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A Level Chemistry Year 1 > Periodic Table > Flashcards

Periodic Table Flashcards

1
Q

What is the modern periodic table arranged in?

A

By order of proton number

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2
Q

What are Döbereiner triads?

A

The groups that were made when, in 1817, Johann Döbereiner attempted to group similar elements

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3
Q

What did the English chemist, John Newlands, notice when her arranged the elements in order of mass?

A

That elements with similar chemical and physical properties occurred at regular intervals, but every 8th element was different

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4
Q

What did Newlands call his discovery?

A

The law of octaves

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5
Q

What does the Periodic law state?

A

If you arrange the elements in order of increasing atomic number then their chemical and physical properties will repeat in a systematic way

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6
Q

What was different about Mendeleev’s table?

A

He left gaps, so that elements with similar properties were in the same group

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7
Q

Give an example of a case where 2 elements in the periodic table are the wrong way around?

A

Ar —-> K are not in order of atomic mass

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8
Q

All the elements within the same period have the same number of…

A

Shells

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9
Q

What is periodicity?

A

A repeating pattern of physical and chemical properties across a period

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10
Q

Why is the sub-shell 4s before 3d?

A

The 4s sub-shell has a lower energy level

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11
Q

What is the electron configuration of cobalt (27 electrons)

A

1s2 2s2 2p6 3s2 3p6 4s2 3d7

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12
Q

What is the first ionisation energy?

A

The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms

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13
Q

What type of process is the first ionisation energy?

A

Endothermic, as you have to put energy in to remove the 1st electron

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14
Q

Write out the 1st ionisation energy equation of oxygen

A

O —–> O+ + E-

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15
Q

Why do you have to use the gas state symbol when writing out the equation for the first ionisation energy?

A

Because ionisation energies are measured for gaseous atoms

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16
Q

The lower the ionisation energy…

A

The easier it is to form an ion

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17
Q

What is an endothermic process?

A

One that takes in heat

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18
Q

What does a high ionisation energy mean?

A

There’s a strong electrostatic attraction between the electron and nucleus, so more energy is needed to remove the first electron

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19
Q

What 3 things affect ionisation energy?

A

Nuclear charge
Atomic radius
Shielding

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20
Q

How does nuclear charge affect ionisation energy?

A

The more protons there are in the nucleus, the more positively charged it is and so the stronger the attraction for the electrons

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21
Q

How does atomic radius affect ionisation energy?

A

Attraction falls off rapidly with distance, so an electron closer to the nucleus will be more strongly attracted than one that’s further away

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22
Q

How does shielding affect ionisation energy?

A

As the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction

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23
Q

What is the trend in ionisation energy down a group?

A

As you go down a group, ionisation energies decrease (as it gets easier to remove an electron)

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24
Q

Why does ionisation energy decreases down a group?

A

Elements further down a group have extra electron shells, so the atomic radius is larger (reducing attraction to the nucleus). Shielding also increases

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25
Q

What is the trend of ionisation energies as you go along a period?

A

They increase (gets harder to remove the outer electrons)

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26
Q

Why do the ionisation energies increase as you go along a period?

A

Number of protons increases, meaning a larger nuclear charge, so a smaller atomic radii and a stronger attraction

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27
Q

Is there any shielding change as you go along a period?

A

No, as all the extra electrons are at roughly the same energy level

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28
Q

As you go along a period, why is there a drop of ionisation energy between Group 2 and 3?

A

Due to the outer electrons in Group 3 being in a p-orbital rather than an s-orbital

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29
Q

As you go along a period, why is there a drop of ionisation energy between Group 5 and 6?

A

Due to electron repulsion

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30
Q

Why does electron repulsion cause a drop of ionisation energy between Group 5 and 6?

A

For the Group 5, the outer electron is being removed from an orbital containing 1 electron, whereas for the Group 6, the outer electron is being taken from an orbital containing 2 electrons (which repel each other). This repulsion makes it easier to remove an electron

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31
Q

What is successive ionisation energy?

A

The energy needed to remove 1 mole of each subsequent electron from each ion in 1 mole of positively charged gaseous ions

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32
Q

What is the equation for the second ionisation of Oxygen?

A

O+ —-> O2+ + e-

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33
Q

Explain the trend in successive ionisation energies

A

Within each shell, successive ionisation energies increase because electrons are being removed from an increasingly positively charged ion (so there’s less repulsion between the remaining electrons). There are big jumps in ionisation energy every time a new shell is broken into

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34
Q

What are macromolecular structures?

