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CH10133 > atomic structure & properties > Flashcards

atomic structure & properties Flashcards

1
Q

definition of ‘atom’

A

smallest unit quantity of an element; may exist alone or in chemical combination with another element.

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2
Q

charge of a proton, neutron and electron

A

+1.602 x 10-19C, 0, -1.602 x 10-19C

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3
Q

mass of a proton, neutron and electron

A

1.673 x 10-27 kg, 1.675 x 10-27 kg, 9.109 x 10-31 kg

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4
Q

definition and symbol of atomic number

A

the number of protons or electrons in an atom - Z.

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5
Q

definition and symbol of mass number

A

the number of protons and neutrons in an atom - A.

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6
Q

definition of ‘isotope’

A

forms of the same element that contains equal numbers of protons but has a different number of neutrons in their nuclei.

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7
Q

what is the more convenient unit for the mass of atoms?

A

atomic mass units (u) .

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8
Q

definition of relative atomic mass:

A

values are calculated by taking a weighted mean of the mass numbers of each isotope, are generally not whole integers.

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9
Q

what was John Daltons part in the history of the atom?

A
  • 1800s
  • argued that all matter is composed of solid and indivisible atoms
  • noted that elements react in the same mass ratios to form a given compound
  • argued that, in a given compound, the atoms are always present in the same ratio - explains why reagents always react together in the same mass ratio
  • Devised the concept of indivisible atoms to explain this, so that each element has a different atom type.
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10
Q

what was JJ Thomson’s part in the history of the atom?

A
  • 1897-1904
  • discovered the electron – first suggestion of subatomic particles.
  • concluded that when a high voltage electric current was passed through a gas at low pressure, negatively charged particles were observed to travel between the electrodes – known as electrons.
  • plum-pudding model arose – ball of +ve charge with electrons in it.
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11
Q

what was Robert Millik’s part in the history of the atom?

A
  • 1910
  • determined the charge on the electron as –1.60 × 10-19 C.
  • observed the rate of fall of charged oil droplets and since Thomson had determined the charge to mass ratio for the electron, the known charge enabled the electronic mass to be determined.
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12
Q

what was Ernest Rutherford’s part in the history of the atom?

A
  • 1909 – 1911
  • disproved the plum-pudding model by trialling the gold-foil experiment.
  • lead to the discovery of the nucleus – showed atoms are composed mostly of empty space.
  • experiment = shot α-particles at a thin metal film, only a few atoms thick.
  • most positively charged particles passed through with minor deflections.
  • a small fraction - about 1/20000 - were deflected by angles greater than 90 ° due to repulsive interactions with a positively charged nucleus with some bouncing back the way they had come.
  • proposed the ‘planetary model’ of the atom
  • atoms have small +ve charged nucleus with electrons floating around it.
  • this proposal implied that there must be a +ve charged particle to balance the –ve electrons.
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13
Q

what was James Chadwick’s part in the history of the atom?

A
  • 1932
  • discovered the neutron – it’d only been predicted previously.
  • presence of these electrically neutral particles was predicted on the basis of charge to mass ratios in ions.
  • particles emitted on bombarding Be and B atoms with α-particles, which are positively charged particles that we now know consist of 2 protons and 2 neutrons.
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14
Q

how can you record the hydrogen emission spectrum? what happens?

A

can be recorded by passing an electric discharge through hydrogen gas – this splits the molecules into atoms and the electrons within the atoms become excited – when the electrons return to the ground state they emit energy.

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15
Q

what is the speed of light?

A

2.998 × 10^8 m s-1

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16
Q

What is frequency? What is its symbol and units?

A

number of wave-crests that passes a fixed-point per second.

units = s^-1 = Hz.

Symbol = ν = Greek lower case ‘n’ = pronounced: nu.

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17
Q

What is velocity? What is its symbol and units?

A

how quickly they pass through space

units = ms^-1

Symbol = c - velocity of light.

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18
Q

What is wavelength? What is its symbol and units?

A

the distance between any given point and the same point in the next wave cycle.

units = m

Symbol = λ = pronounced: lambda.

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19
Q

equation involving velocity, wavelength and frequency:

A

velocity of light = wavelength x frequency

c = λ x ν

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20
Q

What did Max Planck propose?

