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CHM135 > Quantum Theory and Atomic Structure > Flashcards

Quantum Theory and Atomic Structure Flashcards

1
Q

define electromagnetic radiation (= electromagnetic energy or radiant energy)

A

energy propagated by perpendicular electric and magnetic fields that increase and decrease in intensity as they move through space as waves.

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2
Q

define wavelength ( λ)

A

distance between identical points on adjacent waves, such as from the crest/trough of one wave to the crest/trough of the next

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3
Q

what are the units for wavelength?

A

usually metres or nanometers, nm= 10^-9 metres or pico metres pm=10^-12

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4
Q

define frequency

A

the number of complete waves per second that pass through a given point

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5
Q

what are the units for frequency?

A

s^-1 = 1Hz

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6
Q

define the speed of a wave

A

the distance it moves per unit time (m/s)

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7
Q

state the equation relating, frequency, wavelength, and speed of light

A

v λ = c

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8
Q

state the 7 parts of the electromagnetic spectrum

A

highest frequency:
gamma rays
x rays
ultraviolet
visible
infrared
microwave
radio waves

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9
Q

define the amplitude

A

the height of the crest of the wave or the depth of the trough

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10
Q

for an electromagnetic wave, the amplitude is related to the

A

intensity of the radiation, or its brightness in the case of visible light

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11
Q

dimmer light

A

lower amplitude, less intense

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12
Q

brighter light

A

higher amplitude, more intense

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13
Q

what is the similarity between all electromagnetic waves?

A

they travel at the same speed through a vacuum

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14
Q

state the 2 distinctions between energy and matter

A
  1. refraction and dispersion
  2. diffraction and interference
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15
Q

refraction

A
  • when a light wave passes from one medium into another at an angle other, the speed of the wave changes
  • if the wave strikes the boundary between media at an angle other than 90, the change in speed causes a change in direction, and the wave continues at a different angle
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16
Q

dispersion

A

white light separates (disperses) into its component colours when it passes through a prism or another reframing object because each incoming wave is refracted at a slightly different angle

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17
Q

how does a particle of matter contrast to a wave of light in terms of refraction?

A

particle of matter does not undergo refraction, it just slows down gradually along a curved path (after entering water)

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18
Q

diffraction (inc diagram)

A

when a wave strikes the edge of an object and bends around it. if the wave passes through a slit as wide as its wavelength, it bends around both edges of the slit and forms a semi-circular wave on the other side of the opening
p294

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19
Q

how does a particle of matter contrast to a wave of light in terms of diffraction?

A

when you throw a collection of particles at a small opening, only some go through the opening.

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20
Q

interference (inc diagram)

A

when waves of light pass through two adjacent slits, the nearby emerging circular waves interact through the process of interference to form a diffraction pattern (p295)

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21
Q

describe the two possible outcomes for a diffraction pattern

A
  • in phase: if the crests of the waves coincide, the interfere constructively - the amplitudes add together to form a brighter region in the diffraction pattern
  • out of phase: if crests coincide with troughs, the interfere destructively - the amplitudes cancel to form a darker region
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22
Q

how does a particle of matter contrast to a wave of light in terms of interference ?

A

particles passing through adjacent openings continue straight paths, some colliding and moving at different angles

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23
Q

what did scientists observe that caused them to question their view of energy?

A
  1. blackbody radiation -> quantum theory of energy
  2. the photoelectric effect -> photon theory of light
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24
Q

explain blackbody radiation

A
  • when an object is heated, it emits radiation of various wavelengths, including visible light
  • the light of maximum intensity that is emitted shifts to shorter wavelengths as the temperature increases (red by cooler, blue by hotter)
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25
Q

why was the blackbody radiation observed confusing?

A

the classical wave model could not explain the relationship between the energy given off by a hot object and the wavelength of the energy emitted

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26
Q

define a blackbody

A

an idealised object that absorbs all the radiation incident on it, eg a hollow cube with a small hole in one wall approximates one.

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27
Q

what did Planck theorise?

