Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemisry 1411 > Electronic Structure And Periodicity > Flashcards

Electronic Structure And Periodicity Flashcards

1
Q

Electromagnetic spectrum

A

-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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Gamma rays

A

-shortest wavelength
(10^-12 m)
-highest frequency
(10^20 Hz)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

X rays

A
  • wavelength = 10^-10 m

- frequency = 10^18 Hz

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Ultraviolet

A
  • wavelength = 10^-8 m

- frequency = 10^16 Hz

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Visible light

A 
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)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Infrared

A
  • wavelength = 10^-6 m

- frequency = 10^14 Hz

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Microwaves

A
  • wavelength = 10^-3 m

- frequency = 10^10 Hz

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Radio waves

A
  • wavelength = 1 m

- frequency = 10^8 Hz

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Frequency

A
  • (v)
  • the # of wave peaks that pass a given point, per unit time
  • 1/s
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Wavelength

A
  • (lambda)
  • the distance from 1 wave peak to the next
  • m
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Amplitude

A
  • (A)

- the height of the wave (from center line to peak)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Frequency and wavelength relationship

A
  • 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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Diffraction and interference

A
  • 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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Photoelectric effect

A

-irritating a clean metal surface with light causes electrons to be ejected from the metal

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Planck’s postulate

A
  • 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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Intensity & frequency

A
  • 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
17
Q

Work function

A
  • 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
18
Q

Atomic line spectra

A
  • 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
19
Q

Bohr model

A
  • 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

20
Q

Balmer-Rydberg equation

A

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

21
Q

De Broglie equation

A
  • 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
22
Q

quantum mechanical model of an atom

A
  • 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
23
Q

Wave function

A
  • orbital
  • found by solving a wave equation
  • characterized by 4 parameters (quantum numbers)
    - n
    - l
    - m(l)
    - m(s)
24
Q

(n)

A

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
25
Q

(l)

A 
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 = f
26
Q

m(l)

A

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)
27
Q

S orbitals

A
  • 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
28
Q

Nodes

A

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
29
Q

P orbitals

A
  • 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
30
Q

D & F orbitals

A

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

31
Q

(Ms)

A

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
32
Q

Pauli exclusion principle

A
  • 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
33
Q

Effective nuclear charge (Zeff)

A

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

34
Q

Degenerate orbitals

A

Orbitals that have the same energy level

Ex: All 3 p orbitals in a given shell

35
Q

Aufbau principle

A
  • lower-energy orbitals must fill before higher-energy orbitals
  • also unites Hind’s rule and the Pauli exclusion principle
36
Q

Hind’s Rule

A

-if 2 or more degenerate orbitals are available, 1 electron goes in each until all are “half full”

37
Q

Anomalous electron configurations

A
  • 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
38
Q

Blocks

A 
S-block = groups 1A & 2A
P-block = groups 3A-8A
D-block = transition metals
F-block = lanthanides & actinides
39
Q

Atomic radius

A
  • 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

Chemisry 1411 (8 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions