1. Chapter 8: Periodic properties of the elements

2. Electron configuration
Electrons occupy orbitals following a set of rules
The number of electrons determine the eventual electronic configuration
Orbital diagrams can be used to visually represent how electrons occupy orbitals
Orbital diagram
Electron configuration

3. Pauli exclusion principle
Quantum numbers can indentify each individual electrons
Pauli exclusion principle – no two electrons can have the same four quantum numbers
Electron configuration
Quantum Numbers
Orbital diagram

4. Aufbau principle
Aufbau principle - When placing more electrons, fill the lowest energy orbitals first
Lower l value, lower energy
s < p < d < f
Lower n value, lower energy
Exceptions once you get to d and f orbitals
4s  3d

5. Diagonal rule

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

7. 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)

8. 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

9. 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 Xe, xenon
1s22s22p63s23p64s23d104p65s24d105p6
Electron configuration of W, tungsten
1s22s22p63s23p64s23d104p65s24d105p66s24f145d4 or
[Xe]66s24f145d4

10. 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
Sometimes it saves energy to half fill or completely fill the lower sublevel, than to have it in the 4s or 5s orbital

11. 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, like a core from an apple

12. Electron configuration for ions
When writing electron configurations for ions, start from the neutral atoms configuration, then remove or add the appropriate amount of electrons dependant on the charge of the ion
When removing electrons, pull out of the highest principle quantum number orbitals first
V = [Ar] 4s23d3
V2+ = [Ar] 4s03d3 = [Ar] 3d3
electrons dependant on the charge of the ion

13. Coulomb’s law: implications
Electron configuration rules derived from interactions of the electrons and the nucleus
Coulomb’s law – potential energy of two charged particles dependant 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

14. Shielding
The nucleus of an atom is positively charged and will attract negatively charged electrons
In multielectron atoms electrons are repelled by the other electrons present
Shielding – electrons preventing other electrons from experiencing full effects of nuclear charge
Effective nuclear charge (Zeff) – The approximate charge an electron will experience
For Li+
3 protons, 3+
2 electrons, 2-
1s2
An incoming electron outside the “shielded” region experiences
3 – 2 = +1 charge
Zeff = +1

15. 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
This is penetration

16. 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

17. 4s has greater penetration than the 3d orbital, enough be a lower energy orbital in comparison

18. Size of an atom
Atomic radius – the radius of an atom (spherical in shape)
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

19. 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 expection, generally staying constant throughout

20. 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

21. 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
diamagnetic
paramagnetic

22. 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 repulsion

23. Isoelectronic ionic trends
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

24. 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

25. IE1 trends

26. IE1 trends

27. IE2,3,4…n trends

28. 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

29. Metallic character

30. The end of Chapter 8