1. Chapter 4: The Periodic Table
Section 3: Trends in the Periodic Table

2. Key Terms Ionization energy – removes an electron from an atom or ion
Electron shielding – when inner electrons cancel some of the + charge of the nucleus & lessen its attraction of outermost electrons
Trend – predictable change in a particular direction

3. Key Terms 2
Atomic radius – depends on volume occupied by atom’s electron cloud
Bond radius – half the distance between nuclei of atoms that are bonded together
Electronegativity – measure of how much an atom in a chemical com-pound can attract electrons

4. Key Terms 3
Electron affinity – change in energy when a neutral atom gains an electron
There are no existing bonds as with electronegativity.

5. Things To Know/Answer
What are the periodic trends in ionization energy, & how are they affected by atomic structures?
Answer these questions for atomic radius, electronegativity, ionic size, electron affinity, melting points (mp) & boiling points (bp).

6. Periodic Trends
All in a group can be explained by electron configurations.
Reactivity in alkali metals rise from top to bottom in group 1.

7. Periodic Trends 2
When one adds enough energy to overcome the attraction between protons & electrons in an atom, it becomes charged (ion) and an electron escapes. See page 133 Figure 16.
A + ionization energy → A+ + e-

8. Ionization Energy
Ionization energy decreases down-ward in a group because the number of energy levels increases downward.
Increasing energy levels have increasing distance from the nucleus where the positive charge that attracts e- in the atom is.

9. Ionization Energy 2
The farther an e- is from the nu-cleus, the less attraction protons there have on the e-.
Also, the higher an e-’s energy level is the more full levels of e-’s there are between it & the nucleus.

10. Ionization Energy 3
These e-’s “in the middle” reduce the positive attraction that extends from the nucleus through the atom (electron shielding).
Outermost electrons are less tightly held for this reason.

11. Ionization Energy 4
Ionization energy increases as you move across a period. See Figure 17 on page 134.
This is because protons increase 1 at a time rightward on a row, but the energy level stays the same.

12. Ionization Energy 5 / Atomic Radius
Increasing e-’s get crowded & repel each other; this counteracts the positive attrac-tion of the nucleus.
Here, electrons can only get so close before increased positive attraction can-not overcome e- to e- repulsion.
For this reason, atomic radius stops decreasing rightward in a row.

13. Atomic Radius 2
Increasing atomic # across a row pro-duces a much bigger rise in positive attraction than rise in distance of e-’s from the nucleus.
This pattern also causes decreased atomic radius left to right on a row.

14. Atomic Radius 3
Rising + attraction also pulls electrons closer to the nucleus.
Added inner energy levels are present downward in groups & add distance from the nucleus.

15. Atomic Radius 4
Electron shielding also blocks rises in + attraction & yields similar attraction down the group.
These effects cause increasing atomic radius downward in a group.

16. Electronegativity
Electronegativity is relative attraction of electrons by nuclei in bonded atoms where they “play tug of war” with their shared electrons.
Linus Pauling, one of America’s most famous chemists, made a scale of electronegativity values.

17. Electronegativity 2
In the scale, he assigned F 4.0 since it attracts electrons in bonds most then Pauling calculated values for other elements relative to this one.
Electronegativity decreases down a group mostly because higher energy levels are farther from the nucleus.

18. Electronegativity 3
Nuclei cannot attract valence electrons in these distant energy levels well.
For this reason, an element like Cs has a nucleus w/ more protons but weak attraction of a valence electron on its 6th energy level.

19. Electronegativity 4
However, an element like Li attracts a valence electron on its 3rd energy level more strongly.
This makes Li more electronegative than Cs.
Electronegativity increases sharply rightward across a period.

20. Electronegativity 5
This trend arises because no change in electron shielding happens across a row since no electrons get added to inner energy levels.
As atomic # rises quickly, nuclear charge does also and can attract bond electrons much more strongly.

21. Electronegativity 6
Adding inner electrons downward in a group increases electron shielding.
This keeps effective nuclear charge mostly the same.

22. Electronegativity 7 Slight drops in electronegativity result.
This is b/c distance from the nucleus is the key factor not nuclear charge.
Slight rises in distance from the nucleus downward in a group have much less effect than boosts in nu-clear charge rightward across a row.

23. Other Periodic Trends
Effective nuclear charge & electron shielding explain most periodic trends including ionic size & electron affinity.
Ionic size follows trends of atomic radii for the same reasons.

24. Other Periodic Trends 2
Metals tend to lose one or more electrons & become cations (+ions); whereas, nonmetals tend to gain e-’s & form anions (-ions).
Electron affinity follows electronega-tivity trends (decrease down a group but increase right across a series) for the same reasons.

25. Other Periodic Trends 3
Mp & bp do not generally rise or fall but reach 2 different peaks as d & p orbitals fill.
For example, Cs has low mp & bp b/c it only has 1 valence electron for bonding; it is far left in 6th period.

26. Other Periodic Trends 4
As the electron # increases across a row, more bonds can form & require more energy to break.
This effect peaks near the middle of d-block elements at W and Re b/c the d orbitals are half filled.

27. Other Periodic Trends 5
Further e-’s pair in d orbitals beyond W & decrease the # of unpaired e-’s that help strengthen bonds between atoms by forming multiple bonds.
More rightward, Hg & Rn have much lower mp & bp b/c d orbitals are full.

28. Other Periodic Trends 6
Past Hg, mp & bp rise again as elec-trons start filling p orbitals until these are half filled.
Beyond half filled status, mp & bp drop again b/c the p orbitals get full and unable to help strengthen bonds.
By Rn, p orbitals are full also so mp & bp are unusually low.