S1.2: The Nuclear Atom + S1.3 Electronic Configurations Flashcards
Basics of the nuclear atom
Atoms -> electrically neutral (protons=electrons)
Masses and charges -> compared bc they are so small
Nucleus:
- proton+neutrons (nucleons)
- dense -> most of mass concentrated
Proton and neutron -> similar mass
Neutron -> neutral
Proton -> +1 charge
Nucleus surrounded by electrons (cloud)
Electron -> small mass (0), relative charge of -1
Electrostatic attraction between the positive nucleus+negatively charged electrons -> hold atom together
What is the atomic number?
Number of protons in the nucleus of a neutral atom
(Z)
Atoms of the same element -> same number of protons -> identifies the element
What is the mass number?
Number of protons and neutrons in the nucleus of an atom
(A)
Atoms of the same element can possess a different number of neutrons in their nuclei -> isotopes
What is an isotope?
Atoms of the same element that have the same atomic number but a different mass number
Relative abundance of isotopes -> effects mass number
Difference in mass -> have different properties (mp, bp, etc.)
What is relative atomic mass?
Weight mean of the naturally occurring isotopes of an element relative to 1/12 of a carbon-12 atom
Relative amount of the naturally occurring isotopes can be obtained from a mass spectrometer
What is the Bohr theory?
Small nucleus of protons and neutrons surrounded by circular electrons
Each shell/energy level can hold a max amount of electrons
Energy levels becomes greater as they get further from nucleus
Electrons fill the energy levels in a specific order
What is a principal energy level/principle quantum shell?
(What is a principal quantum number?)
Principle energy level -> each ‘ring’ around the nucleus is one principle energy level
Principle quantum number -> used to number the energy levels or quantum shells
Each holds a fixed number of electrons
-> lower -> closer to nucleus
-> higher -> further from nucleus
What is a sub-level?
Study of ionization energy and periodic properties -> main energy levels split into sub levels
Principle quantum shells -> sub levels
-> s, p, d
Elements with 57+ electrons -> also f sub-level
Energy of electrons increases: s -> p -> d
Rules and principles of sub-levels: heisenberg’s uncertainty principle
You cannot determine the position+momentum of an electron at the same time
You can’t say where an electron is exactly
Rules and principles of sub-levels: aufbau principle
Electrons enter the lowest available energy level
Rules and principles of sub-levels: Pauli’s exclusion principle
No two electrons can have the same four quantum number
Two electrons can go in each orbital provide they are of the opposite spin
Rules and principles of sub-levels: hund’s rule of maximum multiplicity
Orbitals of equal energy -> electrons try to remain unpaired
2 electrons in 1 orbital -> electrostatic repulsion between them
Separate orbitals -> reduced repulsion -> increased stability
What are orbitals?
Sub-levels -> broken down into (atomic) orbitals
Exists at specific energy levels
electrons can only be found at these level, not between
Orbitals hold up to 2 electrons (so long as they have opposite spin)
What are the different shapes of orbitals?
An orbital is a 3D shape showing you where one is most likely to find an electron because (according to Heisenberg) you cannot say exactly where an electron is only where it might be
S -> spherical
- 1 in every principle level
- increase every level
P -> dumb-bell
- 3 in every principle level
- x, y, z axes (perpendicular to each other)
- lobes become larger/longer with every level
__________(not required)
D -> various (5 in every principle level 3+)
F -> various (7 in every principle level 4+)
Order of filling orbitals
Not always filled in numerical order -> principle energy levels get closer together as you get further from nucleus
-> overlap of sublevels (ex: 4s -> 3d)
What is ground state (in the context of electron configuration)?
Most stable electron configuration of atom
Lowest amount of energy
Achieved by: filling the subshells of energy with the lowest energy first (Aufbau principle)
Orbital in the sub-shells have same energy
Said to degenerate -> energy in x orbital is same as y orbital
Order of sub-shells filled doesn’t follow regular pattern at n=3+
Électron configuration ‘levels’
Principle quantum shell/principle energy level
-> lower principle quantum number fill up first
Sub-shell/sub-level
-> s < p < d < f
-> exceptions: 4s, 3d
Orbitals
S -> 1 orbital -> 2 electrons
P -> 3 orbitals -> 6 electrons
D -> 5 orbitals -> 10 electrons
F -> 7 orbitals -> 14 electrons
What is spin pair repulsion?
Hunds rule?
Pauli’s exclusion principle?
Electrons can be imagined as small spinning charges
Rotate around their own axis clockwise or anti.
