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Alevel Chemistry OCR A > Module 3: The Periodic Table > Flashcards

Module 3: The Periodic Table Flashcards

1
Q

The periodic table
What were the only ways to categorise elements in the 1800s

A

By their physical and chemical properties or their relative atomic mass

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2
Q

The periodic table
In 1817 what did dobereiner attempt to do and what were they called

A

Group similar elements and the groups were called dobereiners triads

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3
Q

The periodic table
What groups did dobereiner create

A

Saw bromine chlorine and iodine had similar characteristic
Realises properties of bromine fell halfway between chlorine and iodine
Found groups of three eg lithium sodium and potassium

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4
Q

The periodic table
What did john newlands do

A

Tried to make a table of elements in 1863and noticed if her arranged elements in order of mass similar elements appeared regularly

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5
Q

The periodic table
What was the law of octaves

A

John newlands found every eighth elements was similar but this pattern broke down on third row

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6
Q

The periodic table
What did dmitri Mendeleev do

A

In 1869 produced a better table which wasn’t far off current table

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7
Q

The periodic table
How did Mendeleev arrange elements

A

By atomic mass and left gaps in the table where the next element didn’t seem to fit so could keep elements with similar chemical properties in same groups

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8
Q

The periodic table
What did Mendeleev predict

A

Properties of undiscovered elements

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9
Q

The periodic table
How is the modern periodic table arranged

A

By increasing proton number

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10
Q

The periodic table
How is the periodic table laid out

A

Into periods (rows) and groups (columns)

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11
Q

The periodic table
What do all elements in same period have and what does this mean

A

Same number of electron shells and means there are repeating trends in physical and chemical properties across periods

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12
Q

The periodic table
What are trends in properties called

A

Periodicity

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13
Q

The periodic table
What do all elements within one groups have and what does this mean

A

All have same number of electrons in outer shell and have similar chemical properties

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14
Q

Ionisation energies
What does it mean if and atom or molecule has been ionised

A

Electrons have been removed

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15
Q

Ionisation energies
Definition for the first ionisation energy

A

Energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms

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16
Q

Ionisation energies
What type of process is it to ionise an atom or molecule and why

A

Have to put energy in so and endothermic process

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17
Q

Ionisation energies
Equation for first ionisation energies

A

X (g) —> X+ (g) + e-

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18
Q

Ionisation energies
Important to rember when writing ionisation energy equations

A

Must use gas symbol (g) because ionisation energies are measured for gaseous atoms

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19
Q

Ionisation energies
What does the lower ionisation energy mean

A

Easier to form an ion

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20
Q

Ionisation energies
What factors affect ionisation energies

A

Nuclear charge
Atomic radius
Shielding

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21
Q

Ionisation energies
How does nuclear charge affect ionisation energy

A

More protons in the nucleus, more positively charged the nucleus is and the stronger attraction for the electrons

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22
Q

Ionisation energies
How does atomic radius affect ionisation energy

A

Attraction decreases rapidly with distance, an electron close to the nucleus will be much more strongly attracted than one further away therefore require more energy to remove

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23
Q

Ionisation energies
How does shielding affect ionisation energy

A

As number of electrons increases, there’s more shells therefore shielding from inner shells and less attraction between nucleus and outer electron

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24
Q

Ionisation energies
What does a high ionisation energy mean

A

Strong attraction between the electron and the nucleus so more energy is ended to overcome the attraction and removed the electron

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25
Q

Ionisation energies
What happens to ionisation energies as you go down a group

A

Decreases

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26
Q

Ionisation energies
What does ionisation energies decrease down a. Group

A

Elements have more electron shells compared to one above. Extra shells means that atomic radius is larger so outer electrons are further away from nucleus which means there a weaker attraction to the nucleus due to shielding from the inner shells

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27
Q

Ionisation energies
What does ionisation energies decreasing down the group give evidence for

A

Electron shells existing

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28
Q

Ionisation energies
What happens to the ionisation energy across a period

A

Increases

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29
Q

Ionisation energies
Why does ionisation energy increase actions a period

A

Number of protons is increasing so positive charge in nucleus increases, electrons are pulled closer to the nucleus making the atomic radius smaller so stronger stronger attraction and more energy required to overcome it

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30
Q

Ionisation energies
What are the two exceptions to ionisation energies increasing across a period

A

First ionisation energy decreases between groups 2 and 3 and between groups 5 and 6

