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A Level Chemistry > Module 3 Section 1: The Periodic Table > Flashcards

Module 3 Section 1: The Periodic Table Flashcards

1
Q

How were chemical categorised in the early 1800s

A

By their physical and chemical properties and by their relative atomic mass

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2
Q

What did Johann Döbereiner do

A

Attempted to group similar elements - called Döbereiner’s triads
Noticed chlorine, bromine and iodine had similar properties and that properties of bromine fell halfway between chlorine and iodine
This was the same for lithium, sodium and potassium
Called these triads

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3
Q

What did John Newlands do

A

Noticed that if he arranged the elements in order of mass, similar elements appeared at regular intervals
Every eighth element was similar
Called this the law of octaves
However, this broke down on the third row

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4
Q

What did Dmitri Mendeleev do

A

Arranged all known elements by atomic mass
Left gaps in the table where the next element didn’t seem to fit
This meant he could keep elements with similar chemical properties in the same group
Also predicted the properties of undiscovered elements that would go in the gaps
When elements were later discovered that had properties that fitted with Mendeleev’s predictions (germanium, scandium and gallium) it showed that he had got it correct

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5
Q

How are elements ordered in the present day periodic table

A

Elements are arranged by increasing atomic number ( proton number )
They are arranged into groups and periods

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6
Q

How are elements arranged in periods

A

All elements within a period have the same number of electron shells
This means that there are repeating trends in the physical and chemical properties of elements across each period

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7
Q

What is periodicity

A

The repeating trend in the properties of elements

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8
Q

How are elements arranged in groups

A

All elements within a group have the same number of electrons in their outer shell
This means they have similar chemical properties
Group number tells you the number of electrons in the outer shell

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9
Q

How to use the periodic table for electron configuration

A

The periodic table can be split into an s-block, d-block, p-block which can show which sub-shells all the electrons go into

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10
Q

What is the first ionisation energy

A

The first ionisation energy is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms

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11
Q

What type of process is ionisation

A

Energy must be put in so it’s an endothermic process

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12
Q

Equation for first ionisation energy

A

O (g) -> O+ (g) + e-
1st ionisation energy = +1314 kJ mol-1

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13
Q

Important notes on ionisation energy

A

You must use the gas state symbol, (g), because ionisation energies are measured for gaseous atoms
Always refer to 1 mole of atoms, as stated by the definition, rather than to a single atom
The lower the ionisation energy, the easier it is to form an ion

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14
Q

What factors effect ionisation energies

A

Nuclear charge
Atomic radius
Electron shielding

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15
Q

How does nuclear charge effect ionisation energy

A

The more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons
Increases ionisation energy

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16
Q

How does atomic radius affect ionisation energy

A

Attraction falls off very rapidly with distance
An electron close to the nucleus will be much more strongly attracted than one further away
Decreases ionisation energy

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17
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 towards the nuclear charge
This lessening of the pull of the nucleus by inner shells of electrons blocking the attraction is called shielding
Decreases ionisation energy

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18
Q

What does a high ionisation energy mean

A

Means that there’s a strong attraction between the electron and the nucleus, so more energy is needed to overcome the attraction and remove the electron

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19
Q

What happens to ionisation energy as you go down a group

A

The ionisation energies generally fall
This means it gets easier to remove outer electrins
This is because:
Elements further down a group have extra electron shells compared to one’s above
The extra shells man that the atomic radius is larger, so the outer electrons are further away from the nucleus, which greatly reduces their attraction to the nucleus

The extra inner shells shield the outer electrons from the attraction of the nucleus
(The positive charge of the nucleus does increase as you go down a group (due to extra protons), but this effect is overridden by the effect of the extra shells)

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20
Q

Why does ionisation energy decreasing as you go down a group provide evidence that electron shells exist

A

A decrease in ionisation energy going down a group supports the Bohr model of the atom

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21
Q

What happens to ionisation energies as you move across a period

A

The ionisation energies increases
This means it gets harder to remove the outer electrons
This is because the nuclear charge increases and atomic radius decreases as the electrons are pulled closer to the nucleus
The extra electrons that elements gain across a period are added to the outer energy level so they don’t provide any extra shielding effect

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22
Q

Why does the first ionisation energy decrease between groups 2 and 3

A

This is due to sub shell structure
The outer electron in group 3 elements is in a p orbital rather than s orbital
A p orbital has a slightly higher energy than an s orbital in the same shell, so the electrons is, on average, to be found further from the s electrons
These factors override the effect of the increased nuclear charge, resulting in the ionisation energy dropping slightly

