Electronic Structure And Periodicity Flashcards
Electromagnetic spectrum
-Includes all forms of electromagnetic radiation (radiant energy)
~gamma rays, X rays, Ultraviolet, Visible light (ROYGBIV), Infrared, microwaves, radio waves
-characterized by frequency, wavelength, and amplitude
Gamma rays
-shortest wavelength
(10^-12 m)
-highest frequency
(10^20 Hz)
X rays
- wavelength = 10^-10 m
- frequency = 10^18 Hz
Ultraviolet
- wavelength = 10^-8 m
- frequency = 10^16 Hz
Visible light
Wavelengths: -Violet = 380nm (3.8x10^-7 m) -Indigo = 420nm -blue = 490nm -green = 520nm -yellow = 600nm -orange = 690nm -Red = 780nm (7.8x10^-7 m) Frequency = ~10^15 Hz (depending on color)
Infrared
- wavelength = 10^-6 m
- frequency = 10^14 Hz
Microwaves
- wavelength = 10^-3 m
- frequency = 10^10 Hz
Radio waves
- wavelength = 1 m
- frequency = 10^8 Hz
Frequency
- (v)
- the # of wave peaks that pass a given point, per unit time
- 1/s
Wavelength
- (lambda)
- the distance from 1 wave peak to the next
- m
Amplitude
- (A)
- the height of the wave (from center line to peak)
Frequency and wavelength relationship
- inversely related
- longer wavelength = lower frequency
- shorter wavelength = higher frequency
- Wavelength x Frequency = speed
- wavelength = speed/frequency
- frequency = speed/wavelength
- speed = the speed of light (2.998x10^8) unless otherwise specified
Diffraction and interference
- diffraction is the bedding of light around an object
- interference occurs when 2 or more waves superpose to form a new wave
- constructive = larger wave
- destructive = smaller or no wave
Photoelectric effect
-irritating a clean metal surface with light causes electrons to be ejected from the metal
Planck’s postulate
- a beam of light behaves as if it were a stream of small particles (photons)
- the energy of the photons is related to their frequency and wavelength
E = hv or E = hc/y
-h= Planck’s constant (6.626x10^-34)
-v = frequency
-c = speed of light (2.998x10^8)
-y = wavelength
-often asks for answer in per-mol basis, not per-photon (as answer gives), multiply per-photon energy by Avagadro’s number (6.02x10^23) to find per-mol energy
- higher frequency & shorter wavelength = higher E
- lower frequency & longer wavelength = lower E
Intensity & frequency
- energy of a photon depends on its wavelength and frequency (not intensity)
- intensity of a light beam is a measure of the number of photons in the beam
- frequency is a measure of the energy of the photons
Work function
- the amount of energy necessary to eject an electron from a metal
- depends on the metal (specific to each)
- lowest for group 1A & 2A elements
Atomic line spectra
- atoms give off light when energetically excited
- light is not continuous and only occurs at certain wavelength
- different atoms produce distinct spectra that are unique to that element
Bohr model
- proposed that electrons move in circular orbits around the nucleus
- each orbit has its own radius which is directly related to energy (n)
- as radius increases, energy increases(E-final)-(E-initial) = hv
- when an electron falls to a lower energy it releases a photon
- when an electron jumps to a higher energy, it has absorbed energy
*an electron cannot reside between orbits
Balmer-Rydberg equation
E = R [(1/m^2) - (1/n^2)]
- E = energy (J)
- R = 2.178x10^-18
- m = final energy level
- n= initial energy level
*can find wavelength or frequency using the energy
De Broglie equation
- perhaps matter is wavelike, as well as particle-like
- double-slit experiment
(Y) = h/mv
- y = wavelength
- h = planck’s constant (6.626x10^-34)
- m = mass
- v = velocity
quantum mechanical model of an atom
- Erwin Shrodinger
- it is impossible to know precisely where an electron is and what path it follows
- the very act of determining an electrons position (requires input of energy) causes the position to change
Wave function
- orbital