A

Another name for giant covalent lattices

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35
Q

Give 2 examples of elements that can form macromolecular structures

A

Carbon and Silicon

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36
Q

Why can carbon and silicon form giant covalent structures?

A

They both form 4 strong, covalent bonds

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37
Q

What is an allotrope?

A

A different form of the same element in the same state

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38
Q

What are the 3 carbon allotropes?

A

Graphite
Diamond
Graphene

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39
Q

How are the carbon atoms arranged in graphite?

A

Sheets of flat hexagons covalently bonded with 3 bonds each (the 4th outer electron of each carbon atom is delocalised)

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40
Q

What are the sheets of hexagons in graphite bonded together by?

A

Weak induced dipole-dipole forces

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41
Q

What does delocalised mean?

A

An electron isn’t attached to a particular atom

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42
Q

Why is graphite used as a dry lubricant?

A

The weak forces between layers can easily be broken, so the sheets can slide over each other

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43
Q

Can graphite conduct electricity?

A

Yes, as the delocalised electrons are free to move along the sheets

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44
Q

Describe the density of graphite?

A

A low density because the layers are far away from each other compared to the length of the covalent bonds

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45
Q

Does graphite have a high or low melting point?

A

High melting point (around 3700K) because of the strong covalent bonds

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46
Q

Is graphite soluble or insoluble?

A

Insoluble in any solvent, as the covalent bonds in the sheets are too difficult to break

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47
Q

What is the shape formed by the carbon atoms in diamond?

A

Tetrahedral, as each carbon is covalently bonded to 4 others

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48
Q

Describe the melting/boiling point of Diamond

A

Very high melting point

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49
Q

Describe the strength of Diamond

A

Extremely hard (which is why it’s used on the tips of diamond-tipped drills)

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50
Q

Describe the thermal conductivity of Diamond

A

It’s a good thermal conductor because vibrations travel easily through it

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51
Q

Can Diamond conduct electricity?

A

No because all the outer electrons are held in localised bonds

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52
Q

Does Diamond dissolve in solvent?

A

No

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53
Q

What are macroscopic properties?

A

Properties that can directly be perceived

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54
Q

What are the properties of Silicon similar to?

A

That of diamond

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55
Q

What is graphene?

A

One layer of graphite (a sheet of carbon atoms in hexagonal shapes)

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56
Q

Why is graphene a 2D compound?

A

It’s one atom thick

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57
Q

Why is graphene thought of as the best electrical conductor?

A

The delocalised electrons are free to move, but because graphene doesn’t have any layers they can move quickly above and below the sheet making good at conducting electricity

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58
Q

Describe the strength of graphene

A

Very strong because the delocalised electrons strengthen the covalent bonds between the covalent bonds

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59
Q

Give 2 uses of graphene

A

High-speed electronics

Aircraft technology

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60
Q

What properties make graphene useful in high-speed electronics?

A

It’s high strength, low mass and good electrical conductivity

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61
Q

Why is graphene used in the screens of phones?

A

It’s flexible and transparent

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62
Q

What is metallic bonding?

A

A lattice of metal cations in a sea of delocalised electrons

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63
Q

What 2 things affect the melting/boiling points of metals

A

The number of delocalised electrons per atom

Size of metal ion and lattice structure

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64
Q

How does the number of delocalised electrons per atom affect the melting/boiling points of metals?

A

The more there are, the stronger the metallic bonds and so the higher the melting and boiling point is

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65
Q

Which has a higher melting point between Na+ and Mg2+?

A

Mg2+ as it has 2 delocalised electrons per atom compared to Na+ having only 1

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66
Q

How does the size of the ion and lattice structure affect the boiling/melting points of metals?

A

A smaller ionic radius will hold the delocalised electrons closer to the nucleus (meaning a larger melting/boiling point)

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67
Q

Are metals malleable or non-malleable?

A

They’re malleable because there an no bonds holding specific ions together, and so the metal ions can slide over each other when the structure is pulled

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68
Q

Describe the thermal conductivity of metals

A

They are good thermal conductors because the delocalised electrons can pass kinetic energy to each other

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69
Q

Describe the electrical conductivity of metals

A

They are good electrical conductors because the delocalised electrons can carry a current

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70
Q

Describe the solubility of metals

A

They are insoluble (except in liquid metals)

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71
Q

Why are metals insoluble?