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23
Q

Why does ionisation energy drop between groups 5 and 6

A

p orbital structure
In the group 5 elements, the electron is being removed from a singly-occupied orbital
In the group 6 elements, the electron is being removed from an orbital containing two electrons
The repulsion between two electrons in an orbital means that electrons are easier to remove from shared orbitals

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24
Q

How remove all electrons from an atom

A

This would be done using successive ionisation energies
Each time you remove an electron, there’s a successive ionisation energy

O (g) -> O+ (g) + e-: first ionisation energy
O+ (g) -> O2+ (g) + e-: second ionisation energy

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25
Q

Why does a graph showing successive ionisation energy show the ionisation energy as increasing

A

Within each shell, successive ionisation energies increase
This is because electrons are being removed from an increasingly positive ion
Also less repulsion amongst the remaining electrons
So the electrons are held more strongly by the nucleus

The big jumps in ionisation energy happen when a new shell is broken into
This is where an electron is being removed from a shell closer to the nucleus

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26
Q

How can successive ionisation graphs tell you what group the element is in

A

Count how many electrons are removed before the first big jump to find the group number

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27
Q

How can successive ionisation graphs tell you the electronic structure of an element

A

Working from right to left, count how many points there are before each big jump to find how many electrons are in each shell, starting with the first

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28
Q

What are giant covalent lattices

A

Huge networks of covalently bonded atoms
(Sometimes called macromolecular structures too)
Carbon atoms can form this type of structure because they can each form four strong, covalent bond

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29
Q

Examples of giant covalent lattices

A

Diamond
Graphite
Graphene

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30
Q

What are allotropes

A

Different forms of the same element in the same state

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31
Q

Structure of diamond

A

Each carbon atom is covalently bonded to four other carbon atoms
The atoms arrange themselves in a tetrahedral shape
It’s got a crystal lattice structure

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32
Q

Properties of diamond

A

Lots of strong covalent bonds
Very high melting point (sublimes at over 3800K)
Very hard - can be used in diamond tipped drills and saws
Vibrations travel easily through the lattice, so it’s a good thermal conductor
Can’t conduct electricity - all outer electrons are held in localised bonds
Won’t dissolve in any solvent

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33
Q

Structure of silicon

A

Silicon (same group as carbon)
Forms crystal lattice structure with similar properties to carbon
Each silicon can form 4 strong covalent bonds

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34
Q

Structure of graphite

A

Weak forces between layers of graphite are easily broken, so the sheets can slide over eachother, can be used in lubricants

The delocalised electrons in graphite aren’t attached to any particular carbon and are free to move along the sheets, so an electric current can flow

Layers are quite far apart compared to the length of the covalent bonds, so graphite is less dense than diamond and is used to make strong, lightweight sports equipment

Strong covalent bonds in the hexagon sheets, graphite has a very high melting point

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35
Q

Properties of graphite

A

Can be used as a lubricant
Conducts electricity
Relatively low density
High melting point (sublimes at over 3900K)
Graphite is insoluble in any solvent, covalent bonds in the sheets are too strong to break

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36
Q

What is graphene

A

A sheet of carbon atoms joined together in hexagons
Sheet is one atom thick, making it a 2D compound

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37
Q

Properties of graphene

A

The delocalised electrons are free to move along the sheet
Without layers, they can move quickly above and below the sheet, making graphene very electrically conductive
The delocalised electrons strengthen the covalent bonds between the carbon atom which makes graphene strong
Single layer of graphene is transparent and very light
High melting and boiling point
Insoluble due to strong covalent bonds

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38
Q

Uses of graphene

A

Can be used in high-speed electronics and aircraft technology
This is due to its high strength, low mass and good electrical conductivity
Flexibility and transparency also allows it to be used for touchscreens

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39
Q

Structure of metals

A

Giant metallic lattice structures
Electrons in outer shell of metal atoms are delocalised- electrons are free to move around the metal
This leaves a positively metal cation
Metal cations are electrostatically attracted to the delocalised negative electrons
They form a lattice of closely packed cations in a sea of delocalised electrons - this is metallic bonded

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40
Q

Properties of metals

A

Number of delocalised electrons per atom affects the melting point
The more there are, the stronger the bonding will be and the high the melting point
So Mg2+ has a higher melting point than Na+

