1. Periodic Trends
There are several important atomic characteristics that show predictable trends.
Atomic radius
Ionization energy
Electron affinity
Electronegativity
Adapted from: staffweb.psdschools.org/rjensen/powerpoints/periodic_trends.ppt

2. Why Do These Trends Exist?
The arrangement of elements in the periodic table results in characteristic changes in the
Effective nuclear charge
and
Shielding effect
going across a period or down a group.

3. Effective Nuclear Charge
Effective nuclear charge is the pull that an electron “feels” from the nucleus
The closer an electron is to the nucleus, the more pull it feels
As effective nuclear charge increases, the electron cloud is pulled in tighter

4. Shielding Effect
As more energy levels are added to an atom, the inner layers of electrons shield the outer electrons from the nucleus
The effective nuclear charge on those outer electrons is less, and so the outer electrons are less tightly held

5. Atomic Radius
Radius is the distance from the centre of the nucleus to the edge of the electron cloud
This is measured as half the distance between the nuclei of 2 bonded atoms

6. Atomic Radius

7. Atomic Radius
Atomic radius increases going down a vertical column or group
This is because each period has an additional electron shell or energy level
The shielding effect increases while effective nuclear charge decreases

8. Atomic Radius
Atomic radius tends to decrease going from left to right within a period
Each step adds a proton and an electron, but the energy level does not change
Effective nuclear charge increases
The increased attraction pulls the electron cloud in, making the atom smaller

9. Atomic Radius
**Heavy arrows point in the direction of an increase **

10. Ionization Energy
The energy required to remove a electron from an atom in the gas phase
The 1st electron (1st Ion. E.) is the easiest to remove, the 2nd (2nd Ion. E.) is harder, etc.
Ionization energy and atomic radius are inversely proportional
Ionization energy is always endothermic; energy is added to the atom to remove the electron

11. Copyright Pearson Prentice Hall

12. Ionization Energy

13. Ionization Energy
Ionization energy decreases going down a group because the valence electrons are further from the nucleus, therefore less energy is required to remove one
The shielding effect also increases as more energy levels are added

14. Ionization Energy
Ionization energy increases moving from left to right across a period
As the number of protons in the nucleus increases, so does the effective nuclear charge; and the energy level does not change

15. Ionization Energy

16. Electron Affinity
Electron affinity is the energy change that occurs when an atom in the gas phase gains an electron
Electron affinity is directly proportional to the effective nuclear charge
Where ionization energy is always endothermic, electron affinity is usually exothermic (but not always)

17. Electron Affinity
Electron affinity is exothermic if there is an empty or partially empty sub-orbital for an electron to occupy
If there are no empty spaces, a new sub-orbital must be created, making the process endothermic (e.g. alkaline earth metals and noble gases)

18. Electron Affinity Electron affinity decreases going down a group
Each time a new energy shell is added, the electrons are further from the nucleus and the effective nuclear charge decreases, while the shielding effect increases

19. Electron Affinity
Electron affinity increases moving left to right across a period
Electron affinity increases as the effective nuclear charge increases

20. Electron Affinity + +

21. Electronegativity
Electronegativity is a measure of an atom’s attraction for another atom’s electrons while in a chemical bond
It uses an arbitrary scale ranging from 0 – 4
Generally, metals are “electron givers” and have low electronegativities
Nonmetals are are “electron takers” and have high electronegativities
The noble gases have 0 electronegativity

22. Electronegativity
Electronegativity decreases going down a group as the effective nuclear charge decreases and the shielding effect increases
Electronegativity increases going from left to right across a period because the effective nuclear charge increases

23. Electronegativity

24. Overall Reactivity This ties all the previous trends together
However, we must treat metals and nonmetals separately
The most reactive metals are the largest, since they are the best electron givers
The most reactive nonmetals are the smallest ones, the best electron takers

25. Overall Reactivity

26. Ionic Radius Positive ions are called cations
Negative ions are called anions
Cations are always smaller than the original atom, because the entire outer shell is removed during ionization
Conversely, anions are always larger than the original atom, because electrons are added to the outer shell

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

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