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Inorganic Chemistry > Discovery of the atom > Flashcards

Discovery of the atom Flashcards

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1
Q

atom

A

smallest unit quantity of an element that may exist alone or in chemical combintion with another element

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2
Q

atomic number

A

[z] - no. of protons or electrons

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3
Q

mass number

A

[A] - no. of protons and neutrons

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4
Q

isotope

A

forms of the same element that contain equal number of protons but different number of neutrons in their nuclei

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5
Q

how many atomic mass units (u) are equal to the mass of 1 atom of carbon-12?

A

12

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6
Q

what is 1u/Dalton (Da) equal to

A

1/12 x mass of 1 atom of carbon-12

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7
Q

how many atoms is equal to 1 mol of carbon-12

A

(6.022 x 10^23) = Avogadro’s constant

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8
Q

relative atomic mass

A

[Ar] - ratio of average mass of atoms of an element in a given sample to 1u

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9
Q

evidence for isotopes

A

mass spectrometry

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10
Q

key dates for model of atom discovery

A

1800s - John Dalton - each element has different atom type

1897 - J.J. Thomson - discovered the electron

1904 - Plum pudding model

1910 - Robert Millikan - charge on electron

1909-1911 - Ernest Rutherford - Gold-foil experiment
-discovery of nucleus - disproved plum pudding model
-proposed planetary model of atom (small positively charged nucleus balanced by charge of electron)

1932 - James Chadwick - discovery of neutron

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11
Q

conclusions from gold-foil experiment

A
  1. most alpha-particles passed through sheet ∴ atom mostly composed of empty space
  2. nucleus = small and dense
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12
Q

discrete frequencies

A

missing characteristic of H when an electric discharge is passed through a sample of hydrogen

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13
Q

emission spectrum of hydrogen

A

recorded by passing an electric discharge through hydrogen gas

splits molecules into atoms

electrons within atoms get excited

when these electrons return to ground state, they emit energy

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14
Q

electromagnetic radiation

A

wave with electric and magnetic properties

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15
Q

frequency

A

no. of wave crests that pass a fixed point per second

symbol = V
units = Hz

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16
Q

what happens to the wavelength as frequency increases?

A

decreases

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17
Q

what happens to the energy of EM radiation as frequency increases?

A

increases

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18
Q

what did Max Planck propose about EM radiation

A

[1900]

-EM radiation could only be emitted or absorbed in quanta (= discretes packets with energy stored in them)

-quanta = photons (units of light)

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19
Q

what does the energy of a photon depend on?

A

the frequency of the radiation

high energy = high frequency (more dangerous i.e. skin damage + UV light) = short radiation wavelength

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20
Q

what did Albert Einstein propose about light?

A

[1905] - light energy was made up of photons

explained the photoelectric effect = the way certain metals release electrons when UV falls on them

each individual photon needs to have enough energy to remove e- from metal (radiation must reach threshold frequency, v)

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21
Q

work function

A

[Φ] = min. energy needed to eject electron from atom

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22
Q

wave-particle duality

A

light can demonstrate both wave-like and particle-like characteristics

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23
Q

Balmer series

A

hydrogen emission frequencies observed in the visible region of EM spectrum (n = 2)

Lyman series - n = 1
Pashcen series - n = 3
Brackett series - n = 4
Pfund series - n = 5

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24
Q

emission spectrum

A

coloured lines on dark background

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25
Q

absorption spectrum

A

black lines on coloured background

energy is absorbed from EM radiation continuum to promote an e- from ground state -> excited state

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26
Q

what do the line positions in absorption/emission spectra correspond to?

A

frequency/energy (transitions = quantised i.e. involve fixed amounts of energy)

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27
Q

problem with classical model?

