Module 7.1 - Periodicity (3.1.1 - Spec reference)

Question 1

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

Answer 1

Dmitri Mendeleev

Question 2

What did the gaps in the table represent from Mendeleev?

Answer 2

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

Question 3

How many elements were in Mendeleev’s table?

Answer 3

63

Question 4

What did Mendeleev call Germanium before it was discovered?

Answer 4

Eka - silicon

It was next to silicon in the table

Question 5

How did Mendeleev arrange his table?

Answer 5

By Relative atomic mass(Mr)

Question 6

How are elements arranged in the table today?

Answer 6

By increasing atomic number

Question 7

What is a group?

What does it show?

Answer 7

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

Question 8

What is a period?

What does it show?

Answer 8

Horizontal rows in the periodic table

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

Question 9

Periodicity def

Answer 9

A repeating trend in properties of elements

Question 10

Examples of periodicity in elements

Answer 10

Electron Configuration
Ionisation energy
Structure
Melting points

Question 11

What is the trend across a period?

Answer 11

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

Question 12

Trend down a group

Answer 12

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

Question 13

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

Which groups do they represent?

Answer 13

(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

Question 14

Definition of first ionisation energy

Answer 14

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

Question 15

Factors affecting ionisation energy

Answer 15

Atomic radius
Nuclear Charge
Electron Shielding

Question 16

How atomic radius affects ionisation energy

Answer 16

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)

Question 17

How nuclear charge affects ionisation energy

Answer 17

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

Question 18

How electron shielding affects ionisation energy

Answer 18

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

Question 19

Difference in first and second ionisation energies

Answer 19

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

Question 20

Def of second ionisation energy

Answer 20

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

Question 21

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

Answer 21

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

Question 22

What can successive ionisations show?

Answer 22

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

Question 23

Trend in first ionisation energy down a group

Answer 23

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

Question 24

Trend in first ionisation energy across a period

Answer 24

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

Question 25

Sub shell trends in first ionisation energy

Answer 25

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

Question 26

Nitrogen and oxygen sub shell trends

Answer 26

H

Question 27

Structure of metallic bonding

Answer 27

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

Question 28

Def of metallic bond?

Answer 28

The strong electrostatic attraction between cations and delocalised electrons

Question 29

Structure of metals

Answer 29

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

Question 30

Properties of metals

Answer 30

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

Question 31

Why metals have electrical conductivity

Answer 31

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)

Question 32

Why metals have high MP/BP

Answer 32

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

Question 33

Solubility in metals

Answer 33

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

Question 34

Properties of giant covalent structures

Answer 34

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

Question 35

High MP of giant covalent

Answer 35

Contain strong covalent bonds | So more energy needed to break these strong covalent bonds

Question 36

Solubility of giant covalent structures

Answer 36

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

Question 37

Electrical conductivity in giant covalent structures

Answer 37

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

Question 38

Exceptions to conductivity in giant covalent structures

Answer 38

Graphene Graphite Both made of carbon

Question 39

Features of graphene

Answer 39

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

Question 40

Structure of graphite

Answer 40

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

Question 41

Periodic trend in melting points across periods 2 and 3

Answer 41

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

Question 42

What does the sharp decrease in melting points show?

Answer 42

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

Question 43

Difference between giant and simple molecular structures

Answer 43

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

Question 44

Where does the trend in melting points across a period continue

Answer 44

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