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A Level Chemistry Ocr > Module 3 Chapter 7 > Flashcards

Module 3 Chapter 7 Flashcards

1
Q

What did Dmitri Mendeleev do

A

He is one of the main chemists credited with the development of the periodic table

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2
Q

How did Mendeleev organise the periodic table

A

He organised elements by atomic mass
He placed elements into groups of which they shared similar chemical properties with
He left gaps where no known elements fitted with the chemical properties of the group

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3
Q

What were the properties of the periodic table in the past

A

60 elements ordered by atomic mass
groups with similar properties
Properties had to fit, elements swapped and gaps left
Assumed incorrect atomic masses and some not discovered
Predicted properties of missing elements using group trends

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4
Q

What are the properties of the periodic table now

A

118 elements
Vertical groups and horizontal periods
Elements arranged in increasing atomic number

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5
Q

How are elements arranged in the periodic table

A

They are arranged in increasing atomic number

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6
Q

What are the properties of groups in the periodic table

A

In each group the elements have atoms with the same number of electrons in the outer shell and similar properties.

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7
Q

What does the period show

A

The number of the period gives the highest energy shell in an elements atom

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8
Q

What is periodicity

A

The repeating trend of properties across the period

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9
Q

What determines the chemistry of an element

A

A lot of the chemistry is determined by electron configuration, particularly the highest energy outer shell.

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10
Q

What are some trends in periodicity

A

Ionisation energy
Structure
Melting points

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11
Q

What is the trend of the highest energy sub shell where the electrons are held as you go across period 2, 3 and 4

A

Across period 2: 2S —> 2P
Across period 3: 3S —> 3p
Across period 4: 4S —> 3D —> 4P

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12
Q

How many groups are in the periodic table

A

There are 18 groups in the periodic table

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13
Q

What is Group 15 called

A

The Pnictogens

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14
Q

What is group 16 called

A

The Chalogens

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15
Q

What is group 1 called

A

The alkali metals

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16
Q

What is group 2 called

A

The alkaline earth metals

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17
Q

What are groups 3-12 called

A

The transition metals

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18
Q

what is group 17 called

A

The halogens

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19
Q

What is group 18 called

A

The noble gases

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20
Q

What is the trend of reactivity as you go down Group 1

A

The reactivity increases as you go down the group

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21
Q

Why does the reactivity increase as you go down group one

A

The atomic radius increases, so the force of attraction between the nucleus and electrons in the outer shell is weaker, therefore it looses its electrons to form ions more easily and is therefore more reactive.

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22
Q

What does ionisation energy measure

A

It measures how easily and atom can lose electrons to form positive ions

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23
Q

What is first ionisation energy

A

The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.

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24
Q

What are the three factors affecting first ionisation energy

A
  1. Atomic radius
  2. Nuclear charge
  3. Electron shielding
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25
Q

Where is the first electron lost in an atom

A

It is lost from the highest energy level and therefore experiences the least attraction from the nucleus.

26
Q

How does atomic radius affect ionisation energy

A

The greater the distance from the nucleus and outer shell the less nuclear attraction from the nucleus. As the force of attraction decreases with increasing distance.

27
Q

How does nuclear charge affect first ionisation energy

A

The more protons there aer in the nucleus of an atom the greater the attraction between the nucleus and outer electrons.

28
Q

How does electron shielding affect ionisation energy

A

They are negatively charged, the inner-shell electrons repel the outer shell electrons. This repulsion leads to a shielding effect reducing the attraction.

29
Q

Which of the three factors has the greatest affect on ionisation energy

A

Atomic radius

30
Q

What is second ionisation energy

A

The energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.

31
Q

Why is the second ionisation energy always greater than the first

A

Because more energy is required to remove the second electron as the nuclear charge increases as there are less electrons attracted to the same number of protons.

32
Q

What is a key trend in the periodic table

A

The change from metals to non-metals from left to right across each period.

33
Q

Where does the change from metals to non-metals occur

A

On a diagonal line from the top of group 13 to the bottom of group 17.

34
Q

What do elements close to the diagonal line show

A

Show in between properties and are called semi metals or metalloids.

35
Q

What type of bonding is present in metals

A

Metallic bonding

36
Q

What is metallic bonding

A

it is a unique type of bonding where the outer shell electrons leave their shell to form a lattice of metal cations surrounded by delocalised electrons. Where the metallic bond is between the delocalised electrons and positively charged metal cations.