A

Because of the strength of the metallic bonds

72
Q

What is charge density?

A

Amount of charge in relation to the size of the ion

73
Q

Describe the trends of boiling/melting points across Periods 2 and 3

A

A general increase between the 1st and 4th elements in the period, but then decrease from 4th to the 8th

74
Q

Explain the trends in boiling/melting points across Periods 2 and 3

A

As you go along the period, the types of structure changes from giant metallic to giant covalent and then to a simple molecular lattice

75
Q

For metals across a period, why does the boiling/melting point increase?

A

The metal-metal bonds get stronger because the metal ions have a greater charge, an increasing number of delocalised electrons and a smaller atomic radii

76
Q

Which has a greater charge density between a small and a large ions with the same charge?

A

The small ion

77
Q

Why do B, C and Si have the highest melting and boiling point in their period?

A

They have strong covalent bonds linking their atoms together, which need a lot of energy to break

78
Q

Why do simple molecular structures have low melting and boiling points?

A

The melting and boiling points depend on the strength of the London forces (not how strong the covalent bonds are), so these forces are weak and are easily overcome

79
Q

More……..in a molecules mean stronger London forces

A

Atoms

80
Q

Why do the noble gases have the lowest boiling and melting point?

A

They are monatomic, so the London forces are very weak

81
Q

What block of the periodic table are the alkaline earth metals in?

A

S-block

82
Q

Where are the alkaline earth metals found?

A

Group 2

83
Q

What is the trend in reactivity as you go down Group 2?

A

Increases

84
Q

Why does the reactivity increase down Group 2?

A

When Group 2 elements react, they lose electrons forming positive ions (cations). The easier it is to remove an electron, the more reactive the element. So because ionisation energy decreases down a group, the reactivity increases

85
Q

What is the electronic structure of Mg?

A

1s2 2s2 2p6 3s2

86
Q

What is the electronic structure of Mg2+?

A

1s2 2s2 2p6

87
Q

Write an ionic equation for when Mg reacts

A

Mg ——> Mg2+ + 2e-

88
Q

What is the product when Group 2 elements react with water?

A

Metal Hydroxide and Hydrogen

89
Q

Write the equation for the reaction of Mg with H2O

A

Mg + 2H2O —–> Mg(OH)2 + H2

90
Q

What is the pH of the reaction of a Group 2 element with water?

A

pH 12 or 13, as the metal hydroxide that forms dissolves in water releasing Hydroxide ions (making it strongly alkaline)

91
Q

Why do ionisation energies decrease as you go down a Group?

A

Because of the increasing size of the atomic radii and the increasing shielding effect

92
Q

What is the general formula of the reaction between the Group 2 metal ‘M’ and Water?

A

M + 2H20 —-> Ca2+ + 2OH- + H2

93
Q

What is the general formula of the reaction between the Group 2 metal ‘M’ and Oxygen?

A

2M + 02 —–> 2MO

94
Q

What is the product when Group 2 metals burn in oxygen?

A

A solid white oxide

95
Q

Why are Group 2 metals known as Alkaline Earth Metals?

A

Because their oxides form alkaline solutions (pH12-13)

96
Q

What is the product of the reaction between the oxides of Group 2 metals and water?

A

They form metal hydroxides, which are very alkaline and dissolve

97
Q

What is an alkali?

A

A base that is soluble in water

98
Q

Why is Magnesium Oxide an exception to the other oxides of Group 2 metals?

A

It only reacts slowly and the Hydroxide isn’t very soluble

99
Q

What is the trend down Group 2 for the strength of the alkaline solutions formed?

A

As you go down the group, the oxides form more strongly alkaline solutions

100
Q

Why do the oxides form more strongly alkaline solutions as you go down Group 2?

A

Because the hydroxides get more soluble

101
Q

What is the general formula of the reaction between the Group 2 metal ‘M’ and dilute acids (this case being HCl)?

A

M + 2HCl —-> MCl2 + H2

102
Q

What are many compounds of group 2 metals used for?

A

For neutralising acids

103
Q

What is CaOH used for?

A

In agriculture to neutralise acid soils

104
Q

What are Mg(OH)2 and CaCO3 used for?

A

Used in indigestion tablets (used to neutralise stomach acids)

105
Q

What do halogens exist as?

A

Diatomic molecules

106
Q

What colour is fluorine at room temperature?

A

Pale yellow

107
Q

What colour is chlorine at room temperature?

A

Green

108
Q

What colour is bromine at room temperature?

