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6 Electronic Structure of Atoms - Coggle Diagram
6 Electronic Structure of Atoms
electron configuration
quantum numbers
principle quantum number
energy level on which the orbital resides
n
1, 2, 3, 4...
angular momentum quantum number
shape of the orbital
l
0 ~ n-1/ s, p, d, f...
magnetic quantum number
three-dimensional orientation of the orbital
ml
-l ~ l
spin quantum number
s
+1/2 or -1/2
orbitals
s
l=0
spherical
radius increases with the value of n
p
l=1
have two lobes with a node between them
d
l=2
energies of orbitals
degenerate
orbitals on the same energy level have the same energy
as the number of electrons increases, the repulsion between them increases
in many electron atoms, orbitals on the same energy level are no longer degenerate
electron configurations
shows the distribution of all electrons in an atom
components
a number denoting the energy level
principle quantum number
a letter denoting the type of orbital
sub-shell
a superscript denoting the number of electrons in those orbitals
number of electrons
principles
l
Aufbau principle
the manner in which electrons are filled in the atomic orbitals of an atom in its ground state
m
Pauli exclusion principle
two or more identical particles with half-integer spins cannot occupy the same quantum state within a quantum system simultaneously
n
Heisenberg uncertainty principle
it is impossible for us to know simultaneously both the exact momentum of the electron and its exact location in space
s
Hund's rule
every orbital in a sublevel is singly occupied before any orbital is doubly occupied
special cases
potassium
1s2 2s2 2p6 3s2 3p6 4s1
4s is below the 3d in terms of its energy
chromium and copper
Cr
[Ar]3d5 4s1
half-filled d5
Cu
[Ar]3d10 4s1
fully filled d10
periodic table
we fill orbitals in increasing order of energy
different blocks on the periodic table correspond to different types of orbitals
waves
wavelength(λ)
the distance between corresponding points on adjacent waves
frequency(v)
the number of waves passing a given point per unit of time
electromagnetic radiation
c = λv = 3.00*10^8m/s
quantized energy and photons
the wave nature of light does not explain how an object can glow when its temperature increases, Max Planck explained it by assuming that energy comes in packets called quanta
Einstein used this assumption to explain the photoelectric effect
E = hv
Plank's constant h = 6.626 * 10^-34Js^-1
quantum mechanics
Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated
the square of the wave equation, Ψ^2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time
line spectra and the Bohr model
for toms and molecules one does not observe a continuous spectrum, only a line spectrum of discrete wavelengths is observed
Bohr model
electrons in an atom can only occupy certain orbits
electrons in permitted orbits have specific, "allowed" energies; these energies will not be radiated from the atom
energy is only absorbed or emitted in such a way as to move an electron from one "allowed" energy state to another
E = hv
ΔE = -RH(1/nf^2 - 1/ni^2)
Louis de Broglie posited that if light can have material properties, matter should exhibit wave properties
λ= h/mv
the uncertainty principle
(Δx) (Δmv) >= h/4π