A
  • in 1900, Max Planck proposed that EM radiation could only be emitted or absorbed in packets or quanta of radiation, now called photons.
  • energy of a photon is proportional to its frequency.
  • High energy means high frequency.
  • High frequency mean short radiation wavelength.
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21
Q

Equations involving Plancks constant, energy, frequency, and wavelength.

A

energy = Plancks constant x frequency

E = hv

energy = (Plancks constant x velocity of light) / wavelength

E = (hc) / λ

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22
Q

what is Planck’s constant?

A

6.626 × 10^-34 J s.

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23
Q

What is the photoelectric effect?

A
  • simply: when UV radiation strikes a metal surface and electrons are ejected.
  • explanation: Each individual photon needs to have enough energy to remove an electron from the metal surface
  • electrons ejected when the frequency of the UV radiation is above a certain threshold - specific to the metal.
  • once this threshold is passed electrons are ejected - regardless of the intensity of the radiation, though high intensity = more electrons emitted.
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24
Q

What part did Albert Einstein play in the photoelectric effect?

A
  • in 1905, Albert Einstein explained the photoelectric effect using a quantum approach.
  • Einstein reasoned that electrons can only be ejected from a surface if incoming photons transfer a minimum value of energy to atoms on the metal surface.
  • if a photon doesn’t have enough energy - an electron will not be ejected regardless of the intensity of the radiation, as none of the individual photons has enough energy to eject the electron.
  • once above the threshold, the excess energy of the photon is converted into the kinetic energy of the ejected electron
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25
Q

what is the workfunction? symbol and definition.

A

Φ

Name given to the threshold (minimum) energy of an electron (in a single atom).

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26
Q

What does the theory of wave-particle duality of light suggest?

A

observation of diffraction suggests that EM radiation can behave as a wave.

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27
Q

What role did Thomas Young play in the theory of wave-particle duality of light?

A
  • in the early 19th century, Thomas Young demonstrated that when light passes through 2 closely spaced slits, each slit gives rise to a circular wave and these waves interfere with each other to give a diffraction pattern consisting of a series of bright and dark lines.
  • some experiments provide evidence of light consisting of waves, and others provide evidence of light consisting of particles - Light has wave–particle duality - so we treat light as waves when it is useful to do so, and as particles when it is useful to do so.
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28
Q

What is constructive interference?

A

occurs when the peak of the first wave coincides with the peak of the second wave, so we can add the amplitudes of the 2 waves.

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29
Q

What is destructive interference?

A

when the peak of the first wave coincides with the trough of the second, so the 2 waves cancel out

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30
Q

What is a continuous spectrum?

A

when visible light passes through a prism, it is split into the component colours - resulting pattern is a continuous spectrum - as it contains an unbroken distribution of all frequencies.

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31
Q

Who was the first person to identify a pattern in the atomic spectrum of hydrogen?

A

Johann Balmer

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32
Q

What pattern did Johann Balmer find? (expression)

A

v = (1/4) - (1/n^2)

where n = 3,4,5,…

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33
Q

What is the Rydberg equation?

A

v = Rh ((1 / n1^2)-(1/n2^2))

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34
Q

What is the Rydberg constant?

A

3.29 × 1015 Hz.

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35
Q

For the Lyman series what region of the EM spectrum is it? What value is n1 and n2?

A
  • UV
  • n1=1
  • n2 = 2,3,4,…
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36
Q

For the Balmer series what region of the EM spectrum is it? What value is n1 and n2?

A
  • visible
  • n1 = 2
  • n2 = 3,4,5,…
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37
Q

For the Paschen series what region of the EM spectrum is it? What value is n1 and n2?

A
  • IR
  • n1 = 3
  • n2 = 4,5,6…
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38
Q

For the Brackett series what region of the EM spectrum is it? What value is n1 and n2?

A
  • IR
  • n1 = 4
  • n2 = 5,6,7…
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39
Q

emission spectra colours?

A

bright lines against a dark background.

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40
Q

absorption spectra colours?

A

black lines against a coloured background.

41
Q

What did Niels Bohr contribute to the atom?

