1. The Periodic Table
The how and why

2. Atomic Size First problem where do you start measuring.
The electron cloud doesn’t have a definite edge.
Chemists get around this by measuring more than 1 atom at a time.

3. Atomic Size }
Radius
Atomic Radius = half the distance between two nuclei of a diatomic molecule (homo-nuclear molecule).

4. Trends in Atomic Size Influenced by two factors. Energy Level
Higher energy level is further away.
The effective charge from the nucleus
The greater the nuclear charge reaching the valence electrons the closer these electrons are pulled in.

5. Group trends H Li Na K Rb As we go down a group
Each atom has another energy level, so the atoms get bigger.
There are more levels in the kernel and therefore greater shielding of valence electrons (weaker attraction). Li Na K Rb

6. Periodic Trends Na Mg Al Si P S Cl Ar
As you go across a period the radius gets smaller.
Valence electrons are in the same energy level.
Greater nuclear charge gets to the valence electrons – same shielding e- but greater nuclear charge.
Outermost electrons are closer.
Row 3 elements all have 10 shielding electrons;
+11 (nuclear charges) Na Mg Al Si P S Cl Ar

7. Rb K
Overall Na Li
Atomic Radius (nm) Kr Ar Ne H 10
Atomic Number

8. Ionization Energy
The amount of energy required to completely remove an electron from a gaseous atom.
Removing one electron makes a +1 ion.
The energy required is called the first ionization energy.

9. Ionization Energy
The second ionization energy is the energy required to remove the second electron.
Always greater than first IE.
The third IE is the energy required to remove a third electron.
Greater than 1st of 2nd IE.

10. Symbol First Second Third
HHeLiBeBCNO F Ne

11. Symbol First Second Third
HHeLiBeBCNO F Ne

12. What determines IE
The greater the effective nuclear charge the greater IE.
Less distance from nucleus increases IE
Filled and half filled sublevels have lower energy, so removing them raises the IE.
Shielding: effective blocking of nuclear charge weakens the attraction of valence electrons.

13. Shielding
The electron on the outside energy level has to look through all the other energy levels to see the nucleus

14. Shielding
The electron on the outside energy level has to look through all the other energy levels to see the nucleus.
A second electron has the same shielding.

15. Group trends As you go down a group first IE decreases because
The electron is further away.
More effective shielding.

16. Periodic trends
All the atoms in the same period have the same energy level.
Same shielding.
Increasing nuclear charge
So IE generally increases from left to right.
Exceptions at full and 1/2 fill sublevels.

17. He has a greater IE than H. no shielding greater nuclear charge
First Ionization energy
Atomic number

18. further away (outweighs greater nuclear charge)
Li has lower IE than H
more shielding
further away (outweighs greater nuclear charge) H
First Ionization energy Li
Atomic number

19. greater nuclear charge removing an electron from a full s sublevel He
Be has higher IE than Li
same shielding
greater nuclear charge
removing an electron from a full s sublevel
First Ionization energy H Be Li
Atomic number

20. greater nuclear charge He
B has lower IE than Be
same shielding
greater nuclear charge
removing an electron from a partially filled p sublevel
First Ionization energy H Be B Li
Atomic number

21. First Ionization energyHe
First Ionization energy H C Be B Li
Atomic number

22. First Ionization energyHe N
First Ionization energy H C Be B Li
Atomic number

23. First Ionization energyHe N
First Ionization energy H C O Be
Breaks the pattern because in N the electron gets removed from a ½ filled p sublevel B Li
Atomic number

24. First Ionization energyHe F N
First Ionization energy H C O Be B Li
Atomic number

25. Both have full valence shells Ne has more shielding Greater distance
Ne has a lower IE than He
Both have full valence shells
Ne has more shielding
Greater distance F N
First Ionization energy H C O Be B Li
Atomic number

26. Na has a lower IE than Li Both are s1 Na has more shieldingHe Ne
Na has a lower IE than Li
Both are s1
Na has more shielding
Greater distance F N
First Ionization energy H C O Be B Li Na
Atomic number

27. First Ionization energy
Atomic number

28. Driving Force Full Energy Levels are very low energy.
Noble Gases have full energy levels.
Atoms behave in ways to achieve noble gas configuration (become isoelectronic).

29. 2nd Ionization Energy
For elements that reach a filled or half filled sublevel by removing 2 electrons 2nd IE is lower than expected.
True for s2
Alkaline earth metals form +2 ions.

30. 3rd IE Using the same logic s2p1 atoms have an low 3rd IE.
Atoms in the aluminum family form +3 ions.
2nd IE and 3rd IE are always higher than 1st IE!!!

31. Electron Affinity
The energy change associated with adding an electron to a gaseous atom.
Easiest to add to group 17 (7A).
Gets them to full energy level becomes stable – releases a large amount of energy.
Increase from left to right atoms –metals are losers so to gain an electron will only bring about a small increase in stability (if any) – small energy change.
Decrease as we go down a group.

32. Ionic Size Cations form by losing electrons.
Cations are smaller than the atom they come from.
Metals form cations.
Cations of representative elements have noble gas configuration.

33. Ionic size Anions form by gaining electrons.
Anions are bigger than the atom they come from.
Nonmetals form anions.
Anions of representative elements have noble gas configuration.

34. Configuration of Ions Ions always have noble gas configuration.
Na is 1s22s22p63s1
Forms a +1 ion - 1s22s22p6
Same configuration as neon.
Metals form ions with the configuration of the noble gas before them - they lose electrons.

35. Configuration of Ions
Non-metals form ions by gaining electrons to achieve noble gas configuration.
They end up with the configuration of the noble gas after them.

36. Group trends Li+1 Na+1 K+1 Rb+1 Cs+1 Adding energy level
Ions get bigger as you go down.
Li+1
Na+1
K+1
Rb+1
Cs+1

37. Periodic Trends N-3 O-2 F-1 B+3 Li+1 C+4 Be+2
Across the period nuclear charge increases so they get smaller.
Energy level changes between anions and cations.
N-3
O-2
F-1
B+3
Li+1
C+4
Be+2

39. Size of Isoelectronic ions
Iso - same
Iso electronic ions have the same # of electrons (same configuration)
Al+3 Mg+2 Na+1 (Ne) F-1 O-2 and N-3
all have 10 electrons
all have the configuration 1s12s22p6

40. Size of Isoelectronic ions
Positive ions have more protons than electrons so they are smaller (strong attraction for electrons).
N-3
O-2
F-1 Ne
Na+1
Al+3
Mg+2

41. Electronegativity

42. Electronegativity
The tendency for an atom to attract electrons to itself when it is chemically combined with another element.
How fair it shares.
Big electronegativity means it pulls the electron toward it.
Atoms with large negative electron affinity have larger electronegativity.

43. Group Trend
The further down a group the farther the electron is away and the more electrons an atom has – the lower the electronegativity value.

44. Periodic Trend Metals are at the left end.
They let their electrons go easily
Low electronegativity
At the right end are the nonmetals.
They want more electrons.
Try to take them away.
High electronegativity.

45. Ionization energy, electronegativity
Electron affinity INCREASE

46. Atomic size decreases, shielding constant
Atomic size increases
Shielding increases
Ionic size decreases