3.1.1 Periodicity

Question 1

Define group

Answer 1

columns of elements with the same number of electrons in their outer shells
- similar chemical properties

Question 2

Define period

Answer 2

horizontal row of elements with the same number of electron shells
- Repeating trends in physical and chemical properties

Question 3

Define periodicity

Answer 3

repeating trends in properties

Question 4

What do the blocks of the periodic table tell us ?

Answer 4

the sub-shell with the highest energy electrons
- the WIDTH is the same as the number of electrons in the outer sub-shell

Question 5

Define ionisation energy

Answer 5

the energy required to remove one electron from each atom in a mole of gaseous atoms producing one mole of gaseous ions

X (g) -> X+ (g) + e-

Question 6

Define successive ionisation energy

Answer 6

removing additional electrons from each ion in a mole of gaseous ions

Question 7

What happens to the ionisation energy as more electrons are lost ?

Answer 7

• the protons remain constant
• so there is a greater proton to electron ratio
• so there is increases nuclear attraction between the nucleus and outer electrons
• increasing the ionisation energy

Question 8

What three factors affect ionisation energy ?

Answer 8

1. Nuclear charge (number of protons)
2. Atomic / ionic radius
3. Extent of shielding in lower energy shells

Question 9

How does ionisation energy change down the group ?

Answer 9

• DECREASES
• Down the group electrons are added to a new shell further from the nucleus so atomic radius increases
• Shielding (of electrons in lower energy shells) increases
• Nuclear charge increases
• Increased shielding outweighs increased nuclear charge
• so there is weaker attraction between the nucleus and the outermost electron
• less energy needed to remove it which lowers the ionisation energy

Question 10

How does ionisation energy change across the period ?

Answer 10

• INCREASES
• Nuclear charge increases
• Atomic radius decreases (due to the increased nuclear charge)
• Shielding is constant (as electrons are added to the outer shells)
• increased attraction between outermost electron and nucleus
• so more energy is needed to remove it which increases the first ionisation energy

Question 11

Explain the ionisation energy dip between magnesium-aluminum and berrylium-boron

Answer 11

• outer electron in the 3p orbital of aluminum is further away from the nucleus than the two electrons in the 3s orbital of magnesium
• so there is less nuclear attraction
• so less energy is required to remove it
• so aluminum has a lower ionisation energy than magnesium

Question 12

Explain the ionisation energy dip between phosphorus-sulfur and nitrogen-oxygen

Answer 12

• sulfur has an extra electron so there is electron-electron repulsion in the 3p orbital of sulfur
• so the electron requires less energy to remove
• so sulfur has a lower first ionisation energy

Question 13

How can you establish the number of shells from an ionisation energy graph ?

Answer 13

number of jumps shows the gaps between shells (so three jumps = four shells)

Question 14

How can you establish the number of outer electrons in an ionisation energy graph ?

Answer 14

count how many electrons are removed before the first big jump (This also tells us the group number)

Question 15

Define metallic bond
What structure do metals have ?

Answer 15

the strong electrostatic attraction between a lattice of positive ions and a sea of delocalised electrons (delocalised electrons are in the outermost shell)
- Giant metallic lattice

Question 16

Describe and explain electrical conductivity in metals

Answer 16

• good electrical conductors
• as delocalised electrons (in the outer shell) can move and carry charge

Question 17

Describe and explain solubility in metals

Answer 17

• insoluble
• due to the strong metallic bonds

Question 18

Describe and explain thermal conductivity in metals

Answer 18

• good thermal conductors
• as delocalised electrons can move and carry kinetic energy

Question 19

Describe and explain the trend in melting point in metals

Answer 19

• more **delocalised electrons **(higher charge)
• stronger metallic bonding
• higher melting point

Size of metal ion and lattice structure also affect melting point - smaller radius = more attraction so higher melting point

Question 20

Define giant covalent lattice

Answer 20

huge network of covalently bonded atoms
- very high melting and boiling points as a lot of energy is needed to break the strong covalent bonds

Question 21

Describe the structure and properties of diamond
- what other element forms a very similar structure to diamond ?

What is the shape ?

Answer 21

Giant covalent lattice (each carbon is covalently bonded to four other carbon atoms)
* very high melting point
* very hard
* good thermal conductor due to the stiff lattice
* non-conductive as all four outer electrons are involved in covalent bonding
* will not dissolve (strong covalent bonds between carbon)
SILICON IS SIMILAR

tetrahedral

Question 22

Describe the structure and properties of graphite

Answer 22

Giant covalent lattice (carbon atoms in flat hexagonal layers** covalently bonded** to three others - BUT weak intermolecular forces between layers - induced dipole dipole)
* HIGH melting point due to strong covalent bonds in hexagon sheets
* Slippy due to weak intermolecular forces between layers
* Conductive due to delocalised electrons
* Insoluble
* Less dense than diamond as layers are far apart

Question 23

Describe the structure and properties of graphene

Answer 23

Single sheet of carbon atoms that is one atom thick
* conductive as delocalised electrons are free to move and carry charge
* Very strong as delocalised electrons strengthen bonds between carbon atoms
* transparent and light as it is only one layer
* insoluble due to strong covalent bonds

Question 24

Define simple molecular lattice

Answer 24

a few atoms covalently bonded togethet

Question 25

Describe and explain melting and boiling points in simple molecular structures

What about noble gases ?

Answer 25

• low melting/boiling point
• as you only have to overcome the induced dipole-dipole interactions
• which need minimal energy to overcome
• Larger molecule = stronger induced dipole-dipole forces = higher melting and boiling point

Noble gases
- exist as single atoms so have very weak induced dipole-dipole forces
- so have very low melting and boiling points

Question 26

Describe the what affects the melting and boiling point across the period of:
- metals
- giant covalent lattice
- simple molecular structures
- noble gases

Answer 26

• Metals - more delocalised electrons and smaller atomic radius = stronger metallic bonds = increased melting and boiling point
• Giant covalent lattice - more covalent bonds = increased melting and boiling point
• Simple molecular structures - larger molecule = stronger induced dipole-dipole forces
• Noble gases - all exist as single atoms so have very weak intermolecular forces and low melting/boiling points