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Chemistry > Unit 3.3 - Chemistry of the p-block > Flashcards

Unit 3.3 - Chemistry of the p-block Flashcards

1
Q

Periodicity

A

How properties change across the periodic table

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2
Q

Name 2 compound with more than 8 electrons in their outer shell

A

Phosphorus (V) chloride —> PCL5
Phosphorus (V) fluoride —> PF5

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3
Q

What can we notice about the phosphorus compounds PCL5 and PF5?

A

There are 10 electrons around the central phosphorus atom

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4
Q

What’s the name for having more than 8 electrons in an outer shell?

A

Expanding the octet

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5
Q

Which elements are the only ones which can expand their octets?

A

Fluorine
Oxygen
Chlorine

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6
Q

What can fluorine, oxygen and chlorine expand their octets with?

A

Highly electronegative elements

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7
Q

Why are fluorine, oxygen and chlorine able to expand their octets?

A

They have large electronegativity values and have a very strong tendency to attract electrons to themselves

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8
Q

What does Florine, oxygen and chlorine expanding their octets occur with?

A

Elements in the 3rd row of the periodic table (Na —> Ar)

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9
Q

What does fluorine, oxygen and chlorine expanding their octets NOT occur with?

A

The elements in the 2nd row of the periodic table (Li—>Ne)

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10
Q

Why are fluorine, oxygen and chlorine able to expand their octets with the elements in the 3rd row of the periodic table?

A

As the elements in the 3rd row can expand their octet of electron as vacant 3d orbitals of suitable energy are available for bonding, allowing them to hybridise

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11
Q

How come Florine, oxygen and chlorine are able to expand their octet with elements from the 3rd row of the periodic table but not the second?

A

The energy difference between the 3s, 3p and 3d orbitals are small enough to overcome when bonding to strongly electronegative elements

With the 2nd row elements, the energy difference is too large to be overcome and no orbitals are available for expanding the octet

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12
Q

What happens to the gap between subshells as you get further from the nucleus?

A

Decreases

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13
Q

What does a smaller gap between subshells lead to?

A

Smaller energy differences

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14
Q

What can happen to small energy differences?

A

Can be overcome when bonding to strongly electronegative elements

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15
Q

Valency

A

Amount of bonds they can form

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16
Q

Maximum valency of second row elements

A

4 pairs

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17
Q

Maximum valency of 3rd row elements

A

Up to 7 pairs

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18
Q

Octet expansion

A

The ability of some atoms to use d-orbitals to have more than 8 electrons in their valence shell

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19
Q

Which orbitals are used to expand octets?

A

d-orbitals

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20
Q

What is the maximum covalent bonds that can be formed by boron?

A

3 (this is electron deficient)

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21
Q

What is the maximum covalent bonds that can be formed by carbon?

A

4

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22
Q

What is the maximum covalent bonds that can be formed by nitrogen

A

3 (this is also has 1 lone pairs)

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23
Q

What is the maximum covalent bonds that can be formed by oxygen?

A

2 (this is also has 2 lone pairs)

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24
Q

What type of orbitals do elements in period 3 and below have and why is this important?

A

d orbitals
More bonds can be formed

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25
Q

How is PF5 formed?

A

An s-electron is promoted to a d-orbital
These then hybridise forming 5 sp^3d orbitals (each having the same energy)

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26
Q

Trends in bonding across the p-block

A

Ionic to covalent

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27
Q

Trend in bonding down the p-block

A

Bonding becomes more ionic (metallic)

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28
Q

Metallic character

A

The tendency of an element to lose electrons and form positive ions or cations

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29
Q

Trend in metallic character across the p-block

A

Metallic character decreases

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30
Q

Trend in metallic character down the p-block

A

Metallic character increases

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31
Q

What’s related to metallic character and how?

A

Electronegativity
As electronegativity increases, the metallic character decreases and vice versa

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32
Q

Where does amphoteric character occur mainly?

A

In the middle of the period table, here the elements Be, Zn, Al, Ga, In, Sn and Pb are found

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33
Q

Elements in group 2 which show amphoteric behaviour

A

Beryllium

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34
Q

Elements in group 3 which show amphoteric behaviour

A

Aluminium, gallium, indium

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35
Q

Elements in group 4 which show amphoteric behaviour

A

Tin, lead

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36
Q

Transition elements which show amphoteric behaviour

A

Zinc

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37
Q

Amphoteric behaviour

A

The oxide and hydroxide will react with acids and bases to form salts

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38
Q

What are oxides and hydroxides of group 1 and 2?

A

Basic

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39
Q

What does the fact that the oxides and hydroxides of group 1 and 2 are basic mean?

A

They only react as acids

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40
Q

How does an element behave if it’s demonstrating amphoteric behaviour?

A

As either an acid or a base

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41
Q

If something reacts with an acid, how does it behave?

A

As a base

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42
Q

If something reacts with a base, how does it behave?

A

As an acid

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43
Q

Describe the electronegativity values of elements that show amphoteric character

A

Intermediate

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44
Q

Describe the bonding of elements that show amphoteric character

A

It’s changing in character from ionic to covalent

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45
Q

Describe whether the elements that show amphoteric character are metals or non-metals

A

On the region of the periodic table where the elements are changing from being metals to non-metals

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46
Q

What are all of the elements which show amphoteric characters?

A

Metals with relatively high electronegativity

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47
Q

What type of bonding do elements that demonstrate amphoteric character have?

A

Ionic or covalent bonding

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48
Q

How does a hydroxide form in the amphoteric character equations?

A

By the addition of sodium hydroxide solution to a salt solution of the metal

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49
Q

Aluminium hydroxide showing basic behaviour equation

A

Al(OH)3 + 3H+ —> Al3+ + 3H2O

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50
Q

Aluminium hydroxide showing acidic behaviour equation

A

Al(OH)3 + OH- —> [Al(OH)4]-

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51
Q

Tetrahydroxe aluminate (III)

A

[Al(OH)4]-

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52
Q

Lead hydroxide showing basic behaviour equation

A

Pb(OH)2 + 2H+ —> Pb2+ + 2H2O

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53
Q

Lead hydroxide showing acidic behaviour equation

A

Pb(OH)2 + 2OH- +H2O —> [Pb(OH)4]2-

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54
Q

Tetrahydroxo plumbate (II)

A

[Pb(OH)4]2-

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55
Q

What happens if lead (II) hydroxide is reacted with nitric acid?

A

An aqueous solution of lead (II) nitrate is formed

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56
Q

What happens when lead (II) hydroxide is reacted with sodium hydroxide?

A

An aqueous solution of sodium plumbate is formed

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57
Q

What do we need in an question if we have a metal oxide and a base?

A

Water to balance it out

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58
Q

Equation for lead oxide showing basic behaviour

A

PbO + 2H+ —> Pb2+ + 2H2O

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59
Q

Equation for lead oxide showing acidic behaviour

A

PbO + 2OH- + H2O —> [Pb(OH)4]2-

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60
Q

Which elements react with sodium hydroxide solution violently to form a salt and release hydrogen gas?

