1. Orbitals and the Periodic Table
Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental)
s orbitals
d orbitals
p orbitals
f orbitals

2. General Periodic Trends
Atomic and ionic size
Ionization energy
Electronegativity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.

3. Atomic Size Size goes UP on going down a group. Why?
Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.
Size goes DOWN on going across a period.

4. Atomic Size Size decreases across a period Why?
The increase in the positive charge from the protons. Each added electron feels a greater and greater + charge because the protons are pulling in the same direction, where the electrons are scattered.
Large
Small

5. Which is Bigger?
Na or K ?
Na or Mg ?
Al or I ?

6. Ion Sizes Does the size go
up or down when losing an electron to form a cation?

7. Ion Sizes CATIONS are SMALLER than the atoms from which they come.Li +
, 78 pm
2e and 3 p
Forming a cation.
Li,152 pm
3e and 3p
CATIONS are SMALLER than the atoms from which they come.
The electron/proton attraction has gone UP and so size DECREASES.

8. Ion Sizes
Does the size go up or down when gaining an electron to form an anion?

9. Ion Sizes Forming an anion.-
, 133 pm
10 e and 9 p
F, 71 pm
9e and 9p
Forming an anion.
ANIONS are LARGER than the atoms from which they come.
The electron/proton attraction has gone DOWN and so size INCREASES.
Trends in ion sizes are the same as atom sizes.

10. Trends in Ion Sizes
Figure 8.13

11. Which is Bigger?
Cl or Cl- ?
K+ or K ?
Ca or Ca+2 ?
I- or Br- ?

12. Put these in order of increasing size.
Na+1
Mg+2
Al+3
P-3
S-2
Cl-1 Ar
K+1
Ca+2

13. Answer
Al+3
Mg+2
Na+1
Ca+2
K+1 Ar
Cl-1
S-2
P-3
Increasing size

14. Ionization Energy
IE = energy required to remove an electron from an atom (in the gas phase).
Mg (g) kJ ---> Mg+ (g) + e-
This is called the FIRST ionization energy because we removed only the OUTERMOST electron
Mg+ (g) kJ ---> Mg2+ (g) + e-
This is the SECOND IE.

15. Successive ionization energy

16. Explaining the trends in ionization energy

17. Trends in Ionization Energy
IE increases across a period. Why?
Because the positive charge increases and electrons are being put at about the same distance from the nucleus.
Metals lose electrons more easily than nonmetals.
Nonmetals lose electrons with difficulty (they like to GAIN electrons).

18. Trends in Ionization Energy
IE decreases down a group
Because size increases (Shielding Effect)

19. Which has a higher 1st ionization energy?
Mg or Ca ?
Al or S ?
Cs or Ba ?

20. Electronegativity, 
 is a measure of the ability of an atom in a molecule to attract electrons to itself.
Concept proposed by
Linus Pauling

21. Periodic Trends: Electronegativity
In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity increases UP a group of elements.
In a period: More protons, while the energy levels are the same, means atoms can better attract electrons. So, electronegativity increases RIGHT in a period of elements.

22. Electronegativity

23. Which is more electronegative?
F or Cl ?
Na or K ?
Sn or I ?

24. How is electronegativity used?
To predict the nature of the bond between two atoms
Is the bond polar or nonpolar?
The larger the difference, the more likely that the bond is ionic
NOTE: A large difference in electronegativity does not equate with a strong bond!!!