Q

How do london forces arise

A

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
Q

Anomalous properties of water

A

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
Q

Bond angle in CH4

A

109.5

87
Q

Bond angle in NH3

A

107

88
Q

Bond angle in H2O

A

104.5

89
Q

How does lone pairs affect the mp

A

The more the lone pairs, the higher the melting point

90
Q

What colour is anhydrous copper sulfate

A

White

Blue when hydrated

91
Q

Dative covalent bond

A

Both electrons have been donated by one atom

92
Q

Why does iodine have a higher bp than fluorine

A

More electrons
Greater LF’s
More energy needed to overcome

93
Q

All lattices are …

A

Giant

94
Q

Graphite

A

Giant covalent lattice w/ layers
Has delocalised electrons
Soft - layers can slide over each other and has LF

95
Q

Why does pure water not conduct

A

No free charge carriers

96
Q

Bonding in NH4 +

A

4 bonding pairs (one dative)

Tetrahedral

97
Q

Why is the 1st IE lower for Al

A

e- released from 3p which has a higher energy than 3s

98
Q

Why is the 1st IE lower for S

A

Pair of e- removed from 3p

Repulsion between paired e-

99
Q

Graphene

A

One layer of graphite

100
Q

How are atoms vaporized before going into a mass spectrometer

A

Dissolved in a volatile solvent

101
Q

How does a mass spec work

A

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
Q

Shape of 5 bonded pairs

A

Bipyramidal

103
Q

Polar bonds

A

C-O
C-N
N-O
S-F

104
Q

Drawing trig. bipyramidal

A

Three on the plane
One doing back into the plane
One coming out of the plane

105
Q

Angle in trigonal bipyramidal

A

90 and 120

106
Q

What happens to the bond angle when you add a lp

A

Decreases by 2.5

107
Q

s block elements

A

Gp 1 and 2

108
Q

p block elements

A

Gp 13 to 18

109
Q

d block elements

A

Transition metals

110
Q

f bock elements

A

Lanthanoids and actanoids

111
Q

Why is there a drop in IE between Be and B

A

2p has a higher energy than 2s so outermost e- is easier to remove

112
Q

Why is there a drop in IE between N and O

A

Both outermost e- in 2p

In O paired e- in 2p repel each other so its easier to remove the outermost e-

OCR Chemistry A COPY (16 decks)

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