A

couldn’t account for line positions on emission/absoption spectra (classical model = when everything is known + defined)

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28
Q

Bohr model

A

[Neils Bohr, 1913]

e- travel in circular orbits around nucleus

e- hend in orbitals by attractive electrostatic forces within the nucleus

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29
Q

Bohr model - energy of orbits

A

each orbit = fixed / quantised

zero energy = taken as the point where e- is removed (∴ energy levels all lower than 0)

lowest energy state - n = 1 (closest to nucleus)

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30
Q

Rutherford (planetary model) vs Bohr model

A

Rutherford = predicted unstable atoms whereby e- lose energy as they travel (they should collapse in nucleus)

Bohr = orbits have specific energies + e- can only gain or lose energy by changing orbit

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31
Q

successes of Bohr model

A
  1. worked well for hydrogen
  2. explained the Rydberg formula for the emission lines of the H spectrum
  3. quantum numbers + quantinisation introduced orbits with fixed energies
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32
Q

failures of Bohr model

A
  1. useless for anything except hydrogen
  2. didn’t explain why only certain orbits were allowed
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33
Q

what did the mathematics of the Bohr model lead to?

A

radius of hydrogen = 0.529 A (= 1r)

34
Q

what did Louis de Broglie propose?

A

1924

all matter had wave-like properties (formula relates matter to its wave)

led to amended version of Bohr model with e- moving in a wave-like motion about the nucleus

each e- orbital was considered to be a fixed number x wavelength

35
Q

Davisson-Germer experiment

A

1925

shone e- beam at nickel crystal + recorded pattern

pattern had areas of light + dark due to constructive/deconstructive interference of e- beam = only relevant to waves = diffraction pattern

able to observe this due to wavelength of e- beam was close to spacing between nickel atoms

[Conclusions]
1. e- have both particle and wave-like properties (= wave-particle duality)

  1. classical model (certainty) replaced with probability/likelihood
36
Q

Heisenberg Uncertainty principle

A

objects that exhibit wave-like properties, one cannot measure their position AND momentum at same time

Heisenberg determined lower limit to uncertainty = Planck’s constant (h/4π)

37
Q

why was de Broglie’s model incorrect?

A

based on electron particles moving in a wave-like motion

38
Q

wavefunction

A

[𝚿] = a mathematical function that varies with position

39
Q

Schrodinger wave equation

A

used to determine energy levels for e-

uses idea of an e- acting as a wave that alters with position

40
Q

issues with Schrodinger’s wave equation

A
  1. can only be used for 1-electron systems where it gives all energy levels possible for that electron
  2. 𝚿 not measurable
41
Q

how did Born propose?

A

an interpretation for SWE whereby 𝚿 could be related to a measurable quantity, 𝚿^2

equates to electron density

for each 𝚿, SWE treats e- as wave that is delocalised ∴ one cannot say exactly where an e- is located at any given time

possible to calculate probability of an e- being in certain volume of space

42
Q

what does 𝚿2 represent?

A

electron density per unit volume

take small vol. of space 𝜏

d𝜏 = how e- density varies as one moves through space in small volume amounts (= get orbital shapes)

43
Q

radial node

A

when R(r) = 0

44
Q

angular node

A

when [Y(θ,Φ) = 0]

45
Q

R(r)2

A

electron density at specific point in space as a function of distance (r) from the nucleus only i.e. probability of finding electron at a specific point

46
Q

problem with R(r)2

A

doesn’t take into account amount of space available for the e- (rdf = more useful)

47
Q

radial distribution function (rdf)

A

probability of finding an electron in a spherical shell of thickness (dr) at a distance (r) around the nucleus

48
Q

max. of rdf

A

[4πr2R(r)2) = most probable distance (r) from the nucleus of finding an e-

49
Q

explain why 2s fills before 2p

A

rdf(max) 2s > 2p (node in 2s)

2s penetrates 2p orbital (2s e- can get closer to nucleus than 2p e-)

closer to nucleus an e- can get, the lower the energy will be

50
Q

what happens to the probable distance of finding an e- from the nucleus as the no. of radial nodes increases?

A

distance increases

51
Q

angular wavefunction for s-orbitals

A

constant

52
Q

boundary surface

A

represents 95% probability of where electron density is located for a given orbital

53
Q

wavefunctions of multi-electron systems

A

approximation can be viewed as the product of the wavefunctions from each single electrons (based on those for hydrogenic species)

54
Q

Aufbau principle

A

e- enter and fill lower-energy orbitals before filling higher-energy orbitals

55
Q

Pauli’s exclusion principle

A

no 2 e- in same atom can be in same quantum state (i.e. no 2 electrons can have same set of 4 quantum numbers)

56
Q

Hund’s rule of multiplicity

A

when degenerate, electrons will enter orbitals singly + only pair up when orbitals are half-full

avoids e- repulsion and maximises exchange energy (K) - higher K = more stable

57
Q

why does nickel have the electronic configuration [Ar] 3d8 4s2 and not [Ar] 3d10 4s0?