The size of the metal ion and lattice structure also affects the melting point
A smaller ionic radius will hold the delocalised electrons close to the nuclei

As there are no bonds holding specific ions together, the metal ions can slide past eachother when the structure is pulled, so metals are malleable and ductile

Delocalised electrons can pass kinetic energy to eachother, making metals good thermal conductors

Metals are good electrical conductors because the delocalised electrons can move and carry a charge,

Metals are insoluble, except in liquid metals, because of the strength of the metallic bond

41
Q

What are simple molecules

A

Simple molecular structures contain only a few atoms
E.g. oxygen (O2), chlorine (Cl2), phosphorus (P4)

42
Q

Why do simple molecular structures have low melting and boiling points

A

The covalent bonds between the atoms in the molecules are very strong
The melting and boiling points of simple molecular substances depend upon the strength of the induced dipole-dipole forces between their molecules
These intermolecular forces are weak and easily overcome, so these elements have low melting and boiling points

43
Q

Why do larger molecules have high melting and boiling points

A

More atoms in a molecule mean stronger induced dipole-dipole forces
E.g. in period 3 sulfur is the biggest molecule (S8), so it’s got high melting and boiling points than phosphorus or chlorine

44
Q

Why do noble gases have very low melting and boiling points

A

They exist as individual atoms (monatomic), resulting in very weak induced dipole-dipole forces

45
Q

Why does melting and boiling points change as you across a period

A

As you go across a period, the type of bond formed between the atoms of an element changes
This affect the melting and boiling points of the element

46
Q

What does the graph showing melting and boiling points across a period show

A

Melting and boiling points increase from the metals to the group 4 elements (strong metallic or covalent bonds)
Sharp decrease into group 5 (simple molecules)
Stays low up to group 0 (monatomic)

47
Q

What happens to the melting and boiling point of metal as you go across the period (according to the graph)

A

Melting and boiling points increase across the period
The metallic bonds get stronger as the ionic radius decreases and the number of delocalised electrons increase
Resulting in stronger metallic bonds

48
Q

What happens to the melting and boiling point of group 4 elements as you go across the period (according to the graph)

A

The elements with giants covalent lattice structures (C and Si) have strong covalent bonds linking all their atoms together
A lot of energy is needed to break these bonds

49
Q

What happens to the melting and boiling point of simple molecules as you go across the period (according to the graph)

A

Elements that form simple molecular structures have only weak intermolecular forces to overcome between their molecules
They have low melting and boiling points

50
Q

What happens to the melting and boiling point of noble gases as you go across the period (according to the graph)

A

Noble gases (neon and argon) have the lowest melting and boiling points in their periods
They are held together by the weakest forces

51
Q

Properties of ionic compounds, give examples

A

E.g. NaCl, MgCl2
High Mp and Bp
Solid at STP (standard temperature and pressure)
Does not conduct electricity as a solid - ions are held firmly in place
Conducts electricity as a liquid - ions are free to move
Soluble in water

52
Q

Properties of simple molecular (covalent) substances, give examples

A

E.g. CO2, I2, H2O
Low Mp and Bp - have to overcome induced dipole-dipole forces or hydrogen bonds, not covalent bonds
Sometimes solid, usually liquid or gas at STP (water is liquid as it has hydrogen bonds)
Does not conduct electricity as a solid
Does not conduct electricity as a liquid
Solubility depends on how polar the molecule is

53
Q

Properties of giant covalent lattices, give examples

A

E.g. diamond, graphite, graphene
High Mp and Bp
Solid at STP
Does not conduct electricity as a solid (except graphite and graphene)
Will generally sublime as a liquid
Insoluble in water

54
Q

Properties of metallic substances, give examples

A

E.g. Fe, Mg, Al
High Mp and Bp
Solid at RTP
Conducts electricity as a solid - delocalised electrons
Conducts electricity as a liquid - delocalised electrons
Insoluble in water

55
Q

What ions do group 2 elements make

A

All have 2 electrons in their outer shell (s2)
Lose their two outer electrons to form 2+ ions
Ions have a structure of a noble gas

