1. Periodic Trends

2. Periodic Trends Atomic Radius Ionic Radius Ionization Energy
Electron Affinity
Electronegativity

3. Atomic Size or Radius
The distance from the center of the nucleus to the edge of the atom.
© 1998 LOGAL

4. 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.

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

6. 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
(blocking effect?)

7. Shielding
The electron on the outermost energy level has to look through all the other energy levels to see the nucleus.
Second electron has same shielding, if it is in the same period

8. 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

9. Periodic Trends Na Mg Al Si P S Cl Ar
As you go across a period, the radius gets smaller.
Electrons are adding to the SAME energy level. The effect of adding 1 proton to the nucleus overpowers the effect of adding 1 electron to the cloud. Nuclear charge increases, and pulls electrons in.
Therefore, the atoms get SMALLER. Na Mg Al Si P S Cl Ar

10. Atomic Radius
Atomic Radius
Increases to the LEFT and DOWN

11. Examples
Which atom has the larger radius?
Be or Ba
Ca or Br Ba Ca

12. Trends in Ionic Size Cations form by losing electrons.
Cations are smaller that the atom they come from.
Metals form cations.
Cations of representative elements have noble gas configuration.

13. Ionic size Anions form by gaining electrons.
Anions are bigger that the atom they come from.
Nonmetals form anions.
Anions of representative elements have noble gas configuration.

14. Configuration of Ions Ions always have noble gas configuration.
Na is: 1s22s22p63s1
Forms a 1+ ion: 1s22s22p6
Same configuration as neon.
Metals form ions with the configuration of the noble gas before them - they lose electrons.

15. Configuration of Ions
Non-metals form ions by gaining electrons to achieve noble gas configuration.
They end up with the configuration of the noble gas after them.

16. S or S2- Al or Al3+ S2- Al Examples
Which particle has the larger radius?
S or S2-
Al or Al3+
S2- Al

17. Ionization Energy Removing one electron makes a 1+ ion.
The energy required to remove one electron from a neutral atom in the gaseous state.
Removing one electron makes a 1+ ion.
The energy required to remove the first electron is called the first ionization energy.

18. Ionization Energy First Ionization Energy
trend is opposite of atomic radius
Increases UP and to the RIGHT

19. Ionization Energy Down a group:
Electrons are located in higher energy levels, farther from the nucleus, so they are easier to remove.
Also, the “shielding effect” is increasing (more electrons from lower energy levels blocking nuclear charge), making it even easier to remove an electron.
Across a period:
Electrons are located in the SAME energy level. Nuclear charge is increasing.
Therefore, electrons are harder to remove.

20. Ionization Energy Successive Ionization Energies
Large jump in I.E. occurs when a CORE e- is removed.
Mg 1st I.E kJ
2nd I.E. 1,445 kJ
Core e- 3rd I.E. 7,730 kJ

21. Ionization Energy Successive Ionization Energies
Large jump in I.E. occurs when a CORE e- is removed.
Al 1st I.E kJ
2nd I.E. 1,815 kJ
3rd I.E. 2,740 kJ
Core e- 4th I.E. 11,600 kJ

22. Ionization Energy
The 2nd I.E. is always higher than the 1st, the 3rd I.E. is always higher than the 2nd, and so on.
Why? Removing each electron decreases the “shielding effect” around the nucleus. Nuclear charge increases, making it harder to remove another electron.

23. Ionization Energy
The noble gases have extremely high ionization energies.
Why? Due to the stability that comes with filled s and p orbitals.
Therefore, elements that have achieved noble gas configurations (by gaining or losing electrons) also have very high I.E.’s.

24. Examples
Which atom has the higher 1st I.E.?
N or Bi
Ba or Ne N Ne

25. Electron Affinity
The energy given off when an electron is gained by a neutral atom.

26. Electron Affinity Electron Affinity trend is same as ionization energy
Increases UP and to the RIGHT

27. Electron Affinity
Energy is usually given off when an atom gains an electron.
These values for E.A. are negative.
Sometimes, an atom must be “forced” to accept an electron through the addition of energy.
These values for E.A. are positive.
Adding a second electron to an already negative ion is always more difficult, so 2nd E.A.’s are always positive!

28. Electron Affinity Down a group:
Electrons are located in higher energy levels, farther from the nucleus. Nuclear charge is decreased, “shielding effect” is increased.
Therefore, electrons add with greater difficulty
Across a period:
Electrons are located in the SAME energy level. Nuclear charge is increasing.
Therefore, electrons add more easily

29. The group most likely to…
LOSE electrons?
Alkali metals (group 1)
Lowest ionization energy!
GAIN electrons?
Halogens (group 17)
Highest (most negative) electron affinity!

30. Electronegativity
The ability of an atom in a chemical compound to attract electrons.
trend is same as electron affinity

31. Electronegativity Fluorine is the most electronegative element!
Fluorine is assigned an electronegativity value of 4.0 on the Pauling Scale.
All other elements are assigned values relative to this.

32. Electronegativity
In general, elements that tend to lose electrons to form positive ions (metals) have lower e-neg’s.
Elements that tend to gain electrons to form negative ions (nonmetals) have higher e-neg’s.