S1.2: The Nuclear Atom + S1.3 Electronic Configurations

Question 1

Basics of the nuclear atom

Answer 1

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

Question 2

What is the atomic number?

Answer 2

Number of protons in the nucleus of a neutral atom
(Z)

Atoms of the same element -> same number of protons -> identifies the element

Question 3

What is the mass number?

Answer 3

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

Question 4

What is an isotope?

Answer 4

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.)

Question 5

What is relative atomic mass?

Answer 5

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

Question 6

What is the Bohr theory?

Answer 6

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

Question 7

What is a principal energy level/principle quantum shell?
(What is a principal quantum number?)

Answer 7

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

Question 8

What is a sub-level?

Answer 8

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

Question 9

Rules and principles of sub-levels: heisenberg’s uncertainty principle

Answer 9

You cannot determine the position+momentum of an electron at the same time
You can’t say where an electron is exactly

Question 10

Rules and principles of sub-levels: aufbau principle

Answer 10

Electrons enter the lowest available energy level

Question 11

Rules and principles of sub-levels: Pauli’s exclusion principle

Answer 11

No two electrons can have the same four quantum number
Two electrons can go in each orbital provide they are of the opposite spin

Question 12

Rules and principles of sub-levels: hund’s rule of maximum multiplicity

Answer 12

Orbitals of equal energy -> electrons try to remain unpaired

2 electrons in 1 orbital -> electrostatic repulsion between them

Separate orbitals -> reduced repulsion -> increased stability

Question 13

What are orbitals?

Answer 13

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)

Question 14

What are the different shapes of orbitals?

Answer 14

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+)

Question 15

Order of filling orbitals

Answer 15

Not always filled in numerical order -> principle energy levels get closer together as you get further from nucleus
-> overlap of sublevels (ex: 4s -> 3d)

Question 16

What is ground state (in the context of electron configuration)?

Answer 16

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+

Question 17

Électron configuration ‘levels’

Answer 17

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

Question 18

What is spin pair repulsion?
Hunds rule?
Pauli’s exclusion principle?

Answer 18

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

Question 19

Writing electron configuration: full electron configuration or shorthand

Answer 19

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

Question 20

Ions and electron configuration notation

Answer 20

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

Question 21

Electron configuration and the periodic table

Answer 21

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

Question 22

What are the exceptions to the aufbau pinrciple?

Answer 22

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

Question 23

What is the electromagnetic spectrum?

Answer 23

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

Question 24

What is frequency and wavelength?
(+ their relationship and relation to velocity)

Answer 24

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)

Question 25

What is a continuous spectrum?
What is a line spectrum?

Answer 25

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

Question 26

The emission spectra: explained

Answer 26

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

Question 27

Line emission spectra

Answer 27

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

Question 28

How does which energy level an electron fall to effect the type of wave it produces?

Answer 28

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

Question 29

Put in order from biggest to smallest:

Principle energy level/principle quantum shell

Orbitals

Sub-levels/sub-shells

Answer 29

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