Unit 1: electronic structured and periodic properties of elements

What is 𝛎 (nu)

symbol for frequency

does electromagnetic radiation travel at different speeds?

no! all electromagnetic radiation travels at 3 × 10^8 m/s or the speed of light.

whats the speed of light

3 × 10^8 m/s = C

3 characteristics of waves 

• lengths ( λ lambda)

• Amp (height)

• Frequency ( 𝛎 Nu )

how is wavelength and frequency related?

C = λν, where c is the speed of light

constructive interference

when the amps of two waves combine- making extra stretched wave

destructive interference

when tops and bottoms of 2 seperate waves cancel out leaving a line

Planck’s theory

that energy can only be absorbed or released from atoms in certain amounts - basically energy releases in these little bursts called quanta

Quantum def:

Smallest amount of energy that can be emitted/absorbed as electromagnetic radiation

what is Plancks constant ( h)

= 6.626 × 10 ^ -34 js

relationship between energy and frequency

E = hv

energy = plancks Constant x Nu

this describes that energy comes in little bursts proportional to its frequency.

what does the photoelectric effect describe?

• when the surface of a metal is hit by electromagnetic radiation it ejects electrons 

• however in order for this to happen the photons in the electromagnetic waves must have enough energy 

• based on the derived equation 

E = h(c/λ) ,

• the shorter wavelength something has the more energy it will have - therefore the more electrons it will eject

• the maximum wavelength for emitting electrons is 550 nm ( which is equivalent to green light) 

Bohr Model 3 main ideas

electrons are confined to specific energy states ( orbits)

1)only orbits of specific radii are permitted for electrons in an atom( electrons must stick to that radii??)

2) an electron in a permitted orbit has a specific energy

3)energy is emitted or absorbed by an electron as it moves from one energy state to another

what happens as electrons move between energy states?

• as electrons move to lower energy states light is emitted

• as electrons move up energy states light is absorbed

Matter Waves

De Brogile’s theory that if light can have material properties- matter should exhibit wave properties 

basically matter has waves - for most objects however the wavelength they emit is so small its pointless 

leading to this eq: λ= h/mv 

where h is plancks constant ( I assume)

m is mass, and v is velocity 

Uncertainty principle

Since everything ( materials and light) have both material and wave properties it sets a limit on how precisely we can know the location and momentum of an object

TLDR : cant find the momentum and location of an object at the same time ( measuring one will interfere with the other ) - obv this applies to subatomic particles not like everyday objects

formula involving uncertainty principle

(uncertainty of x )(uncertainty of mv) ≥ h/4pi

what is Ψ?

Ψ wave function ( symbol being psi) describes the behaviour of a quantum mechanical object ( too advanced for me but what is relevant→) Ψ² gives electron density - a region of high electron density = high probability of finding an electron

basically there are orbitals ( not fixed orbits) where electrons are most likely to be found

What do the three quantum numbers tell u ( theory based)

the three quantum numbers tell you where the orbital is , what shape it is and how its oriented. 

What is the principle quantum number n?

n is a positive integer - that describes the main energy lvl( or shell), basically how far the orbital is to the nucleaus

n also correlates to the amount of sub shells that are in that energy lvl 

a greater n value correlates to being in a shell further away from the electron nucleaus and therefore higher in energy

What is the angular momentum quantum number 𝓁?

𝓁 tells u the shape of the orbital.

maximum possible value for 𝓁 is n -1 ( 𝓁 can be less than the number quite frequently tho)

and each 𝓁 value correlates to a letter that describes the shape of the orbital. 

𝓁: 0 1 2 3 

   s p d f 

What is the magnetic quantum number M (subscript 𝓁) ?

correlates to one orbital with a unique orientation is space ( each orbital with a different orientation is a seperate orbital) 

maximum value depends on 𝓁 ( thats why 𝓁 is in subscript) , m𝓁 can take on values from -𝓁 to 𝓁. So basically the greater the 𝓁 value, the more complex the shape and the more possible orientations available ( or possible m𝓁 values)

how can u figure out fuckass number of nodes/node types?

total nodes = n-1

however within the total nodes there are two types, angular and radial

angular nodes define the shape - ( having one sphere, vs 2 , vs a clover shape) and the amount of angular nodes is determined by 𝓁

radial nodes are whatever is left over: basically (n-1) - 𝓁 = amount of radial nodes, since in many cases 𝓁 < n - 1

the radial nodes appear in a certain radius from the nucleus, cutting the orbital up into rings kinda

• overall shape tho ( like amount of circles around the nucleaus ) is determined by the angular nodes, then the leftover radial nodes cut the orbital up into rings…

how can u tell where nodes are on diagram?

cause the color of the orbital changes as it passes thru a node ( goes from neg to pos if nodes are 0) color of orbital changes as sign changes

whats the funky d orbital? 

dz² , its a nodal cone 

4th quantum number? M (subscript S)

spin magnetic quantum number - shows wether electron is spin up or spin down values include +1/2 for spin up and -1/2 for spin down.

