Discovery of the atom Flashcards
atom
smallest unit quantity of an element that may exist alone or in chemical combintion with another element
atomic number
[z] - no. of protons or electrons
mass number
[A] - no. of protons and neutrons
isotope
forms of the same element that contain equal number of protons but different number of neutrons in their nuclei
how many atomic mass units (u) are equal to the mass of 1 atom of carbon-12?
12
what is 1u/Dalton (Da) equal to
1/12 x mass of 1 atom of carbon-12
how many atoms is equal to 1 mol of carbon-12
(6.022 x 10^23) = Avogadro’s constant
relative atomic mass
[Ar] - ratio of average mass of atoms of an element in a given sample to 1u
evidence for isotopes
mass spectrometry
key dates for model of atom discovery
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
conclusions from gold-foil experiment
- most alpha-particles passed through sheet ∴ atom mostly composed of empty space
- nucleus = small and dense
discrete frequencies
missing characteristic of H when an electric discharge is passed through a sample of hydrogen
emission spectrum of hydrogen
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
electromagnetic radiation
wave with electric and magnetic properties
frequency
no. of wave crests that pass a fixed point per second
symbol = V
units = Hz
what happens to the wavelength as frequency increases?
decreases
what happens to the energy of EM radiation as frequency increases?
increases
what did Max Planck propose about EM radiation
[1900]
-EM radiation could only be emitted or absorbed in quanta (= discretes packets with energy stored in them)
-quanta = photons (units of light)
what does the energy of a photon depend on?
the frequency of the radiation
high energy = high frequency (more dangerous i.e. skin damage + UV light) = short radiation wavelength
what did Albert Einstein propose about light?
[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)
work function
[Φ] = min. energy needed to eject electron from atom
wave-particle duality
light can demonstrate both wave-like and particle-like characteristics
Balmer series
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
emission spectrum
coloured lines on dark background
absorption spectrum
black lines on coloured background
energy is absorbed from EM radiation continuum to promote an e- from ground state -> excited state
what do the line positions in absorption/emission spectra correspond to?
frequency/energy (transitions = quantised i.e. involve fixed amounts of energy)
problem with classical model?
couldn’t account for line positions on emission/absoption spectra (classical model = when everything is known + defined)
Bohr model
[Neils Bohr, 1913]
e- travel in circular orbits around nucleus
e- hend in orbitals by attractive electrostatic forces within the nucleus
Bohr model - energy of orbits
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)
Rutherford (planetary model) vs Bohr model
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
successes of Bohr model
- worked well for hydrogen
- explained the Rydberg formula for the emission lines of the H spectrum
- quantum numbers + quantinisation introduced orbits with fixed energies
failures of Bohr model
- useless for anything except hydrogen
- didn’t explain why only certain orbits were allowed
what did the mathematics of the Bohr model lead to?
radius of hydrogen = 0.529 A (= 1r)
what did Louis de Broglie propose?
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
Davisson-Germer experiment
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)
- classical model (certainty) replaced with probability/likelihood
Heisenberg Uncertainty principle
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)
why was de Broglie’s model incorrect?
based on electron particles moving in a wave-like motion
wavefunction
[𝚿] = a mathematical function that varies with position
Schrodinger wave equation
used to determine energy levels for e-
uses idea of an e- acting as a wave that alters with position
issues with Schrodinger’s wave equation
- can only be used for 1-electron systems where it gives all energy levels possible for that electron
- 𝚿 not measurable
how did Born propose?
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
what does 𝚿2 represent?
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)
radial node
when R(r) = 0
angular node
when [Y(θ,Φ) = 0]
R(r)2
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
problem with R(r)2
doesn’t take into account amount of space available for the e- (rdf = more useful)
radial distribution function (rdf)
probability of finding an electron in a spherical shell of thickness (dr) at a distance (r) around the nucleus
max. of rdf
[4πr2R(r)2) = most probable distance (r) from the nucleus of finding an e-
explain why 2s fills before 2p
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
what happens to the probable distance of finding an e- from the nucleus as the no. of radial nodes increases?
distance increases
angular wavefunction for s-orbitals
constant
boundary surface
represents 95% probability of where electron density is located for a given orbital
wavefunctions of multi-electron systems
approximation can be viewed as the product of the wavefunctions from each single electrons (based on those for hydrogenic species)
Aufbau principle
e- enter and fill lower-energy orbitals before filling higher-energy orbitals
Pauli’s exclusion principle
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)
Hund’s rule of multiplicity
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
why does nickel have the electronic configuration [Ar] 3d8 4s2 and not [Ar] 3d10 4s0?
4s orbital has lower energy than 3d - when ionised, 4s electrons lost first
rule about electron configurations in ions vs compounds
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
Cr electronic configuration + reasoning
predicted = [Ar] 4s2 3d4
actual = [Ar] 4s1 3d5
actual = more exchange energies = more stable
screening
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)
what happens to Zeff as Z increases?
increases for a given PQN
the increase in Z is not cancelled out by the addition of another e
covalent radius of a non-metallic element
1/2 internuclear separation of neibouring atoms of same element in a molecule
metallic radius
1/2 the experimentally determined distance between nuclei of nearest neighbouring atoms in the solid state
ionic radius
measure of ion size - related to distance between neighbouring cations + anions
trend in atomic radii down a group
[increases]
due to increase in PQN
increase in radial nodes
rdf max moves further from nucleus as ‘n’ increases
trend in atomic radii across a period
[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
Van der Waals radius
1/2 internuclear distance of closest approach between 2 atoms of the same type in different molecules
ionisation energy (Ie)
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-
ionisation enthalpy
standard enthalpy change per mole to remove an electron from an atom in the gas phase
electron affinity (Ae)
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)
ion size trends - cations
smaller radii as they have closed outer shells with smaller PQN
fewer radial nodes
rdf max = closer to nucleus
ion size trends - anions
larger - e- being added to an already partially filled shell
e- repulsion not offset by addition of proton to nucleus
Zeff = smaller
trends - increase of Ie across period (L -> R)
reflects on increase in Zeff
trends - B < Be
P e- easier to remove (2p1 in B) than 2s e- in Be
trends - O < N
easier to remove paired e- than unpaired e- (exchange energy not as diluted for N relative to O)
electronegativity trends - L->R
increases due to increase in Zeff
electronegativity trends - down group
increases due to increase in PQN which casues outer e- to be screened from nuclear charges
ionic bonding
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
covalent bonding
electrons shared between atoms
occurs between elements with small electronegativity differences
ΔX < 2
metallic bonding
electrostatic attraction between cations and mobile valence e-