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31
Q

Ionisation energies
Why is there a drop in ionisation energy between groups 2 and 3 as an expection to period trend

A

Due to subshell structure

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32
Q

Ionisation energies
Why is there a drop in ionisation energy between groups 2 and 3 due to subshell structure

A

Outer electron in group 3 is in a p orbital rather than and s orbital
A p orbital is slightly higher energy than an s orbital in same shell so electron on average is further from nucleus
P orbital also has additional shielding
These factors override effect of increasing nucleus charge

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33
Q

Ionisation energies
Why is there a drop in ionisation energy between groups 5 and 6 as an exception to periodicity

A

Due to p orbital repulsion

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34
Q

Ionisation energies
Why is there a drop in ionisation energies between group 5 and 6 due to p orbital repulsion

A

In group 5 elements the electron is being removed from a single occupied orbital
In group 6 elements electron is removed from an orbital with 2 electrons
The repulsion between the two electrons in an orbital means that they’re easier to remove

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35
Q

Ionisation energies
What are successive ionisation energies

A

Each time you remove and electron after the first time and can remove all electrons to leave only the nucleus

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36
Q

Ionisation energies
What is the equation for the second ionisation energy

A

X+ (g) —> X 2+ (g) + e-

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37
Q

Ionisation energies
What does a graph of successive ionisation energies provide evidence for

A

Shell structure of atoms

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38
Q

Ionisation energies
Why do successive ionisation graphs show evidence for shell structure

A

Within each sha;; successive ionisation energies increase because electrons are being removed from and increasingly positive ion and there’s also less repulsion amongst the examining electrons so electrons are held strongly by the nucleus
Big lumps in ionisation energies happens when a new shell is broken into

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39
Q

Ionisation energies
How can you tell which groups and Lenten is from from a successive ionisation energy graph

A

Couldn’t how many electrons are removed before first big group

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40
Q

Ionisation energies
How can successive ionisation energy graphs be used to predict the electronic structure of an elements

A

Working from right to left count how many pints there are before each big jump to find how many electrons are in each shell

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41
Q

Structure bonding and properties
Examples of giant covalent lattices

A

Diamond graphite and graphene and silicon

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42
Q

Structure bonding and properties
What are giant covalent lattices

A

Huge network of covalently bonded atoms

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43
Q

Structure bonding and properties
How many bonds can each carbon atom make

A

4

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44
Q

Structure bonding and properties
What are allotropes and example

A

Different form so the same element in the same state eg carbon has graphite carbon and graph even

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45
Q

Structure bonding and properties
What is the bonding and structure in diamonds

A

Each carbon is Covalently bonded to four other carbon atoms and they arrange themselves is a tetrahedral shape -crystal lattic structure

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46
Q

Structure bonding and properties
Properties of diamonds

A

Very high melting point (sublimes)
Extremely hard
Good thermal conductor because vibration Eason;ly travel through the stiff lattice
Can’t conduct electricity
Won’t dissolve

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47
Q

Structure bonding and properties
Reason for properties in diamonds

A

String covalent bonds

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48
Q

Structure bonding and properties
Structure and bonding in silicon

A

Forms a crystal lattice structure with similar properties to carbon each silicon forms four throng covalent bonds

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49
Q

Structure bonding and properties
Structure and bonding in graphite

A

Carbon atoms arranged in sheets of flats hexagons each carbons atom covalently bonded with three bonds each
Sheets bonded together with weak London forces
Fourth electron of each atom delocalised between the sheets

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50
Q

Structure bonding and properties
What do the weak forces between the layers mean for graphite

A

Layers easily broken so the beets can slide over each other
Graphite feels slippery and issued as a lubricant and in pencils

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51
Q

Structure bonding and properties
Why can graphite conduct electricity

A

Delocalised electrons aren’t attached to any particular carbon atom and are free to move

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52
Q

Structure bonding and properties
Why is graphiteused to Kate strong lightweight sports equipment

A

Layers are quite far apart compared to the length of covalent bond so graphite is less dense

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53
Q

Structure bonding and properties
Why is graphite have a high melting point

A

Strong covalent bonds in hexagon sheets

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54
Q

Structure bonding and properties
Why is graphite insoluble

A

Covalent bonds too strong to break

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55
Q

Structure bonding and properties
What is graphene

A

Sheet of carbon atoms joined together in hexagons
One atoms thick each carbon atoms makes three bonds

