Electrons in atom

Ionisation Energy

Atomic radius

Open Ended

What is definition of 1st ionisation energy?

Open Ended

Write the equation of 3rd ionisation energy of calcium

The big gap in ionisation energies can determine the valence electrons of an atom and the group

Electronic configurations

Multiple Choice

Which subshell has 5 orbitals?

• s

• p

• d

• f

Periodic patterns of atomic and ionic radii

Atomic radius

The atomic radius increases down any group. This is because, going down the group, each successive element has one more shell of electrons which is further from the nucleus.
Although there is also an increasing nuclear charge going down the group, the increased effect of inner shell electrons shielding the outer shell electrons is more important.
The atomic radius decreases across any period. This is because the number of protons, and therefore the nuclear charge, increases by one with each successive element. The number of electrons also increases by one but the extra electron added goes into the same (outer) energy level. This means that the shielding does not change significantly. The greater attractive force of the increased nuclear charge on the outer shell electrons pulls them closer to the nucleus.

The ionic radius increases down any group (for ions with the same charge).

Positively charged ions are smaller than their original atoms because they have lost their outer shell electrons.

Negatively charged ions in Groups 15 to 18 are larger than their original atoms because each atom will have gained one or more extra electrons into their outer energy level.ncreased

Patterns in ionisation energies in the Periodic Table

There is a general increase in IE1 across a period.
This applies to Period 1 (hydrogen and helium), Period 2 (lithium to neon) and also to other periods.

As you go across a period the nuclear charge increases. But the electron removed comes from the same shell. The force of attraction between the positive nucleus and the outer negative electrons increases across the period because:
• the nuclear charge increases
•the distance between the nucleus and the outer electron remains reasonably constant
•the shielding by inner shells remains reasonably constant.

​Patterns across a period

There is a rapid decrease in ionisation energy between the last element in one period and the first element in the next period.
The IE1 for lithium is much smaller than the IE1 for helium. Helium has two electrons. These are in the first quantum shell. But lithium has three electrons.
The third electron must go into the next quantum shell further away from the nucleus.

The force of attraction between the positive nucleus and the outer negative electrons decreases because: •the distance between the nucleus and the outer electron increases
•the shielding by inner shells increases
•these two factors outweigh the increased nuclear charge.constant.

There is a slight decrease in IE1 between beryllium and boron.
Although boron has one more proton than beryllium, there is a slight decrease in IE1 on removal of the outer electron.
Beryllium has the electronic structure 1s² 2s² and boron has the electronic structure 1s² 2s² 2p¹.
The fifth electron in boron must be in the 2p sub-shell, which is slightly further away from the nucleus than the 2s sub-shell.
There is less attraction between the fifth electron in boron and the nucleus because:
•the distance between the nucleus and the outer electron increases slightly
•the shielding by inner shells increases slightly
•these two factors outweigh the increased nuclear charge.

There is a slight decrease in IE1 between nitrogen and oxygen. Oxygen has one more proton than nitrogen and the electron removed is in the same 2p sub-shell. So, you might think that IE1 would increase.
However, the spin-pairing of the electrons plays a part here.
the electron removed from the nitrogen is from an orbital that contains an unpaired electron.
The electron removed from the oxygen is from the orbital that contains a pair of electrons. The extra repulsion between the pair of electrons in this orbital results in less energy being needed to remove an electron. So, IE1 for oxygen is lower, because of spin-pair repulsion.

The first ionisation energy decreases as you go down a group in the Periodic Table.
For example, in Group 1 the values of IE1 are:
Li = 519 kJ mol⁻¹
Na = 494 kJ mol⁻¹
K = 418 kJ mol⁻¹
Rb = 403 kJ mol⁻¹
As you go down the group, the outer electron removed is from the same type of orbital but from a successively higher principal quantum level: 2s for lithium, 3s for sodium and 4s for potassium.
Although the nuclear charge is increasing down the group there is less attraction between the outer electron and the nucleus because:
the distance between the nucleus and the outer electron increases
the shielding by complete inner shells increases
these two factors outweigh the increased nuclear charge

​Patterns down a group

Open Ended

Why is Oxygen has lower 1st ionisation energy than nitrogen