A.
Patterns in element properties were recognized
1.
Dobereiner (1780-1849) found “triads” of similar elements: Cl, Br, I
2.
Newlands suggested in 1864 that elements should be arranged in
“octaves” because similarities occurred every 8 th element
B.
The Modern Periodic Table
1.
The German Meyer (1830-1895) and Russian Mendeleev (1834-1907) independently developed the current arrangement of elements
2.
Mendeleev predicted the properties of “missing” elements

3.
Mendeleev’s Periodic Table a.
Left blank spaces for the elements that he predicted would be discovered b.
Arranged elements by mass, instead of by atomic number c.
Scientific Law: predictive; Not a Theory: no explanation why d.
Quantum Mechanics is the Theory that explains the Periodic Table

1. The atomic number of an element tells us how many protons or e- it has
2. As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals.
2p y
2p z
H configuration 1s 2s 2p x
1s 1 ↑
He 1s 2 ↑ ↓
Li 1s 2 2s 1 ↑ ↓ ↑
Be
B
C
N
1s
1s
1s
1s
2
2
2
2
2s
2s
2s
2s
2
2
2
2
2p
2p
2p
1
2
3
↑ ↓ ↑ ↓
↑ ↓ ↑ ↓ ↑
↑ ↓ ↑ ↓ ↑ ↑
↑ ↓ ↑ ↓ ↑ ↑ ↑
O
F
Ne
1s
1s
1s
2
2
2
2s
2s
2s
2
2
2
2p
2p
2p
4
5
6
↑ ↓ ↑ ↓ ↑ ↓ ↑ ↑
↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑
↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓

3. Hund’s Rule : The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals.
C configuration 1s 2s 2p x
1s 2 2s 2 2p 2
2p
↑ ↓ ↑ ↓ ↑ ↑ y
N 1s 2 2s 2 2p 3
2p
↑ ↓ ↑ ↓ ↑ ↑ ↑ z
O 1s 2 2s 2 2p 4 ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↑
4. We can use the previous Noble Gas as an abbreviation to indicate filled inner orbitals a. Na = 1s 2 2s 2 2p 6 3s 1 b. Cl = [Ne]3s 2 3p 5 or [Ne]3s 1 c. Ca = [Ar]4s 2 d. Rb = [Kr]5s 1

5.
Complications with Multi-electron Atoms a.
-
Coulomb’s Law for relationship between charged particles
Like charges (q
1
, q
2
E

4 q
1 q
 o
2 r for two electrons) repel, weakens as r increases
Opposite charges attract (nucleus attraction to an electron)
Attraction increases as q increases: 2 + nucleus more attractive than 1 + b.
Shielding = reduced charge attraction due to electrons that are closer c.
Penetration = ability to get closer to nucleus than inner electrons

B.
1.
Valence electrons are electrons in the outermost shell only
2.
All other electrons are referred to as core electrons
3.
Only the valence electrons are involved in reactivity and bonding a.
O has 6 valence electrons b.
Na has 1 valence electron c.
Cl has 7 valence electons
4.
Elements in the same group (column) of the periodic table have the same number of valence electrons (and hence, reactivity and bonding) a.
Li (2s 1 ), Na(3s 1 ), K(4s 1 ) b.
F(2s 2 2p 5 ), Cl(3s 2 3p 5 ), Br(4s 2 4p 5 )

1.
From energy ordering, 4s fills before 3d
2.
3d can hold 10 e- (5 d-orbitals)
3.
Transition elements have e- in d-orbitals a.
Sc = [Ar]4s 2 3d 1 b.
Mn = [Ar]4s 2 3d 5 c.
Zn = [Ar]4s 2 3d 10
4.
Not all transition elements follow simple order; the reasons are complex.
You may have to look up the e- configuration a.
Cr = [Ar]4s 1 3d 5 b.
Cu = [Ar]4s 1 3d 10 c.
?Zn
2+ = [Ar]4s 2 3d 8 d.
?Zn
2+ = [Ar]4s 0 3d 10 e.
Paramagnetic : having unpaired electrons (attracted to magnetic field) f.
Diamagnetic : having all paired electrons (weakly repelled by magnetic field)