A

Aluminium, zinc and beryllium

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61
Q

What do alumnimum, zinc and beryllium do when reacting with sodium hydroxide?

A

React violently to form a salt and release hydrogen gas

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62
Q

What is it that usually reacts with NaOH to release hydrogen gas?

A

Acids

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63
Q

What is an additional characteristic property of amphoteric behaviour?

A

Aluminium, zinc and beryllium reacting with sodium hydroxide solution violently to form a salt and release hydrogen gas

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64
Q

Equation for the reaction between aluminium and hydroxide

A

2Al + 2OH- + 6H2O —> 2[Al(OH)4]- + 3H2

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65
Q

Which elements does the violent reaction with sodium hydroxide NOT occur with?

A

Tin and lead

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66
Q

What is the only amphoteric group II element?

A

Beryllium

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67
Q

What is the only thing really that beryllium has in common with the other group II elements?

A

It has 2 electrons in the s-orbital

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68
Q

How is beryllium different to the other group 2 elements?

A

It’s oxide and hydroxide are amphoteric

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69
Q

How are many of beryllium’s compounds?

A

Covalent

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70
Q

Example of a covalent compound contains beryllium

A

BeCl2

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71
Q

Why are many of beryllium’s’ compounds covalent?

A

A consequence of its small size and comparatively large electronegativity which favours the formation of covalent bonds

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72
Q

Why does beryllium favour the formation of covalent bonds?

A

Small size and comparatively large electronegativity

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73
Q

What are the properties of beryllium closer to than the properties of other group II elements?

A

Closer to those of aluminium

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74
Q

Alkali

A

A soluble base

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75
Q

Silicon oxide (IV)

A

SiO2

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76
Q

Phosphorus pentoxide

A

P2O5

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77
Q

Describe the elements on the right hand side of the periodic table

A

Acidic

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78
Q

How can we investigate amphoteric behaviour in a practical?

A

By adding aqueous sodium hydroxide solutions to a number of aqueous salt solutions

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79
Q

Method for investigating amphoteric behaviour by adding aqueous sodium hydroxide solutions to a number of aqueous salt solutions

A

1.) place 4cm^3 of the zinc sulfate solution in a test tube
2.) to this solution, add sodium hydroxide solution drop wise until a precipitate forms
3.) divide the precipitate into 2 equal amounts
4.) to the first portion, add nitric acid slowly until it is in excess
5.) to the second portion, add sodium hydroxide solution slowly until it is in excess
6.) note all observations
7.) repeat the experiment with lead (II) nitrate and aluminium chloride

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80
Q

What do zinc, lead and aluminium all have in common?

A

All show amphoteric behaviours

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81
Q

What was the observation for adding the initial sodium hydroxide to zinc sulfate, lead (II) nitrate and aluminium chloride?

A

White precipitate

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82
Q

Why do we add the initial sodium hydroxide in the investigating amphoteric behaviour experiment?

A

To get the elements which show amphoteric behaviour (zinc, lead (II) and aluminium) to be as hydroxides

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83
Q

What’s the observation when adding both excess nitric acid and excess sodium hydroxide to the zinc, lead (II) and aluminium hydroxides?

A

Precipitate dissolves
Colourless solution

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84
Q

What word do we always use to describe a “clear” solution?

A

Colourless

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85
Q

What’s happened to a precipitate when it’s no longer visible?

A

Dissolved

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86
Q

Why do we add both excess nitric acid and excess sodium hydroxide in the investigating amphoteric behaviour experiment?

A

To see the hydroxides of zinc, lead (II) and aluminium react in the same way with both acids and bases = amphoteric behaviour

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87
Q

equation for the reaction between lead (II) ions and hydroxide ions

A

Pb ^2+ (aq) + 2OH- (aq) –> Pb(OH)2 (s)

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88
Q

equation for the reaction between lead (II) hydroxide and sodium hydroxide solution

A

Pb(OH)2 + 2OH- –> [Pb(OH)4]^2- (aq)
(sodium is a spectator ion)

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89
Q

equation for the reaction between aluminium ions and hydroxide ions

A

Al^3+ (aq) + 3OH- (aq) –> Al(OH)3 (s)

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90
Q

equation for the reaction between aluminium hydroxide and sodium hydroxide solution

A

Al(OH)3 + OH- (aq) –> [Al(OH)4]- (aq)

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91
Q

how do we identify complex ions?

A

placed in square brackets

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92
Q

what type of ion is one in square brackets?

A

a complex one

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93
Q

[Pb(OH)4]2-

A

tetrahydroxo plumbate (II)

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94
Q

[Al(OH)4]-

A

tetrahydroxo aluminate (III)

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95
Q

what do the roman numerals represent in the complex ions equations?

A

the oxidation state of the lead/aluminium

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96
Q

tetrahydroxo plumbate (II)

A

[Pb(OH)4]2-

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97
Q

tetrahydroxo aluminate (III)

A

[Al(OH)4]-

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98
Q

latinised name of lead

A

plumbum

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99
Q

latinised name of aluminium

A

alum

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100
Q

how do we explain why metals form amphoteric oxides and hydroxides?

A

their electronegativity values are intermediate (lie between a metal and a non-metal

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101
Q

what happens to an aluminium metal when exposed to oxygen?

A

forms an oxide spontaneously

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102
Q

why would there be no initial reaction when trying to react an aluminium metal with something?

A

aluminium forms an oxide spontaneously when exposed to oxygen –> need to get through the oxide layer first

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103
Q

what are the group II elements we will study the most?

A

boron and aluminium

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104
Q

boron electronic configuration

A

1s^22s^22p^1

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105
Q

aluminium electronic configuration

A

[Ne]3s^23p^1

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106
Q

what do both boron and aluminium have in common?

A

both have 3 electrons in their outer shell

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107
Q

oxidation state of boron and aluminium in their compounds

A

+3

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108
Q

general formula of group III halides

A

MX3
(x is the halide)

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109
Q

structure of boron (III) chloride or boron fluoride

A

trigonal planar

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110
Q

number of electrons around boron and aluminium

A

6

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111
Q

what can we see about boron and aluminium due to the fact that they have 6 electrons around them?

A

they’re short of an octet = electron deficient

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112
Q

what does it mean if something is short of an octet?

A

is electron deficient

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113
Q

can boron and aluminium gain a pair of electrons readily?

A

yes

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114
Q

how come boron and aluminium can gan a pair of electrons readily?

A

they have strong electron acceptor properties

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115
Q

what are the electron acceptor properties of aluminium chloride reflected in?

A

the ready formation of the dimer Al2Cl6

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116
Q

what is Al2Cl6?

A

a dimer

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117
Q

what is the dimer Al2Cl6 made up of?

