1. The Periodic Table
The how and why

2. History Mid 1800 - molar masses of elements were known (Cannizarro).
Russian scientist Dmitri Mendeleev taught chemistry in terms of properties.
Wrote down the elements in order of increasing mass.
Found a pattern of repeating properties.

3. Mendeleev’s Table-1870
Grouped elements in columns by similar properties.
Found some inconsistencies
the properties were more important than the mass, so switched order.
Found some gaps.
Must be undiscovered elements.
Predicted their properties before they were found.

4. The Modern table Order is by increasing atomic number (Mosley).
Elements are still grouped by properties.
Similar properties are in the same column.
Seaborg moved bottom two rows out(~1944)
Noble gas elements added later.
Mendeleev didn’t know about - nonreactive.

5. Horizontal rows are called periods
There are 7 periods

6. Vertical columns are called groups.(18)
Elements are placed in groups by similar properties.
Also called families

7. The elements in these groups are called the representative elements (main groups)1 18 2 13 14 15 16 17

8. The center group is called the transition elements
These are called the inner transition elements and they belong here

9. Group 1 are the alkali metals
Group 2 are the alkaline earth metals

10. Group 17 is called the Halogens
Group 18 are the noble gases

11. S- block Alkali metals all end in s1
Alkaline earth metals all end in s2
(could include He but it fits better later)
He has the properties of the noble gases.

12. Transition Metals -d block

13. The P-block p1 p2 p3 p4 p6 p5

14. F - block f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13
inner transition elements f1 f5 f2 f3 f4 f6 f7 f8 f9
f10
f11
f12
f14
f13

15. 1 2 3 4 5 6 7 4f 5f
f orbitals start filling at 4f

16. D orbitals fill up after previous energy level so first d is 3d even though it’s in row 4.1 2 3 4 5 6 7 3d

17. Electron Configurations repeat
The shape of the periodic table is a representation of quantum mechanics.
For example:
At the end of each period/row the element is very unreactive. **Noble Gas**
This elec. config. is stable (called full).
basis for kernel (shortcut).

18. Why?
If two atoms approach each other, it is the electron clouds that interact.
Most importantly, the outside orbitals. (valence electrons)
The orbitals fill up in a regular pattern.
The electron configuration repeats.
The properties of atoms repeat.
Called periodicity.

19. Valence Electrons Electrons at the highest energy level.
Includes s + p orbitals only.
Never more than 8.

20. Atomic Size/Radius First problem - where do you start measuring?
The electron cloud doesn’t have a definite edge, but the nucleus can be located easily.
Measure more than 1 atom at a time.

21. Atomic Size }
Radius
Atomic Radius = half the distance between two nuclei of a diatomic molecule.

22. Trends in Atomic Radius
Influenced by two factors.
Energy Level
Higher energy level is further away.
Charge on nucleus
More charge pulls electrons in closer.

23. Group trends H Li Na K Rb As we go down a group
Each atom has another energy level,
So the atoms get bigger. Li Na K Rb

24. Period Trend Na Mg Al Si P S Cl Ar
As you go across a period the radius gets smaller.
Same energy level.
More nuclear charge.
Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

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

26. 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.

27. 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 or 2nd IE.

28. Symbol First Second Third
HHeLiBeBCNO F Ne

29. What determines IE nuclear charge Distance from nucleus
Filled and half filled orbitals have more stability
Shielding

30. Shielding
The electron(s) on the outer energy level has to look through all the other energy levels to see the nucleus.
More levels = less attraction

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

32. Period 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 full orbitals.

33. First Ionization energyHe
He has a greater IE than H.
same shielding
greater nuclear charge H
First Ionization energy
Atomic number

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

35. First Ionization energyHe Ne
Ne has a lower IE than He
Both are full,
Ne has more shielding
Greater distance F N
First Ionization energy H C O Be B Li
Atomic number

36. 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

37. First Ionization energy
Atomic number

38. Driving Force
Full Energy Levels have lower potential energy. (means stability increases)
Noble Gases have full orbitals.
Atoms behave(react) in ways to achieve noble gas configuration.

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

40. Ionic size Anions form by gaining electrons.
Anions are negatively charged.
Anions are bigger that the atom they come from.(more repulsion)
Nonmetals form anions.
Anions of representative elements have noble gas configuration.

41. Configuration of Ions Ions 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
Metals lose electrons

42. 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.
Nonmetals gain electrons.

43. 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

44. 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

45. Size of Isoelectronic ions
Iso - same
Iso electronic ions have the same # of electrons
Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3
all have 10 electrons
all have the configuration 1s22s22p6

46. Size of Isoelectronic ions
Positive ions have more protons so they are smaller.
N-3
O-2
F-1 Ne
Na+1
Mg+2
Al+3

47. Electronegativity

48. Electronegativity
The tendency for an atom to attract a pair of electrons to itself when it is bonded with another atom.
How fair it shares.
Big electronegativity means it pulls electrons toward it.
Values are not given for noble gases.

49. Group Trend Electronegativity decreases as you go down a group
Reason: More occupied energy levels (shielding)

50. Period Trend Metals tend to release electrons easily
Low electronegativity
Nonmetals tend to pull electrons in.
High electronegativity.
Reason: Nuclear charge

52. Ionization energy, electronegativity
Electron affinity INCREASE

53. Atomic size increases, shielding constant
Ionic size increases