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56
Q

Structure bonding and properties
Why can graphene conduct electricity

A

Delocalised electronics are free to move along the street
Can move quickly above and below the sheet

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57
Q

Structure bonding and properties
Why is graphene extremely strong

A

Delocalised electrons strengthen the covalent bonds between the carbon atoms

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58
Q

Structure bonding and properties
Physical properties of graphene

A

Single layer makes it transparent and extremely light

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59
Q

Structure bonding and properties
Potential applications of graphene

A

High speed electronics and aircraft technology due to high strength low mass and good conductivity
Touchscreens due to flexibility and transparency

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60
Q

Structure bonding and properties
Structure and bonding in giant metallic lattice strictures

A

Delocalised electrons electrostatically attracted yo the metal positive ions

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61
Q

Structure bonding and properties
Melting point factors in giant metallic structures

A

Number of delocalised electrons, the more there are the stronger the bonding there will be and the higher the melting point eg MG2+ has 2 electrons per atom whereas NA+ only has 1 electron per atom to MG will have a higher melting point
Size of metal ions also has impact smaller atomic radius means delocalised electrons closer to nucleus

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62
Q

Structure bonding and properties
Why are giant metallic structure ductile and malleable

A

No bonds holding specific ions together so metal ions can slide past each other when structure is pulled

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63
Q

Structure bonding and properties
Why are giant metallic stricture good thermal conductors

A

Delocalised electrons can pass kinetic energy to each other

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64
Q

Structure bonding and properties
Why are metals good conductor

A

Delocalised electrons can move freely and carry charge

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65
Q

Structure bonding and properties
Why are most metals insoluble

A

Because of strength of metallic bonds

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66
Q

Structure bonding and properties
What affects melting and boiling points across a period

A

The type of bonding

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67
Q

Structure bonding and properties
How do you explain graph with melting or boiling point across a period

A

Describe type of bonding

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68
Q

Group 2- alkaline earth metals
What ion do group 2 elements form

A

2+

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69
Q

Group 2- alkaline earth metals
How many electrons are in the outer shell of grp 2 elements

A

2 subshell (s2)

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70
Q

Group 2- alkaline earth metals
What stricture do group 2 ions have

A

Noble gas

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71
Q

Group 2- alkaline earth metals
What does reactivity do as you go down the group

A

Increases

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72
Q

Group 2- alkaline earth metals
Why does reactivity increase as you go down croup 2

A

As you go down ionisation energies decrease due to increasing atomic radii and shielding effect from inner shells

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73
Q

Group 2- alkaline earth metals
What happens to group 2 element when they react

A

They’re oxidised
From oxidation’s state 0 to +2
Lose 2 electrons

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74
Q

Group 2- alkaline earth metals
What do groups 2 metals react with water to form

A

Metal hydroxide and hydrogen gas

75
Q

Group 2- alkaline earth metals
Equation for group 2 nd water

A

M (s). + 2 H2O (l) —> M(OH)2 (aq) H2 (g)

76
Q

Group 2- alkaline earth metals
Group2 metals how they react with water

A

I Be -doesn’t react
I Mg- VERY slowly
I Ca - steadily
I Sr- fairly quickly
V Ba-rapidly

77
Q

Group 2- alkaline earth metals
What do group 2 burn with oxygen to form

A

Solid white oxides

78
Q

Group 2- alkaline earth metals
Reaction with oxygen equation

A

2 M (s) + O2 (g) —> 2 MO (s)

79
Q

Group 2- alkaline earth metals
What do groups 2 react with dilute acid o form

A

Salt and hydrogen

80
Q

Group 2- alkaline earth metals
Equation for groups 2 with dilut HCl

A

M (s) + 2 HCL (aq) -> MCl2 (aq) + H2(g)

81
Q

Group 2- alkaline earth metals
Reactivity of group 2 with acids as you go down the group

A

Get more vigorous

82
Q

Group 2- alkaline earth metals
What are group 2 oxides and hydroxides

A

Bases and mostly soluble in water therefore alkalis

83
Q

Group 2- alkaline earth metals
Why are groups 2 hydroxides strongly alkaline

A

Hydroxide ions when in aqueous

84
Q

Group 2- alkaline earth metals
Exception for solubility and alkaline in group 2

A

Magnesium oxide

85
Q

Group 2- alkaline earth metals
Trend for group 2 oxides alkaline

A

Oxides form more strongly alkaline solutions as you go down the group because hydroxides get more soluble

