1. Fill in the arrows on the blank periodic table with trends using your graphs made during last class.
Periodic Trends

2. Short hand e- configurations
To save time (and energy) electron configurations can be written in a shortened version.
Rather than write 1s22s22p63s23p64s23d104p2 we can abbreviate the configuration:
[Ar] 4s23d104p2
Here’s how…

3. Short-hand electron configurations
1. Identify the element
Ge (Germanium)
2. Find the last noble gas in the configuration
Ar (argon)
3. Write the symbol in parentheses
[Ar]
4. Starting at the next element, finish the electron configuration
[Ar] 4s23d104p2

4. Practice Write the short-hand for : p (1s22s22p63s23p3) Te K
[Ne] 3s23p3 Te
[Kr] 5s24d105p4 K
[Ar] 4s1

5. To find Valence electrons
Because valence electrons are only in the s and p sublevels, you can count the representative groups (or Group A) elements.
For example: Cl is in group 7A, therefore it has 7 valence electrons
Find the valence electrons for: Se
In group 6A, 6 valence electrons Cs
In group 1A, 1 valence electron Al
In group 3A, 3 valence electrons

6. Atomic Radius
Atomic Radius - radius of the atom, an indication of the atom's volume.
Looking at your graph, what is the trend down groups?
Across periods?

8. Atomic Radius
Period - atomic radius decreases as you go from left to right across a period.
Why?
Stronger attractive forces in
atoms (as you go from left to
right) between the opposite
charges in the nucleus and
electron cloud cause the atom to
be 'sucked' together a little tighter.

9. Atomic Radius Group - atomic radius increases as you go down a group.
Why?
There is a significant jump in the size of the nucleus (protons + neutrons) each time you move from period to period down a group.
Additionally, new energy levels of elections clouds are added to the atom as you move from period to period down a group, making the each atom significantly more massive, both is mass and volume.

10. Atomic Radius Decreases
Atomic Radius Increases

11. Atomic Radius Practice
3rd period element with the largest radius Na
6th period element with the smallest radius Rn
Element with the smallest radius He
Which is larger; Sr, Cs, Rb or Ca? Cs

12. Ionization Energy
Ionization Energy - the amount of energy required to remove the outmost electron.
Looking at your graph, what is the trend going down a group?
Across a period?

14. Ionization Energy
Period - ionization energy increases as you go from left to right across a period.
Why?
Elements on the right of the chart want to take others atom's electron (not given them up) because they are close to achieving the octet.
This means it will require more energy to remove the outer most electron.
Elements on the left of the chart would prefer to give up their electrons so it is easy to remove them, requiring less energy (low ionization energy).

15. Ionization Energy
Group - ionization energy decreases as you go down a group.
Why?
The shielding affect makes it easier to remove the outer most electrons from those atoms that have many electrons (those near the bottom of the chart).

16. Atomic Radius Decreases
Ionization Energy Increases
Atomic Radius Increases
Ionization Energy Decreases

17. Ionization Energy Practice problems
Element with the highest ionization energy: He
Which has a lower ionization energy; I, Cl, F, Br? I
Element in the 2nd period with the highest ionization energy: Ne

18. Atomic Radius of ions
Negatively charged ions are larger than their neutral atoms
More electrons without adding protons has more repulsion between outer electrons
Positively charged ions are smaller than their neutral atoms
Loses a shell when loses valence electrons

19. Ionization “bumps” in graph
Compare Li with Be
Why does Be have a lower IE than Li?
Consider the electron configuration and Aufbau diagrams

20. Electronegativity
Electronegativity - tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element.
Looking at your graph, what is the trend as you go down a group?
Across a period?

22. Electronegativity
Period - electronegativity increases as you go from left to right across a period.
Why?
Elements on the left of the period table have 1 -2 valence electrons and would rather give those few valence electrons away (to achieve the octet in a lower energy level) than grab another atom's electrons. As a result, they have low electronegativity.
Elements on the right side of the
period table only need a few
electrons to complete the octet, so
they have strong desire to grab
another atom's electrons.

23. Atomic Radius Decreases
Ionization Energy Increases
Electronegativity Increases
Atomic Radius Increases
Ionization Energy Decreases

24. Electronegativity Group - electronegativity decreases as you
go down a group.
Why?
Elements near the top of the period table have few electrons to begin with; every electron is a big deal. They have a stronger desire to acquire more electrons. Elements near the bottom of the chart have so many electrons that losing or acquiring an electron is not as big a deal.

25. Electronegativity
This is due to the shielding affect where electrons in lower energy levels shield the positive charge of the nucleus from outer electrons resulting in those outer electrons not being as tightly bound to the atom.

26. Atomic Radius Decreases
Ionization Energy Increases
Electronegativity Increases
Atomic Radius Increases
Ionization Energy Decreases
Electronegativity decreases

27. Electronegativity Practice Problems
Element with the lowest electronegativity: Fr
4th period element with the highest electronegativity: Br
Which has a lower electronegativity: Li, Be, Na, or Mg? Na

28. Reactivity
Reactivity: refers to how likely or vigorously an atom is to react with other substances.
This is usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom's electrons (electronegativity) because it is the transfer/interaction of electrons that is the basis of chemical reactions.

29. Metal Reactivity
Period - reactivity decreases as you go from left to right across a period.
Group - reactivity increases as you go down a group
Why?
The farther to the left and down the periodic chart you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity.

30. Metal Reactivity Decreases
Metal Reactivity Increases

31. Non-metal Reactivity
Period - reactivity increases as you go from the left to the right across a period.
Group - reactivity decreases as you go down the group.
Why?
The farther right and up you go on the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electron.

32. Metal Reactivity Decreases
Non-metal Reactivity increases
Metal Reactivity Increases
Non-metal Reactivity decreases

33. Reactivity Practice Metal with the lowest reactivity:
Nonmetal with the highest reactivity:
F (noble gases are stable and have a full octet)
Which has a higher reactivity: Se, S, As, P? S
Which has a lower reactivity: Cd, In, Hg, Tl? In