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My AS Level Chemistry - OCR A > Module 7.1 - Periodicity (3.1.1 - Spec reference) > Flashcards

Module 7.1 - Periodicity (3.1.1 - Spec reference) Flashcards

1
Q

Who crated one of the first periodic tables with gaps in it?

A

Dmitri Mendeleev

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2
Q

What did the gaps in the table represent from Mendeleev?

A

They represented elements that he predicted would have specific properties
These predictions were based on other elements in the same group

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3
Q

How many elements were in Mendeleev’s table?

A

63

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4
Q

What did Mendeleev call Germanium before it was discovered?

A

Eka - silicon

It was next to silicon in the table

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5
Q

How did Mendeleev arrange his table?

A

By Relative atomic mass(Mr)

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6
Q

How are elements arranged in the table today?

A

By increasing atomic number

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7
Q

What is a group?

What does it show?

A

Vertical columns of elements
Elements in groups have the same number of outer shell electrons
Group number gives the number of outer shell electrons

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8
Q

What is a period?

What does it show?

A

Horizontal rows in the periodic table

Number of the period gives the number of the highest energy electron shell of the element’s atom

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9
Q

Periodicity def

A

A repeating trend in properties of elements

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10
Q

Examples of periodicity in elements

A

Electron Configuration
Ionisation energy
Structure
Melting points

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11
Q

What is the trend across a period?

A

For each period, the s - and p - sub shells are filled in the same way
A periodic pattern

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12
Q

Trend down a group

A

All elements in that group have the same number of outer shell electrons

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13
Q

What are the old and new group numbers now for the table?

Which groups do they represent?

A 
(Old)1 - 1 - alkali metals
(Old)2 - 2 - alkaline earth metals 
Groups 3-12 are transition metals
(Old)7 - 17 - halogens
(Old)8- 18 - Noble gases
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14
Q

Definition of first ionisation energy

A

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

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15
Q

Factors affecting ionisation energy

A

Atomic radius
Nuclear Charge
Electron Shielding

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16
Q

How atomic radius affects ionisation energy

A

Greater the distance between nucleus and outer shell electrons the less the nuclear attraction
(Force of attraction falls off sharply with increasing distance, so atomic radius has a large effect)

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17
Q

How nuclear charge affects ionisation energy

A

More protons in nucleus gives greater attraction between nucleus and outer shell electrons

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18
Q

How electron shielding affects ionisation energy

A

Electrons are negatively charged, so inner shell electrons repel outer shell electrons
This reduces attraction between nucleus and outer shell electrons

19
Q

Difference in first and second ionisation energies

A

Second ionisation energy is larger than first
When one electron is lost, other electrons have a stronger nuclear attraction to unchanged number of protons
Therefore electrons pulled closer to nucleus, so atomic radius decreases
So more attraction to nucleus so more energy needed to remove the remaining electrons

20
Q

Def of second ionisation energy

A

Amount of energy required to remove on electron from each atom of one mole of gaseous 1+ ions to form one mole of gases out 2+ ions

21
Q

What do large spikes in ionisation energies in successive ionisations show?

A

Suggest that electron must have been removed from a new electron shell
Due to atomic radius being smaller, so nuclear attraction increases

22
Q

What can successive ionisations show?

A

Number of electrons in the outer shell
Group of the element in the periodic table
The identity of an element

23
Q

Trend in first ionisation energy down a group

A

Atomic radius increases
More inner shells so shielding increases
Nuclear attraction on outer shells decreases
Therefore first ionisation energy decreases

24
Q

Trend in first ionisation energy across a period

A 
Nuclear charge increases(same no. of protons whilst electrons increases)
Same shell:similar shielding
Nuclear attraction increases
Atomic radius decreases
So first ionisation energy increases
25
Q

Sub shell trends in first ionisation energy

A

First ionisation energy decreases when a new sub shell is filled
E.g. Be and B marks 2p sub shell being filled
2p sub shell has a higher energy than 2s sub shell
So 2p electron is easier to remove than one of 2s electrons in Be
So first I.E in B is less than Be

26
Q

Nitrogen and oxygen sub shell trends

A

H

27
Q

Structure of metallic bonding

A

Each atom donates its negative outer-shell electrons to a shared pool of electrons, which are delocalised(spread out) throughout the whole structure
The cations left behind consist of the nucleus and the inner electron shells of the metal atom

28
Q

Def of metallic bond?

A

The strong electrostatic attraction between cations and delocalised electrons

29
Q

Structure of metals

A

Cations are fixed in position, maintaining the structure and shape of the metal
Delocalised electrons are mobile and able to move throughout the structure - only the electrons move

30
Q

Properties of metals

A

Strobe metallic bonds(attraction between positive ions and delocalised electrons
High electrical conductivity
High MP/BP

31
Q

Why metals have electrical conductivity

A

They conduct electricity in solid and liquid states
When voltage applied across a metal, delocalised electrons can move through structure carrying a charge
(In comparison, ionic compounds have no mobile charge carriers)

32
Q

Why metals have high MP/BP

A

Strong electrostatic attraction between the cations and sea of delocalised electrons
So more energy needed to break these forces

33
Q

Solubility in metals

A

Metals do not dissolve
Any interactions lead to a reaction, rather than dissolving
E.g. sodium and water

34
Q

Properties of giant covalent structures

A

High MP/BP
Insoluble in most solvents
Don’t conduct electricity

35
Q

High MP of giant covalent

A

Contain strong covalent bonds

So more energy needed to break these strong covalent bonds

36
Q

Solubility of giant covalent structures

A

Insoluble in most solvents

Covalent bonds holding atoms together in lattice too strong to be broken by interactions with solvents

37
Q

Electrical conductivity in giant covalent structures

A

Don’t conduct electricity
E.g. in diamonds, all four outer shell electrons involved in covalent bonding, so none are available for conducting electricity

38
Q

Exceptions to conductivity in giant covalent structures

A

Graphene
Graphite
Both made of carbon

39
Q

Features of graphene

A

A single layer of graphite
Shape - hexagonal planar layer
Bond angles - 120 - by electron-pair repulsion
Composed of hexagonally arranged carbon atoms linked by strong covalent bonds
Very strong material
Same electrical conductivity as copper

40
Q

Structure of graphite

A

Bond angles - 120
Compose of layers of hexagonally arranged carbon atoms, like a stack of graphene layers
Layers bonded by weak London Forces
Bonding in the hexagonal layer only uses three of carbon’s four outer shell electrons.
Spare electron is delocalised between layers
This means electricity can be conducted as in metals as has a mobile charge carrier

41
Q

Periodic trend in melting points across periods 2 and 3

A

Melting point increases from group 1 to group 4
Sharp decreases in melting point between group 4 and group
Melting pints are comparatively low from group 5 to group 0

42
Q

What does the sharp decrease in melting points show?

A

Marks change from giant to simple molecular structures. Shows start of divide between metals and non-metals

43
Q

Difference between giant and simple molecular structures

A

Giant covalent structures have strong forces to overcome so have high melting points.
Simple molecular structures have weak forces to overcome, so have much lower melting points

44
Q

Where does the trend in melting points across a period continue

A

Across period 2 it is repeated across period 3, and continues across the s- and p- blocks from period 4 downwards

My AS Level Chemistry - OCR A (14 decks)

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