A
  • in 1913, Neils Bohr proposed a new model of the atom.
  • he proposed that electrons travel in a fixed orbit around the nucleus + the electrons in these orbits are held by attractive electrostatic forces – energy associated with each orbit has a fixed value.
  • by absorbing energy = an electron can move from one orbit to another, further from the nucleus.
  • by emitting energy = an electron can move from one orbit to another, closer to the nucleus.
42
Q

What is Bohr’s expression for the energy of an electron in a hydrogen atom?

A

E = (2π^2)(me)(e^4) / (h^2)(n^2)

E = energy of orbit

me = mass of electron

e4 = charge of electron

n2 = quantum number

π , me , e and h = constants

43
Q

What is the simplified version of Bohr’s expression for the energy of an electron in a hydrogen atom?

A

E = -((k) / (n^2))

k = (2π^2)(me)(e^4) / (h^2)

44
Q

Successes of the Bohr model:

A

works well for hydrogen.

explained the Rydberg formula for emission lines of the hydrogen spectrum.

introduced quantum numbers.

Bohr’s model led to the calculation of the radius of hydrogen of 0.529Å.

45
Q

Failures of the Bohr model:

A

useless for anything other than hydrogens.

didn’t explain why only certain orbits are allowed

46
Q

What did Louis De Broglie contribute?

A

In 1924, Louis de Broglie proposed that all matter had wave like properties.

this led to an amended version of Bohr’s model of the atom – where electrons are moving in a wave-like motion around the nucleus.

each orbit was considered to be a fixed number times the wavelength.

wavelength of everyday objects are too small = unnoticeable – but cannot be ignored for subatomic particles.

47
Q

What is de Broglie’s equation?

A

wavelength = Planck’s constant / (mass x velocity)

λ = h / (m x v)

48
Q

What is the Davisson – Germer experiment?

A

directed a beam of electrons at a nickel crystal and observed a diffraction pattern.

pattern had areas of light and dark due to constructive/destructive interference of the electron beam.

constructive/destructive interference = only relevant for waves.

able to observe the pattern because the wavelength of the electron beam was close to the spacing between nickel atoms.

electrons show wave-particle duality

49
Q

What is the Heisenberg uncertainty principle equation?

A

Δp Δq >/ h / 4π

Δp = uncertainty in momentum
Δq = uncertainty in position
>/ = greater than or equal to
h = Planck constant = 6.626 × 10^-34 J s.
50
Q

What is the Heisenberg uncertainty principle? What does it describe?

A

since objects have wave-like properties, one cannot measure their position and momentum simultaneously – known as the uncertainty principle – founded by Werner Heisenberg in the 1920’s.

To determine the position of an electron will mean that the momentum is affected.

To determine the momentum of an electron will mean that the position is affected.

51
Q

What is the wavefunction? Who introduced it? What is it used for?

A

Schrödinger introduced the wavefunction – denoted by psi, ψ.

function with a value that varies with position.

can calculate the wavefunction for an electron using the SWE.

SWE can be used to determine energy levels for electrons.

52
Q

What does a partial differentiation show?

A

a partial differential shows how a function depends on one variable when several are changing

53
Q

What is the Schrodinger wave equation used for?

A

only possible to solve the SWE exactly for two-body problems - in the case of atoms means systems that contain a nucleus and a single electron - only neutral atom for which the equation can be solved is hydrogen.

solutions also possible for single-electron ions such as He+ and Li2+.

54
Q

What did Max Born suggest?

A

Max Born suggested how wavefunctions could be related to a measurable property = ψ^2 = electron density.

the square of the wavefunction ,ψ^2 = proportional to the probability of finding the electron within a small volume of space dτ.

region where ψ^2 is large, = probability of finding the electron is high.

region where ψ^2 is small = probability of finding the electron is low.

each ψ is treated as a wave by SWE that is delocalised – so one cannot say exactly where an electron is located at a given time.

can take a small volume of space τ, Tau.

electron density measured as τ varies, dτ.

55
Q

What does Dτ measure?

A

Dτ = measures how electron density varies as one moves through space in small volume amounts, I.e. at different points.

56
Q

What is Bohr’s radius?

A

Bohr’s radius = 0.529Å = 1 radial unit.

57
Q

What is the principle quantum number (PQN)?

A

the integer n is called the principle quantum number (PQN) - value of n can be any integer from 1 to ∞ - though n = 1 to n = 7 are the most important chemically. (“shell”)

tells you about the energy of the orbital.

