3.1 - Periodicity

Question 1

How are the known elements of the Periodic Table arranged?

Answer 1

In order of increasing proton number.

Question 2

What do all elements in the same period have?

Answer 2

Electron shells.

Question 3

What do all elements in the same group have?

Answer 3

Outer electrons and similar chemical properties.

Question 4

What are elements classified into?

Answer 4

Blocks.

Question 5

What do elements in the same block have?

Answer 5

The outer electrons are in the same type of orbital.

Question 6

What block are the outer electrons in for Groups 1 and 2 in the Periodic Table?

Answer 6

S-block.

Question 7

What block are the outer electrons in for Groups 3 to 0 in the Periodic Table?

Answer 7

P-block.

Question 8

What block are the outer electrons in for the transition metals in the Periodic Table?

Answer 8

D-block.

Question 9

What block are the outer electrons in for the lanthanides and actinides in the Periodic Table?

Answer 9

F-block.

Question 10

What is Periodicity the study of?

Answer 10

Other trends in the Periodic Table linked to the different electronic configurations.

Question 11

What order are the orbitals filled in?

Answer 11

The energy of the orbitals increase from S-F, therefore they are filled in that order.

Question 12

When can the next orbital hold electrons?

Answer 12

When the previous orbital has been filled.

Question 13

When and why does the atomic radius decrease?

Answer 13

Along a period, as the nuclear charge is increased for the same number of electron shells. The outer electrons are pulled in closer to the nucleus, as the increase charge produces a greater attraction.

Question 14

When and why does the atomic radius increase?

Answer 14

Down a group, as an electron shell is added, therefore the distance between the outer electrons and the nucleus is increased. This reduces the power of attraction. More shells increases electron shielding (where the inner shells create a ‘barrier’ that blocks the attractive forces), and this further reduces nuclear attraction.

Question 15

What is the definition of the first ionisation energy?

Answer 15

The minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state. It is measured in kJ mol-1.

Question 16

When and why do successive ionisation energies occur?

Answer 16

When further electrons are removed, this usually requires more energy because the electrostatic forces of attraction increases between the positive nucleus and the negative outer electron. More energy is needed to overcome this attraction, therefore increasing the ionisation energy.

Question 17

What are the trends followed by the first ionisation energy within the Periodic Table?

Answer 17

Along a period, the first ionisation energy increases due to a decreasing atomic radius and greater electrostatic forces of attraction.
Down a group, the first ionisation energy decreases due to an increasing atomic radius and electron shielding, which reduces the effect of the electrostatic forces of attraction.

Question 18

Why do successive ionisation energies increase?

Answer 18

The atomic radius decreases and the attraction between the nucleus and the outer shell electrons is greater.

Question 19

How can the successive ionisation energies indicate the group an element is in?

Answer 19

A large jump between successive ionisation energies indicates that the first electrons are relatively easy to remove, and then there is a jump to show that a lot more energy is required to remove that one, therefore proving that that electron is on a different electron shell.

Question 20

What is the trend for first ionisation energies in Period 2?

Answer 20

The ionisation energies increase along Period 2 because the atomic radius decreases and the nuclear charge increases, therefore causing the outer electrons to be held more strongly.

Question 21

What are the exceptions to the trend in Period 2 and why are they the exceptions?

Answer 21

Boron has a lower first ionisation energy than expected due to the energy difference between the 2s and 2p sub-shells. The electron is being removed from a higher energy level, further from the nucleus, therefore the electron is held less strongly.

Oxygen has a lower first ionisation energy that expected, due to the repulsion within the 2p orbital when two electrons with opposite spins are in the same orbital.This destabilises and therefore allows the electron to be removed more easily.

Question 22

What is the trend for first ionisation energies in Period 3?

Answer 22

The ionisation energies increase along Period 3 because the atomic radius decreases and the nuclear charge increases, therefore causing the electrons to be held more strongly.

Question 23

What are the exceptions to the trend in Period 3 and why are they exceptions?

Answer 23

Aluminium has a lower first ionisation energy than expected due to the energy difference between the 3s and 3p sub-shells. The electron is being removed from a higher energy level, further from the nucleus, therefore it is held less strongly.

Sulfur has a lower first ionisation energy than expected due to the repulsion within the 3p orbital, when two electrons with opposite spins are in the same orbital. this destabilises and allows the electron to be removed more easily.

Question 24

What does metallic bonding consist of?

Answer 24

A lattice of positively charged ions (cations) surrounded by a ‘sea’ of delocalised electrons. Between the oppositely charged particles, there are very strong electrostatic forces of attraction.

Question 25

What does a greater charged ion do in metallic bonding?

Answer 25

It increases the strength of the attractive force, due to more electrons being released into the ‘sea’.

Question 26

What do ions larger in size cause in metallic bonding?

Answer 26

They produce a weaker attraction and their greater atomic radius decreases the charge density.

Question 27

What are the properties of a metallic structure?

Answer 27

Good conductors, as the ‘sea’ of delocalised electrons can move and carry a flow of charge.
Malleable, as the uniform layers of positive ions can slide over each other.
High melting points, as the electrostatic forces of attraction between the cations and delocalised electrons