C7 - Periodicity

Question 1

What is periodicity?

Answer 1

The repeating trend in properties of the elements across each period.

Question 2

Who identified the law of octaves?

Answer 2

John Newlands.

Every 8th element has similar properties, known as the law of periodicity.

Question 3

Who formulated the periodic table?

Answer 3

Dimitri Mendeleev

Question 4

How are the elements of the periodic table arranged?

Answer 4

They have increasing atomic number from left to right.

They are in vertical groups based on their number of electrons on their outer shell and chemical properties

They are in horizontal rows. The period number gives the number of the highest energy electron shell of that element’s atom.

Question 5

Why did Mendeleev leave spaces in his periodic table and what did he predict?

Answer 5

He left gaps for elements which he believed existed but hadn’t been discovered yet. He predicted the existence of elements e.g. gallium.

Question 6

What is ionisation?

Answer 6

A process where atoms become charged ions due to the loss or gain of electrons.

Question 7

What is the first ionisation energy?

Answer 7

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

M -> M+ + e-

Question 8

What factors affect ionisation energy?

Answer 8

Atomic radius - the greater the distance, the weaker the nuclear attraction.

Nuclear charge - the greater the amount of protons, the greater nuclear attraction.

Electron shielding - The more inner-shell electrons present, the more shielding due to repulsion of the electrons so attraction decreases.

Question 9

What is the second ionisation energy?

Answer 9

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

Question 10

Why would a successive ionisation graph show large increases in its energy levels?

Answer 10

Because an electron is being removed from an inner shell.

Removing an electron from a lower energy level requires more energy.
It also shows the filling of different shells.

Question 11

What is the trend of first ionisation energy across a period?

Answer 11

First ionisation energy increases as nuclear charge increases therefore nuclear attraction increases and atomic radius decreases.

Within a period, electrons are in the same shell so there is similar shielding/shielding doesn’t affect the energy.

Question 12

What is the trend of first ionisation energy down a group?

Answer 12

First ionisation energy decreases as atomic radius increases and there are more inner shells so shielding increases.
This means nuclear attraction decreases so ‘I’ energy decreases.

Question 13

How are metals bonded?

Answer 13

By metallic bonding

Question 14

What is the structure of a metal?

Answer 14

Rows of (positive) atoms surrounded by a sea of delocalised electrons. 
This conduct electricity due to these electrons which can move carrying charge. 

Each atom donates it’s outer shell electrons to a shared pool of (delocalised) electrons spread throughout the whole structure.
The (positive) cations left behind consist of the nucleus and inner electrons.

The cations are fixed meanwhile the electrons are mobile.

Question 15

Do metals conduct better when hot or cold?

Answer 15

Cold as the ions are moving less so are less intrusive to the electrons carrying charge.

Question 16

What are allotropes?

Answer 16

Different forms of the element

Question 17

What’s graphene?

Answer 17

A single layer of carbon one atom thick

Question 18

What causes the increase in metallic bond strength from Na to Al?

Answer 18

Charge density - Na+ ions are large with a small charge so have a low charge density.

The number of free electrons - Na has 1 free electron per ion whereas Al has 3 so more attractions must be broken within Al.

Question 19

What is charge density?

Answer 19

The ratio of an ion’s charge to its size

Question 20

What is the structure of silicon?

Answer 20

It has a macro molecular structure similar to a diamond.

Each Si atom has 4 strong covalent bonds.

Question 21

What’s the structure of phosphorus?

Answer 21

It forms a P4 tetrahedral molecule.

It has the second highest boiling point in period 3.

Question 22

What’s the structure of sulfur?

Answer 22

Forms crown shaped molecules. (S8)

They’re the largest in period 3 and have more Van der waal forces so has the highest m/b points.

Question 23

What’s the structure of chlorine?

Answer 23

It forms diatomic Cl2 molecules.

It’s small and has weak Van de waal forces so has low m/b points.

Question 24

What’s the structure of argon?

Answer 24

It’s a monatomic molecule with very weak m/b points.

Question 25

Define the term ‘first ionisation energy’:

Answer 25

It’s the amount of energy required to remove one electron from each atom in one mole of gaseous atoms.

Question 26

Explain why first ionisation energies increase from Li to Ne:

Answer 26

Moving across the period, nuclear charge increase so nuclear attraction increases and atomic radius decreases.
Shielding (in the same period) remains the same and doesn’t affect ionisation energies.

Question 27

Write an equation to represent the second ionisation energy of oxygen:

Answer 27

O+ (g) -> O 2+ (g) + e-

Question 28

Why is the second ionisation energy of oxygen greater than the first?

Answer 28

There are fewer electrons for the same number of protons and the atomic radius is smaller.
Shielding decreases.

Question 29

Solid carbon exists as diamond and graphite.

Explain why it is unnecessary to refer to carbon as either diamond or graphite when discussing their (first) ionisation energies:

Answer 29

Because the ionisation energy is the amount of energy required to remove one electron from each atom in one mole of gaseous atoms and, when gaseous, no covalent bonds exist.

Question 30

What is metallic bonding?

Answer 30

The strong electrostatic attraction between cations and delocalised electrons.