56
Q

Group 2 elements and the electronic structure of the ion they make

A

Be atom: 1s2 2s2
Be ion: 1s2

Mg atom: 1s2 2s2 2p6 3s2
Mg ion: 1s2 2s2 2p6

Ca atom: 1s2 2s2 2p6 3s2 3p6 4s2
Ca ion: 1s2 2s2 2p6 3s2 3p6

57
Q

What happens to reactivity as you go down group 2

A

Reactivity increases
The ionisation energies decrease due to increasing atomic radius and shielding effect
Group 2 elements lose electrons when they react
The easier it is to lose electrons (I.e. the lower the first and second ionisation energy), the more reactive the element, so reactivity increases down the group

58
Q

What happens when group 2 elements react

A

They are oxidised from a state of 0 to +2, forming M2+ ions

M -> M2+ + 2e- (oxidation number goes from 0 to +2)

E.g. Ca -> Ca2+ + 2e-

59
Q

What do group 2 metals produce when they react with water

A

Group 2 elements reacts with water to form an alkaline hydroxide, with the general formula M(OH)2, and hydrogen gas

E.g. M + 2H2O = M(OH)2 + H2
Oxidation number M: 0 -> +2 (ox)
Oxidation number H: +1 -> 0 (red)

Not all hydrogen atoms are reduced

60
Q

What group 2 atom doesn’t react with water

A

Beryllium
Magnesium reacts slowly with cold water and faster with steam

61
Q

What do group 2 metals produce when they react with dilute acid

A

Metal react with acids to produce a salt and hydrogen
E.g. when a metal reacts with hydrochloric acid, you get a metal chloride and hydrogen

M + 2HCl -> MCl + H2

Oxidation number: M 0 -> +2 (ox)
Oxidation number: H -1 -> +1 (red)

62
Q

What do group 2 metals produce when they react with oxygen

A

Group 2 elements react with oxygen to form a white metal oxide with the general formula MO, made up of M2+ and O2- ions

2M + O2 -> MO

Oxidation number M: 0 -> +2 (ox)
Oxidation number O: 0 -> -2 (red)

63
Q

Formula, colour and physical state at RTP for halogens

A

Fluorine: F2, pale yellow, gas
Chlorine: Cl2, green, gas
Bromine: Br2, red brown, liquid
Iodine: I2, grey, solid

64
Q

What molecules do halogens exist as

A

Diatomic molecules joined by a single covalent bond

65
Q

Trend in melting and boiling point of halogens down the group

A

Bp and mp increase down the group
Due to increasing strength of induced dipole-dipole forces
Size and relative mass of atoms also increases
Can be seen as changes to physical state from Cl and I and volatility decreases down the group

66
Q

How do halogens react

A

Gain an electron in their outer shell
Form 1- ions
They are reduced and they oxidise another substance - so they’re oxidising agents

67
Q

What happens to reactivity as you go down group 7 and why does this happen

A

Reactivity decreases
Atomic radii increase - electrons further from nucleus
Outer electrons are more shielded
Overrides increase in charge from nucleus
Makes it harder for larger atoms to attract the electron needed to fill the outer shell

68
Q

How can you see the relative oxidising strengths in halogens

A

Can be seen in displacement reactions with other halide ions

69
Q

What would happen if you put bromine in a potassium iodide solution

A

Bromine displaces the iodide ions (oxidises) giving potassium bromide and iodine

70
Q

What happens when another halogen displaces a halide

A

The colour of the solution changes

71
Q

What colour does the solution go when bromine is aqueous and displaced

A

Yellow in aqueous solution
Orange in organic solution

72
Q

What colour does the solution go when iodine is displaced

A

Orange/ brown in aqueous solution
Purple in organic solution

73
Q

What colour is chlorine water

A

Colourless

74
Q

What can halogen displacement reactions be used for

A

Can be used to identify which halogen or halide is present in a solution

75
Q

Group 7s and what other halides they displace

A

Chlorine: bromide and iodide
Bromine: iodide
Iodine: no reaction with any

76
Q

How to test for halide ions

A

Add dilute nitric acid: removes ions that may interfere with the test
Add silver nitrate solution
A precipitate is formed of silver halide
To make sure add ammonia solution to see how easily the silver halide dissolves (larger ions are more difficult to dissolve)

77
Q

Test results for halide ion test

A

Chloride: white ppt, dissolves in dilute NH3
Bromide: cream ppt, dissolves in conc NH3
Iodide: yellow ppt, insoluble in conc NH3

78
Q

What is the electron configuration of halogens

A

Always ends in -s2 p5

78
Q

How to make the colour change of a halogen displacement reaction more visible

A

This can be made easier to see by adding an organic solvent like hexane
The halogen present will dissolve readily in the solvent which settles as a distinct layer above the aqueous solution

