S1.3 Electron configurations Flashcards
1
Q
how are emission spectra produced
A
by atoms emitting photons when electrons in excited states return to lower energy levels
2
Q
continuous spectrum
A
- in the visible region
- all colours
- all wavelengths
- all frequencies
3
Q
line spectrum
A
- Only has specific wavelengths
- Only has specific frequencies
- Only has specific energies of light
- Tells us that emitted light from atoms can only be at fixed frequencies → quantised
- Electrons can only have certain energy amounts
4
Q
radio waves
A
low energy
long wavelength
low frequency
5
Q
gamma rays
A
high energy
short wavelength
high frequency
6
Q
relationship between frequency and wavelength + the supporting equation
A
Speed of light (c) = vλ → frequency and wavelength are inversely proportional
7
Q
hydrogen line emission spectrum
A
- Provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies
- As energy increases (frequency increases, wavelength decreases), lines converge so the electron is reaching a maximum amount of energy → ionisation energy
- The visible lines correspond to the electron jumping from higher levels to n=2
8
Q
n∞ → n=3
A
- paschen
- infrared region
- low energy
9
Q
n∞ → n=2
A
- balmer
- visible region
- medium energy
10
Q
n∞ → n=1
A
- lyman
- UV region
- high energy
11
Q
how is energy related to frequency and wavelength
A
- Wavelength determines colour
- If constant energy, the same colour will always be emitted
12
Q
principal quantum number (n)
A
- main energy level
- number of energy levels/quantum shells
- can hold a maximum of 2n^2 electrons
13
Q
what is the main energy level divided into
A
- sublevels/subshells of successively higher energies
- s, p, d, f
14
Q
s atomic orbital
A
- spherical
- low energy
- most likely to find electrons
15
Q
p atomic orbitals
A
- 3: x, y, z
- dumbbell
- higher energy than s
- lobes are larger and longer as n increases
16
Q
ground state
A
most stable electronic configuration with the lowest amount of energy
17
Q
Aufbau’s principle
A
- fill the sub shells with lowest energy first (1s)
18
Q
exception to Aufbau’s principle
A
- 3d is higher in energy than 4s
- So 4s is filled first
- When filled, 4s is higher in energy
19
Q
Hund’s Rule
A
- Electrons fill all orbitals singularly first, then in pairs
- Electrons can spin clockwise or anticlockwise
- Those with the same spin repel (spin pair repulsion)
- So → occupy separate orbitals to minimise repulsion
- They then pair up with electrons with spin in the opposite direction
20
Q
Pauli exclusion principle
A
- Orbital must hold 2 electrons with the opposite spin
- Energy to jump to a higher energy orbital is greater than inter-electron repulsion
- So they pair up and occupy lower energy levels first
21
Q
transition metal rules
A
- Fill 4s before 3d
- Lose 4s before 3d
22
Q
exception to Aufbau’s principle: Chromium
A
- 1s2 2s2 2p6 3s2 3p6 3d5 4s1
- [Ar]3d5 4s1
23
Q
exception to Aufbau’s principle: Copper
A
- 1s2 2s2 2p6 3s2 3p6 3d10 4s1
- [Ar]3d10 4s1
24
Q
explanation for copper and chromium
A
it is more energetically favourable if a 3d sublevel is ½ full or full, therefore a 4s electron is used
25
ions rules
- When removing electrons, we remove the highest energy electron first
4s before 3d
- In an emission spectrum, the limit of convergence at higher frequencies corresponds to ionisation
26
limit of convergence
where lines appear to meet, limit is the frequency at which spectral lines converge
27
4 factors affecting ionisation energies
Size of nuclear charge
Atomic radii
Shielding effect of inner electrons
Spin-pair repulsion
28
size of nuclear charge
Greater atomic number = greater nuclear charge
Greater attractive forces between nucleus and outer electron
More energy required to overcome attractive forces
29
atomic radii
Electrons in shells further from nucleus are less attracted to nucleus
So lower ionisation energy
30
shielding effect of inner electrons
Electrons in full inner shells repel electrons in outer shells
Prevents them to feel full nuclear charge
The greater the shielding the lower the ionisation energy
31
spin-pair repulsion
Paired electrons in the same atomic orbital in a subshell repel each other (more than electrons in different atomic orbitals)
Easier to remove an electron, lower ionisation energy
32
general trend in ionisation energies across a period
Increases
Increase in nuclear charge, same number of shells so same distance of outer electrons to nucleus
Reasonably constant shielding
Decreased atomic radius
Outer electron is held more tightly to nucleus → harder to remove
33
trend in ionisation energies down a group
Decrease
Nuclear charge increases
Distance between nucleus and outer electron increases (atomic radii)
Shielding by inner electron increases
Effective nuclear charge decreases as shielding increases
Outer electron is held more loosely to nucleus → easier to remove
34
explain the rapid decrease in IE between last element in a period and 1st element in next period
Increased atomic radii
Increased shielding by inner electrons
2 factors outweigh increased nuclear charge
35
group 2 to 3 dip
drop in the I.E of group 3, there is an extra electron in the 2p orbital which is removed as the p orbital is higher energy than s (easier to remove)
36
group 5 to 6 dip
drop in the I.E of group 6, extra electron paired up with one electron already in one of the 3p orbitals, there is a repulsive force between the 3 paired electrons, so less energy is needed to remove them
37
successive ionisation energies of an element
- give information about its electron configuration
- increase because it is harder to remove electrons from an already positive ion
- as electrons are removed, attractive forces increase, due to decreased shielding and an increase in proton to electron ratio
38
frequency and energy
E = hf
directly proportional
39
frequency, energy, wavelength
f = c/λ
frequency & energy are inversely proportional to wavelength
