Week 5 Lecture 2 - Topic 4 (Atomic energy levels)

Atomic Energy Levels and Quantum Numbers

Electrons occupy specific regions around the nucleus with discrete energy levels.
Atoms absorb energy (e.g., light) to excite electrons to higher energy states.
The energy absorbed is directly related to the energy change during excitation.
When an electron relaxes back to the ground state, it loses the same amount of energy.

Quantum numbers describe the probability of finding an electron in a specific location.
Key quantum numbers:
- Principal Quantum Number (n):
  - Relates to the energy level of the orbital or shell. n = 1, 2, 3,…
  - Indicates the size of the orbital; as n increases, energy and size increase.
- Azimuthal Quantum Number (l):
  - Describes the shape of the orbital.
  - l = 0: s orbital (sphere).
  - l = 1: p orbital (figure eight).
  - l = 2: d orbital (clover).
  - l = 3: f orbital (more complex shape).
- Magnetic Quantum Number (m):
  - Indicates the orientation of the orbital in space.
  - s orbital: one orientation (one "house" or bedroom).
  - p orbital: three orientations (three "houses" or bedrooms) along x, y, and z axis
- Spin Quantum Number (m_s):
  - Relates to the spin of the electron.
  - Electrons can have spin up (+1/2) or spin down (-1/2).

States that no two electrons in an atom can have the same set of quantum numbers.
Each orbital can hold a maximum of two electrons, which must have opposite spins (spin up and spin down).
Energetically unfavorable for two electrons in the same orbital to have the same spin.

When filling orbitals, add one electron to each orbital within a subshell before pairing any electrons.
Electrons are added with the same spin orientation initially for energy reasons.

Electron configuration relates to the periodic table and various properties.
Valence electrons are electrons in the outermost shell of an atom.
Valence electrons influence the reactivity of an atom, which tends to gain or lose them to achieve a full outer shell.

As you go from left to right across the periodic table:
- The amount of protons in the nucleus increases.
- This increases the positive charge in the nucleus (nuclear charge).
- The attraction to electrons increases, pulling them closer.
- The atomic radii decreases.
- Shielding occurs due to the inner core electrons which partially shield the outermost electrons from the full nuclear charge, so the outermost electrons experience less of the nuclear charge.
As you go down the periodic table; atomic size increases as the principal quantum number increases.

Count the number of electrons in the electron configuration.
Find the element on the periodic table with that number of protons (atomic number).
If the number of electrons doesn't match the number of protons, it's an ion.
Determine the charge of the ion based on the difference between protons and electrons.
Example: Electron configuration with 18 electrons.
Argon (Ar) has 18 electrons (and 18 protons), so it could be argon.
If argon isn't an option, consider ions.
Sulfur (S) has 16 protons. To have 18 electrons, it would need a 2- charge (S^{2-}

Transition metals (d-block elements) lose electrons from the s orbital before the d orbital.
Example: Iron (Fe)
- Neutral iron has 26 electrons.
- Iron(II) (Fe^{2+}) has 24 electrons.
- Iron(III) (Fe^{3+}) has 23 electrons.
Two exceptions for filling orbitals in the d-block:
- Copper (Cu)
- Chromium (Cr)

As the principal quantum number (n) increases, atomic size also increases.
Core electrons shield valence electrons from the full nuclear charge.
The effective nuclear charge is the net positive charge experienced by valence electrons, considering the shielding effect.

Across a row:
- Nuclear charge increases due to increasing protons in the nucleus.
- Increased attraction to electrons and effective nuclear charge that draws electrons closer.
- Atomic radius decreases with an increasing effective nuclear charge from left to right across a period.
Down a group:
- Nuclear charge gets larger because more neutrons are being added.
- Effective nuclear charge does not increase as much, because there are an increased amount core electrons that are shielding the outermost electrons.
- Atomic radius can get larger, even though the nuclear charge tends to increase the net effect, because more and more of the attraction is negated by shielding.

Describes the atomic size
Determined by taking distance between two carbon atoms that are bonded to each other, and the raidus would be half of that distnace
In metals, it's called the metallic radius

First ionization energy has a general trend with how much energy is required to start the reaction to do it. The energies increase as we go up, or you could say it gets easier as you go down. They need less energy because the electrons are further along the nucleus as they increase. Therefore, they experience less attraction of the nucleus with larger nucleus.
By remembering those trends of atomic size, working out other predictable trends work. As atoms get larger, it is easier to take out the first amount of electron. So it is more efficient when the atomic size gets bigger at removing it from the first electron.
It gets increasingly harder to remove more and more electrons from each shell as you go
Eventually when the specific starts losing electrons, the configuration will reach a noble gas form, which is super stable and hard to remove and break the shell. Carbon is a prime example.
It will require high amount of intense energy for all other forms of shell type electrons, the most popular being the Noble Gas configuration, which needs a lot of energy to keep going and force its way to continue going and forcing the electron to deplete and go after another one and another one each at a time.