A

2 aluminium chloride molecules have combined

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118
Q

what type of bonds form to form an Al2Cl6 dimer?

A

co-ordinate bonds

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119
Q

how are co-ordinate bonds formed in the Al2Cl6 dimer?

A

by each aluminium atom accepting a lone pair of electrons from one of the chlorine atoms in the other molecule

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120
Q

why is the Al2Cl6 dimer able to form?

A

aluminium has a vacant 3d orbital
chlorine has a lone pair of electrons

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121
Q

is aL2cL6 electron deficient?

A

no

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122
Q

bonding in AlCl3

A

covalent

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123
Q

bonding in AlF3

A

ionic

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124
Q

bonding in Al2O3

A

ionic

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125
Q

properties of AlCl3 with its covalent bonding

A

sublimes at 180 degrees celcius (fairly low temperature)
decomposes in water
exists as dimers, Al2Cl6, when solid and as monomers
AlCl3 in the gaseous state

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126
Q

state of AlCl3

A

gaseous

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127
Q

when does AlCl3 exist as Al2Cl6??

A

when solid and as monomers

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128
Q

properties of AlF3 with its ionic bonding

A

melts at around 1300 degrees celcius
dissolves readily in water
molten alumnium fluoride is a good conductor of electricity

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129
Q

properties of Al2O3 with its ionic bonding

A

melts at 2100 degrees celcius
insoluble in water
conducts electricity when molten

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130
Q

which bonding is exhibited between aluminium and oxygen and fluorine and why?

A

ionic
large electronegativity

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131
Q

with which elements does aluminium exhibit ionic bonding and why?

A

with oxygen and fluorine
large electronegativity

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132
Q

what type of bonding does a large electronegativity usually lead to?

A

ionic

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133
Q

which type of bonding happens with less electronegative elements?

A

covalent

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134
Q

apart from with oxygen and fluorine, when else does aluminium exhibit ionic bonding?

A

when hydrated with water

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135
Q

why does ionic bonding occur in aluminium when hydrated with water?

A

large hydration energy owing to the very small size of the aluminium ion

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136
Q

list some aluminium elements that exhibit ionic bonding

A

AlF3
Al2O3
[Al(H2O)6]^3+

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137
Q

list some aluminium elements that exhibit covalent bonding

A

Al2Cl6
AlBr3
Al2H6

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138
Q

give an example of a donor-acceptor compound

A

ammonia-boron trifluoride

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139
Q

ammonia-boron trifluoride equation

A

NH3BF3

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140
Q

when is ammonia-boron trifluoride formed?

A

when ammonia reacts with boron trifloride, forming a co-ordinate bond

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141
Q

how are donor-acceptor compounds formed?

A

when one atom, having a lone pair of electrons to donate, reacts with another atom, which is electron deficient (i.e - has an empty orbital to accept the electron)

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142
Q

what type of bonding is in a donor-acceptor compound?

A

coordinate

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143
Q

what does each central atom have around it in a donor-acceptor compound?

A

an octet

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144
Q

why is ammonia-boron trifloride able to form?

A

nitrogen in ammonia: has a lone pair of electrons which is available for bonding
boron in boron trifluroride: has only 6 electrons in its outer shell
therefore, boron attains a full shell of electrons in forming a co-ordinate bond with the lone pair of electrons from the nitrogen

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145
Q

which group element is nitrogen?

A

group V

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146
Q

describe boron

A

electron deficient group III element

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147
Q

how many electrons does aluminium have in its outer shell in aluminium chloride and what does this mean?

A

6
is electron deficient

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148
Q

what happens when aluminium chloride reacts with a chloride ion?

A

aluminium will readily accept a lone pair of electrons from a chloride ion to form the complex ion tetrachloroaluminate (III)

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149
Q

[AlCl4]-

A

tetrachloroaluminate (III)

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150
Q

tetrachloroaluminate (III) ion

A

[AlCl4]-

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151
Q

equation for the reaction between aluminium chloride and chloride ions

A

AlCl3 + Cl- ⇌ AlCl4-

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152
Q

why will aluminium readily accept a lone pair of electrons from a chloride ion to form the complex ion tetrachloroaluminate (III)?

A

since aluminium in aluminium chloride only has 6 electrons in its outer shell (is electron deficient)

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153
Q

What leads to aluminium chloride being an industrially important catalyst?

A

the affinity of AlCl3 for chlorine species

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154
Q

what does the affinity of aluminium chloride for chlorine species lead to?

A

it’s an industrially important catalyst

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155
Q

examples of industrially important catalysts that aluminium chloride forms

A

AlCl3 - catalysed Friedel-crafts reactions. there is no need to heat when carrying out reactions = saves energy
Low melting temperature ionic liquids, which contain the chloroaluminate (III) ion, [AlCl4]- are being developed as “clean technology” catalysts for the polymerisation of alkenes. the ionic liquid contains the chloroaluminate (III)) ion and a large organic cation.

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156
Q

why is boron nitride of interest to us?

A

as the B-N bond is similar to the C-C bond

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157
Q

which 2 words can be used to describe boron nitride and its comparisons to carbon?

A

isoelectronic (with the elemental forms of carbon)
isomorphism (occurs between the 2 species)

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158
Q

isoelectronic

A

the same number of electrons or the same electronic structure

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159
Q

isomorphism

A

similar chemical composition and exist in the same crystalline form

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160
Q

the same number of electrons or the same electronic structure

A

isoelectronic

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161
Q

similar chemical composition and exist in the same crystalline form

A

isomorphism

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162
Q

how many electrons do both C-C and B-N bonds have between the 2 atoms?

A

12

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163
Q

describe the structure of boron nitride

A

giant covalent

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164
Q

what are the giant covalent structures of boron nitride?

A

hexagonal
cubic

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165
Q

what is hexagonal boron nitride isomorphic with?

A

graphite

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166
Q

what is hexagonal boron nitride with graphite?

A

isomorphic

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167
Q

similarities between hexagonal boron nitride and graphite

A

high melting and boiling points
softness
insoluble

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168
Q

why do both hexagonal boron nitride and graphite have high melting and boiling points?

A

each atom forms 3 or 4 covalent bonds which require a lot of heat energy to overcome

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169
Q

why are both hexagonal boron nitride and graphite soft?

A

weak van der waal forces allow them to slide over each other

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170
Q

what does the feature that both graphite and hexagonal boron nitride are soft allow them to be used as?

A

a lubricant

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171
Q

how come hexagonal boron nitride and graphite can be used as a lubricant? explain this

A

they’re both soft
weak van der waal forces allow them to slide over each other

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172
Q

why are both hexagonal boron nitride and graphite insoluble?

A

no charges particles to react with the permanent dipole of water molecules

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173
Q

differences between hexagonal boron nitride and graphite

A

hexagonal boron nitride is an electrical insulator, whereas graphite is an electrical conductor
atoms in adjacent layers in hexagonal boron nitride are in register. in graphite, they’re out of register
electronegativity difference between boron and nitrogen = polar B-N bonds. no electronegativity difference between the carbon atoms in graphite (all atoms are the same).

