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

2. Atomic Size Size goes UP on going down a group.
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.

3. Atomic Size
Size decreases across a period owing to 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 ANIONS are LARGER than the atoms from which they come.-
, 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. Mg+ (g) + 1451 kJ ---> Mg2+ (g) + e-
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.

13. Trends in Ionization Energy
IE increases across a period because the positive charge increases.
Remember, low energy means easier. Removing an electron from Li than F
Metals lose electrons more easily than nonmetals.
Nonmetals lose electrons with difficulty (they like to GAIN electrons).

14. Trends in Ionization Energy
IE increases UP a group
Harder to remove an electron closer to the nucleus
Because size increases (Shielding Effect)

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

16. Electron Affinity
Definition - the energy change associated with the addition of an electron
The sign of the electron affinity can be confusing.
When an atom becomes less stable upon gaining an electron, its potential energy increases, which implies that the atom gains energy as it acquires the electron. In such a case, the atom's electron affinity is positive.
An atom with a negative electron affinity is far more likely to gain electrons.

17. Periodic Trend: Electron Affinity

18. ELECTRON AFFINITY: ACROSS A PERIOD
Electron affinities becoming increasingly negative from left to right. (Easier to add e- to Cl than Na)
Just as in ionization energy, this trend conforms to and helps explain the octet rule. The octet rule states that atoms with close to full valence shells will tend to gain electrons.
Such atoms are located on the right of the periodic table and have very negative electron affinities, meaning they give off a great deal of energy upon gaining an electron and become more stable.
Be careful, though: the noble gases, located in the extreme right hand column of the periodic table do not conform to this trend. Noble gases have full valence shells, are very stable, and do not want to add more electrons: noble gas electron affinities are positive.
Similarly, atoms with full subshells also have more positive electron affinities (are less attractive of electrons) than the elements around them.

19. ELECTRON AFFINITY: DOWN A PERIOD
DECREASES AS YOU GO DOWN (easier to add to Na than to K)
Electron affinities change little moving down a group, though they do generally become slightly more positive (less attractive toward electrons).
This is because there is less attractions between the electrons you are trying to add and the nucleus because there is more distance between the electron shell and nucleus
The biggest exception to this rule are the third period elements, which often have more negative electron affinities than the corresponding elements in the second period. For this reason, Chlorine, Cl, (group VIIa and period 3) has the most negative electron affinity.

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. Summary of Periodic Trends

25. Periodic Trends Summary
As Move to the right
Reason?
As Move down
Atomic Radius
Gets Smaller
Effective Nuclear Charge Increases with no change in shielding
Gets Larger
More energy levels filled, more shielding between nucleus and electrons
Ionization Energy
Energy Increases (harder to remove)
Octet rule: easier to remove unpaired e-
Coulomb’s Law (has greater Effective Nuclear Charge)
Energy Decreases (easier to remove)
-Electrons are farther and have more shielding from Nucleus and easier to remove according to Coloumb’s Law
Electron Affinity
Energy Decreases (easier to add)
Octet rule: easier to add e- to almost full orbitals
Greater Effective nuclear charge, more attraction to add electron
Energy Increases (harder to add)
Electrons are farther and have more shielding from nucleus and will have smaller attractive forces to nucleus
Electro-negativity
Decreases (Easier to attract electrons)
Effective nuclear charge increases=more attracting to bonding electrons
Increases (Harder to attract electrons)
- Electrons are farther away, more shielding, less attraction to nucleus by bonding electrons