1. Topic 2 & 3 :Atomic Structure & Periodicity Modified from Scheffler
IB Chemistry
Topic 2 & 3 :Atomic Structure & Periodicity
Modified from Scheffler

2. Ionization energy is defined as the amount of energy required to remove a mole of electrons from a mole of gaseous atoms of a particular element.
E(g)  E+(g) + e-

3. Trends in ionization energy occur in the Periodic Table
Trends in ionization energy occur in the Periodic Table. Ionization energy decreases down a group and increases across a period.

4. Trends in ionization energy occur in the Periodic Table
Trends in ionization energy occur in the Periodic Table. Ionization energy decreases down a group and increases across a period.
WHY?

5. Effective nuclear charge is the net positive charge felt by an electron in an atom.
The basics of electrostatics imply that each and every electron is attracted to each and every proton in the nucleus and repelled by every other electron. However . . .
Electrons between the valence electron and the nucleus provide a shielding effect – weakening the electrostatic force on the valence electron.

6. Across a Period:
shielding remains constant
atomic number increases so effective nuclear charge increases
ionization energy increases

7. Down a Group:
shielding increases AND atomic number increases
effective nuclear charge does not change significantly
valence electrons further from nucleus
so weaker electrostatic force and lower ionization energy

8. This diagram shows how 1st ionization energy decreases down a group and that trends in ionization energy also occur across a period. Look at the H, Li, Na, K, Rb, Cs values! .

9. Looking at just the trend across the 1st period, what does the graph imply?
The theory is . . .
Across a period, number of p+ increases so effective nuclear charge increases.
As a result, the valence electrons are more strongly held, and are closer to the nucleus (radius decreases)
This does not explain the drop in ionization energy (decreased stability) observed between Be and B and between N and O.

10. Be B NEW IDEA – suborbitals (or subshells)
Within a given energy level (shell), there are different subshells that electrons can occupy that have slightly different energy levels
3d ___ ___ ___ ___ ___
4s ____
3p ___ ___ ___
3s ___
2p ___ ___ ___
2s ___
1s ___ Be
3d ___ ___ ___ ___ ___
4s ____
3p ___ ___ ___
3s ___
2p ___ ___ ___
2s ___
1s ___ B

11. Be B NEW IDEA – suborbitals (or subshells)
Within a given energy level (shell), there are different subshells that electrons can occupy that have slightly different energy levels
3d ___ ___ ___ ___ ___
4s ____
3p ___ ___ ___
3s ___
2p ___ ___ ___
2s ___
1s ___ Be
3d ___ ___ ___ ___ ___
4s ____
3p ___ ___ ___
3s ___
2p ___ ___ ___
2s ___
1s ___ B

12. N O NEW IDEA – suborbitals (or subshells)
Within a given energy level (shell), there are different subshells that electrons can occupy that have slightly different energy levels
3d ___ ___ ___ ___ ___
4s ____
3p ___ ___ ___
3s ___
2p ___ ___ ___
2s ___
1s ___ N
3d ___ ___ ___ ___ ___
4s ____
3p ___ ___ ___
3s ___
2p ___ ___ ___
2s ___
1s ___ O

13. N O NEW IDEA – suborbitals (or subshells)
Within a given energy level (shell), there are different subshells that electrons can occupy that have slightly different energy levels
3d ___ ___ ___ ___ ___
4s ____
3p ___ ___ ___
3s ___
2p ___ ___ ___
2s ___
1s ___ N
3d ___ ___ ___ ___ ___
4s ____
3p ___ ___ ___
3s ___
2p ___ ___ ___
2s ___
1s ___ O

14. Ionization energy trends
Down a group : ionization energy decreases
- ENC constant but atoms larger so easier to ionize
Across a period : ionization energy increases
- increasing ENC therefore smaller size (e- closer to nucleus)
so harder to ionize

15. Explaining the “dips” – support for s and p orbital model
Be to B “dip”
- because s shields p and lowers ENC
N to O “dip”
- because repulsions between electron pair in first full orbital (experimental evidence supporting Aufbau and Hund)