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174
Q

how come the bonds between B and N in boron nitride are polar?

A

electronegativity difference

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175
Q

how come hexagonal boron nitride is an electrical insulator whilst graphite is an electrical conductor even though they’re isomorphic?

A

hexagonal boron nitride only has localised electrons due to the large electronegativity difference between the nitrogen and boron atoms = no free electrons
in graphite, electrons are evenly delocalised (non-bonding electrons)

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176
Q

uses of hexagonal boron nitride

A

used in electronics as a substrate for semi-conductors
ceramics
microwave windows
catalyst carrier in fuel cells and batteries
single layers can be wrapped to create nanotubes

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177
Q

how are nanotubes created?

A

single layers of hexagonal boron nitride are wrapped

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178
Q

which property of hexagonal boron nitride allow them to be used in electronics as a substrate for semi-conductors?

A

electrical insulator

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179
Q

what can hexagonal boron nitride be used for due to it being an electrical insulator?

A

can be used in electronic as a substrate for semi-conductors

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180
Q

what is cubic boron nitride isomorphic with?

A

diamond

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181
Q

what is cubic boron nitride with diamond?

A

isomorphic

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182
Q

similarities between cubic boron nitride and diamond

A

hard
strong
extremely high melting and boiling points
insoluble
electrical insulator
Tetrahedral arrangement of atoms

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183
Q

why are both cubic boron nitride and diamond hard?

A

each atom is bonded to 4 others with a strong covalent bond = structure held together in a rigid 3D structure

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184
Q

why do both cubic boron nitride and diamond have extremely high melting and boiling points?

A

each atom has 3 or 4 covalent bonds which require a lot of heat energy to overcome

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185
Q

why are both cubic boron nitride and diamond electrical insulators?

A

no delocalised electrons in the structure to carry voltage

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186
Q

differences between cubic boron nitride and diamond

A

cubic boron nitride isn’t as hard as diamond = preferred for grinding certian materials
cubic boron nitride is more stable due to the lack of electronegativity difference between the atoms
diamond can react with transition metals like iron and above 800 degrees celcius can react with air to form CO2
Cubic boron nitride is a good thermal conductor

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187
Q

why is cubic boron nitride more stable than diamond?

A

due to the lack of electronegativity difference between the atoms

188
Q

what can diamond do above 800 degrees celcius?

A

can react with air to form CO2

189
Q

uses of cubic boron nitride

A

wear-resistant coating
industrial abrasive
cutting tools

190
Q

How many oxidation states do p-block elements usually show?

A

2

191
Q

What does the higher oxidation state of p-block elements correspond to?

A

The group number

192
Q

What is the lower oxidation state of p-block elements usually?

A

2 and lower

193
Q

Oxidation state of group 3 elements

A

+3 and +1

194
Q

Oxidation state of group 4 elements

A

+4 and +2

195
Q

Oxidation states of group 5 elements

A

+5 and +3

196
Q

Inert pair effect

A

The increasing reluctance, as you move down the group, of the s2 pair of electrons in the binding level to become involved in bonding

197
Q

What does the inert pair effect cause in groups 3, 4 and 5?

A

The lower oxidation states of the element are more stable the lower down the group due to the inert pair effect

198
Q

What happens to the metallic character of group 4 elements as the atomic number increases?

A

Increases

199
Q

What happens to the bonding as the atomic number increases down group 4?

A

Covalent bonding
Metallic bonding

200
Q

What happens as the atomic number increases down group 4 to the lower oxidation states?

A

Increase in stability

201
Q

Which oxidation states increase in stability down group 4?

A

Lower oxidation states

202
Q

Non-mets of group 4

A

Carbon and silicon

203
Q

Semi-metal of group 4

A

Germanium

204
Q

Metals of group 4

A

Tin and lead

205
Q

Describe the electronic structure of group 4 elements

A

4 electrons in the outer shell of their atoms
2 in the s-sub-shell, 2 in the p sub-shell

206
Q

2 allotropic forms of carbon

A

Diamond and graphite

207
Q

What structure do both of the allotropes of carbon have (diamond and graphite)?

A

Giant covalent structure

208
Q

Describe the structure, appearance, melting point and conductivity of silicon

A

Structure similar to that of diamond
Black crystalline solid
High melting point
Semi-conductor

209
Q

Describe the appearance and melting point of lead and tin

A

Soft shiny metals
Fairly low melting points

210
Q

Metallic character

A

The tendency of an element to lose electrons and form positive ions or cations

211
Q

What can explain the change form non-metallic to metallic character as the atomic number increases down group 4?

A

Electronegativity

212
Q

Why do the elements in group 4 change or non-metallic to metallic as the atomic number increases?

A

On going down the groups there is a decrease in electronegativity

213
Q

What does a decrease in electronegativity down group 4 lead to?

A

A change from non-metallic to metallic character

214
Q

Main oxidation states of carbon

A

+4
+2

215
Q

Main oxidation state of silicon

A

+4

216
Q

Main oxidation states of tin

A

+4
+2

217
Q

Main oxidation states of lead

A

+4
+2

218
Q

What happens to the stability of the +4 oxidation state down group 4?

A

Decreases

219
Q

What happens to the stability of the +2 oxidation state down group 4?

A

Increases

220
Q

Why does the stability of the +4 oxidation state decrease down group 4, and the +2 oxidation state increases in stability?

A

Due to the inert pair effect

221
Q

What happens in terms of oxidation states down groups 3, 4 and 5?

A

Increasing tendency for the lower oxidation state to become more stable than the group valency

222
Q

What’s the outer electronic structure of all of the elements in group 4?

A

ns^2npx^1npy^1
Where n varies from 2 (for carbon) and 6 (for lead)

223
Q

What does the oxidation state of +4 imply in group 4?

A

Shows that all of the outer electrons are directly involved in bonding

224
Q

What happens to the s^2 pair in group 4 as you go down the group?

A

There’s an increased tendency for it not to be involved in bonding

225
Q

What does it lead to if there’s a decreased tendency for a pair of electrons not to be used in bonding?

A

Has a lower oxidation state

226
Q

Explain the inert pair effect

A

As you get closer to the bottom of the group, there’s an increasing tendency for the s^2 pair not to be used in bonding. This leads to the element having a lower oxidation state —> the inert pair effect

227
Q

What happens to the oxidation state when going from gallium to indium to thallium and why?

A

+1 state becomes more stable
Inert pair effect

228
Q

What happens to the oxidation state when going from tin to lead and why?

A

+2 state becomes more stable
Inert pair effect

229
Q

What happens to the oxidation state when going from arsenic, antimony to bismuth and why?

