U1.3: Electronic Configurations Flashcards
What is the trend of the EM Spectrum and what does it consist of
RW, MW, IR, VL, UV, XR, GR
Increasing energy and frequency
decreasing wavelength
What is the difference between a continuous spectrum and a line spectrum
Continuous spectrum: Shows all the frequencies of light
Line spectrum: Shows specific frequencies of light
What is the trend of the hydrogen emission spectrum
- Made up of discrete energy levels containing electrons that have a particular energy
- The energy levels converge as they successively increase hence, becoming higher in energy
- Electrons can either be in an excited (absorbing energy) or relaxed (emitting energy) state
Why do lines converge at higher energy levels
The difference between the energy levels is too little hence, it reaches the converging limit.
What is the converging limit
After this point, the atom has no control over the electron hence, the atom turns into an ion.
What are higher and lowest energy level states described as
Higher energy levels: Excited states
Lowest energy level: Ground state (all elements are initially at a ground state)
How can you add energy to the hydrogen atom
- Heat
- Electricity
- Light
What happens when you add energy to the hydrogen atom
The electron may jump from the inner shell to the outer shell as the electron can only exist on those shells
What does it mean when an electron gets excited
Once the electron goes to an outer shell
What are the 2 main words to describe energy levels
Concentric: They’re within one another
Converging energy levels: The further they get from the nucleus, the closer they get to one another
What happens when the electron moves from one energy level back down to the energy level beneath it
It releases a little amount of energy according to ROYGBIV.
What happens when an electron gets so excited that it hits n=infinity
An ion is made as the electron gets ripped off
What is the trend in energy sub-levels?
s<p<d<f
How to write the electronic configuration
1s
2s, 2p
3s, 3p,3d
4s,4p,4d,4f
5s,5p,5d,5f
—————-
2, 6, 10, 14
When drawing the electronic configuration of an ion, where is the electron gained/taken off from?
The highest energy level (so following s<p<d<f and from the greatest number like 4 or 3)
What is an orbital
Where the probability of finding an electron is maximum
What is the pauli exclusion principle
Each orbital box is shown with a half-upward arrow and a half-downward arrow
What are degenerate orbitals
Orbitals with the same energy level
What is Aufbau’s principle
Lowest available energy orbitals fill before higher energy orbitals do
What is Hund’s rule
Every degenerate orbital in a sublevel’s singly occupied before any orbital’s doubly occupied.
All electrons in singly occupied orbitals have the same spin
What 2 elements are the exceptions to Aufbau’s principle
Copper and Chromium
How do you write condensed formulas of compounds
Find the closest noble gas and write the rest after the noble gas is written like: [Insert Noble Gas]
What is the ionisation energy
The minimum energy required to eject an electron from 1 mol of gaseous atoms to form 1 mol of gaseous 1+ ions under standard conditions
What is shielding
Where inner electrons reduce the attractive force between the positively charged nucleus and outer electrons
What is the trend on the first ionisation energy
- Decreases down the groups - number of sublevels increases, valence electrons are more greatly shielded from the pull of the nucleus by the inner electrons
- Increases across the periods - number of protons in nucleus increases hence, outermost electrons are held closer to the nucleus by the increased nuclear charge and the shielding effect remains constant as number of inner electrons doesn’t change
What is the trend for period 2 elements in terms of ionisation energy? (B and Br)
- electron in 2p subshell of Boron is slightly further from nucleus
- hence, there is more shielding in the 2s subshell
- therefore, electron is easier to remove and so IE decreases
What is the trend for group 15 and group 16 elements in terms of ionisation energy?
N: 1s2 2s2 2p3
O: 1s2 2s2 2p4
- Paired electrons in the same 2p orbital repel more than unpaired electrons it’s slightly easier to remove
In what conditions is a sublevel most stable
Half filled, fully filled and empty
What is the number for planck’s constant
6.63 x 10^-34
What is the 3 common trends of ionisation energies (IE)?
- As the charge of the gaseous atom / ion increases, IE increases
- As shielding increases, IE decreases
- As atomic radius increases, IE decreases
What are the periodic trends of the ionisation energy?
- IE increases across the whole periodic table
- IE increases up the groups
- IE increases going left to right for the periods
Why does the IE increase
Nuclear charge increases and there is similar shielding thus, there is a greater force of attraction between v.e- and nucleus thus, IE increases
What are successive ionisation energies?
They measure the energy required to eject electrons from an ion/atom
Describe the photons released when the electron jumps back to the n=1, n=2 and n=3 energy level
Jump back to the n=1 - UV photon
jump back to the n=2 - VL photon
jump back to the n=3 - IR photon
What is the electronic configuration of Cr and Cu
Cu - [Ar] 4s1 3d10
Cr - [Ar] 4s1 3d5
When do we talk about Electron-Electron repulsion?