1.
Representative Elements (main group): fill s and p orbitals (Na, Al, N)
2.
Transition Elements : fill d orbitals (Fe, Co, Ni)
3.
Lanthanide and Actinide Series (inner transition elements): fill 4f and 5f orbitals (Eu, Am, Es)

E.
1.
Chemical properties are largely determined by number of valence electrons
2.
Quantum Mechanics explains filling of shells and sub-shells to give charges
3.
We now know more about why ions have specific charges
4.
Transition metal charges are variable in part because of d/s energy similarity

A.
Ionization Energy (I.E.) = The quantity of energy required to remove an electron from the gaseous atom or ion.
1.
X(g) X + + e-
2.
It is easiest to take away the first e-, more energy needed for others
3.
Greater charge makes it harder to extract e-
4.
First e- comes from the outermost (weakest bound) orbital
5.
I. E. increases from left to right on periodic table because larger elements have larger +/- attraction for electrons
6.
I. E. decreases down a group because outer e- becomes more weakly bound

7.
Exceptions to Ionization Energy Trends a.
1s electrons shield the 2p orbital more than they shield 2s orbital b.
I. E. actually decreases from column 2 to column 3 c.
An electron is repelled by the other electron in its same orbital d.
Nitrogen I. E. is greater than Oxygen I. E.
Nitrogen: 1s(↑↓) 2s(↑↓) 2p (↑ ) (↑ ) (↑ )
Oxygen: 1s(↑↓) 2s(↑↓) 2p (↑↓) (↑ ) (↑ )
8.
Additional Ionizations (2 nd and 3 rd Ionization Energies) a.
First ionization energies are always lower than 2 nd b.
Relative difference has to do with filled Noble Gas Configurations

B.
Electron Affinity (EA) = energy change when e- is added to a gas atom
1.
X(g) + e- -------> X (g)
2.
EA generally become more negative (exothermic) from left to right a.
Interactions between the +nucleus and the added electrons is favorable b.
The larger the nucleus, the better the interaction
3.
Exceptions to the trend is a function of electron—electron interactions a.
N- doesn’t form because N has 1s 2 2s 2 2p 3 configuration b.
The extra electron would have to go into an occupied 2p orbital c.
C- does form because it has an empty orbital: 1s 2 2s 2 2p 2 d.
This phenomenon doesn’t always hold true. O- can form 1s 2 2s 2 2p 4
4.
EA diminishes down a group a.
Added e- is farther from the nucleus, so less interaction is observed b.
The difference in EA is small and has exceptions (small F 2p orbital)

C. Atomic Radius: half the distance between the nuclei in a molecule consisting of identical atoms.
1.
Easy to understand for diatomic X
2 molecules
2.
Can be calculated from known radii of XY molecules
3.
Decreases from left to right on periodic table because +/- attraction becomes greater as we add protons to the nucleus across a row
4.
Increases down a group because more e- and larger shells filled

5.
Electrons in the same shell don’t shield each other from the nucleus a.
This gives rise to the decrease in atomic radius across a row b.
Each subsequent element adds a proton (increasing attraction to nucleus) c.
Each subsequent element adds an electron (no shielding since same shell)
6.
Core electrons effectively shield outer electrons a.
This gives rise to the large increase in size down a column b.
Effective nuclear charge is reduced due to shielding c.
Z eff
= Z actual
– Shielding
7.
Transition Elements a.
Atomic Radius of transition metals changes little across the row b.
As number of protons goes up, the electron being added is to a core shell c.
Z eff is almost constant as: Z act
+ 1, Shielding +1

D.
1.
Cations are smaller than neutral atoms because more attraction per electron
2.
Anions are larger than neutral atoms because less attraction per electron

Metals : tend to give up electrons to form cations found on the lower left
Nonmetals : tend to gain electrons to become anions found on the upper right
Metalloids : have properties of both metals and nonmetalsdepending on the conditions