1. Still all about + and – charges!
Periodic Trends
Still all about + and – charges!

2. Which of the following atoms is the largest?Mg Ca Sr
How the fluorine should I know?

3. Which of the following atoms is the largest?Mg Al Cl

5. Periodic for a Reason
It is a “periodic” table because of the “periodic trends” that make it up.
We saw that the electron configurations are periodic (s-block, p-block, d-block, etc.) and I suggested that the Chemistry associated with the atoms followed the electron configuration.

6. Other Periodic Properties
Electron configuration isn’t the only periodic trend in the table.
The other important trends are all rationalizable based on the most important trend – atomic size!

7. Atomic Size
Suppose we want to compare the atomic radius of two atoms on the periodic table, for example Na and K.
Which would you think is bigger?

9. What does an atom look like?
Bohr Model

10. The Bohr Model So, what determines the size of the atom? e- n p p
Nucleus e- n p p n e- n p n e-
So, what determines the size of the atom?

11. The size is determined by…
…where the last electron lies.
So, which is bigger Na or K?
K – it is Na with a whole extra shell of electrons!

13. Na vs. K Na: 1s22s22p63s1 K: 1s22s22p63s23p64s1
This is a general trend. As you go down a column in the periodic table, the atomic radius increases.

14. What about the rows?
Suppose I ask the same question about Na and Mg; which is larger?

16. Na vs. Mg Na: 1s22s22p63s1 Mg: 1s22s22p63s2 Does this help us any?
They have the same valence shell (n=3). The same highest orbital (3s).
Does this mean they are the same size? 

17. Those pesky charges What are the differences between Na and Mg?
One extra electron
One extra proton
Does that help?

18. The Bohr Model - Na
11 p
12 n
full e-

19. The Bohr Model - Mg
12 p
12 n
full e-

20. The Bohr Model – Na vs. Mg
12 p
12 n
full e-
11 p
12 n
full e-

21. Na vs. Mg Same outer shell of the electrons (- charge).
More + charges in nucleus.
What do you think happens?
Mg is actually slightly smaller than Na due to the extra + charge in the nucleus pulling the – electrons in closer!

22. Periodic Trend – atomic radius
This pair of observations describes the general trend of atomic size:
Larger as you go down a column (large effect)
Smaller as you go across the row (small effect)
Note: There are exceptions due to special cases (1/2 full orbitals, etc.)

24. Rank from Largest to smallest
Sn, In, Ga, As
In Sn As Ga
In Sn Ga As
In Ga Sn As

25. Periodic Trend – atomic radius
Both of these trends are related to the charged species:
Larger as you go down a column (more electron shells – adding layers)
Smaller as you go across the row (stronger attraction between more + protons and the outer electrons)

26. Other Trends
Understanding the trend in atomic radius and keeping the charge issues in mind make it easy to understand and predict some other periodic trends.

27. Ionization Energy
Ionization energy is the amount of energy required to remove an electron from an atom:
Na + energy  Na+ + e-
(You simply raise the electron from n=valence to n=∞)

28. Ionization Energy Na + energy  Na+ + e-
If you are going to remove an electron, what is the relevant issue?
Charge – what a surprise!
What Charge?
The nuclear charge! - At least in part

29. Na vs. Mg Compare the Ionization Energy of Na to that of Mg.
Which do you think would be larger?
Na or Mg
Why?
They have the same outer orbital, but Mg has a larger nuclear charge (sound familiar) – Mg should have the larger ionization energy!

30. Larger Ionization Energy
Sodium
Magnesium
I have no frigging idea

31. Na vs. K Compare Na to K, what is relevant? K is bigger than Na – why?
Because K has more shells and larger radius. Electrons are farther from nucleus.
Same argument for ionization energy. Electron, farther away, less attracted to nucleus, smaller ionization energy!

32. Ionization Energy vs. Atomic Radius
It is the same arguments, with the same results.
Ionization energy has a trend that tracks the radius. Smaller atoms, larger ionization energy!

33. Electron Affinity
Electron Affinity is complementary to ionization energy. Rather than give up an electron, the atom receives it:
Na + e- → Na-
What will determine if an atom wants an electron?
Attraction for the nucleus – so electron affinity will also track atomic radius

34. Electron Affinity Na vs. K
K will have a smaller electron affinity – K is larger, a new electron is farther away, less attracted to nucleus
Na vs. Mg
Na will have a smaller electron affinity – same radius, smaller charge, less attraction for the electron.

36. Electronegativity
Electronegativity is the ability of an atom to attract electrons to itself. (Kind of like electron affinity, but on a different scale)
Electronegativity is important in predicting whether a bond is ionic or covalent.
Electronegativity will have the same trend as electron affinity.