37
Q

What is a property of the delocalised electrons

A

They are mobile and free to move throughout the entirety of the solid metal.

38
Q

What are the properties of metals

A

High electrical conductivity
High melting and boiling points
Insoluble in water

39
Q

Why are metals insoluble in water

A

There would be a reaction instead of dissolving so therefore they are insoluble in water.

40
Q

Why do metals have high melting points

A

High temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic forces of attraction between the delocalised electrons and the cations. As a result they have a high melting and boiling points.

41
Q

What are the forces present in simple covalent structures

A

They have strong covalent bonds between atoms within a molecule. However, they have very weak intermolecular forces between individual molecules.

42
Q

Can simple covalent molecules conduct electricity and why

A

They can’t conduct electricity as they don’t have any free to move charged particles, so a charge can’t be carried and a current can’ flow.

43
Q

What is the melting and boiling points of simple covalent molecules

A

They have a low melting and boiling points as there are weak intermolecular forces between molecules, so have a low melting and boiling points.

44
Q

What are the physical properties of simple covalent molecules

A

Soft and waxy in the solid state
Brittle
Depends on polarity whether it is soluble or not.

45
Q

What are properties of giant covalent structures

A

They have high melting and boiling points as they have lots of strong covalent bonds that require lots of energy to break.
Insoluble in almost all solvents
Don’t conduct electricity except graphite and graphene

46
Q

Why are giant covalent structures insoluble in all solvents

A

The covalent bonds holding atoms together in the lattice are too strong to be broken by interactions with solvents.

47
Q

What are the key properties of diamond

A

It is very hard
Has a high, melting and boiling points
Can’t conduct electricity
Not soluble in water

48
Q

What is the bonding present in diamond

A

Each carbon atom is bonded to 4 other carbon atoms.
This forms a tetrahedral arrangement.
And as there are lots of these strong covalent bonds between carbons they are very hard and have a high melting and boiling point.
It has no free to move charged particles so can’t conduct electricity.
Not polar, so therefore aren’t soluble in water.

49
Q

What are the key properties of graphite and graphene

A

Is a lubricant
Has delocalised electrons so can conduct electricity
Is soft

50
Q

What is the bonding in graphite and graphene

A

Carbon forms giant covalent structures based on planar hexagonal layers
Each carbon is bonded to 3 carbon atoms leaving 1 delocalised electron.
This results in it having a bound angle of 120* due to electron pair repulsion.
The delocalised electron allows it to conduct electricity
It has weak intermolecular forces between layers, allowing for layers to slide over each other making it a lubricant.

51
Q

What is graphene

A

It is one layer of graphite

52
Q

What is the key property of SiO2

A

Transparent and brittle

53
Q

What is the bonding in SiO2

A

Each oxygen is bonded to 2 silicons and each silicon is bonded to 4 oxygen atoms.

54
Q

What is the trend in melting points across period 2 and 3

A

The melting point increases from group 1 to group 14
There is a sharp decrease in melting points between groups 14&15
Melting points are relatively low from group 15 to 18

55
Q

What is the sharp decrease in melting points due to

A

A change from giant to simple molecular structures.

56
Q

What is the general trend in ionisation energy with ionisation number

A

Ionisation energy increases with ionisation number. However, there will be a big increase in ionisation energy when there is a decrease in the principal quantum number.

57
Q

Why does the ionisation energy fall from nitrogen to oxygen

A

There is a paired electron in one of the 2P orbitals in oxygen whereas there isn’t in nitrogen. These electrons repel each other. Making it easier to remove an electron from an oxygen atom than a nitrogen atom. Therefore the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen.

58
Q

What is the trend in first ionisation energies

A

They decrease down a group

59
Q

What is the general trend in ionisation energies across a period

A

There is a general increase in first ionisation energies across period 2 and 3

60
Q

What are the 2 exceptions to this general trend

A

A decrease in ionisation energy from beryllium to boron
A decrease in ionisation energy from nitrogen to oxygen.

61
Q

Why is there a dip in ionisation energy between beryllium and boron

A

It is because boron’s highest energy level electron in in 2P whereas beryllium’s highest energy level electron is 2S. The 2P subshell is at a higher energy level so therefore it is easier to remove than boron’s electron in 2S, so as a result the first ionisation energy of boron is less than the first ionisation energy of beryllium.

A Level Chemistry Ocr (12 decks)

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