1. Structure & Properties of Matter
Lesson # 3: Electron Configurations

2. Multi-Electronic Atoms
The Bohr-Rutherford model of the atom works well for the first 20 elements, up to calcium, but not so well for the transition metals or higher.
The Quantum Mechanical model better explains how electrons are arranged in atoms with multiple electrons.

3. Example - Helium
Consider helium with 2 electrons in its 1s orbital (when n = 1, l = 0, ml = 0). Within this atom, there is kinetic energy of the electrons as they move about the nucleus, potential energy of attraction between the nucleus and electrons, and potential energy of repulsion between the two electrons. The repulsion is what makes it very difficult to pinpoint the exact location of the electrons, and occurs in any atom with more than one electron in it.

4. Example - Sodium
For example, a sodium atom has 2 electrons in the first shell, 8 in the second, and 1 in the third. That outer electron feels attraction to the nucleus but repulsion from inner electrons, which causes it to not be as tightly bound.
This leads to the fact that Ens < Enp < End < Enf, etc. The s shell is always filled first as it takes the least energy to fill it, and moves upwards from there, with p orbitals at a slightly higher energy level than s, even with the same quantum number.

5. The Aufbau Principle
It is possible to show how electrons are arranged in atoms at the ground state. This is called an electron configuration.
The Aufbau Principle states that an atom is “built up” by progressively adding electrons, and are added so that they assume their most stable condition by filling the lowest available energy orbitals before filling higher energy orbitals.

6. Aufbau Principle Hydrogen: 1s1 2p 2s 1s
We can write in both the electron configuration using values for n and l, and we can also draw energy level diagrams, which are a more visual representation of electron arrangement (without getting into that crazy diagram I drew on the board the other day).
Hydrogen: 1s1 2p 2s 1s

7. Examples
Helium: 2p 2s 1s
Lithium:

8. Examples
Beryllium: 2p 2s 1s
Boron:

9. Examples
Carbon: 2p 2s 1s
Nitrogen:

10. Examples
Oxygen: 2p 2s 1s
Fluorine:

11. Examples
Neon: 2p 2s 1s
Sodium:

12. Examples
Sodium Ion: 2p 2s 1s
Fluoride Ion:

13. Aufbau Principle Summary
Use the periodic table to determine the number of electrons in the atom or ion
Assign electrons by main energy level and then by sublevel, using electron configurations or energy level diagrams.
Hund’s rule states that a particular set of orbitals that have the same energy, the lowest energy configuration is one with the maximum number of unpaired electrons – which is why we kept them unpaired and distributed evenly until the had to be filled up (which is also why we did this with Bohr Rutherford diagrams in grade 9 and 10).
Fill each sublevel before starting with the next. Continue until all electors are assigned.
For ions, either remove or add the appropriate amount of electrons to the diagram.

15. Anomalous Electron Configurations
When we get into atoms that have a d shell, the Aufbau principle breaks down a bit.
For example, potassium, which is in period 3 on the periodic table (n = 3, so l = 2,1,0), has s, p and d shells.
You would expect an electron configuration of:
Instead of that one electron placed in the 3d shell, we have found by experiment that the electron fills the 4s shell:
This makes sense as its properties are very similar to both lithium and sodium, which also have one electron in their outermost s shell.
In summary, the (n+1)s subshell is filled before the nd orbital.

16. Anomalies
Transition metals are the first to start filling the d shell.
Scandium:
It can’t fill the 4d shell because the 3d is empty.
Even weirder is one like Chromium.
It should be:
Experimentally we find that it is actually:
Experimentation has show that unfilled subshells are less stable than half-filled and filled subshells, and that unfilled subshells have higher energy.
This seems to affect the d and f orbitals more than any other.
Copper is another exception, seen as [Ar] 4s1 3d10.

17. Explaining Ion Charges
Neutral cadmium has an electron configuration of:
It’s 4th shell is completely full, and there are 2 electrons in the 5s shell.
Losing those would make it stable, hence it has a charge of 2+.
Neutral lead is [Xe] 6s2 4f14 5d10 6p2.
It’s 4th and 5th shell are completely full.
If it loses the two 6p electrons it would have a charge of 2+, and if it lost both the two 6p and 6s electrons, it would have a charge of 4+.

18. Explaining Magnetism
Based on magnetism associated with electron spin and the presence of several unpaired electrons, an initial explanation for magnetism is that unpaired electrons cause magnetism.
However, the presence of several unpaired electrons may account for some magnetism (paramagnetic: weakly magnetic) but not for strong magnetism (ferromagnetic: strongly magnetic).
For an atom to be ferromagnetic, the atoms must be small & closely packed together, capable of orienting themselves in a magnetic field. The theory is that each atom acts like a little magnet. This only occurs in certain atoms.

19. Explaining Magnetism
These atoms influence each other to form groups (domains) in which all of the atoms are oriented with their north poles in the same direction. Ferromagnetism is based on the properties of a collection of atoms, rather than just one atom.
Paramagnetism is also explained as being due to unpaired electrons within substances where domains do not form. In other words, paramagnetism is based on the magnetism of individual atoms.

20. Explaining Periodic Trends
1. Atomic Radius
Atomic radius is half the distance between the nuclei of two adjacent atoms, or in metals it is the distance between atoms in a crystal, and for molecules it is between two atoms chemically bonded together.
As n increases, there is a higher probability of finding electrons farther from the nucleus, so the atomic radius increases with increasing n.

21. Atomic Radius
The effective nuclear charge, Zeff is difference between an electron’s attraction to the nucleus and its repulsion from other electrons. As Zeff increases, electrons are attracted more strongly, so the size of the atom decreases. As Zeff decreases, there is a reduced force of attraction, and the size increases.
The exception is transition metal elements. Because of the large shielding from the d electrons, Zeff remains fairly constant for all transition metals in a period.

22. 2. Ionization Energy
Ionization energy is the energy required to remove an electron from a ground state atom in the gaseous state.
In order to remove this electron, energy is needed to overcome the force of attraction that is exerted on the electron by the nucleus.
In general, ionization energy increases across and period and decreases down a group. This is the inverse of atomic radius.

23. 2. Ionization Energy
Across a period, the atomic radius decreases because Zeff increases. The increased effective nuclear charge of each successive element increases the attractive forces between the nucleus and valence electrons, therefore more energy is needed to remove one such electron.
There are some exceptions to the rule, as explained by those anomalous electron configurations discussed above. If removing an electron produces half filled or completely filled orbitals, it will be more stable than going to a configuration with an empty orbital (so ionization energy would be slightly higher than expected).

24. 3. Electron Affinity
Electron affinity is a change in energy that accompanies the addition of an electron to an atom in the gaseous state.
Energy is typically released when the first electron is added because it is attracted to the atom’s nuclear positive charge. Adding another in, however, becomes more difficult as energy must be absorbed in order to overcome the electrostatic repulsions and add another electron to a negative ion.
Electron affinity is actually much harder to predict than ionization energy or size, and is beyond the scope of the course.