Pauli’s exclusion principle

no 2 electrons can have the same set of 4 quantum numbers, therefore if they exist in the same orientiaon (m𝓁) one has to be spin up and one has to be spin down

• an example of Paulis exclusion principle would be 3 elections in one box ( or 3 elections in the same sub shell orientation?)

Hunds Rule:

for orbitals of the same energy ( n), the lowest energy is attained when the number of electrons with the same spin is maximized. Filling all the spin ups first will help with stability - minimizes electron-electron repulsion.

What orbital does D start with?

D starts with 3d, even though it’s in the 4 line.

what orbital does f start with?

f starts with orbital 4 even though its in line 6

2 Anomaly electron configurations ( memorize)

• Cr [ Ar ] 4s^1 3d^5 ( all shells are half filled for stability)

• Cu [ Ar ] 4s^1 3d^10 ( all shells are again half filled for stability)

Aufbau Principles 

1) lower energy orbitals (n) fill with electrons first’

2) Any orbital can hold up to 2 electrons. 2 electrons in the same orbital must have opposite spins - Pauli exclusion principle 

3)if 2 or more orbitals with the same energy are available, one electron fills each shell until are all half filled ( improves stability ) - Hands Rule

what is orbital digram? ( notation)

what is sub shell notation?

draw different orientations for d orbital

electron configuration of cations and anions?

Anomalous Electron Configurations

Chromium (Cr) - [Ar]4s¹3d⁵ (3d gets fully filled up before 4s cause its more stable)

Copper (Cu) - [Ar]4s¹3d¹⁰ (3d fully filled cause its more stable)

Atomic Size ( how does it vary across periodic table)

Atoms get larger as u go down ( adding shells)

Atoms get smaller as you go right (Zeff gets larger)

Biggest atom?

Cesium

Smallest atom

Helium

What is Zeff?

charge experienced by an electron from the nucleaus ( attraction), taking into account the shielding effect from the inner electrons 

( the charge it is experiencing isn’t as strong because of the other inner electrons blocking it)

Zeff equation

Zeff = z - s ( atomic number(#of protons) - screening constant )

what is screening constant?

screening constant is portion of nuclear charged shielded from the valence electrons by inner electrons near atom

• s is given by # of core electrons in an atom

how zeff changes across perodic table 

Zeff increase as we go to the right ( cause more protons pulling on the electrons ) - therefore the atom is tighter ( protons pull electrons tighter) and atom is smaller 

Zeff gets slightly larger as we go down ( slightly) 

BUT fully filled d and f orbitals should be treated as core electrons - they screen nucleaus 

Are cations larger or smaller than ogs?

cations are smaller, since they have one less electron , the zeff increases ( I think??) and therefore atom gets tighter/smaller

are anions bigger or smaller than og?

bigger! electrons get added to largest orbital, total electron - electron repulsion increases so they space out more

If u have a series of ions with the same amount of electrons ( eg, 0^-2, F-,Ne,Na+, etc) how do their size differ?

ones that are cations ( have more protons) will be smaller

ionization energy?!?!?

minimum energy required to remove an electron - larger ionization energy = harder to remove electron

I1, I2… ( ionization energy)

basically the idea that I1 is the energy to remove an electron from an atom, I2 is the energy required to remove another electron from teh same atom ( I2 will always be greater)

how does ionization energy change across periodic table

ionization energy increases as u go right - cause as Zeff gets larger its harder to remove an electron ( harder to remove electron from a tighter electron )

ionization energy exceptions 

• Boron, Aluminum ( removing the first p electron because prev s electrons are closer to the nucleaus )

• (oxygen, sulphur) removing fourth p electron due to electron - electron repulsion?? 

where does ionization energy change across product table

ionizionization energy increases as we go right ( with a higher zeff its harder to remove), and decreases slightly as we go down

elements with most and least ionization energy

most = helium least = caesium

how electron configuration of ions work

• start with og atom

• remove or add to the most accessible orbital

• electrons are removed from orbital with the highest n value first ( so 4s before 3d)

electron affinity

energy change when an atom gains an electron ( measures how much an atom wants an extra electron) - energy is usually released when an atom adds an electron

unfavourable electron affinity examples

atoms dont want an electron when it is going intoa new sub shell or pairing electrons in p orbitals ( new sub shell is too much work - they were stable before) ( pairing p electrons leads to highg electron - electron repulsion)

Anything that ends at the end of s orbital - moving into p is too much work

Nitrogen specifically cause the shell doesnt want to pair more p electrons?

every Nobel gas but neon?? which still has low electron affinity

negative electron affinity??

negative electron affinity means it really wants an electron ( sounds counterintuitive)

but if Ea is negative it means the atom gaining an electron releases energy so its favourable

if Ea is positive it means it requires energy to gain an electron so its not favourable