A
  • an atom changes its energy state by emitting/absorbing one or more quanta (energy packet of fixed quantity)
  • the energy of the emitted/absorbed radiation is equal to the difference in the atom’s energy states.
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28
Q

planck’s equation and explain a more popular alternative

A

ΔE = E(emitted/absorbed radiation) = Δnhv

atoms can change energy only by integer multiples of hv, so the smallest energy change occurs when Δn=1:

ΔE = hv
or ΔE = hc/λ

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29
Q

h

A

Planc’s constant

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30
Q

v

A

frequency of radiation

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31
Q

E

A

energy of radiation

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32
Q

energy is directly proportional to ——- and indirectly proportional to ——–

A

frequency; wavelength

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33
Q

describe the photoelectric effect

A

when monochromatic light of sufficient frequency shines on a metal plate, a current flows (p297)

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34
Q

two confusing features that led scientists to the idea of the photoelectric effect

A

presence of a threshold frequency:
- wave theory associates the energy of light with its amplitude (intensity) not frequency (color)
- it thus predicts that an electron would break free when it absorbed enough energy from light of any color
- for current to flow, the light shining on the metal must have a minimum frequency (diff metals = diff frequencies)

absence of a time lag:
- wave theory predicts that with dim light there would be a time lag before the current flows as the electrons would have to absorb enough energy to break free
- current flows the moment light of the min frequency shines on the metal, regardless of the light’s intensity

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35
Q

describe photon theory

A

Einstein proposed that light is particulate, quantised into photons - tiny bundles of energy.

E(photon) = hv = ΔE(atom)

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36
Q

how does photon theory explain the two features of the photoelectric effect?

A

threshold frequency:
- the intensity (brightness) of a beam of light is related to the number of photons, but not to the energy of each.
- a photon of a certain minimum energy must be absorbed to free an electron from the surface

absence of time lag:
- an electron breaks free when it absorbs a photon of enough energy
- current is weak in dim light cos fewer photons of enough energy can free fewer electrons per unit time, but some current flows as soon as light of sufficient energy/freq strikes the metal

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37
Q

above the minimum frequency of light, the kinetic energy of the ejected electron increases with….

A

light frequency

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38
Q

increased light intensity increases …… but not…

A

the number of ejected electrons; their kinetic energy

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39
Q

describe how the emission of light from atoms was confusing for scientists

A

the passage of electricity through gas of atoms causes atoms to emit light.
- scientists expected to observe continuous light (of all wavelengths)
- instead, they only observed discrete frequencies in the visible part of the spectrum

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40
Q

give a roadmap of new ideas of light and matter, from before 1900 to after

A

before 1900:
light: wave character - eg wavelength, frequency, diffraction
matter: particulate character - eg mass, position

experiments:
- photoelectric effect
- diffraction by electrons
- atomic line spectra

wave-particle duality:
light: both wave and particulate character
matter: both particulate and wave character

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41
Q

describe the Schrodinger model of the atom

A

H^ψ = Eψ (wave equation)
- has many solutions psi

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42
Q

what does ψ psi mean?

A

‘wavefunction’ or ‘orbital’
= a mathematical function describing the shape of a wave

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43
Q

what does ψ^2 mean?

A

the probability of finding the electron at any point about the nucleus

44
Q

each orbital is described by 3 quantum numbers, and electrons are described by a quantum number:

A
  1. principal quantum number (n) - size of the orbital
  2. angular momentum quantum number (l) - shape of the orbital
  3. magnetic quantum number (ml or m) - orientation of the orbital
  4. electron spin quantum number (ms or s) - +1/2 or -1/2, up or down
45
Q

principal quantum number (n)

A

size of orbital
n = 1, 2, 3

46
Q

angular momentum quantum number (l)

A

shape of orbital
l = 0 (s), 1 (p) … n-1

47
Q

magnetic quantum number (ml or m)

A

orientation of the orbital
m = -l, -l+1, …, l-1, l

48
Q

describe s orbitals

A

spherical, l=0

49
Q

1s orbital

A
  • radial probability distribution plot is highest slightly out from the nucleus
50
Q

2s orbital

A
  • has two regions of higher electron density
  • radial probability distribution of the more distant region is higher than that of the closer one
  • between the two regions is a spherical node
51
Q

why is the radial probability distribution of the more distant 2s region higher than that of the closer one?

A

because the sum of psi^2 for it is taken over a much larger volume

52
Q

what happens at the node?

