1. Chapter 5 Periodicity and the Electronic Structure of Atoms

2. Electron Spin
experiments by Stern and Gerlach showed a beam of silver atoms is split in two by a magnetic field
the experiment reveals that the electrons spin on their axis
as they spin, they generate a magnetic field
spinning charged particles generate a magnetic field
if there is an even number of electrons, about half the atoms will have a net magnetic field pointing “North” and the other half will have a net magnetic field pointing “South”

3. Electron Spin and the Pauli Exclusion Principle
spin quantum number describes how the electron spins on its axis
clockwise or counterclockwise
spin up or spin down
spins must cancel in an orbital
paired
Electrons have spin, which gives rise to a tiny magnetic field and to a spin quantum number (ms).
Another way to think about it is that two electrons in an orbital must be spin-paired.
Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers.

4. Pauli Exclusion Principle
no two electrons in an atom may have the same set of 4 quantum numbers
therefore no orbital may have more than 2 electrons, and they must have with opposite spins
knowing the number orbitals in a sublevel allows us to determine the maximum number of electrons in the sublevel
s sublevel has 1 orbital, therefore it can hold 2 electrons
p sublevel has 3 orbitals, therefore it can hold 6 electrons
d sublevel has 5 orbitals, therefore it can hold 10 electrons
f sublevel has 7 orbitals, therefore it can hold 14 electrons

5. Electron Configurations of Multielectron Atoms
Electron Configuration: A description of which orbitals are occupied by electrons
Degenerate Orbitals: Orbitals that have the same energy level—for example, the three p orbitals in a given subshell
Ground-State Electron Configuration: The lowest-energy configuration
Aufbau Principle (“building up”): A guide for determining the filling order of orbitals

6. Electron Configurations of Multielectron Atoms
Rules of the aufbau principle:
Lower-energy orbitals fill before higher-energy orbitals.
An orbital can hold only two electrons, which must have opposite spins (Pauli exclusion principle).
If two or more degenerate orbitals are available, follow Hund’s rule.
Hund’s Rule: If two or more orbitals with the same energy are available, one electron goes into each until all are half-full. The electrons in the half-filled orbitals all have the same value of their spin quantum number.
Electrons repel each other. Electrons in different orbitals (and thus different spatial regions) will experience fewer repulsive forces than electrons placed into the same orbital.

7. Electron Configurations of Multielectron Atoms

8. Electron Configurations of Multielectron AtomsH:
1s1
1 electron
s orbital (l = 0)
n = 1 He C Mg
Electron configurations show the distribution of the electrons between the subshells.

9. Electron Configurations of Multielectron Atoms
Orbital-Filling
Diagram H:
1s1 1s
He:
1s2
Li:
1s2 2s1
Orbital-filling diagrams show the distribution of the electrons between the orbitals.
One “line” for each orbital. An s subshell has only one orbital. N:
1s2 2s2 2p3

10. Electron Configurations of Multielectron Atoms
Shorthand
Configuration
Na:
1s2 2s2 2p6 3s1
Ne configuration
Think of it as a mathematical substitution.

11. Electron Configurations of Multielectron Atoms
Shorthand
Configuration
Na:
1s2 2s2 2p6 3s1
[Ne] 3s1 P:
1s2 2s2 2p6 3s2 3p3
Ne configuration

12. Anomalous Electron Configurations
Expected
Configuration
Actual
Configuration
Cr:
[Ar] 4s2 3d4
[Ar] 4s1 3d5
Cu:
[Ar] 4s2 3d9
[Ar] 4s1 3d10

13. Electron Configurations and the Periodic Table

14. Electron Configurations and the Periodic Table
Valence Shell: Outermost shell
Li:
2s1
Na:
3s1
Cl:
3s2 3p5
For main-group atoms, the column number (using the U.S. standard numbering system) is equivalent to the number of valence electrons.
Br:
4s2 4p5

15. Electron Configurations and Periodic Properties: Atomic Radii
Atomic radii increase down a column because successively larger valence-shell orbitals are occupied.
Atomic radii decrease from left to right because the effective nuclear charge increases.
Column
Radius
Row
Radius

16. Trend in Atomic Radius – Main Group
Different methods for measuring the radius of an atom, and they give slightly different trends
van der Waals radius = nonbonding
covalent radius = bonding radius
atomic radius is an average radius of an atom based on measuring large numbers of elements and compounds
Atomic Radius Increases down group
valence shell farther from nucleus
effective nuclear charge fairly close
Atomic Radius Decreases across period (left to right)
adding electrons to same valence shell
effective nuclear charge increases
valence shell held closer

17. Effective Nuclear Charge
in a multi-electron system, electrons are simultaneously attracted to the nucleus and repelled by each other
outer electrons are shielded from full strength of nucleus
screening effect
effective nuclear charge is net positive charge that is attracting a particular electron
Z is nuclear charge, S is electrons in lower energy levels
electrons in same energy level contribute to screening, but very little
effective nuclear charge on sublevels trend, s > p > d > f
Zeffective = Z - S

18. Orbital Energy Levels in Multielectron Atoms
Effective Nuclear Charge (Zeff): The nuclear charge actually felt by an electron
Zeff = Zactual − Electron shielding