IBDP Chemistry HL Cheat sheet - Structure 1.3 - Electron configurations | IBDP Chemistry Cheatsheets

Emission spectra form when electrons in excited states return to lower energy levels and emit photons.
A larger energy difference between levels produces a higher-energy photon.
Towards violet light, wavelength decreases while frequency and photon energy increase.
Towards red light, wavelength increases while frequency and photon energy decrease.
A continuous spectrum contains an uninterrupted range of wavelengths; a line spectrum contains only specific wavelengths.

Each orbital has a defined energy for a particular electron configuration and chemical environment.
An orbital can contain a maximum of two electrons with opposite spin, represented as $\uparrow\downarrow$ .
An $s$ sublevel contains $1$ orbital, $p$ contains $3$ , $d$ contains $5$ and $f$ contains $7$ .
The Aufbau principle fills available orbitals from lower to higher energy.
Hund’s rule fills equal-energy orbitals singly before electron pairing occurs.
The Pauli exclusion principle requires two electrons occupying the same orbital to have opposite spins.

At increasingly high frequency, spectral lines approach a limit of convergence.
This limit corresponds to the minimum energy required to ionize the atom from the relevant state.
If frequency is given, calculate photon energy using $E=hf$ .
If wavelength is given, use $c=\lambda f$ , so $E=\dfrac{hc}{\lambda}$ .
For a molar first ionization energy, use $\mathrm{IE}=E_{\text{photon}}N_A$ .
Convert from $\mathrm{J,mol^{-1}}$ to $\mathrm{kJ,mol^{-1}}$ by dividing by $1000$ when required.

Successive ionization energies measure the energies required to remove electrons one after another from the same species.
Values normally increase because each electron is removed from an increasingly positive ion.
A large jump indicates that the next electron belongs to a lower main energy level closer to the nucleus.
Count the electrons removed before the first major jump to determine the number of valence electrons.
For main-group elements, the number of valence electrons can therefore be used to deduce the group.

The hydrogen emission spectrum consists of discrete lines rather than a continuous range.
Each line corresponds to an electron making a transition between two discrete energy levels.
Transitions ending at $n=1$ , $n=2$ and $n=3$ form separate groups of spectral lines.
Within each group, the lines become progressively closer together as the energy levels converge at higher energies.
The names assigned to the different hydrogen spectral series are not required.

Each main energy level is divided into sublevels, labelled $s$ , $p$ , $d$ and $f$ .
These sublevels are generally of successively higher energy within the same main energy level.
An orbital is a region of space with a high probability of finding an electron.
An $s$ orbital has a spherical shape.
The three $p$ orbitals have dumbbell shapes oriented along three mutually perpendicular directions.

Deduce electron configurations for atoms and ions up to $Z=36$ using orbital-filling principles.
A full electron configuration shows every occupied sublevel, for example sodium as $\mathrm{1s^2,2s^2,2p^6,3s^1}$ .
A condensed configuration replaces inner electrons with the preceding noble-gas core, giving sodium as $\mathrm{[Ne],3s^1}$ .
For ions, first determine the new total number of electrons and then write the configuration.
When forming transition-metal ions, electrons are removed from the highest main energy level first.
Chromium is $\mathrm{[Ar],3d^5,4s^1}$ and copper is $\mathrm{[Ar],3d^{10},4s^1}$ ; both are required exceptions.

First ionization energy generally increases across a period because nuclear charge increases while electrons are added to the same main energy level.
First ionization energy generally decreases down a group because the outer electron is farther from the nucleus and experiences greater shielding.
A drop can occur when the electron removed occupies a higher-energy sublevel, making removal easier.
Another discontinuity can occur when electrons pair in an orbital because increased electron–electron repulsion makes one easier to remove.
Use electron configurations to explain both the overall trend and any discontinuities.

Can you explain how electron transitions produce line emission spectra?
Can you relate colour, wavelength, frequency and photon energy qualitatively?
Can you use $2n^2$ to determine the maximum number of electrons in a main energy level?
Can you recognize an $s$ orbital and the three orientations of $p$ orbitals?
Can you apply the Aufbau principle, Hund’s rule and Pauli exclusion principle to atoms and ions up to $Z=36$ ?
Can you explain the chromium and copper exceptions and write full and condensed configurations?
Can you calculate first ionization energy from a convergence wavelength or frequency and interpret successive ionization energy data?

IBDP Chemistry HL Cheat Sheet - Structure 1.3 - Electron configurations