A

+3 state becomes more stable
Inert pair effect

230
Q

Stability of CO2

A

Stable

231
Q

Stability of PbO2

A

Strong oxidising agent

232
Q

Equation for the reaction between lead oxide and hydrochloric acid

A

PbO2 + 4HCl —> PbCl2 + Cl2 + 2H2O

233
Q

What type of bonding predominates in the +4 state of group 4 oxides?

A

Covalent bonding

234
Q

When does covalent bonding predominate in group 4 oxides?

A

In the +4 state

235
Q

Bonding of group 4 oxides in the +2 state

A

Tendency to be ionic

236
Q

Proof that there’s a tendency to be ionic in the +2 state of group 4 oxides

A

+2 compounds of tin and lead contain the ions Sn^2+ and Pb^2+
e.g - SnCl2 and PbCl2

237
Q

What happens to +4 compounds in group 4’s stability down the group and why?

A

Decreases down the group
Are covalent
The oxidising power of the state increases down the group

238
Q

What happens to the stability of +2 compounds in group 4 down the group and why?

A

In the +2 stat, the compounds are more ionic than in the +4 state and their stability increases down the group (i.e - the reducing power of the state decreases)

239
Q

Which group 4 oxides do we need to be aware of?

A

Carbon monoxide
Carbon dioxide
Lead (II) oxide
Lead (IV) oxide

240
Q

Carbon monoxide:
Formula
Structure
Appearance/25°C
Redox behaviour
Acid-base behaviour

A

CO
Simple molecular
Colourless gas
Reducing agent
N/A

241
Q

Carbon dioxide:
Formula
Structure
Appearance/25°C
Redox behaviour
Acid-base behaviour

A

CO2
Simple molecular
Colourless gas
Stable
Weak acid

242
Q

Lead (II) oxide:
Formula
Structure
Appearance/25°C
Redox behaviour
Acid-base behaviour

A

PbO
Ionic lattice
Yellow solid
Stabl
Amphoteric

243
Q

Lead (IV) oxide:
Formula
Structure
Appearance/25°C
Redox behaviour
Acid-base behaviour

A

PbO2
N/A
Dark brown solid
Strong oxidising agent
N/A

244
Q

Why are CO and CO2 gas at room temperature?

A

Weak forces between separate molecules

245
Q

Why are lead (II) oxide and lead (IV) oxide solid at room temperature?

A

Low ionisation energy = can form positive ions
Oxides are ionic with strong attractions between ions
Solids

246
Q

Oxides of Carbon to learn

A

Carbon monoxide
Carbon dioxide

247
Q

What is carbon monoxide?

A

A very poisonous gas

248
Q

Why is carbon monoxide so dangerous?

A

Reacts in preference to oxygen with haemoglobin in the blood

249
Q

What is carbon monoxide send as?

A

A reducing agent industrially - extraction of iron in the blast furnace (reduction of iron (III) oxide to iron)

250
Q

What happens to carbon monoxide during the extraction of iron in the blast furnace?

A

Is oxidised to CO2

251
Q

Reaction in the blast furnace

A

Fe2O3 + 3CO —> 2Fe + 3CO2

252
Q

What is carbon monoxide easily oxidised from and to and why is this good?

A

+2 to +4
+4 is more stable

253
Q

Describe CO2 in water

A

Fairly soluble

254
Q

Is CO2 acidic? Explain

A

Weakly acidic
Is partially hydrolysed by water -> about 1% of the molecules

255
Q

HCO3- (aq)

A

Carbonic acid

256
Q

Equation for the reaction between CO2 and water

A

CO2 (g) + H2O (l) ⇌ H+ (aq) + HCO3 - (aq)

257
Q

What does CO2 react with alkalis for?

A

To form salts

258
Q

How does CO2 form salts?

A

Reacts with alkalis

259
Q

What type of salts does CO2 react with alkalis to form?

A

Carbonates
Hydrogen-carbonates

260
Q

CO2 reaction with sodium hydroxide

A

CO2 + NaOH —> NaHCO3

261
Q

CO2 reaction with 2 sodium hydroxides (concentrated)

A

CO2 + 2NaOH —> Na2CO3 + H2O

262
Q

What do carbonates and hydrogencarboantes react with acids to produce?

A

CO2

263
Q

When do carbonates and hydrogencarbonates release CO2?

A

When reacting with acids

264
Q

What does CO2 do to limewater?

A

Turns milky

265
Q

Lime water

A

Calcium hydroxide

266
Q

What turns limewater milky?

A

CO2

267
Q

Which reaction is used to test for acids?

A

When carbonates and hydrogencarbonates react with acids, it releases CO2
The CO2 turns limewater milky

268
Q

Equation for the reaction between Ethanoic acid and sodium hydrogencarbonate

A

CH3COOH + NaHCO3 —> CH3COONa + CO2 + H2O

269
Q

Oxidation state of carbon dioxide

A

+4

270
Q

What is the most stable, carbon monoxide or carbon dioxide?

A

Carbon dioxide

271
Q

Lead (II) oxide formula

A

PbO

272
Q

Lead (II) oxide at room temperature

A

Stable

273
Q

Lead (II) oxide description

A

Ionic yellow solid

274
Q

Lead (I) oxide in water

A

Insoluble

275
Q

Oxidation state of lead (II) oxide

A

+2

276
Q

Reducing or oxidising properties of lead (II) oxide

A

None - stable

277
Q

Acid-base behaviour of lead (II) oxide

A

Amphoteric
(Forms salts with acids and alkalis)

278
Q

Equation for the reaction between lead (II) oxide and an acid

A

PbO + 2H+ —> Pb2+ + H2O

279
Q

Equation for the reaction between lead (II) oxide and an alkali?

A

PbO + 2OH- + H2O —> [Pb(OH)4]^2-

280
Q

Description of lead (IV) oxide

A

Dark brown solid

281
Q

Lead (IV) oxide formula

A

PbO2

282
Q

Redox behaviour of PbO2

A

Powerful oxidising agent

283
Q

Example of PbO2 being a powerful oxidising agent

A

Will oxidise chloride ions to chlorine. The lead (IV) is reduced to lead (II) chloride which is found as white precipitate.

284
Q

How is lead (II) chloride found?

A

White precipitate

285
Q

What has to be done to the reagents in the reaction between PbO2 and chloride ions?

A

Have to be heated

286
Q

Equation for the reaction between lad (IV) oxide and hydrochloric acid

A

PbO2 + 4HCl —> PbCl2 + Cl2 + 2H2O

287
Q

Oxidation state of PbO2

A

+4

288
Q

What is lead easily reduced from and to in PbO2 and why?