A

Red-Brown

109
Q

What colour is iodine at room temperature?

A

Grey

110
Q

What state is fluorine at room temperature?

A

Gas

111
Q

What state is chlorine at room temperature?

A

Gas

112
Q

What state is bromine at room temperature?

A

Liquid

113
Q

What state is iodine at room temperature?

A

Solid

114
Q

What is the trend in boiling/melting points you go down Group 7?

A

Increase

115
Q

Why does the boiling/melting point increase as you go down Group 7?

A

Due to the increasing strength of the London forces, as the number of electrons increases when the size and relative mass of the atom increases

116
Q

Why does the states change from gas to solid as you go down the top 4 elements in Group 7?

A

Because of the increasing boiling point

117
Q

What does it mean if a substance is volatile?

A

It has a low boiling point

118
Q

Are halogen atoms oxidised or reduced?

A

Reduced

119
Q

What forms when halogens are reduced?

A

Halide ions

120
Q

Why are halogens oxidising agents?

A

Because when they are reduced, they oxidise another substance

121
Q

What is the trend in atomic radii as you go down Group 7?

A

The atomic radii increases, so the outer electrons become further away from the nucleus

122
Q

What is the trend in reactivity as you go down Group 7?

A

They become less reactive

123
Q

Why does the reactivity decreases as you go down Group 7?

A

The atomic radii increases so the outer electrons get further away from the nucleus. The outer electrons are also shielded more from the attraction of the nucleus, because there are more inner electrons

124
Q

What 2 ions will Chlorine displace?

A

Br- and I-

125
Q

What ion will Bromine displace?

A

I-

126
Q

What can you add to a Group 7 displacement reaction to make the colour change stand out more?

A

Any organic compound such as Hexane

127
Q

What colour is iodine in the presence of an organic compound?

A

A violet/pink

128
Q

What colour is bromine in the presence of an organic compound?

A

Orange/red

129
Q

What colour is chlorine in the presence of an organic compound?

A

Very pale yellow/green

130
Q

What are the 6 steps for testing for halide ions?

A
  • If you’re testing solid substances, dissolve each of them in separate water
  • Use a pipette to add about 3cm^3 of each solution to 3 different test tubes
  • Label each test tube
  • Prepare a table to record your results, leaving room to record what precipitate forms
  • Use a pipette to add a few drops of nitric acid to remove unwanted ions
  • Add a few drops of silver nitrate solution and observe the colour of the precipitate formed to see which ions are present
131
Q

What colour is the precipitate when silver nitrate is added to chloride ions?

A

A white precipitate

132
Q

What colour is the precipitate when silver nitrate is added to bromide ions?

A

A cream precipitate

133
Q

What colour is the precipitate when silver nitrate is added to iodide ions?

A

A yellow precipitate

134
Q

Describe the 5 steps of the further test if you can’t tell which ion is present

A
  • Using a pipette, add dilute ammonia to each of the silver halide solutions
  • Stir the mixture using a glass rod and record whether any precipitates dissolve in the ammonia
  • If they haven’t dissolved, add concentrated ammonia to the solution
  • Stir the mixture using a glass rod and record whether any of the precipitates dissolve
135
Q

What happens to the silver chloride precipitate when dilute ammonia is added?

A

Precipitate dissolves

136
Q

What happens to the silver bromide precipitate when dilute ammonia is added?

A

It doesn’t dissolve

137
Q

What happens to the silver bromide precipitate when concentrated ammonia is added?

A

Precipitate dissolves

138
Q

What happens to the silver iodide precipitate when dilute ammonia or concentrated ammonia is added?

A

It doesn’t dissolve

139
Q

What is disproportionation?

A

When a single element is oxidised and reduced simultaneously

140
Q

The halogens undergo disproportionation when they react with what?

A

Cold dilute alkali solutions such as KOH or NaOH

141
Q

What is the ionic equation of the disproportionation reaction between a halogen (X) and NaOH?

A

X2 + 2OH- —-> XO- + X- + H2O

142
Q

In the reaction X2 + 2OH- —-> XO- + X- + H2O (where X is a halogen) what is the oxidation number of X in XO-?

A

+1

143
Q

What is the full equation of the disproportionation reaction between a halogen (X) and NaOH?

A

X2 + 2NaOH —-> NaXO + NaX + H2O

144
Q

What is Sodium Chlorate (I) used for?

A

Household bleach

145
Q

How do you form Sodium Chlorate (I)?