58
Q

What is the secondary quantum number?

A

the second quantum number, l, is known as the secondary quantum number - value of l can be any integer from 0 to (n – 1) - though values of 0, 1, 2, and 3 are important chemically. (“orbital number”).

Gives the types of orbitals possible for a given ‘n’ and their shapes.

also referred to as the angular, azimuthal, or orbital quantum number.

59
Q

What is the magnetic quantum number?

A

the third quantum number, ml, is called the magnetic quantum number – values can be any whole number from –l - +l. (“orientations for orbitals”)

60
Q

What is an atomic orbital?

A

a region of spaced defined by a wavefunction.

61
Q

What does this symbol represent: θ?

A

θ = theta (angular coordinate) - deviation from the z axis.

62
Q

What does this symbol represent: ϕ?

A

ϕ = phi (angular coordinate) - deviation from the x axis in the xy plane.

63
Q

Expression that involves the wavefunction as a function of r, θ and ϕ :

A

ψ(x,y,z) = ψ(r, θ, ϕ) = R(r) x Y(θ, ϕ)

64
Q

What does R(r) represent?

A

radial wavefunction: contains information about what happens to the wavefunction as the distance from the nucleus increases

65
Q

What does Y(θ, ϕ) represent?

A

angular wavefunction: contains information about the shape of the orbital.

66
Q

What is a node?

A

Node: a point at which a wavefunction is zero - this occurs in the radial wavefunction, R(r), it is called a radial node.

67
Q

Description of the 2s orbital radial wavefunction and any potential nodes?

A

radial wavefunction for the 2s orbital: R(r) starts with a +ve value, goes through 0 becoming -ve, and then tends towards zero at high values of r.

node = spherical shape - at this value of r, the wavefunction is 0 (regardless of the direction from the nucleus).

68
Q

Description of the 2p orbital radial wavefunction and any potential nodes?

A

radial wavefunction = 0 at the nucleus, increases to a maximum value, then decreases towards 0 at high values of r.

so there is a node at the nucleus -> present in both the radial and angular wavefunctions.

69
Q

How many radial nodes does the 3s orbital have?

A

2

70
Q

How many radial nodes does the 3p orbital have?

A

1 radial node for r > 0

71
Q

How many radial nodes do these have: 1s, 2p, 3d, and 4f orbitals?

A

they are the first of their types, so have no radial nodes for r > 0

72
Q

What does this function represent: R(r)2dτ?

A

it gives the probability of finding an electron within the volume dτ at a distance r from the nucleus - equivalent to the electron density at that point.

73
Q

What is thy notation used for the radial distribution function?

A

4πr^2R(r)^2

74
Q

What is the radial distribution function?

A

gives the probability density for an electron to be found anywhere on the surface of a sphere located a distance r from the proton.

related to the probability of finding an electron in a spherical shell of radius r and thickness dr.

75
Q

What are the main effects on the functions when multiplying R(r)^2 by 4πr^2?

A
  • firstly, the maxima are pushed further from the nuclei since 4πr^2 increases with r.
  • secondly, the RDF for all orbitals is 0 at r = 0, since multiplying R(r)^2 by zero equals zero
76
Q

What type of wavefunction is related to the shape of atomic orbitals?

A

angular wavefunction, Y(θ, ϕ)

77
Q

What does a boundary surface represent?

A

A boundary surface represents a 95% probability of where electron density is located for a given orbital.

95% because R(r) never touches the r axis, which means that R(r) only reaches ‘0’ at infinity – so cannot plot 100%.

78
Q

What is an angular node?

A

nodal plane which arises from the angular wavefunction.

  • Either planes or cones.
  • Does not depend on r.
79
Q

What is the orbital approximation?

A

states that the wavefunction for an N-electron atom, ψ, is equal to the product of N single-electron wavefunctions.

80
Q

What does the fourth quantum number relate to?

A

fourth quantum number, ms - relates to the spin of the electron in the orbital.

81
Q

What is the Aufbau Principle?

A

involves building up the electronic structure of an atom by filling the lowest energy orbitals first – used to determine the electronic configuration.

82
Q

What is the Pauli exclusion principle?

