Electronic Structure and Periodic Table Flashcards
Orbital structure of hydrogen atom
The hydrogen electron orbits the nucleus, and the electron exists in a spherical probability cloud around the nucleus.
Principal quantum number n
The number “n” defines what shell the electron is in. The higher “n” shells indicate higher energy.
Number of electrons per orbital
There are n^2 squared orbitals per shell, and 2 electrons per orbital, so 2n^2 electrons per shell.
Ground state
Electrons are normally in their ground state. Excited electrons come down to ground state via release of energy
Excited state
When electrons in their ground state absorb energy, they get promoted to excited states
Absorption spectra
The absorption spectrum shows what wavelengths of light are absorbed, and looks like black lines on a rainbow background. They correspond to the emission spectrum in pattern.
Emission spectra
The emission spectrum shows what wavelengths of light are emitted. They look like colored lines on a black background. The spectrum shifts to a slightly longer wavelength.
Quantum number l
The number “l” is the angular momentum for quantum numbers ranging from 0 to n-1. “l” tells us whether it’s in the s (l=0) subshell, p (l=1) subshell, d (l=2), subshell, or f (l=3) subshell.
How many electrons exist in each subshell
The s subshell holds 1 orbital, the p subshell holds 3 orbitals, the d subshell holds 5 orbitals, and the f subshell holds 7 orbitals. Each hold a maximum of 2 electrons per orbital.
Quantum number m
The number “m” is the magnetic quantum number that range from “-l” to “l”, including zero.
Quantum number s
The number “s” is the quantum spin number, which is either +0.5 or -0.5
Common names and geometric shapes for orbitals s, p, d
For “s”, this equates to one orbital, “l” = 0, and a spherical shape. For “p”, this equates to 3 orbitals, “l” = 1, and a fission shape. For “d”, this equates to 5 orbitals, “l” = 2, and a clover-like shape.
Arrangement of subshells
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d
Orbital diagrams
Each subshell gets a row where each +0.5 spin and each -0.5 spin are an arrow.
Aufbau principle
Shells and subshells of lower energy get filled first.
Hund’s rule
When you fill a subshell with more than 1 orbital, you must fill each orbital with a single electron with the same spin.
Pauli Exclusion Principle
2 electrons in the same orbital must be of different spins
Bohr atom
Electron orbiting nucleus in a circular orbit and the larger “n” values have a larger orbiting radii.
Effective nuclear charge
The nuclear charge minus the shielding charge. The higher the efective nuclear charge, the more stable it is. It increases for outer electrons as you go left to right in the periodc table. Shielding electrons are those that stand between the nucleus and the electron. They are lower in energy.
Alkali metals
Group 1. Single valence electron, low ionization energy, very reactive, more reactive as you go down because of increasing radii. Most commonly found in +1 oxidation state
Alkaline earth metals
Group 2. Two valence electrons, low ionization energy, quite reactive, more reactive as you go down, and most commonly found in the +2 oxidation state
Halogens
Group 7. Seven valence electrons (2 s subshell and 5 p subshell). High electron affinity. Very reactive. More reactive as you go up because of decreasing radii. Most commonly found in -1 oxidation state.
Noble gases
Group 8. Eight valence electrons. High ionization energy and low electron affinity. Inert. Have oxidation state of 0.
Transition metals
High conductivity due to loosely bound outer d electrons. Varied oxidation states, but always positive.