A

Mix chlorine gas and cold, dilute aqueous sodium chloride

146
Q

What is the full equation for the reaction to make bleach?

A

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

147
Q

What is chlorine’s oxidation state in NaClO?

A

As ClO- is a chlorate ion, the chlorine has a +1 oxidation number

148
Q

What type of reaction is it when you react chlorine and water?

A

A disproportionation reaction

149
Q

What are the 2 products when you react chlorine and water?

A

Hydrochloric acid and Chloric (I) acid

150
Q

Why is chlorine used water treatment?

A

It kills disease causing microorganisms as well as preventing algae growth and removes discolouration

151
Q

What is the main risk of using chlorine in water treatment?

A

Chlorine is very toxic

152
Q

What is the affect on your health if you breath in chlorine?

A

It irritates the respiratory system

153
Q

What are the health risks with liquid chlorine?

A

Liquid chlorine on the skin or in eyes causes severe chemical burns

154
Q

When chlorine reacts with the organic compounds in the water, it forms chlorinated hydrocarbons. What are the health risks with this?

A

The chlorinated hydrocarbons are carcinogenic (cancer causing)

155
Q

Compare the health risks of treated water and untreated water

A

The risk of causing cancer is small compared to a cholera epidemic from untreated water

156
Q

What are 2 alternatives to chlorine for treating water?

A

Ozone

UV light

157
Q

Name an advantage of using ozone in water treatment

A

It’s a strong oxidising agent, meaning it’s great at killing microorganisms

158
Q

What are the 2 disadvantages of using ozone in water treatment?

A

It’s very expensive to produce

It’s short half-life in water means the treatment isn’t permanent

159
Q

How does UV light kill microorganisms in water treatment?

A

It damages their DNA

160
Q

What is the chemical formula of ozone?

A

O3

161
Q

What is a false positive result?

A

A result that suggests an ion is present, when it actually isn’t

162
Q

How do false positive results occur when testing for ions?

A

If there are ions present that will interfere in the test

163
Q

What order should you test for halides, carbonates and sulfates to prevent false positive results?

A

Carbonates —> Sulfates —> Halides

164
Q

How do you test for carbonates?

A

Add a dilute strong acid (e.g. nitric acid) to unknown solution. If carbonates are present, CO2 will be released (which turns limewater cloudy)

165
Q

What is the ionic equation for the reaction between carbonates and acid?

A

CO3 2- + 2H+ —-> CO2 + H2O

166
Q

How do you test for sulfates?

A

Add a few drops of Barium Nitrate (Ba(NO3)2) to a solution of the unknown substance. If you get a white precipitate, then it’s Barium Sulfate so it contains sulfate ions

167
Q

Are sulfates soluble or insoluble in water?

A

Soluble, except for Barium Sulfate

168
Q

What reaction do you have to avoid when testing for sulfates?

A

Sodium Sulfide reacting with Barium Nitrate solution

169
Q

Why do you have to avoid the reaction of Sodium Sulfide reacting with Barium Nitrate solution when testing for sulfates?

A

It forms Barium Sulfite and Sodium Nitrate. Barium Sulfite is also a white precipitate, which will confuse results for test for sulfates

170
Q

When testing for carbonates, why do you add a dilute strong acid?

A

The acid will react with any carbonates and sulfites present, making sure they don’t interfere with the test for sulfates and cause a false positive result

171
Q

What is the ionic equation for the test for sulfates?

A

SO4 - + Ba 2+ —> BaSO4

Sulfate Anion + Barium Cation —> Barium Sulfate

172
Q

Why do you test for halides after sulfates?

A

Because silver sulfate is insoluble. If you tested for halides before sulfates, then any precipitate could be a false result as it could be silver sulfate rather than silver halide

173
Q

How do you test for halides?

A

Add nitric acid, and then silver nitrate solution

174
Q

What will happen in the test if chloride, bromide or iodide are present?

A

A precipitate will form:

Chloride = White
Bromide = Cream
Iodide = Yellow
175
Q

How do you test of ammonium ions?

A

Use a damp piece of red litmus paper. If there’s ammonia present, the paper will turn blue

176
Q

Ammonium Chloride + Sodium Hydroxide —> ?

A

Ammonia + Water + Sodium Chloride

177
Q

When testing for ammonium ions, why does the litmus paper have to be damp?

A

So the ammonium gas can dissolve and make the colour change

A Level Chemistry Year 1 (9 decks)

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