A

the probability of finding the electron drops to 0

53
Q

what is the result of the 2s orbital being larger than the 1s?

A

an electron in the 2s spends more time farther from the nucleus than it does when it occupies the 1s

54
Q

3s orbital

A
  • three regions of high electron density and two nodes
  • highest radial probability is at the greatest distance from the nucleus
55
Q

describe p orbitals

A

l=1
- has two regions (lobes) of high probability, one on either side of the nucleus
- nucleus lies at the nodal plane of this dumbbell shaped orbital

56
Q

what do the three possible ml values of p orbitals -1,0,+1 refer to?

A

three mutually perpendicular orientations associated with the x, y, and z axis

57
Q

a d orbital has any one of —— orientations

A

five

58
Q

describe d orbitals

A
  • 4 of the 5 d orbitals have 4 lobes with two mutually perpendicular nodal planes between them and the nucleus at the junction of the lobes
  • 3 of these orbitals lie in the xy, xz, and yz planes, with their lobes between the axes, and are called dxy, dxz, and dyz orbitals.
  • a fourth, the dx^2-y^2 orbital, also lies in the xy plane, but its lobes are along the axes
  • the 5th d orbital, the dz^2, has two major lobes along the z axis, and a donut shaped region girdling the centre
59
Q

describe f orbitals

A

l=3; seven orbitals with seven orientations

60
Q

radial probability distribution for s orbitals

A
61
Q

radial probability distribution for a p orbital

A
62
Q

radial probability distribution for a d orbital

A
63
Q

what is the function of the Rydberg equation?

A

predicts the position and wavelength of any line in a given series

64
Q

state the Rydberg equation

A

1/ λ = R ( 1 /n2,1 – 1 /n2,2 )
- λ is the wavelength of the line, n1 and n2 are positive integers with n2>n1, and R is the Rydberg constant

65
Q

what does the value n1 determine?

A

the region of the electromagnetic spectrum in which the lines occur

66
Q

n1= 1
Rydberg equation

A

(n2=2, 3, 4…) for the lines in the ultraviolet series

67
Q

n1 = 2
Rydberg equation

A

(n2 = 3, 4, 5…) for the lines the visible series

68
Q

n1 = 3
Rydberg equation

A

(n2= 4, 5, 6…) for the lines in the infrared series

69
Q

describe main features of the Bohr model for the hydrogen atom

A
  • the H atom has only certain energy levels called stationary states - electrons cannot be between states
  • the lower the quantum number value (n), the smaller the radius of the orbit, the closer the electron is to the nucleus, and the lower the energy level
  • E(photon) = ΔE(atom) = E(final) - E(initial)
70
Q

when does the electron of a H atom move to an outer (higher energy) orbit?

A

if an H atom absorbs a photon whose energy equals the difference between lower and higher energy levels

71
Q

when does an atom emit a photon?

A

if an H atom in a higher energy level returns to a lower energy level, the atom emits a photon whose energy equals the difference between the two levels

72
Q

Why is an atomic spectrum not continuous?

A

an atom’s energy is not continuous, but rather has only certain states

73
Q

when an electron drops from outer orbits to the n=1 orbit,

A

photons with the energy of UV radiation are emitted, creating the ultraviolet series of lines

74
Q

when an electron drops to the n=2 orbit,

A

photons of lower energy visible radiation are emitted, creating the visible series of lines

75
Q

when electrons drop to the n=3 orbit,

A

photons of even lower energy infrared radiation are emitted, resulting in the infrared series of lines

76
Q

what is the main limitation of the Bohr model?

A

fails for atoms with more than one electron - the electron-electron repulsions and additional nucleus-electron attractions that are present create much more complex interactions. also, electrons do not move in fixed, defined orbits.

77
Q

Bohr’s work leads to an equation for calculating the energy levels of an atom:

A

E = -2.18 x 10^-18 (Z^2/n^2)

78
Q

how to find the difference in energy between two levels mathematically using Bohr’s equation

A

ΔE = E(final) - E(initial) = -2.18 x 10^-18 (1/n^2(final) - 1/n^2(initial))

79
Q

how to find the wavelength of a spectral line mathematically using Bohr’s equation

A

sub ΔE into Bohr’s model to derive Rydberg’s equation (p302)

80
Q

the wavelength of any particle of mass m moving at speed u can be calculated using the equation

A

de Broglie wavelength eq:
λ = h/mu

81
Q

what does de Broglie’s wavelength imply?

