Electronic Structure (+ ionisation energy) A1

Question 1

What is the number of electrons each shell can hold?

Answer 1

Shell 1 = 2
Shell 2 = 8
Shell 3 = 18
Shell 4 = 32

Question 2

What is an atomic orbital?

Answer 2

A region within an atom that can hold up to two electrons with opposite spin.

Question 3

Describe s-orbitals

Answer 3

• Spherical shape
• Can up up to 2 electrons
• Every shell has 1 s-orbital

Question 4

Describe p-orbitals

Answer 4

• Dumb-bell shape
• Each shell from n=2 contains 3 p-orbitals
• Each p-orbital can hold up to 2 electrons, 6 in a sub-shell

Question 5

Describe d-orbitals

Answer 5

• Every shell from n=3 contains 5 d-orbitals
• Each d-orbital contains 2 electrons, 10 in a sub-shell

Question 6

Describe f-orbitals

Answer 6

• Every shell from n=4 contains 7 f-orbitals
• Each f-orbital contains 2 electrons, 14 in a sub-shell

Question 7

What is the order of filling for electrons?

Answer 7

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p

Question 8

What are the rules when drawing box diagrams for shells?

Answer 8

• Up to two electrons (arrows) per box.
• If 2 arrows in a box, they must be facing opposite directions.
• If filling an empty sub-shell, fill one arrow per box first.

Question 9

When removing electrons:
Ions - remove electrons in reverse filling order
Transition metals - ?

Answer 9

Remove electrons in reverse filling order but the 4s is lost before the 3d.

Question 10

What is first ionisation energy?

Answer 10

The energy required to remove 1 mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

Question 11

What are the 3 factors that affect ionisation energy?

Answer 11

1) Atomic radius (the greater the atomic radius, the lower the ionisation energy).
2) Nuclear Charge (the greater the nuclear charge the higher the ionisation energy).
3) Electron shielding (the more inner shells the lower the ionisation energy).

Can you remember why?

Question 12

As you go down a group, first ionisation energy ______ because:

Answer 12

Decreases

1) Atomic Radius increases
2) There are more inner electron shells so Shielding increasing.
3) So nuclear attraction of the outer electrons decreases.
4) So less E required to remove outer electron (First IE decreases)

Question 13

As you go across a period, first ionisation energy ______ because:

Answer 13

Increases

1) Nuclear Charge increases
2) The electrons are in the same shell so the Shielding is similar.
3) Nuclear attraction increases (so Atomic Radius decreases)
4) So more E required to remove outer electron (first IE increases)

Question 14

the higher the energy (level) of the sub-shell, the ____ energy needed to remove electron.

Answer 14

Lower

Question 15

Why does B have a lower ionisation energy than Be?

Answer 15

Because it is in a higher energy sub-shell.

Question 16

Group 1 and 2 are known as the __-block
Transition metals are known as the __-block
Group 3,4,5,6,7,8 are known as the __-block

because…

Answer 16

s
d
p

That is the last sub-shell that was filled for that element.

Question 17

Why does O have a lower ionisation energy than N?

Answer 17

Because the paired 2p electrons in O repel each other, making one of them easier to remove than a 2p electron from N. (refer to box diagram)

Question 18

Why are there large jumps in successive ionisation energies? (as number of ionisation increases for an element)

Answer 18

Large jumps correspond to moving electrons into a ‘deeper shell’

e.g. for group 2 element, IE1, IE2 (big jump) IE3 IE4 IE5 etc

Question 19

What is periodicity?

Answer 19

The repeating trend in properties on the periodic table.

Question 20

Oxygen is in group 6. How would successive ionisation energies show this?

Answer 20

Large increase between 6th and 7th ionisation energy values.