L5 - Electronic Config

Electrons and Electron Configuration

Recap of yesterday's lecture on the location of electrons in an atom.
Electrons reside in a cloudy sphere around the nucleus.
Chemical reactivity is driven by valence electrons, which are the outermost electrons with higher energy states.
Valence electrons can be shared, transferred, or accepted in chemical reactions.
Bohr's model proposed concentric rings around the nucleus with specific electron occupancies (2, 8, 18, 36, etc.).
Bohr's model is not entirely correct.
A better approximation is electron configuration (arrangement of electrons) with occupancies of 2, 8, 8, etc.

Electron configuration: Arrangement of electrons around and within an atom, influencing energy and accessibility.
Electrons exhibit both particle and wave behaviors.
Bohr's model is a stepping stone but not accurate.
Bohr's model introduced the concept of energy levels occupying lowest energy level first, with electrons closest to the nucleus and energy levels increasing with distance from the nucleus (N=6 has a higher energy level).
Spectra involve electron transitions between energy levels.

Shells and Subshells (Orbitals): Different energy levels where electrons reside.
- Orbitals are just another term for subshells. It does not mean orbiting.
Subshell Labels: S, P, D, (and F, though less time spent on it).
Key Information:
- Electron capacity of each subshell
- Characteristics of each subshell
- Filling order of subshells guides valence electron count.

Electron configuration is like a map of where electrons are in an atom.
Explains phenomena like excited atoms emitting colors when energy is applied and then relaxing back down to ground state, emitting a frequency of energy corresponding to the color.

Briefly reviewed the Borale's equation from the last lecture relating electrons wavelength to momentum.
Electrons are quantized, possessing specific energies.
Wave-like properties of electrons include frequency, wavelength, and amplitude.
Amplitude: The height of a wave from the zero line to its peak, also described as a wave function.
Orbitals are wave functions, describing where electrons sit.
Wave functions are solutions to Schrodinger's equation.
Amplitude represents electric field strength; its square relates to intensity or brightness.
Amplitude^2 = Intensity of brightness
The square of the electron wavelength gives a description of where electron density exists around the nucleus, giving a probability of the electrons position.

Wave function describes spatial distribution of an electron within an atom.
- (s) kind is spherical.
- (p) kind is two lobes where nucleus is in the center.
S, P, D, and F orbitals have different shapes, with S being spherical and P resembling two lobes. More complex shapes for D and F orbitals.
The S, P, D, and F orbitals dictate how many electrons each can hold.
No orbiting occurs in atomic orbitals; the shapes represent probability.

Fixed harmonic notes correspond to fixed electron orbitals.
Guitar string example: changing finger position (bound distance) creates different standing waves (quantized notes), analogous to electrons having quantized possibilities for energy levels.
Electron standing waves:
- One-dimensional on a guitar string.
Electrons are three-dimensional standing waves considering X, Y, and Z coordinates.

Electrons fill subshells (S, P, D, F) based on energy levels.
Periodic table organization corresponds to these subshells.
Lowest energy levels are filled first (1s).
Hydrogen and helium occupy the 1s orbital.
S orbitals hold two electrons.
Lithium and beryllium occupy the 2s orbital.
2s is the (n=2) orbital.
After filling 1s, electrons go to 2s and then 2p orbitals.
Example: After completing 2s, electrons occupy 2p orbitals (2(p^6)), then 3s (3(s^2)), and 3p (3(p^6)).
Larger elements' configurations deviate from expected order due to energy levels. 4s is lower energy than 3d.

The example used in lecture has two electrons in each of the one s's (1(s^2)), two in each of the two s's (2(s^2)), six in that p's (2(p^6)), two in that three s's (3(s^2)), six in that three p's (3(p^6)), two electrons into 4s's (4(s^2)) and one electron into three d's (3(d^1)).
Valence electrons are electrons in the highest energy shell, not subshell.
In the example provided, highest energy shell is (n = 4), so there are two valence electrons.
Despite electrons occupying the 3d subshell after 4s, 3d electrons are considered core electrons.
Energy levels across shells have significant gaps, while subshells within a shell have smaller gaps.

Atomic emission spectra result from energy input, causing electrons to jump to higher levels.
Ground state: All electrons are in their lowest possible energy states.
Excited state: An electron has absorbed energy and moved to a higher energy level.
Relaxation/Emission: When an electron falls back to a lower energy level, it emits energy, producing emission spectra.
Distinct lines in spectra correspond to possible electron transitions.
Electron configurations can represent excited states (e.g., 1s2 2s2 2p6 3p1).