A

From +4 to +2
+2 is more stable

289
Q

Method used to identify Pb2+ ions and results

A

Pb2+ solution and I- ions (usually KI is used)
= bright yellow precipitate of lead (II) iodide

290
Q

Equation for the reaction between lead ions and iodide ions

A

Pb2+ (aq) + 2I- (aq) —> PbI2 (s)

291
Q

Group 4 halides to remember

A

Tetrachloromethane
Silicon (IV) chloride
Lead (I) chloride

292
Q

Tetrachloromethane:
Formula
Bonding
Structure
Appearance
Reaction with water

A

CCl4
Covalent
Simple molecular
Colourless liquid
No reaction - forms a separate layer

293
Q

Silicon (IV) chloride:
Formula
Bonding
Structure
Appearance
Reaction with water

A

SiCl4
Covalent
Simple molecular
Colourless liquid
Violent reaction

294
Q

Lead (II) chloride:
Formula
Bonding
Structure
Appearance
Reaction with water

A

PbCl2
Ionic
Ionic lattice
White solid
Insoluble - cold water
Soluble- hot water

295
Q

How com Tetrachloromethane has a simple molecular structure?

A

Covalent
Low boiling point

296
Q

Why can’t Tetrachloromethane react with water?

A

Carbon cannot expand is outet to form a coordinate bond with a lone pair of electrons from the water molecule
Carbon has no empty d levels of suitable energy available for bonding

297
Q

How come silicon (IV) chloride is able to react violently with water?

A

Has empty d levels of suitable energy available and the tetra slide reacts with water violently, forming co-ordinate bonds

298
Q

Reaction between silicon (IV) chloride and water equation

A

SiCl4 + 2H2O —> SiO2 + 4HCl

299
Q

What does SiO2 (product of the reaction between solid ion (IV) chloride and water) appear as?

A

White precipitate

300
Q

How does HCl show up after the reaction between silicon (IV) chloride and water?

A

White fumes

301
Q

Structure of silicon (IV) chloride

A

Tetrahedral (like all chlorides)

302
Q

Appearance of lead (II) chloride

A

White ionic solid

303
Q

Describe the solubility of lead (II) chloride in water

A

Cold water = insoluble
Hot water = partially soluble

304
Q

Equation for the reaction between lead (II) chloride and water

A

H2O
PbCl2 (s) ——> Pb^2+ (aq) + 2Cl- (aq)

305
Q

Natural source of silicon

A

Sand (silicon dioxide)

306
Q

Describe lead as a metal

A

Dense

307
Q

What is lead used in?

A

Roofing

308
Q

What does a +2 oxidation state imply?

A

2 electrons used in bonding, 2 not

309
Q

What does something being stable imply?

A

Unreactive

310
Q

What does something being unstable imply?

A

Reactive

311
Q

CO2 reacting with dilute NaOH

A

NaHCO3

312
Q

CO2 reacting with concentrated NaOH

A

Na2CO3

313
Q

How do we know if NaOH is concentrated?

A

“2” in front

314
Q

Why do carbonates and hydrogencarbonates do in the body?

A

Control blood pH

315
Q

Under which conditions alone can we extract elements by reduction?

A

With elements less reactive than carbon

316
Q

What must we do to an equation if we have anything reacting as an acid?

A

Must include water to balance it out

317
Q

When must we include water in these equations in this unit?

A

If we have anything reacting as an acid

318
Q

What must we back any answers up with?

A

Equations

319
Q

How come silicon (IV) chloride has a violent reaction with water but Tetrachloromethane doesn’t?

A

Silicon is a bigger atom than carbon
Can surround more atoms around it

320
Q

What are group 7 elements otherwise known as?

A

The halogens

321
Q

How are the halogens always written?

A

Diatomically

322
Q

Are the halogens metals or non-metals?

A

Non-metals

323
Q

Are the halogens similar to one another?

A

Yes - they show gradual changes only with increasing atomic mass

324
Q

What type of bonds do the halogens form with metals?

A

Ionic

325
Q

What type of bonds do the halogens form with non-metals?

A

Covalent bonds

326
Q

Structure of halogens

A

Simple molecular

327
Q

Change in atomic radius down the group of halogens + explanation

A

Atomic radius increases down the group
More shells present

328
Q

Describe the melting + boiling points of the halogens + explain this

A

Low melting and boiling points
Diatomic molecules and a simple molecular structure

329
Q

What happens to the melting and boiling points of the halogens down the group? Explain this

A

Increases
On going down the group, the size increases with a subsequent increase in the Van der Waal’s forces (induced dipole-induced dipole)
It’s these forces that are broken when the halogens are changing state

330
Q

Type of Van der Waal forces in halogens + why

A

Induced dipole-induced dipole
Non-polar covalent molecules

331
Q

Explain the formation of the induced-dipole induced-dipole Van der Waal forces

A

Slight uneven distribution of electrons = temporary dipole
Causes a chain of induced dipole-induced dipole bonds

332
Q

In which halogen are the induced-dipole induce-dipole bonds the most significant and why?

A

Iodine
Largest molecule = most electrons

333
Q

States of the different halogens + explanation

A

F2, Cl2 = gases
Br2 = liquid
I2 = solid

On going down the group, the Van der Waal forces increase, increasing the melting points down the group

334
Q

What happens to electronegativity down the group of halogens?

A

Decreases

335
Q

What happens down group 7 when the electronegativity decreases down the group?

A

Decrease in reactivity from F to I

336
Q

What happens to the electrode potentials values down group 7?

A

The positive value of the electrode potential decreases

337
Q

What does the decrease in the positive value of the electrode potentials down group 7 reflect?

A

The decrease in reactivity and oxidising power of the elements down the group

338
Q

What are all of the halogens?

A

Oxidising agents

339
Q

What’s different about the oxidising properties of the halogens?

A

Have different strengths as oxidising agents

340
Q

What happens to the oxidising power of the halogens down group 7?

A

Decreases

341
Q

How can we measure the halogens ability to remove electrons from other species?

A

Measure using the standard electrode potentials for the halogens

342
Q

Most powerful oxidising agent of group 7

A

Cl2

343
Q

Most powerful reducing agent of group 7

A

I-

344
Q

In the reaction below, what is the oxidising agent and what is the reducing agent?
Cl2 (g) + 2e- —> 2Cl- (aq)
<—

A

Cl2 - oxidising agent
Cl- - reducing agent

345
Q

Chlorine physical appearance at room temperature

A

Green gas

346
Q

Bromine physical appearance at room temperature

A

Orange liquid

347
Q

Iodine physical appearance at room temperature

A

Black/purple solid

348
Q

Which halogen is a green gas at room temperature?

A

Chlorine

349
Q

Which halogen is an orange liquid at room temperature?

A

Bromine

350
Q

Which halogen is a black/purple solid at room temperature?

A

Iodine

351
Q

What can all of the halogens do at room temperature?

A

Produce gases

352
Q

What can bromine easily do at room temperature and what word describes this?

A

Easily vaporises at room temperature
Volatile

353
Q

What can iodine do at room temperature and what does this mean?

A

Sublime
Straight from solid to gas

354
Q

Which halogen can sublime at room temperature?

A

Iodine

355
Q

What halogen is volatile (easily vaporises at room temperature)?

