1. Learning Intention: To explain periodic trends associated with ionisation energy

2. By the end of this section, I will be able to:
Success Criteria
By the end of this section, I will be able to:
Define the term ionisation energy
Describe the first ionisation energy for an atom by writing balanced ion electron equations.
Explain trends in first ionisation energies across a period and down a group.
Use electron configurations to explain exception to the general periodic trends for ionisation energy.

3. Ionisation energy
Definition:
Ionisation energy is defined as the energy required to remove the least tightly bound electron from 1 mole of gaseous atom/ion.

4. Balanced ion-electron equation: X(g)  X+ (g) + e- ΔH = +ve
The first ionisation energy is the energy required to remove the first valence electron from one mole of gaseous atoms.
Balanced ion-electron equation:
X(g)  X+ (g) + e ΔH = +ve
e.g. Li (g) Li+ (g) e- ΔH = +519 kJ/mol 

5. 1st IE < 2nd IE < 3rd IE < ……
Second IE – energy used to remove the second electron from a mole of gaseous ion.
This is always greater than the first IE as electrons are removed from a positive ion and the attractive force between the nucleus and remaining valence electrons is stronger.
1st IE < 2nd IE < 3rd IE < ……
Na (g) Na+(g) + e- Δ H = +119kJ/mol
Na+ (g) Na2+(g) + e- Δ H = +1090kJ/mol
Na2+ (g) Na3+(g) + e- Δ H = +1652kJ/mol   

6. The higher the ionisation energy, the stronger the force of attraction between an atom’s nucleus and it’s valence electrons.
This force of attraction is influenced by factors such as:
Number of protons (nuclear charge found in an atom)
The distance between an atom’s nucleus and it’s valence electrons
And the shielding of valence electrons by inner electrons.
The value is always a positive number as energy is required to remove electrons
– it is an endothermic reaction.

7. Periodic Trend down a group
As we go down the periodic table ionisation energy decreases.
This is because as the atomic number increases more energy levels and electrons are added. The size of the atom increases.The electron lost is further away from the nucleus, and there is more shielding from the nuclear attraction by the increased number of inner electrons ( effective nuclear charge decreases). Electrons are less tightly held so less energy required to remove an electron

8. 1st Ionisation Energy Gp 2
Ionisation Energy Trends
As we go down the group in the periodic table ionisation energy decreases. This is because more energy levels are added, the electron lost is further away from the nucleus,and there is more shielding from the nuclear attraction by the increased number of inner electrons.( nuclear charge
decreases).

9. 1st Ionisation Energy P3
As we go across a period in the periodic table ionisation energy increases. This is because as atomic number increases there are more protons in the nucleus and more electrons are added to the same shell so there is a stronger nuclear attraction on electrons. Electrons are more tightly held so more energy is required to remove an electron. It becomes more difficult to remove an electron from a non-metal than a metal in the same period (row).

10. PeriodicTrend across a period
As we go across the periodic table ionisation energy increases.
This is because as atomic number increases there are more protons in the nucleus and the electrons are added to the same energy level. The size of the atom decreases so there is a stronger attraction on electrons ( nuclear charge increases). Electrons are more tightly held therefore more energy is required to remove an electron
It becomes more difficult to remove an electron from a non-metal than a metal in the same period (row).

11. The trend across the period is not always regular.
When we go from group 2 to 13 the IE decreases.
Be B
1s2, 2s s2, 2s2, 2p1
B has a lower first IE as the first e- is removed from a partially filled p orbital where as in Be the first e- is removed from a fully filled s orbital which is closer to the nucleus. This requires more energy.

12. When we go from group 15 to 16 the IE decreases.
N O
1s2, 2s2, 2p s2, 2s2, 2p4
O has a lower first IE than N since the first e- is removed from a partially filled p orbital where as in N the first e- is removed from a half-filled p orbital which is more stable and requires more energy.

13. Ionization energy vs. atomic number
04/10/99

14. All Elements

15. Successive Ionisations

17. Successive Ionisations

18. Subsequent Ionisation Energy’s
Second and third ionisation energies are always greater than the first as it takes more energy to remove an electron from an increasingly positive ion. However, if electrons are being removed from the same shell the increase is not as great.
1st
2nd
3rd
4th
Na 2,8,1
484
4560
6940
9540
Mg 2,8,2
736
1450
7740
10500
Ne 2,8
2080
3950
6150
9290
Removing one electron from sodium requires 484 kJ mol-1. It becomes Na+. But to remove a subsequent electron from a shell closer to the nucleus takes 4560 kJ mol-1. So Na2+ isn’t a natural occurance.
But, removing two electrons from magnesium is possible as Mg has two valence electrons. The ionisation energy increases if you are trying to remove the third. For neon, it is relatively difficult to remove the first electron because its electrons are stable.

20. Atomic Radii
atomic radii is half the average distance between neighbouring atom nuclei.

22. Atomic Radii Gp 2
As we go down a group of the periodic table the atomic radii increases.
This is because we are adding an extra shell of electrons.

23. Atomic Radii P3
As we go across the periodic table from left to right the atomic radii decreases.
This is because more electrons are added to the same energy level and each valence electron feels the pull of more and more protons in the nucleus(effective nuclear charge increases), so metals (the left) are larger than non-metals (the right).
Atomic number increases, more proton added and more electrons added to the same energy level, nuclear charge increases, outer energy level pulled towards the nucleus.

24. Electronegativity

25. Ionic Radii Trend
As we go across the periodic table from left to right the ionic radii increases.
This is because
positive ions- electrons (and a whole shell) are lost from metals making them smaller and
Negative ions-valence shells are filled with extra electrons in the non-metals making the cloud bigger as there is increased electron – electron repulsion.

26. Ionic Radii

27. Electronegativity This measures how strongly each atom in a molecule attracts the electrons involved in bonding. Fluorine is the most electronegative atom followed by oxygen, nitrogen and chlorine.

28. Trend across a period
As we go across the periodic table left to right electronegativity increases (ie atoms hold their electrons more tightly).
This is because atoms are smaller in size which means the electrons are closer to the nucleus so they are feeling the pull of protons more strongly. This allows the nucleus to attract the bonding electrons easily.

29. Trend down a group
As we go down a group of the periodic table electronegativity decreases (ie atoms cannot hold onto their own valence electrons as well). This is because
Atoms are getting bigger in size so the valence electrons involved in bonding are getting further away from the nucleus of the atom.

30. The outside electrons are being shielded from the pull of the protons by all the inner layer of electrons
There are more layers of electrons so they are further away from the nucleus and again feel less pull from the protons
The nucleus is unable to attract the valence electrons easily.

31. Though the nucleus gets more positive as we go down the group since more protons are added, the distance factor is more influential in determining the value of the electronegativity.

32. Q6. Answer
All are in the same priod so have same number of energy levels or electron shells.
Mg > Cl because Cl has more protons so has highest nuclear charge. Therefore there is a stronger attraction for the valence electrons bringing the valence electrons closer to the nucleus.
Cl- > Cl because although they both have the same nuclear charge (same number of protons) Cl- has an additional electron added to the same valence shell so has increased electron – electron repulsion with same nuclear charge. Effective nuclear charge decreases and valence electrons occupy more space. Hence Cl- is larger than Cl.