Same spin -> repel
-> spin-pair repulsion
Repulsion between electrons -> but still occupy same region of space in orbitals
-> energy required to jump > inter-electron repulsion
Hunds rule:
Electrons with same spin -> will occupy separate orbitals in same subshell to minimize repulsion
Pauli’s exclusion principle:
An orbital can only hold 2 electrons
MUST have opposite spin
Writing electron configuration: full electron configuration or shorthand
Full: describes the arrangement of ALL electrons from 1s +
Shorthand: electron configuration incudes using the symbol of the nearest preceding Nobel gas to account from however many electrons are in that noble gas
Rest goes as normal
Ions and electron configuration notation
Neg ions -> adding electrons to outer subshell
Pos ions -> removing electrons from outer subshell
Transitions metals:
- fill 4s sub-shell before 3d
- also loses electrons from 4s before 3d
Electron configuration and the periodic table
Periodic table -> 4 main blocks
S block -> valance electron in s orbital
- alkali/alkali earth metals
P block -> valance electrons in p orbital
- non-metals
D block -> valance electrons in d orbital
- transition metals
F blocks -> valance electrons in f orbital
What are the exceptions to the aufbau pinrciple?
Chromium:
Cr -> (Ar)3d5, 4s1 not (Ar)3d4, 4s2
Copper:
Cu -> (Ar)3d10, 4s1 not 3d9, 4s2
Due to the real electron configurations being more energetically favorable
-> promote electron -> atoms achieve full/half full d-sub shell
What is the electromagnetic spectrum?
Range of frequencies that cover all electromagnetic radiation and their respective wavelengths and energy
Divided in bands/regions
Shows relationship between frequency, wavelength and energy
All travel at same speed -> speed of light -> 3 x 10^8 m/s
What is frequency and wavelength?
(+ their relationship and relation to velocity)
Frequency (v): how many waves pass per second
Wavelength (λ): distance between 2 consecutive peaks on the wave
Frequency and wavelength -> inversely proportional
c (velocity) = λ (wavelength) x v (frequency)
What is a continuous spectrum?
What is a line spectrum?
Continuous (of visible region) -> all the colors of the spectrum
Line (of visible) -> only shows certain frequencies
-> can tell us that the emitted light from atoms can be certain fixed frequencies -> quantized
The emission spectra: explained
Electrons move rapidly around nucleus in energy shell -> energy increase -> jump to high energy level
Process reversible -> when electron return to original level -> emit energy
- same frequency of energy, just emitted not absorbed
Energy emitted -> mix of frequencies -> correspond to the many possibilities of electron jumps
Each element -> unique emission spectrum (not continuous) -> discrete lines at particular wavelengths that correspond to differences between energy levels
Emitted energy -> visible region -> analyzed by passing through diffraction grating
Line emission spectra
Packets of energy -> quanta
Each line -> specific value
- suggest electron can only possess a limited choice of allowed energies
Discrete lines which converge towards high energy end of spectrum -> electron reaching max amount of energy
- max -> ionization energy (become ion)
Electron jump down to n=2 -> visible
How does which energy level an electron fall to effect the type of wave it produces?
Families of lines correspond to electrons jumping from higher -> Lower levels
N -> 3+ -> infrared
N -> 2 -> visible
N -> 1 -> UV
visible emission spectrum is due to electrons falling to the n=2 level from higher levels
- n = 3 → n = 2 → 656nm
- n = 4 → n = 2 → 486 nm
- n = 5 → n = 2 → 434 nm
- n = 6 → n = 2 → 410 nm
-> visible with electron spectrometer
Put in order from biggest to smallest:
Principle energy level/principle quantum shell
Orbitals
Sub-levels/sub-shells
Principle energy level/principle quantum shell
-> principle quantum number
-> 1 ring on bohrs atom diagram
Sub-levels/sub-shells
-> s, p, d (,f)
Orbitals
-> can hold 2 electrons of opposite spin
-> electrons only found at these levels
What is a mass spectrometer/mass spectrometry?
A device which can be used to find the percentage abundance of the isotopes of an element
Mass spectrometry -> used to show the identity of isotopes
-> now used to calculate molecular masses and characterize new compounds
What are the basic processes of mass spectrometry?
Sample is vaporized
Sample is ionized -> gaseous atoms bombarded by electrons -> positive ions
Ions are charged -> accelerated by electric field
Ions deflected by electric/magnetic field
Each ion produces signal -> detected as mass-to-charge ratio (m/e or m/z)
Electric or photographic methods
-> large m/z -> deflected the least
-> small m/z -> deflected the most
Mass spectrum produced/portrays different ions in order of their m/z value
How does mass effect the amount of deflection?
Radius of path -> depend on the value of the m/z ratio
Ions or heavier isotopes -> larger m/z -> larger curve
Due to most ions having 1+ charge -> amount of separation depends on MASS
When ion has 2+ charge -> deflected more -> m/z value HALVED