79
Q

Example of a disproportionation reaction involving halogens in real life

A

Cl2(aq) + 2NaOH(aq) = NaClO(aq) + NaCl(aq) + H20
Chlorine molecule has oxidation number of 0
Then forms molecules with Cl with an oxidation number of +1 and -1
NaClO is used as household bleach

80
Q

What is the product when chlorine is mixed with sodium hydroxide

A

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

81
Q

When do halogens undergo disproportionation

A

When halogens react with cold dilute alkali solutions they undergo disproportionation

82
Q

General equation for halogens and cold dilute alkali

A

X2 + 2NaOH -> NaXO + NaX + H2O
Oxidation number of X:
Reactants: 0
NaXO: +1
NaX: -1

83
Q

Chemical name for bleach

A

NaClO: sodium chlorate

84
Q

How is chlorine used to kill bacteria in water

A

When chlorine is mixed with water it undergoes disproportionation to make a mixture of hydrochloric acid and chloric acid
Aqueous chloric acid ionises to make chlorate(I) ions (ClO-)
Chlorate(I) ions kill bacteria so adding chlorine to water can make it safe to drink or swim in

85
Q

How is chlorine used to clean water in the UK

A

Chlorine kills disease-causing microorganisms
Some chlorine remains in the water and prevents re infection further the supply
Prevents the growth of algae, eliminating bad tastes and smells, and removes discolouration caused by organic compounds

86
Q

Risks of chlorine to treat water

A

Chlorine gas is harmful if it’s breathed in
Irritates the respiratory system and liquid chlorine on the skin or eyes causes severe chemical burns

Untreated water contains organic compounds (e.g. from decomposing plants)
Chlorine reacts with these to form chlorinated hydrocarbons which can be carcinogenic

However, the risk of untreated water outweighs this e.g. a cholera epidemic

87
Q

Ethical concerns of chlorine water treatment

A

We do not get a choice about having our water chlorinated
Some called this forced mass medication

88
Q

Alternatives to chlorine

A

Ozone (O3) - strong oxidising agent, which makes it good at killing micro
It is expensive to produce and it’s short half life in water means it isn’t permanent

UV light - kills microorganisms by damaging their DNA, but it is ineffective in cloudy water and it won’t stop the water being contaminated further down the line (like O3)

89
Q

How do the oxides and hydroxides of group 2 metals react

A

Oxides of group 2 metals react steadily with water to form metal hydroxides
These dissolve in water
These solutions are alkaline due to the OH- ions

90
Q

How is magnesium oxide an exception to the trend in group 2 oxides

A

Magnesium oxide is an exception as it only reacts slowly
The hydroxide isn’t very soluble

91
Q

Trend with group 2 oxide reactions

A

The oxides form more strongly alkaline solutions as you go down the group
This is because the hydroxides get more soluble as atomic radius increases so ionic bonds are weaker

92
Q

How can Group 2 metals be used in real life

A

Group 2 (alkaline earth metals) are used commonly to neutralise acids e.g:
Calcium hydroxide is used to neutralise acidic soils in agriculture
Magnesium hydroxide and calcium carbonate are used as antacids in indigestion tablets

93
Q

How to test for carbonates

A

Add dilute HCl to the sample and CO2 will be released
CO3 2- + 2H+ -> CO2 + H2O
Can use limewater which turns cloudy when CO2 is bubbled through it

94
Q

How to test for sulfates

A

Add dilute HCl and BaCl2
White ppt will appear
Ba2+ (aq) + SO4 2- (aq) -> BaSO4 (s)

95
Q

How to test for halides

A

Add nitric acid and silver nitrate
Silver chloride: white ppt, dissolves in dilute NH3
Silver bromide: cream ppt, dissolves in conc NH3
Silver iodide: yellow ppt, insoluble in conc NH3

96
Q

How test for ammonium compounds
Include an equation

A

Add NaOH to the sample and warm the mixture
Use damp red litmus paper, it will turn blue if ammonia is present
NH4+ (aq) + OH- (aq) -> NH3 (g) + H2O (l)

97
Q

What order to do tests

A
  1. Carbonate, if no CO2 ->
  2. Test for sulfate, if no ppt ->
  3. Test for halides
98
Q

What colour are group 2 compounds in solution

A

Group 2 compounds are colourless

A Level Chemistry (15 decks)

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