16. Electron Configurations and the Periodic Table
So far, we have seen how the subshell model provides and explanation for the patterns in ionization energy we see in the periodic table.
You have also seen how to write electron configurations
Example CALCIUM  1s22s22p63s23p64s2
Principle energy level subshell # of e-
Calcium can also be written shorthand as:
[Ar]4s2

17. The organization of the Periodic table correlates directly to electron structure

18. Condensed electron configurations – for example the electron configuration of bromine can be written [Ar] 4s23d104p5
Read questions carefully – many IB questions require you to write the FULL electron configuration

19. Electron configuration of ions:
In general, electrons will be removed from orbitals (ionization) in the reverse order that the orbitals were filled. In other words, electrons vacate higher energy orbitals first.
The exception: TRANSITION METAL IONS
When these ions form, electrons are removed from the valence shell s orbitals before they are removed from valence d orbitals when transition metals are ionized.
For example: Cobalt has the configuration [Ar] 4s2 3d OR [Ar] 3d7 4s2
The Co2+ and Co3+ ions have the following electron configurations.
Co2+: [Ar] 3d7 Co3+: [Ar] 3d6

20. You are responsible for configurations up to Z = 54 (Xe)
You are responsible for configurations up to Z = 54 (Xe). The table works well for this with the exception of Cr and Cu

21. Chromium’s configuration is:
[Ar]4s13d5
Copper’s configuration is:
[Ar]4s13d10
These configurations are energetically more stable than the expected arrangements. KNOW THEM!

22. Successive ionization
energy data supports the electron configuration model
169988
3rd
7732.7
2nd
1450.7
1st
737.7

23. Trends in Atomic Size
First problem: Where do you start measuring from?
The electron cloud doesn’t have a definite edge.
They get around this by measuring more than 1 atom at a time.

24. Atomic Size }
Radius
Atomic Radius = half the distance between two nuclei of a diatomic molecule.

25. Trends in Atomic Size Influenced by three factors: 1. Energy Level
Higher energy level is further away.
2. Charge on nucleus
More charge pulls electrons in closer.
3. Shielding effect
e <-> e repulsion

26. Group trends As we go down a group...H
As we go down a group...
each atom has another energy level,
so the atoms get bigger. Li Na K Rb

27. Periodic Trends As you go across a period, the radius gets smaller.
Electrons are in same energy level.
More nuclear charge.
Outermost electrons are closer. Na Mg Al Si P S Cl Ar

28. Rb
Overall K Na Li
Atomic Radius (nm) Kr Ar Ne H 10
Atomic Number

29. Ionic Radius Cations are always smaller than the original atom.
The entire outer PEL is removed during ionization.
Conversely, anions are always larger than the original atom.
Electrons are added to the outer PEL.

30. Cation Formation
Effective nuclear charge on remaining electrons increases.
Na atom
1 valence electron
11p+
Remaining e- are pulled in closer to the nucleus. Ionic size decreases.
Valence e- lost in ion formation
Result: a smaller sodium cation, Na+

31. Anion Formation
A chloride ion is produced. It is larger than the original atom.
Chlorine atom with 7 valence e-
17p+
Effective nuclear charge is reduced and the e- cloud expands.
One e- is added to the outer shell.

32. Electronegativity
Definition: how much an atom will hog the electrons when it’s bonded to another atom (unfair sharing)
Trend: increases from left to right within a period, and decreases from top to bottom within a group.

33. Why increase from left to right
Why increase from left to right? Because there are more protons added to the nucleus so the electrons get pulled more tightly. A smaller atomic radius and a larger + charge in the nucleus means atoms will attract bonded electrons more strongly
Why decrease from top to bottom? Because there are more energy levels. This makes the valence electrons (bonding electrons) farther away from the nucleus. Furthermore, the valence electrons are shielded or repelled (see below) by inner energy shells filled with electrons.