86
Q

Group 2- alkaline earth metals
What are group 2 compounds often used for

A

Neutralising acids
Eg calcium hydroxide used in agriculture to neutralise acidic soil
Eg magnesium hydroxide and calcium carbonate used in indigestion tablets

87
Q

Group 7-the halogens
what are all group 7 elements

A

Diatomic molecules

88
Q

Group 7-the halogens
What colour and state is fluorine

A

Pale yellow gas

89
Q

Group 7-the halogens
What colour and state is chlorine

90
Q

Group 7-the halogens
Colour and state of bromine

A

Red brown liquid

91
Q

Group 7-the halogens
Colour and state of iodine

A

Grey solid

92
Q

Group 7-the halogens
What happens to the melting and boiling point as you go down the group

93
Q

Group 7-the halogens
Why do melting and boiling points increase as you go down

A

Due to increasing strength of London forces as the size and relative mass of atoms increases

94
Q

Group 7-the halogens
How do the physical properties show the trend of melting and boiling points down the group

A

The congaed of physical state from chlorine has to iodine solid- volatility decreases

95
Q

Group 7-the halogens
What does volatile mean

A

Substance that have low boiling point

96
Q

Group 7-the halogens
What happens to reactivity down the group

A

Get less reactive

97
Q

Group 7-the halogens
What happens to halogens when they react

A

React by gaining and electron to form a 1- ion which means they’re reduced

98
Q

Group 7-the halogens
What type of agents are they

A

Oxidising agents

99
Q

Group 7-the halogens
Why does it get less reactive down the group

A

The atomic radii increases so outer electrons are further away from the nucleus, outer electron also shielded more form th attraction of the positive nucleus because of inner shields
Harder for larger atoms to attract the electron needed to form an ion

100
Q

Group 7-the halogens
What happens to the relative oxidising strengths

A

Decrease down the group

101
Q

Group 7-the halogens
How can there relative oxidising strengths be seen

A

In their displacing reaction with halide ions

102
Q

Group 7-the halogens
Full equation for Bromine reaction with potassium iodide

A

Br2 (aq) + 2 KI (aq) —> 2 KBr (aq) + I2 (aq)

103
Q

Group 7-the halogens
Ionic equation for bromine after with potassium iodide

A

Br2 (aq) + 2 I- (aq) —> 2 Br- (aq) + I2 (aq)

104
Q

Group 7-the halogens
What happens in the displace ment reaction to show the oxidising strengths

A

Colour changes

105
Q

Group 7-the halogens
How can the colour changes be easily seen in displacement with halide ion

A

By shaking mixture with and organic solvent such as cyclohexane

106
Q

Group 7-the halogens
WHaat will happen to the halogen that’s present which organic solvent is added

A

Settles as a distance layer above the aqueous solutions

107
Q

Group 7-the halogens
What colour is KCl (aq)

A

Colourless

108
Q

Group 7-the halogens
What colour is KBr (aq)

A

Colourless

109
Q

Group 7-the halogens
What colour is KI (aq)

A

Colourless

110
Q

Group 7-the halogens
What colour is chlorine water / Cl2 (aq)

A

Pale Green

111
Q

Group 7-the halogens
What colour is bromine after / Br2 (aq)

112
Q

Group 7-the halogens
Wat colour is iodide solution/ I2 (aq)

113
Q

Group 7-the halogens
What is the colour of KCl (aq) and Cl2 (aq) in aq

A

No reaction

114
Q

Group 7-the halogens
What is the colour of KCl (aq) and Cl2 (aq) in organic solution

A

No reaction

115
Q

Group 7-the halogens
What is the colour of KBr (aq) and Cl2 (aq) in aq and what does it show us present

A

Yellow
Br2

116
Q

Group 7-the halogens
What is the colour of KBr (aq) and Cl2 (aq) in organic. And what does it show is present

A

Orange
Br2

117
Q

Group 7-the halogens
What is the colour of KI (aq) and Cl2 (aq) in aq and what does it show is present

A

Orange brown
I2

118
Q

Group 7-the halogens
What is the colour of KI (aq) and Cl2 (aq) in organic and what does it show is ppresent