A

in 1925, Wolfgang Pauli proposed that no 2 electrons in an atom could have the same 4 quantum numbers – known as the Pauli exclusion principle – any particular orbital is defined by 3 quantum numbers ( n, l, ml ) - consequence of the Pauli exclusion principle is that a max of 2 electrons can be in any orbital, these 2 electrons must have different spins and spins are paired.

83
Q

What is Hund’s rule?

A

when filling a set degenerate orbitals, electrons are added with parallel spins to different orbitals rather than pairing 2 electrons in one orbital.

84
Q

2 reasons why Hund’s rule applies are?

A
  • parallel electrons make less electron-electron repulsion – compared with paired electrons.
  • increased stability from parallel electrons = exchange energy.
85
Q

What is exchange energy?

A

Exchange energy is the energy released when two or more electrons with the same spin-exchange their positions in the degenerate orbitals of a subshell

86
Q

What effect does the exchange energy have on the stability?

A

the more exchanges possible (higher values of K), the more stable the arrangement.

87
Q

Why is Scandium an anomaly in electronic configuration?

A

3d orbitals are lower in energy than the 4s orbital.

would expect it to be [Ar]3d3 but is actually [Ar]3d1 4s2

favourable to fill the 4s orbital = less electron-electron repulsions – repulsions are greater when all the electrons are in the d orbitals (more compact than the 4s orbital) - this reduction in repulsion overcomes the cost of occupying the higher energy orbital and leads to lower overall energy.

88
Q

Why is Chromium an anomaly in electronic configuration?

A

would expect it to be [Ar]3d4 4s2 but is actually [Ar]3d54s1

difference in electron configuration explained by exchange energy.

exchange energy is higher with 5 parallel electrons in the 3d54s1 than the 4 parallel electrons in the 3d44s2.

differences in exchange energy has a bigger effect than the electron-electron repulsions.

89
Q

What is the effective nuclear charge?

A

If an electron is far from the nucleus, then at any given moment, many of the other electrons will be between that electron and the nucleus.
Hence the electrons will cancel a portion of the positive charge of the nucleus and thereby decrease the attractive interaction between it and the electron farther away.
As a result, the electron farther away experiences an effective nuclear charge ( Zeff ) that is less than the actual nuclear charge Z

90
Q

What is the order of penetration with the orbitals?

A

s-orbitals > p-orbitals > d-orbitals.

91
Q

Definition of atomic radius?

A

atomic radius: half the distance between the nuclei of neighbouring atoms in the pure element.

92
Q

Definition of Van der Waals radii?

A

Van der Waals radii: radius for an atom is half the distance between non-bonding nuclei of neighbouring atoms in the structure if the solid – considerably larger than covalent radii.

93
Q

Main factor in explaining the general increase in atomic radii down a group?

A

down a group, outer electrons lie in higher energy orbitals – where radial distribution function lie further from the nucleus.

94
Q

Main factor in explaining the general decrease in atomic radii across a period?

A

Zeff increases pulling the electron cloud closer to the nucleus.

95
Q

What is f-block contraction (Lanthanoid contraction) ?

A

F-orbitals are poor screeners of electrons from the nuclear charge – because f-orbitals are diffuse.

So in the case of elements in the range of Z=57-71, there is a greater than expected decrease in observed radii.

In the d-block, the effect is that while 4d elements have larger atomic radii than 3d elements, the poor screening effect of the f-orbitals means that the 4d and 5d elements in the same group have similar radii

96
Q

Definition of Ionic radius?

A

determined from the internuclear distance in ionic solids – e.g. sum of anionic and cationic radii.

97
Q

Why is the cationic radius is smaller than its atomic radius?

A

They have closed outer shells with a smaller PQN than the element.

So an element with a PQN = n, the ion will have electrons in shells up to PQN = n-1.

This means that orbitals have fewer radial nodes than those in shell n.

98
Q

Definition of ionisation energy?

A

ionisation energy: the energy change when an electron is removed from an atom of the element in the gas phase. E.g. First ionisation energy : X (g) -> X- (g) + e-

99
Q

Definition of Electron gain energy?

A

Electron gain energies: energy change that occurs when an electron is attached to an atom in the gas phase. E.g. X (g) + e- -> X- (g)

CH10133 (5 decks)

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