- Only if the elements are in the same period and if the elements are non-metals
In the line spectra, what does the furthest line show?
Highest wavelength
What does the line emission spectrum of hydrogen look like?
In terms of decreasing wavelength
How do you know which species will be deflected the most in a mass spectrometer?
The species with the highest negative charge (because that species gains electrons) and the lowest possible atomic mass
State, giving a reason, whether the chlorine atom or the chloride ion has a larger radius.
Cl- due to more electron-electron repulsion
Outline why the chlorine atom has a smaller atomic radius than the sulfur atom.
Cl has a greater nuclear charge therefore causing a stronger pull on the outer electrons.
What are the steps involved in Mass Spectrometry?
VIADD
- Vaporization
- Ionisation
- Acceleration
- Deflection
- Detection
why would the relative atomic mass of an element be greater if the other element has a greater atomic number
The element with the greater atomic mass has more neutrons
Define first ionisation energy and state what’s meant by the term periodicity (2 marks)
1st ionisation energy:
- The minimum energy required to remove 1 mole of e-
- from 1 mole of gaseous atom
- to form 1 mole of gaseous ion with a 1+ charge
Periodicity:
- Repeating pattern of chemical and physical properties
Define second ionisation energy (1 mark)
Draw the equation if only one mark e.g.:
- M(g)+ –> M(g)2+ + e-
How is the successive ionisation energy data for Na related to its electron configuration
- Na: 1s2 2s2 2p6 3s1
- One of the biggest difference is between the 1st and 2nd IE
- The 2nd electron to be removed is removed from inner shell, closer to nucleus, less shielded, more attracted to the nucleus therefore, it’s much more difficult to remove
- There is also only one outer shell electron as it’s in group 1.
How to read the successive ionisation energy data
- Goes from backwards (outer electrons to inner electrons)
if electron config was 2.8.4, first 4 electrons would be the valence electrons and the first big jump would be from the 4th and 5th IE as the electron would be much more difficult to remove.
Why is there a sharp increase in IE between the ____ and _____ ionisation energies?
- 2 electrons are removed from inner shell
- they are much closer to the nucleus
- this means they are much lesser shielded
- and that the 2 electrons are more attracted to the nucleus therefore, they are difficult to eject
Modal answer for IE generally increasing across a period reasons
Across a period,
- # of protons increases as the nuclear charge increases
- Shielding remains relatively constant as e- fill in the same shell
- Net attraction of outer shell e- to nucleus gets stronger
Therefore, 1st IE generally increases and atomic radii decreases
Modal answer for IE generally decreasing down a group reasons
Down a group,
- There are more protons as the nuclear charge increases
- There are more e- shells
- As atomic radii increases, shielding increases
- Overall net attraction between outer shell e- and the nucleus gets weaker
Therefore, 1st IE decreases and atomic radii as well as size increases
Limitation of the Bohr Model
- doesn’t explain spectral lines of atoms with more than 1 electrons hence, the nature of electrons needs to be reconsidered
- doesn’t show orbitals
What were the 2 assumptions made about the atom from the gold atom experiment
- With 4He2+ ion passing straight through the gold atom –> proved most of the gold atom contains empty space
- Very few 4He2+ deviating largely from their path –> proved the nucleus of the atom is positively charged
How to conevrt from kJ/mol to J/atom?
- IE / 6.02 x 10^23
- x1000
which Cu or Cu2+ has a greater IE?
- 2+ charge means greater charge hence, ion with 2+ charge has a greater IE
Where do the lines on the line spectra converge?
- Converge at higher energy, higher frequency and lower wavelengths
What does “occupied” mean in an aotm
- Fully filled
Equation representing second ionisation energy for Cu
Cu+(g) –> Cu2+ (g) + e-
Describe the relationship between frequency and wavelength
- Inverse Relationship
if an electronic config for one element is [Ar] 3p1 and the other is [Ar] 3s2, why would the 1st IE of the 1st element be less than the 2nd element?
- element with [Ar] 3p2 electronic config. has a higher stability than the other element thus, a greater energy is required to remove the e- from the s orbital than it does to remove the e- from the p orbital in the other element
if an electronic config for one element is [Ar] 3p3 and the other is [Ar] 3p4, why would the 1st IE of the 1st element be less than the 2nd element?
- element with [Ar] 3p3 electronic config. has a partially filled p orbital thus, a higher stability than the other element therefore, a greater energy is required to remove the unpaired e- than it does to remove the paired e- from the p orbital in the other element as there is increased electron-electron repulsion between the paired electrons