119
Q

Group 7-the halogens
What is the colour of KCl (aq) and Br2 (aq) in aq

A

No reaction

120
Q

Group 7-the halogens
What is the colour of KCl (aq) and Br2 (aq) in ornagic

A

No reaction

121
Q

Group 7-the halogens
What is the colour of KBr (aq) and Br2 (aq) in aq

A

No reaction

122
Q

Group 7-the halogens
What is the colour of KBr (aq) and Br2 (aq) in
Organic

A

No reaction

123
Q

Group 7-the halogens
What is the colour of KI (aq) and Br2 (aq) in aq and what does it show is present

A

Orange brown
I2

124
Q

Group 7-the halogens
What is the colour of KI (aq) and Br2 ( aq) in organic and what doe it show is present

125
Q

Group 7-the halogens
What is the colour of KCl (aq) and I2 (aq) in aq

A

No reaction

126
Q

Group 7-the halogens
What is the colour of KCl (aq) and I2 (aq) in aq

A

No reaction

127
Q

Group 7-the halogens
What is the colour of KBr (aq) and I2 (aq) in aq

A

No reaction

128
Q

Group 7-the halogens
What is the colour of KBr (aq) and I2(aq) in organic ic

A

No reaction

129
Q

Group 7-the halogens
What is the colour of KI (aq) and I2 (aq) in aq

A

No reaction

130
Q

Group 7-the halogens
What is the colour of KI (aq) and I2 (aq) in organic

A

No reaction

131
Q

Group 7-the halogens
What else are the displacement reactions useful for

A

Identifying solutions

132
Q

Group 7-the halogens
How are displacement reaction useful for identifying solutions

A

Halogen will displace a halide from the solution if the halide is below it in the periodic table

133
Q

Group 7-the halogens
What will chlorine (Cl2) displace

A

Bromide (Br-) and iodide (I-)

134
Q

Group 7-the halogens
Ionic equation from chlorine displace bromide

A

Cl (aq) + 2 Br- (aq) —> 2 Cl- (aq) + Br2 (aq)

135
Q

Group 7-the halogens
Ionic equation for chlorine displacing iodide (I-)

A

Cl (aq) + 2 I- (aq) —> 2 Cl- (aq) + I2 (aq)

136
Q

Group 7-the halogens
What does bromine displace

A

Iodide (I-)

137
Q

Group 7-the halogens
Ionic equation for bromine displacing h iodide

A

Br2 (aq) + 2 I- (aq) -> 2 Br- (aq). +I2 (aq)

138
Q

Group 7-the halogens
What does iodine displace

A

No reaction with F-, Cl- or Br-

139
Q

Group 7-the halogens
What is used to test for halides

A

Silver nitrate solution

140
Q

Group 7-the halogens
How do you test for halides

A

First add dilute nitric acid to remove ions that might interfere with the yeast then add silver nitrate solution and a precipitate is formed of the silver halide

141
Q

Group 7-the halogens
Ionic equation of halide with silver nitrate solutions

A

Ag+ (aq) + X- (aq) —> AgX (s)
( X is Cl Br or I)

142
Q

Group 7-the halogens
What does the colour of the precipitate determine

A

Which halide it is

143
Q

Group 7-the halogens
What colour does chloride Cl- go with silver nitrate

A

White precipitate

144
Q

Group 7-the halogens
What colour does bromide Br- go with silver nitrate

A

Cream precipitate

145
Q

Group 7-the halogens
What colour does Iodide I- go with silver nitrate

A

Yellow precipitate

146
Q

Group 7-the halogens
How can you be extra sure after the silver nitrate test

A

Add ammonia solution as they have different solubility

147
Q

Group 7-the halogens
Ammonia solution added to chlorine after silver nitrate

A

Dissolved in dilute ammonia solution

148
Q

Group 7-the halogens
Ammonia solution added to bromide and solver nitrate

A

Dissolves in concentrated ammonia solution

149
Q

Group 7-the halogens
Ammonia solution added to iodide after silver nitrate

A

Insoluble in concentrated ammonia solutropn

150
Q

Group 7-the halogens
What do they undergo with alkalis

A

Disproportionation

151
Q

Group 7-the halogens
What conditions are needed fro them to disproportionate

A

Cold dilute alkali solution

152
Q

Group 7-the halogens
What can make bleach

A

Chlorine gas and cold dilute aqueous sodium hydroxide

153
Q

Group 7-the halogens
Equation for making bleach

A

2NaOH (aq) + Cl2 (g) —> NaClO (aq) + NaCl (aq) + H2O

154
Q

Group 7-the halogens
What are some uses of bleach ( sodium chlorate (l) )