- found by solving a wave equation
- characterized by 4 parameters (quantum numbers)
- n
- l
- m(l)
- m(s)
(n)
Principal quantum number
- a positive integer
- relates directly to size and energy of orbital
- for atoms with more than 1 electron, the energy level of an orbital depends on both “n” and “l”
- as n increases, the number of allowed orbitals increases and those orbitals become larger (allows an electron to be further from the nucleus)
- the energy of an electron in the orbital increases as quantum number n increases
(l)
Angular momentum quantum number
- defines 3D shape of the orbital
-can have any integral value from 0 to n-1
- so, within each shell there are “n” different shapes for orbitals
0 = s
1 = p
2 = d
3 = fm(l)
Magnetic quantum number
- defines spatial orientation of the orbital with respect to a standard set of coordinate axis
- can have any integral value from -l to l (L not 1)
S orbitals
- all are spherical
- size and energy increases in successively higher shells (1s<2s<3s…)
* larger orbitals are higher in energy because their electrons have a higher probability of being found further from the nucleus, and it takes energy to separate a positive (nucleus) & negative (electron) charge
Nodes
A surface with 0% probability of finding an electron
- corresponds to zero amplitude sections of a wave
- either side of a node corresponds to an either (+) or (-) wave phase
P orbitals
- Dumbbell shaped
- consist of 2 identical lobes on either side of the nucleus separated by a planar node which cuts trough the nucleus
* each of the 2 lobes represent opposing wave phases (+ & -)
* only lobes of the same phase can interact to form covalent bonds - 3 p-orbitals per shell
* each oriented along 1 of the 3 major axis (px,py,pz) giving - higher “p” shells are larger and extend farther rom the nucleus
D & F orbitals
D-orbitals:
- 4 of the 5 are cloverleaf shaped with 4 lobes and 2 planar nodes passing through the nucleus (dxy, dyz, dxz, dx^2-y^2)
* the 5th has 2 lobes on the z axis (1 up, 1 down), with a spherical (donut) region on the x/y plane (dz^2) - alternating lobes have opposing phases
- all orbitals in a given shell have the same energy level (degenerate)
F-orbitals:
- 8 lobes separated by 3 nodal planes through the nucleus
(Ms)
Magnetic spin quantum number
- relates to property called electron spin
- spinning charge gives rise to tiny magnetic field - can be either +1/2 (up arrow) or -1/2 (down arrow)
- independent of other quantum numbers
Pauli exclusion principle
- No 2 electrons in an atom can have the same 4 quantum numbers
* an orbital can hold 2 electrons which must have opposite spin signs
Effective nuclear charge (Zeff)
Zeff = Zactual - electron shielding
- the difference in energy between subsets results from electron repulsion
- the repulsion of outer-shell electrons by inner-shell electrons
- outer shell electrons are pushed away and held less tightly by the nucleus
*a higher Zeff corresponds to lower energy
Degenerate orbitals
Orbitals that have the same energy level
Ex: All 3 p orbitals in a given shell
Aufbau principle
- lower-energy orbitals must fill before higher-energy orbitals
- also unites Hind’s rule and the Pauli exclusion principle
Hind’s Rule
-if 2 or more degenerate orbitals are available, 1 electron goes in each until all are “half full”
Anomalous electron configurations
- have to do with unusual stability of half-filled and fully-filled sub shells in some atoms
* [Ar]4s1 3d5 - Involves the transfer of an electron from one shell to another which decreases the total energy of the atom by decreasing electron repulsion
Blocks
S-block = groups 1A & 2A P-block = groups 3A-8A D-block = transition metals F-block = lanthanides & actinides
Atomic radius
- determined as half the distance between the nuclei of 2 identical atoms when they are bonded together
* increases top to bottom
* decreases left to right
