3.1.1 Periodicity

Question 1

How is the periodic table set out?

Answer 1

In order of increasing atomic number, Groups and periods.

Question 2

Atomic radius definition

Answer 2

Total distance from an atoms nucleus to the outer most orbital

Question 3

Ionisation energy definition

Answer 3

Amount of energy required to remove one mole of gaseous electrons from one mole of gaseous atoms.

Question 4

Electronegativity definition

Answer 4

The ability to attract the bonding electron pair

Question 5

Electrical conductivity definition

Answer 5

the ability to carry an electrical current/ allow for the flow of electrons.

Question 6

Melting point definition

Answer 6

The temperature at which a substance turns from a solid to a liquid

Question 7

Boiling Point definition

Answer 7

temperature at which a substance turns from a liquid to a gas

Question 8

First ionisation energy definition

Answer 8

the energy required to remove one mole of gaseous electrons from one mole of gaseous atoms to form one mole of gaseous ions. Under standard conditions

Question 9

3 Main factors effecting first ionisation energy

Answer 9

Atomic radius, nuclear charge, electron shielding.

Question 10

how does atomic radius affect ionisation energy

Answer 10

The larger the radius the less energy required to remove the outermost electron

Question 11

how does nuclear charge affect ionisation energy

Answer 11

the higher the nuclear charge the higher the amount of energy required to remove the outermost electron

Question 12

how does Electron shielding affect ionisation energy

Answer 12

the higher the shielding the effect the lower the amount of energy required to remove the outermost electron

Question 13

Why isn’t the graph of ionisation linear

Answer 13

When atoms only have 1 electron in their highest energy orbital it is easier to remove than the atom with a lower atomic number.

Question 14

what is successive ionisation energies

Answer 14

The energy required to remove the outer most electron

Question 15

Why does successive ionisation energies require more energy along the periods

Answer 15

As the electrostatic force of attraction increased between the nucleus and the outermost electron as electrons are removed.

Question 16

How many ionisation energies can an atom have

Answer 16

as many as electrons it has

Question 17

How are ionisation energies used to indicate the group of an element

Answer 17

a large jump between successive ionisation energies indicates the group of the element.

Question 18

What occurs in metallic bonding, (4 points)

Answer 18

Each atom donates outer shell electrons,
positive cations remain fixed maintaining structure and shape,
Delocalised electrons are free to move around,
Billions of metal ions held together in giant metallic lattice.

Question 19

Why does metallic bonding have strong electrostatic forces of attraction

Answer 19

Strong attraction between positive cations and delocalised electrons

Question 20

Why does metallic bonding have high electrical conductivity

Answer 20

When a voltage is applied the delocalised electrons are able to carry a charge through the structure

Question 21

Why does metallic bonding have high melting and boiling points.

Answer 21

Large amounts of energy are required to overcome strong electrostatic forces of attraction between cations and delocalised electrons

Question 22

Why is metallic bonding non soluble

Answer 22

Any interactions between polar solvents and cations cause reactions rather than dissolving

Question 23

The 3 atoms that can create giant covalent bonds

Answer 23

Carbon, boron and silicon

Question 24

The shape and bond angle of diamond and silicon carbide and why.

Answer 24

Tetrahedral, 109.5°, as each atom is bonded to 4 other atoms

Question 25

Why do giant covalent structure have a high melting and boiling point

Answer 25

The large amount of covalent bonds requires lots of energy to break

Question 26

Why are giant covalent structures non soluble in almost all solvents

Answer 26

As the covalent bonds are too strong to be broken by solvent interactions

Question 27

Why are giant covalent structures not electrical conductors (bar graphite)

Answer 27

All outer shell electrons are used for bonding meaning non are delocalised/available for electrical conduction.

Question 28

How come graphite conduct electrical charge

Answer 28

As spare delocalised electrons occupy the space in between the graphene layers.

Question 29

What happens to the melting and boiling point from Na to Al

Answer 29

It increases as metallic bond strength increases

Question 30

2 reasons for metallic bond strength to increase

Answer 30

1. Charge density, the ratio of cations to atomic radius
2. Number of free electrons, more free electrons = more attractions

Question 31

Name of silicon’s structure

Answer 31

Macromolecular

Question 32

What and why is the melting and boiling point of silicon

Answer 32

Very high as, each atom is bonded to 4 other silicon atoms by strong covalent bonds.

Question 33

What and why is the melting/ boiling point and structure of phosphorus

Answer 33

P4 as each atom is bonded to 3 others with strong covalent bonds, but weak London forces mean low melting and boiling point

Question 34

What and why is the melting/boiling point and structure of sulfur

Answer 34

Each atom is bonded to 2 others creating a ring, S8. Low boiling and melting point as strong covalent bonds but weak London forces.

Question 35

What and why is the melting/boiling point and structure of chlorine

Answer 35

Forms a diatomic molecule, low melting and boiling point as weak London forces. Simple molecule

Question 36

What and why is the melting/boiling point and structure of argon (noble gases)

Answer 36

Monatomic structure very weak intermolecular forces, very low melting + boiling point.

Question 37

Describe the trend in atomic radius across period 3

Answer 37

Decreases due to increasing nuclear charge

Question 38

general trend in first ionisation energy across period 3

Answer 38

increases with a dip at aluminium and sulfur

Question 39

Why does first ionisation energy across period 3 dip at aluminium

Answer 39

It’s outermost electron is the only one in the 3p orbital meaning it is further from the nucleus and easily lost

Question 40

Why does first ionisation energy across period 3 dip at sulfur

Answer 40

sulfur outermost electron is the 4th in the 3p orbital meaning it shares an orbital with another electron of opposite spin. As these repel each other it is lost easierly.

Question 41

Explain the trend in melting point from sodium to magnesium

Answer 41

increases due to increasing metallic bond strength due to increases charge density and number of free electrons.