A

Water treatment
Bleach paper and textiles
Cleaning

155
Q

Group 7-the halogens
What happens when you mix chlorine with water

A

Undergoes disproportionation

156
Q

Group 7-the halogens
Eqatrion from chlorine wit water

A

Cl2 (aq) + H2O (l) <==> HCl (aq) + HClO (aq)

157
Q

Group 7-the halogens
What is the mixture after chlorine and water

A

Hydrochloric acid and chloric (l) acid

158
Q

Group 7-the halogens
What does aqueous chloric (l) ionise too and how does it treat water

A

Ionises to chlorate(l) ions also called hypochlorite ions which kill bacteria making water safe to drink and swim in

159
Q

Group 7-the halogens
Strength of using chlorine to water

A

Kills bacteria

160
Q

Group 7-the halogens
Risks of chlorine to water

161
Q

Group 7-the halogens
Alternatives of chlorine to water

A

Ozone O3 (string oxidising agent kills microorganisms)
Ultra violent light- kill’s microorganisms damaging there DNA

162
Q

Testing for ions
How do you test for carbonate ions

A

Add dilute HCl or HNO3

163
Q

Testing for ions
What will the result be form the carbonate test of the carbonate ions are present

A

Carbon dioxide will be released and effervescence

164
Q

Testing for ions

Testing for ions
Ionic equation for carbonate test

A

CO3 2- (s) + 2H + (aq) —> CO2 (g) + H2O (l)

165
Q

Testing for ions
What is the test for sulfate ions

A

Add dilute HCl or HNO3 followed by barium chloride solution BaCl2

166
Q

Testing for ions
Why is barium chloride used to test for sulfates

A

Most sulfates are soluble in water but barium sulfates is insoluble

167
Q

Testing for ions
What will be the result to show the presence ofsulfate ions

A

White precipitate will be formed

168
Q

Testing for ions
Ionic equation for sulfates test

A

Ba 2+ (aq) + SO4 2- (aq) —> BaSO4 (s)

169
Q

Testing for ions
What is the test for halide ions

A

Add nitric acid followed by silver nitrate

170
Q

Testing for ions
Results if chloride ion present

A

White precipitate and soluble in ammonia

171
Q

Testing for ions
Result if bromide ion present

A

Cream precipitate and only soluble in concentrated ammonia

172
Q

Testing for ions
Results if iodide ions are present

A

Yellow precipitate and not soluble in ammonia

173
Q

Testing for ions
What is the test for ammonium ions

A

Add sodium hydroxide and gently warm

174
Q

Testing for ions
Result for if ammonium ions are present

A

Damp red litmus paper will turn blue

175
Q

Testing for ions
Ionic equation for ammonium test

A

NH4 + (aq) + OH - (aq) —> NH3 (g) + H2O(l)

176
Q

Testing for ions
What order should you plan tests

A

If present o ammonium, then carbonate , then sulfates, then halides

177
Q

Periodicity
What happens to atomic radius across periods

178
Q

Periodicity
Why does atomic radius decrease across groups

A

Outer electrons in same shell so same amount of shielding but more protons in nucleus
Stronger attraction between nucleus and outer shell electrons so pulled closer to nucleus

179
Q

Periodicity
What happens to electronegativity across periods

180
Q

Periodicity
Why does electronegativity increase across periods

A

More protons in nucleus
Smaller atomic radius
Stronger attraction between nucleus and pair of electrons in the covalent bond

181
Q

Periodicity
What happens to melting and boiling points across periods

182
Q

Periodicity
How is variance in melting and boiling points across periods explained

A

By type of bonding and structure

183
Q

Periodicity
Explain trend in melting and boiling point across period using period 3 (long)

A

{Na,Mg,Al} mp,bd increases because metallic bonds, strong attraction between metal ions and delocalised e-, Al>Mg>Na because higher charge ion,more delocalised e-, smaller ions
{Si} mp,bp highest on period as giant covalent structure, strong covalent bonds require lots of energy to break
{P4,S8,Cl2} mp,bp increase from P to S then decrease because simple covalent molecules with weak London forces between molecules, S8>P4>Cl2 because bigger molecules have bigger molecules bigger electrons and larger IMF
{Ar} lowest as very weak London forces because monatomic atoms

Alevel Chemistry OCR A (16 decks)

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