A

Bromine

356
Q

What type of reaction is the reaction which occurs between a halogen and a halide ion?

A

A redox reaction

357
Q

How are redox reactions explained?

A

Using electrode potentials

358
Q

Electrode potential

A

The tendency of the electrode to attract electrons to itself (to be reduced)

359
Q

What do the electrode potentials tell us about redox reactions?

A

The trend in the oxidising power of the elements

360
Q

What will a halogen displace in a displacement reaction?

A

A halogen lower in the group form one of its ions

361
Q

Chlorine displacement reaction with bromide ions equation

A

Cl2 + 2Br- —> 2Cl- + Br2

362
Q

Explain why chlorine is able to oxidise bromide ions using electrode potential values

A

The electrode potential value for the 1/2Cl2/Cl- half cell is more positive than the value for the 1/2Br2/Br- half cell,so chlorine will be able to oxidise the bromide ions

363
Q

Which ions is bromine able to oxidise and which isn’t it able to?

A

Oxidise iodide ions
Cannot oxidise chloride ions

364
Q

What happens to the oxidising power of the halogens down the group and why?

A

Oxidising power decreases own the group
The positive value for the electrode potential decreases

365
Q

Chlorine and iodide ions displacement equation

A

Cl2 + 2I- —> 2Cl- + I2

366
Q

Chlorine reaction with…
Chloride ions
Bromide ions
Iodide ions

A

(No reaction)
Yellow orange Br2 formed
Black I2 formed

367
Q

Bromine reaction with…
Chloride ions
Bromide ions
Iodide ions

A

No reaction
No reaction
Black I2 formed

368
Q

Iodine reaction with…
Chloride ions
Bromide ions
Iodide ions

A

No reaction for all

369
Q

Describe iodine in water

A

Fairly insoluble

370
Q

What does iodine dissolve in since it’s fairly insoluble in water?

A

In aqueous potassium iodide

371
Q

What happens when iodine is dissolved in aqueous potassium iodide?

A

Forms a red-brown solution, containing the ion I3-

372
Q

How do we form a red-brown solution containing the ion I3-?

A

Iodide dissolved in aqueous potassium iodide

373
Q

1st stage of the reaction between sodium chloride (or any chloride) with concentrated sulphuric acid

A

Hydrogen halide is formed in a displacement reaction

374
Q

Sodium chloride and concentrated sulphuric acid equation

A

NaCl + H2SO4 —> NaHSO4 + HCl

375
Q

HCl formed in the displacement reaction between sodium chloride and concentrated sulphuric acid

A

Hydrogen chloride gas, HCl
(Not hydrochloric acid!)

376
Q

What does hydrogen chloride gas form as?

A

White misty fumes

377
Q

Why does hydrogen chloride gas form white misty fumes?

A

Reacts with water vapour in the air

378
Q

What can also be observed apart from misty fumes when sodium chloride reacts with sulphuric acid? Why?

A

Effervescence
Gas is released

379
Q

What type of reaction is that between sodium chloride and sulphuric acid?

A

Displacement
Exothermic

380
Q

Explain the second stage of the reaction between sulphuric acid and sodium chloride

A

The extent of further reaction depends on the reducing power of the halide ion

381
Q

What happens to the reducing power of the halides down the group + why?

A

Increases down the group
The electrode potential becomes less positive

382
Q

When does the second reaction occur between sulphuric acid and sodium chloride?

A

When the reducing power of the halide ion is high enough, increasing the tendency for the ion to be oxidised

383
Q

First stage of the reaction between concentrated sulphuric acid and sodium bromide

A

Produces HBr gas

384
Q

Equation for the reaction between sodium bromide and sulphuric acid

A

NaBr + H2SO4 —> NaHSO4 + HBr

385
Q

Hydrogen bromide gas

A

HBr

386
Q

How does hydrogen bromide gas appear?

A

White misty fumes

387
Q

What’s the second stage of the reaction between sodium bromide and sulphuric acid? Explain

A

Further reaction with HBr
Redox reaction
HBr is oxidised
H2SO4 is reduced

388
Q

Equation for the second reaction between sodium bromide and sulphuric acid (the HBr reacting further)

A

2HBr + H2SO4 —> SO2 ++ Br2 + 2H2O

389
Q

Observations during the further reaction of HBr gas after reacting sodium bromide and sulphuric acid

A

Brown/orange fumes of the Br2 gas
SO2 —> pungent smelling, acidic gas

390
Q

Describe SO2

A

Pungent smelling, acidic gas

391
Q

1st reaction for the addition of sodium iodide to sulphuric acid

A

Produced HI gas

392
Q

Equation for the reaction between sodium iodide and sulphuric acid

A

NaI + H2SO4 —> NaHSO4 + HI

393
Q

Hydrogen chloride gas

A

HCl

394
Q

Hydrogen iodide gas

A

HI

395
Q

What does HI form as?

A

White misty fumes

396
Q

Describe the second stage of the reaction between sodium iodide and sulphuric acid

A

Further reaction with HI
Redox reaction
HI —> oxidised
H2SO4 —> reduced

397
Q

Reduction of H2SO4 equation

A

H2SO4 —> SO2 + S + H2S

398
Q

Equation for the reaction between HI and sulphuric acid

A

2HI + H2SO4 —> SO3 + I2 + 2H2O

399
Q

Observations for the reaction between HI and sulphuric acid

A

Dark purple fumes/brown solid as I2 forms
SO2 = pungent smell
When sulphur is reacted further i.e - H2SO4 —> SO2 + S + H2S, a H2S yellow solid may be seen, with a smell of rotten eggs

400
Q

Describe H2S

A

Solid yellow
Smells of rotten eggs

401
Q

Which is the only halide to react with sulphuric acid to produce a rotten egg smell?

A

Iodine

402
Q

What smell does iodide form when reacting with sulphuric acid?

A

Rotten eggs

403
Q

If there’s no rotten egg smell when reacting halogens with sulphuric acid, what can we deduce?

A

It’s not iodide, so it must be one of the other halides

404
Q

What does how chlorine reacts with sodium hydroxide depend on?

A

The conditions

405
Q

Which type of reaction occurs between chlorine and sodium hydroxide no matter the conditions?

A

Disproportionation reaction

406
Q

Disproportionation reaction

A

Same element is oxidised and reduced during the reaction

407
Q

What happens to chlorine during a disproportionation reaction?

A

Undergoes self oxidation-reduction

408
Q

Which ions are formed when chlorine reacts with cold, dilute aqueous sodium hydroxide?

A

Chlorate (I) ion, ClO-
Chloride ion

409
Q

Under which conditions does chlorine reactions with sodium hydroxide form the chlorate (I) ion, ClO- and the chloride ion?

