1. Periodic Table Groups

2. Group 1: Alkali Metals All metals, except for hydrogen
React violently and quickly
Form ions with a charge of +1
One valence electron

3. Group 2: Alkaline-Earth Metals
Very reactive
Harder and denser than group 1
Form ions with a charge of +2
Two valence electrons

4. Group 3-12: Transition Metals
Malleable (flat sheets)
Ductile (made into a wire)
Lusterous (shiny)
Cannot tell their valence electrons from looking at their group #

5. F-Block: Rare Earth Metals
Plentiful in Earth’s crust, but not found in large quantities at a time
Many are man-made
Many are radioactive

6. Group 17: Halogens Contains elements in each state of matter
Extremely reactive
Form diatomic molecules
Form ions with a -1 charge
Seven valence electrons

7. Group 18: Noble Gases Extremely unreactive Do not form ions, stable
All gases
Eight valence electrons, except He

8. Periodic Trends Atomic radius Electronegativity Ions & Ionic Radius
Ionization energy
Reactivity &
Metallic properties

9. Atomic Radius
Distance from the nucleus to the outer valence electron

10. Radius
increases down a group since we add energy levels.

11. Effective Nuclear Charge-
Strong force pulling the electrons toward the nucleus (like a magnet)
This is because of the positive pull from the protons in the nucleus on the negative electrons
As you go LR, you add 1 p+ and 1 e- increasing the attractive force

12. radius decreases from L R b/c
the more e- the more pull from the p+ in the nucleus making the atom SMALLER

15. Valence (outer) electrons1 8 2 3 4 5 6 7
Re-emphasize why an element will more likely lose or gain an electrons, Octet Rule

17. Backwards again!!! UGH!
Atoms that become Cations lose electrons to become positive and smaller
Atoms that become Anions gain electrons to become negative and larger

18. Common charges – oxidation #’s1+ 2+ 3+ 4± 3- 2- 1-
Metals LOSE
Nonmetals GAIN
Re-emphasize why an element will more likely lose or gain an electrons, Octet Rule

19. Ionic Radius – size of the atom after it
gains or loses an electron
Cation- smaller than its neutral atom (because it loses electrons)
Anion- larger than its neutral atom (because it gains electrons)

20. Neutral atom is in gray
Cation (+) is in red - notice it is SMALLER than the neutral
Anion (-) is in blue - notice it is BIGGER than the neutral

21. 5 4 Let’s Practice! Name Symbo l At # e- in neutral atom e- in ion
ox #
cation
anion neither or both
Group
Period
Valence e- 5 4

23. (new flap) Electronegativity
The tendency for an atom to attract electrons to itself when it is chemically combined with another element.

24. Tug of war for shared electrons
Electronegativity
L VE of electrons
Tug of war for shared electrons
Pair of Shared
Electrons

25. Noble gases (group 18) are already stable with 8 electrons in their outer shell (NO LOVE!!)
Floozy Fluorine is the most electronegative element.
Electronegativity increases from L R
Electronegativity decreases going down

26. Electronegativity

29. First Ionization Energies (Group)
As you go down a group IE decreases since you’re adding energy levels (easier to steal)

30. First Ionization Energies (Period)
Since the effective nuclear charge increases as you go across the period (left to right), electrons are held tighter by the nucleus and require more energy to be removed (high IE = harder to steal).

31. First Ionization Energies

33. 2nd Ionization Energy
Energy required to remove a second electron from the outermost electrons.

34. Ionization Energies

36. Metallic Properties
Having properties of a metal ie, luster, malleability, ductility
Increase down a group.
Increase across a period Right to left
Francium Fr- most metallic

37. Metallic Properties

38. Reactivity
The state and degree to which a substance will react with another substance. N O E

39. Energy released when an atom gains an electron

40. Electron
The electron is the most important subatomic particle in determining physical and chemical properties.