1. Chapter 8: Periodic properties of the elements

2. Electron configuration
How electrons fill up atomic orbitals
The number of electrons determine the eventual electronic configuration
Orbital diagrams can be used to visually represent how electrons occupy orbitals

3. Pauli exclusion principle
Quantum numbers can identify each individual electrons
Pauli exclusion principle – no two electrons can have the same four quantum numbers
Only 2 electrons per orbital maximum
Quantum Numbers

4. Aufbau principle
Aufbau principle - When placing more electrons, fill the lowest energy orbitals first
Lower n value, lower energy
1s vs 2s = 1s is lower in energy
Lower l value, lower energy
s < p < d < f

6. Diagonal rule

7. Start from top left (hydrogen)
Follow order of increasing atomic number
6s  then 4f  then 5d
7s  then 5f  then 6d

8. Electron configuration (continued)
Individual orbitals can hold up to two electrons
Different subshells contain different amounts of orbitals
s subshell – one orbital = 2 electrons total (maximum)
p subshell – three orbitals = 6 electrons (maximum)
d subshell – five orbitals = 10 electrons (maximum)
f subshell – seven orbitals = 14 electron (maximum)

9. Hund’s rule
Hund’s rule – when placing more electrons, fill in empty orbitals first with parallel spins, then pair up half filled orbitals with opposite spin electrons
Occurs with degenerate (equal energy) orbitals (ex: 3px, 3py, 3pz)

10. Hund’s rule discussion
In groups, come with up an explanation for Hund’s rule. Why do you fill vacant orbitals first instead of pairing up electrons first?

11. Practice: Electron configuration for Vanadium
atomic number = 23
Number of electrons = 23
Use electrons to fill lower energy orbitals in correct order (refer to chart)
Fill in vacant orbitals first before pairing two electrons up on degenerate orbitals

13. Noble Gas Abbreviation
The electron configuration of the noble gas that precedes the element in question is represented by the noble gas’ bracketed symbol
He = 1s2
C = 1s22s22p2 = [He] 2s22p2
A quicker way to write out electron configurations
Electron configuration of W, tungsten
1s22s22p63s23p64s23d104p65s24d105p66s24f145d4 or
[Xe]6s24f145d4
Noble Gas Abbreviation

15. d-block and f-block exceptions
Sometimes transition metals, lanthanides, and actinides can deviate from the expected rules
Chromium expected valence shell – 4s23d4
Chromium observed valence shell – 4s13d5

16. Valence Electrons Electrons in the outermost (valence) shell
The shell with the highest n value
C = 1s22s22p is highest number
count up electrons in all the valence orbitals
2 electrons in the 2s orbital and 2 electrons in the 2p orbital = 4 valence electrons
P = 1s22s22p63s23p is highest number
2 + 3 = 5 valence electrons
Core electron – the rest of the electrons aside from the valence electrons
The inner electrons

17. Electron configuration for ions
When removing electrons, pull out of the highest principle quantum number orbitals first
V = [Ar] 4s23d3
V2+ = [Ar] 4s03d3 = [Ar] 3d3
Electron configuration for ions

18. Coulomb’s law: Implications
Coulomb’s law – potential energy of two charged particles dependent on magnitude of charges (q1, q2) and the distance between them (r)
If charges are the same ( both positive or both negative) the energy is positive
Since like charges repel each other, the closer they are the more potential energy they have
If charges are opposite (one negative and one positive) the energy will be negative
Opposites attract, the closer they are, the more negative the potential energy
Larger magnitude of q will increase the potential energy
+3 charge will have stronger attraction than a +1 charge

19. Shielding
The nucleus of an atom is positively charged and will attract negatively charged electrons
In multielectron atoms, electrons repel each other
Shielding – electrons preventing other electrons from experiencing full effects of nuclear charge
Effective nuclear charge (Zeff) – The approximate charge an electron will experience

20. Penetration
The actual nuclear charge Z is weakened by the charge shielded by screening electrons S, giving the effective nuclear charge Zeff
If an electron were to “penetrate” inside of the shielded region, it would then experience the full nuclear charge
More experienced charge = drawn closer to the nucleus = lower energy

21. Orbital ordering explained
Radial distribution functions for orbitals solved from the schrödinger equation
When comparing orbitals with the same n value, the 2s has higher penetration than the 2p orbital
Lower l value = higher penetration

22. 4s has greater penetration than the 3d orbital, enough to lower its energy below the 3d orbital

23. Periodic table trends Size of atom Zeff Magnetic properties
Ionic radii
Isolectronic trends
Ionization energy
Electron affinities
Metallic character

24. Size of an atom Atomic radius – the radius of an atom
Non bonding atomic radius or van der Waals radius – determined upon freezing an element into the solid state, and measuring the distance between atom centers using the density
Bonding atomic radius or covalent radius
Nonmetals: half the distance between two of the atoms bonded together
Metals: One-half the distance between two of the atoms next to each other in a crystal of the metal

25. Periodic trend for atomic radius
As we move down a column/family, the radius increases
As we move to the right across a period, the radius decreases
Transition metals are a slight expectation, generally staying constant throughout

26. Zeff trend
Effective nuclear charge will increase going to the right across a period
Core electrons shield better than valence electrons
Going to the right increases proton number and actual charge Z, but the S increases at a lower rate, resulting in a higher Zeff going to the right

27. Magnetic properties of atoms and ions
Unpaired electrons generate a magnetic field due to its spin
Atoms or ions that have unpaired electrons are called paramagnetic
More unpaired electrons make a more paramagnetic atom or ion
Using the electron configuration you can determine the presence of unpaired electrons
If all electrons are paired, then the atom or ion is diamagnetic

28. Ionic radii
When an atom loses or gains electrons, it changes the size of the atom
Upon losing electrons, the atom gets smaller
Cations are smaller than the neutral atom
Upon gaining electrons, the atom gets larger
Anions are larger than the neutral atom
Presence of more electrons causes electron repulsion, that increases size

29. Isoelectronic ionic
Isoelectronic – to have the same number of electrons, also being the same net electron configuration
For isoelectronic ions, the higher the number of protons, the greater the nuclear charge which pulls the electrons in, the smaller the size

30. Ionization energy (IE)
IE – the energy required to remove an electron from an atom or ion in the gaseous state
Removing the first electron is the first ionization energy (IE1)
Removing another electron is the second ionization energy (IE2)
So on and so on
Depending on the electron being removed, and where it is located, this relates to the energy that it will take to remove it

31. IE1 trends

32. IE1 trends

33. IE2,3,4…n trends

34. Electron Affinities (EA)
EA – how easily an atom or ion will accept an electron in the gaseous state
The EA is the energy associated with the gaining of an electron
Values are typically negative since an atom releases heat when it gains an electron
As EA becomes more negative, it becomes more favorable for an atom or ion to gain an electron

35. Metallic character

36. The end of Chapter 8