A

4s orbital has lower energy than 3d - when ionised, 4s electrons lost first

58
Q

rule about electron configurations in ions vs compounds

A

transition metals only populate outermost s-orbitals in their elemental forms

e.g. Ni2+ electronic configuration = [Ar] 3d8 and not [Ar] 3d6 4s2
in compounds, s electrons

59
Q

Cr electronic configuration + reasoning

A

predicted = [Ar] 4s2 3d4
actual = [Ar] 4s1 3d5

actual = more exchange energies = more stable

60
Q

screening

A

in all atoms, there are attractive forces between e- and nucleus

in atoms with >1 e-, there is repulsion to consider

e- repulsions screen an e- from full nuclear charge (Z)

3d > 3p > 3s in terms of screening

= effective nuclear charge (Zeff) whereby Z is reduced by S (screening)

61
Q

what happens to Zeff as Z increases?

A

increases for a given PQN

the increase in Z is not cancelled out by the addition of another e

62
Q

covalent radius of a non-metallic element

A

1/2 internuclear separation of neibouring atoms of same element in a molecule

63
Q

metallic radius

A

1/2 the experimentally determined distance between nuclei of nearest neighbouring atoms in the solid state

64
Q

ionic radius

A

measure of ion size - related to distance between neighbouring cations + anions

65
Q

trend in atomic radii down a group

A

[increases]

due to increase in PQN

increase in radial nodes

rdf max moves further from nucleus as ‘n’ increases

66
Q

trend in atomic radii across a period

A

[decreases]

Zeff increases - pulls e- cloud closer to nucleus
rise to f-block contraction then d-block contraction

f-orbitals = poor screeners (diffuse) -> elements with Z from 57 to 71, there is a greater than expected decrease in radii

d-block => net effect; although 4d > 3d, the poor screening effect of f-orbitals means that 4d/5d elements in the same group have similar radii i.e. Ti < Ze ~ Hf

67
Q

Van der Waals radius

A

1/2 internuclear distance of closest approach between 2 atoms of the same type in different molecules

68
Q

ionisation energy (Ie)

A

energy change on removing an electron from an atom to infinite distance in the gas phase

units = ev

eqns:
1st ionisation energy = M -> M+ + e-
2nd ionisation energy = M+ -> M2+ + e-

69
Q

ionisation enthalpy

A

standard enthalpy change per mole to remove an electron from an atom in the gas phase

70
Q

electron affinity (Ae)

A

reverse of ionisation potential (also in gas phase)

units = eV
eqn: 1st electron affinity = M + e- -> M-

NOTE - always check convention (i.e. if exothermic is +/-) by looking at Ae of F (which the gain of 1e- = exothermic)
no thermodynamic conflict with electron gain enthalpy (exo = always -ve)

71
Q

ion size trends - cations

A

smaller radii as they have closed outer shells with smaller PQN

fewer radial nodes

rdf max = closer to nucleus

72
Q

ion size trends - anions

A

larger - e- being added to an already partially filled shell

e- repulsion not offset by addition of proton to nucleus

Zeff = smaller

73
Q

trends - increase of Ie across period (L -> R)

A

reflects on increase in Zeff

74
Q

trends - B < Be

A

P e- easier to remove (2p1 in B) than 2s e- in Be

75
Q

trends - O < N

A

easier to remove paired e- than unpaired e- (exchange energy not as diluted for N relative to O)

76
Q

electronegativity trends - L->R

A

increases due to increase in Zeff

77
Q

electronegativity trends - down group

A

increases due to increase in PQN which casues outer e- to be screened from nuclear charges

78
Q

ionic bonding

A

electrostatic attraction between a cation and anion, formed from the transfer of electrons from one element to another

occurs between elements with large electronegative differences

ΔX > 2

79
Q

covalent bonding

A

electrons shared between atoms

occurs between elements with small electronegativity differences

ΔX < 2

80
Q

metallic bonding

A

electrostatic attraction between cations and mobile valence e-

Inorganic Chemistry (15 decks)

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