A

When reacting with cold, dilute aqueous sodium hydroxide

410
Q

Ionic equation for the reaction between chlorine and cold, dilute sodium hydroxide

A

Cl2 + 2OH- —> Cl- + ClO- + H2O

411
Q

Alternative equation for the reaction between chlorine and sodium hydroxide, including the sodium ions (spectator ions)

A

Cl2 + 2NaOH —> NaCl + NaClO + H2O

412
Q

chlorate (I) ion

A

ClO-

413
Q

ClO-

A

Chlorate (I) ion

414
Q

What is sodium chlorate (I) used in?

A

Bleach, germicide

415
Q

What is sodium chloride used for?

A

Treating icy roads in the winter
Food preservative

416
Q

How does sodium chloride help treat icy roads in the winter?

A

The impurity lowers the freezing temperature

417
Q

Why is chlorate (I) ion important in bleaches and bactericides?

A

Powerful oxidising agent

418
Q

How is the chlorate (I) ion in domestic bleaches?

A

As an aqueous solution of sodium chlorate (I)

419
Q

What happens to the chlorate (I) ion at higher temperatures?

A

Disproportionation

420
Q

Reaction for the disproportionation of the ClO - ion

A

3ClO- —> 2Cl- + ClO3-

421
Q

What is chlorine used for and wh?

A

Used to kill bacteria in swimming pols and purify domestic water supplies
Strong oxidising power

422
Q

Describe reaction with chlorine and hot, concentrated aqueous sodium hydroxide

A

The chlorate (I) ion disproportionates as it is formed and the chlorate (v) ion is formed with the chloride ion

423
Q

Ionic equation for the reaction between chlorine and hot, concentrated aqueous sodium hydroxide

A

3Cl2 + 6OH- —> 5Cl- + ClO3- + 3H2O

424
Q

Alternative equation for the reaction of chlorine with hot, concentrated sodium hydroxide

A

3Cl2 + 6NaOH —> 5NaCl + NaClO3 + 3H2O

425
Q

Sodium chlorate (V)

A

NaClO3

426
Q

NaClO3

A

Sodium chlorate (V)

427
Q

Wha’s sodium chlorate (V) (NaClO3) used in?

A

Weed killers

428
Q

Why is sodium chlorate (V) used as a weed killer?

A

Removes weeds and prevents new ones from growing
Powerful oxidising agent

429
Q

Problem with sodium chlorate (V)

A

Bad for the environment (as a weed killer)

430
Q

Commercial use of chlorine

A

Used to purify drinking water and swimming pools

431
Q

Equation for the reaction between chlorine and water, showing how it’s used to purify drinking water and swimming pools

A

Cl2 + H2O —> HOCl + HCl
HOCl = oxidising agent

432
Q

What is sodium chloride used for?

A

Enhancing the flavour of food
As a food preservative
For de-icing roads in wintry weather

433
Q

List 3 chlorine compounds and their uses

A

Chlorothene = used to make plastic PVC
Organochlorine compounds have found widespread use as CFC’s, solvents, pesticides, anaesthetics
Sodium chlorate (I) is used in bleach

434
Q

What’s chloroethene used for?

A

To make plastic PVC

435
Q

Uses of organochlorines

A

CFC’s, solvents, pesticides, anaesthetics

436
Q

Commercial uses of bromine

A

Used to make…
Bromide salts = used in photography
Organobromide compounds such as 1,2 dibromoethene (petrol additive)

437
Q

Commercial uses of iodine

A

In alcoholic solution is used to disinfect wounds
Potassium iodide is used in the photographic industry

438
Q

How do we test for halide ions?

A

1.) nitric acid
2.) add silver nitrate solution

= precipitate of filter halide forms

439
Q

Why do we add nitric acid when testing for halide ions?

A

To dilute any OH- ions present

440
Q

Colour of silver halide precipitate with Cl-

A

White

441
Q

Colour of silver halide precipitate with Br-

A

Cream precipitate

442
Q

Colour of silver halide precipitate with I-

A

Yellow precipitate

443
Q

Equation for the reaction between silver nitrate and chloride ions

A

Ag+ (aq) + Cl- (aq) —> AgCl (s)

444
Q

Testing silver halide ions further after nitric acid + observations

A

Add ammonia (NH3)
Cl- = soluble
Br- = mostly insoluble
I- = insoluble

445
Q

What type of analysis are flame tests?

A

Qualitative

446
Q

What are flame tests used for?

A

To identify group I and II cations

447
Q

flame test method

A
  1. Dip a platinum wire/wet wooden splint into concentrated HCl
  2. Dip the top of the wire in the solid to be tested
  3. Tip of the wire is placed in the blue Bunsen flame
  4. Note the colour of the flame
448
Q

Lithium cation flame colour

A

Red

449
Q

Sodium cation flame colour

A

Intense yellow

450
Q

Potassium cation flame colour

A

Lilac

451
Q

Magnesium cation flame colour

A

No colour

452
Q

Calcium cation flame colour

A

Brick red

453
Q

Strontium cation flame colour

A

Crimson

454
Q

Barium cation flame colour

A

Apple green

455
Q

Beryllium showing acidic behaviour equation

A

Be(OH)2 + 2OH- —> [Be(OH)4] 2-
Tetrahydroxo beryllate (II)

456
Q

Tin showing acidic behaviour equation

A

Sn(OH)2 + 2OH- —> [Sn(OH)4]2-
Tetrahydroxo stannate (II)

457
Q

Equation aluminium showing acidic behaviour

A

Al(OH)3 + OH- —> [Al(OH)4]-
Tetrahydroxo laminate (III)

458
Q

When showing basic or acidic behaviour is there always water formed?

A

Basic

459
Q

Electron deficiency

A

Outer shell of electrons is not full (short of an octet)

460
Q

List and explain 4 reaction of Pb2+ ions

A

NaOH
Pb2+ + OH- (aq) —> Pb(OH)2 (s) white ppt

XsNaOH
Precipitate dissolves to give colourless solution
Pb(OH)2 (s) + OH- (aq) —> [Pb(OH)4]^2- (aq)

HCl
Pb2+ (aq) + Cl (aq) —> PbCl2 (s) white ppt

KI
Pb2+ (aq) + 2I- —> PbI2 (s) yellow ppt

461
Q

How does H2SO4 undergo further reaction in the reaction with sodium iodide?

A

HI reduces the sulphur in H2SO4

462
Q

Is it iodine or iodide ions that make a good reducing agent?

A

Iodide ions

463
Q

Why are the oxides of lead both solids?
PbO = yellow solid
PbO2 = dark brown solid

A

They’re ionic with strong attractions between the ions

464
Q

What is the only soluble lead salt?

A

Lead nitrate

465
Q

What occurs during all of the halide ions and concentrated sulphuric acid reactions that we tend to not really mention?

A

White solid forms

466
Q

Why can’t nitrogen not expand its octet?

A

No available d orbitals in the outer shell

467
Q

Describe the observations of the reaction between SiCl4 and water

A

Violent reaction —> white precipitate and steamy fumes

Chemistry (39 decks)

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