A

matter behaves as though it moves in a move. An object’s wavelength is inversely proportional to its mass.

82
Q

give the application for De Broglie’s equation but for photons

A

λ = h/p

where p is the momentum and h Planck’s constant

83
Q

describe Heisenberg’s uncertainty principle

A

For a particle with constant mass m,

Δx * mΔu >/ h/4pi

where Δx is the uncertainty in position, Δu is the uncertainty in speed, and h is Planck’s constant

84
Q

define the Pauli exclusion principle

A

requires each electron in an atom to have a unique set of four quantum numbers; therefore, an orbital can hold no more than two electrons and their spins must be paired (opposite)

85
Q

give 3 ways in which electrostatic interactions determine sub level energies

A
  1. greater nuclear charge lowers sub level energy, making electrons harder to remove
  2. electron-electron repulsions raise sub level energy, making electrons easier to remove. repulsions shield electrons from the full nuclear charge, reducing it to an effective nuclear charge, Z(eff). Inner electrons shield outer electrons very effectively.
  3. penetration makes an electron harder to remove because nuclear attraction increases and shielding decreases. therefore, an energy level is split into sub levels with the energy order s<p<d<f
86
Q

define penetration and give 2 examples of its effects

A

the ability of an electron to be closer to the nucleus
- increases nuclear attraction for a 2s electron over that for a 2p electron
- decreases the shielding of a 2s electron by the 1s electrons

87
Q

define the Aufbau principle

A

electrons will fill orbitals with lowest energies first

88
Q

Hund’s rule

A

orbitals of equal energy become half-filled, with electron spins parallel, before any pairing of spins occurs

89
Q

what are considered valence electrons for main group elements vs transition elements?

A
  • MG: in the outer (highest energy level) only
  • T: (n-1)d electrons are also considered valence electrons
90
Q

describe how to figure out which sub shell will fill first in a many-electron atom

A

draw the orbitals out
arrows diagonal to the left

91
Q

Z(eff) =

A

Z(actual) - electron shielding

92
Q

draw out the s, p, d, and f blocks

A

s d p and lanthanides + actinides are f

93
Q

the elements of a group have similar

A

outer electron configurations and similar chemical behaviour

94
Q

describe two factors affecting atomic size

A
  1. changes in n; as the principal quantum number (n) increases, the probability that outer electrons spend most of their time farther from the nucleus increases as well; atomic size increases.
  2. changes in Z(eff): as effective nuclear charge increases, outer electrons are pulled closer to nucleus; atomic size decreases
95
Q

atomic radius generally —– down a group

A

increases

96
Q

atomic radius generally —– across a period

A

decreases

97
Q

define atomic size

A

half the distance between nuclei of adjacent atoms

98
Q

what is the trend in atomic size for a transition series?

A

size remains relatively constant

99
Q

define first ionisation energy

A

energy required to remove a mole of electrons from a mole of gaseous atoms or ions

100
Q

relationship between first ionisation energy and atomic size

A

inversely related

101
Q

why do successive ionisation energies of an element show a very large increase after all valence electrons have been removed?

A

the first inner (core) electron is in an orbital of much lower energy so is held very tightly

102
Q

define electron afinity

A

the energy involved in adding a mole of electrons to a mole of gaseous atoms or ions

103
Q

the 1s22s22px2 state of a carbon atom is

A

excited

104
Q

describe sub shells for a many-electron atom

A

not degenerate (due shielding)

105
Q

as n increases, the size of the orbital and number of nodes ——

A

increase

106
Q

define shielding and effective nuclear charge

A

Shielding occurs when inner electrons protect or shield outer electrons from the full nuclear attractive force. The
effective nuclear charge is the nuclear charge an electron actually experiences. As the number of inner electrons
increases, shielding increases, and the effective nuclear charge decreases

107
Q

what is penetration?

A

Penetration occurs when the probability distribution of an orbital is large near the nucleus, which results in an increase of the overall attraction of the nucleus for the electron, lowering its energy.

CHM135 (23 decks)

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