1. Unit 8: The Periodic Table Trends

2. Notes goals Understand basic history and arrangement of periodic table
Be able to use the periodic table to predict chemical trends
Be able to use periodic trends to determine the identity of elements

3. History Dmitri Mendeleev (1869, Russian)
Organized elements by increasing atomic mass.
Elements with similar properties were grouped together.
There were some discrepancies.

4. History Henry Mosely (1913, British)
Organized elements by increasing atomic number.
Resolved discrepancies in Mendeleev’s arrangement.

5. Arrangement Period = a single row going across left to right
Numbered from 1 to 7
Group/Family = a column going down top to bottom
Different numbering systems (1-18) or (1A – 8A, etc.)
Elements in the same group have similar physical and chemical characteristics

6. Periodic Trends Effective Nuclear Charge Atomic Radius Ionic Radius
Ionization Energy
Electron affinity
Electronegativity
Overall Reactivity

7. Effective Nuclear Charge
Valence electrons = electrons in outer level
Core electrons = electrons in levels underneath outer level
Valence electrons are:
Attracted to protons in nucleus
Repelled by core electrons which are in the way (shielding)

8. Effective Nuclear Charge
Effective Nuclear Charge = the net attraction of the valence electrons to the nucleus
Greater effective nuclear charge means valence electrons are closer to the nucleus

9. Effective Nuclear Charge (Zeff )
Felt by
Shielding electrons (S)
Effective Nuclear Charge (Zeff)
Atomic Number (Z = #p+)
Zeff = Z S
Inner Core of Shielding e- = screening constant
Attractive Charge felt by valence electrons
Atomic Number (#p+)
{Eff.Nuclear.Charge*}

10. Effective Nuclear Charge (Zeff )
depends on both the fact that valence electrons are both:
attracted to the nucleus
repelled by the other (shielding) electrons.
Effective Nuclear Charge (Zeff.)
Increasing +1
Shielding electrons +8
(0e-) +2 +3 +4 +5 +6 +7
(2e-)
(10e-)

11. Effective Nuclear Charge Trend
Increases to the RIGHT

12. 6.3
Atomic Radius
The atomic radius is one half of the distance between the nuclei of two atoms of the same element when the atoms are joined.
This diagram lists the atomic radii of seven nonmetals. An atomic radius is half the distance between the nuclei of two atoms of the same element when the atoms are joined.

13. Atomic Radius Why larger going down? Why smaller to the right?
Higher energy levels have larger orbitals
Why smaller to the right?
Increased nuclear charge without additional shielding pulls electrons in tighter

14. Atomic Radius 14

15. Atomic Radius 15

17. Atomic Radius Trend
Atomic Radius
Increases to the LEFT and DOWN

18. Ionic Radius Cations (positive ions) Anions (negative ions)
Lose electron
Become smaller than the neutral atom
Anions (negative ions)
Gain electron
Become larger than the neutral atom
Trend is the same as the atomic radius

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

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

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

22. Ionic Radius

23. Ionic Radius Trend Ionic Radius
Increases to the LEFT and DOWN within the cations and the anions
Cations (+)
Anions (-)

24. Ionization Energy
Ionization Energy = Energy required to remove an electron from a neutral atom
Smaller atoms have electrons closer to the nucleus so they are held tighter.
This means a higher ionization energy is required to take an electron away.
There are a few exceptions where where s, p, or d orbitals are filled or half-filled which creates a more stable structure

25. Ionization Energy

26. Ionization Energy Trend
First Ionization Energy
Increases UP and to the RIGHT

27. Electron Affinity What does the word ‘affinity’ mean?
Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ).
Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.

28. Electron Affinity
Increases UP and to the RIGHT

29. Electronegativity
Electronegativity is a measure of an atom’s attraction for another atom’s electrons.
ranges from 0 to 4.
Generally, metals are electron givers and have low electronegativities.
Nonmetals are are electron takers and have high electronegativities.
What about the noble gases?

30. Electronegativity
Increases UP and to the RIGHT

31. Overall Reactivity
This ties all the previous trends together in one package.
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.

32. Overall Reactivity
Your help sheet will look like this:

33. The Octet Rule
The “goal” of most atoms (except H, Li and Be) is to have an octet or group of 8 electrons in their valence energy level.
They may accomplish this by either giving electrons away or taking them.
Metals generally give electrons, nonmetals take them from other atoms.
Atoms that have gained or lost electrons are called ions.