S1.3 (Electron Configuration) Flashcards
Emission Spectra
Range of frequencies or wavelengths of electromagnetic radiation emitted during an electron transition from a higher energy to a lower energy level shown by lines converging at higher energy
Electron Transition
Movement of an electron between energy levels in an atom
electromagnetic Spectrum
Radiowaves<Microwaves<Infra red<Visible Light<UV<X Ray< Gama Ray
highest energy & frequency wave
UV
continuous spectrum
all the freuencies across a range of electromagnetic radiation (rainbow)
Emission line Spectra (How it works)
- energy is applied to a sample of vaporized hydrogen
- only shows specific wavelengths: colored lines on black sheet of paper
*
Why electrons prodcue colored lines
- electron absorbs energy
- n=2 -> n=3
- electron is excited & unstable
- electron releases energy
- n=3 -> n=2
- emitted energy is shown as wave
- line is observed
What is special about the lines on an emission spectra sheet
they converge at high frequencies
Hydrogen Emission Spectrum
The emission line spectrum of hydrogen (proves that electrons stay at specific energy levels)
Hydrogen Emission Spectrum
UV
higher -> n=1
Hydrogen Emission Spectrum
Visible Light
higher -> n=2
Hydrogen Emission Spectrum
IR
higher -> n=3
meaning of n=∞
point where the electron has left the atom thus having an infinite amount of energy
Heisenbergs uncertainty principle
it is impossible to know the postion and momentum of an electron at once
principle quantum numbers
shows main energy levels of an electron (n)
Orbital
A region of space with the highest possibility of finding an electron
How many electrons can each/any orbital hold
2
S orbital
- spehrical
- lowest in energy
- size increases as shell number increases
Total number of orbitals in each energy level
- n=1; 1
- n=2; 4
- n=3; 9
- n=4; 16
P Orbital
- dumbell shaped
- has 3 degenerate orbitals (Py, Px, Pz)
- total of 6 e-
D Orbital
- double dumbel shaped
- 5 orbitals
- can hold a total of 10 e-
F orbital
- 7 orbitals
- total of 14 e-
Total number of electrons in each energy level
- n=1; 2
- n=2; 8
- n=3; 18
- n=4; 32
Sublevels in each energy level
- n=1; 1s
- n=2; 2s,2p
- n=3; 3s,3p,3d
- n=4; 4s,4p,4d,4f
Aufbau Principle
Electrons fill the lowest energy levels first and fill orbitals in order of increasing energy
Hunds Rule
When degenerate orbitals are available, electrons will fill those orbitals singly and with parallel spins before pairing up
Pauli Exclusion Principle
Two electrons in the same orbitals must have opposite spins
Aufpau Principles filling order
1s>2s>2p>3s>3p>4s>3d>4p>5s>4d>5p>6s
Noble Gas Notation example
Calcium (20 e-)
(Ar)4s2
Shortened Spectroscopic Notation/ Noble Gas notation
factors out the part of the notation which overlaps with one of the noble gases
Spectroscopic Notation
Exceptions
- copper
- chromium
they first fill the d orbitals singly rather having them half full
Orbital box Notation
shows electrons and their pairing + spins
Eelectron configurations of ions
- filling= lowest first
- removing= highest first
- except: 4s first in & first out
Ionisation Energy
The amount of energy required to remove one mole of electrons from one mole of gaseous atoms
Why ionisation energy increases across a period
- nuclear charge increases as protons increase
- this means stronger attraction between nucleus and electrons
- energy required to remove the outer electrons increases
Trend in ionisation energies
- ionistaion energy increases across a periods
- ionisation energy decreases going down a group
Why ionisation energy decreases down a group
- number of shells increase, increasing the distance from nucleus and electrons
- attraction between electrons and nucleus decreases => easier to remove electrons
- &shielding from electrons in shells decreases attraction
=> energy required to remove an electron decreases
Convergence limit
en electron is removed from the attraction of the nucleus (ionised)
While knwoing the frequency
ionisation Energy calculations (kJ/mol)
- use E=hf to calculate energy in Joules
- then divide by 1000 for kJ
- then times by avogadros constant (L) for K+kJ per mole
Without knowing the frequency
Ionisation energy calculations (kJ/mol)
- f=c/λ and E=Lhf (per mole) so subtitude f into E=Lhf
- E=Lhc/λ
- to get in in kJ = E= LHC/1000λ
Succesive Ionisation Energies
process of removing successive electrons from an atom or positive ion
Important about the second ionisation energy
It is always significantly larger than the first as you’re trying to remove an electron from a positive ion meaning the attraction to the nucleus is greater
(the more you need to remove the larger the ionisation energy)
Deducing valence electrons based on ionisation energy
the ionisation energy difference which is the largest means that was the last electron in the shell that was being removed
why is there a general increasing value for ionization energy across a period
electrons are being added to the same energy level, but the amount of protons increases
=> this means the attraction between them gets stronger
Things to look out for when comparing ionization energies (especially graph)
- trends across period (ionization increases)
- trend across groups (ionization increases down)
- DRAW ELECTRON CONFIGUARTION and speak about energy levels
What happens to ionisation energy when electrons are paired vs unpaired
- if all are unapired = harder to remove 0 ionisation energy increases
- if some are paired & unpaired= paired ones are easier to remove due to repulsion = decreases ionisation energy
Evidence that electrons exist in fixed energy levels
- emission spectra is discrete lines (if continuous: electrons would be anywherre)
- the lines converge showing different energy levels as the lines are succesively converging
- the differneces in ionization energies shows that different energies are needed to remove from different levels
limitations of hydrogen emission spectra
- does not represent sub levels/orbitals
- only applies to hydrogen atom/atom of one electron
- does not explain why only certain energy levels allowed
- the atom is considered to ber isolated
- does not consider the number of electrons thta can fit into a energy level
- does not consider the probability of finding an electron (orbitals)
emission spectrum described
- series of lines
- only at specific wavelengths
- converge at high frequences
relationship of emission spectrum tp hydrogen atom (outlined)
- electron transition between high energy level to lower energy level
- DRAW DIAGRAM
converting nm to m
multiply by 10xe9
things to consider when speaking about ionisation enrgy trends
- shielding => more shielding = less attaraction
- which orbital/subshell or main energy level it is taking e- out of
- attraction to nucleus (is it already psotive or negatively charged?)
