The Classical Picture of the Atom Flashcards
Law of conservation of mass
A fundamental law of chemistry showing that the total mass of a sealed vessel and its contents are the same before and after a reaction.
Dalton
All matter is composed of solid and invisible atoms, there are as many different types of atom as there are elements, with the atoms of different elements having different masses. Chemical reactions change the way differentiators are grouped together, atoms cannot be created or destroyed.
J.J. Thomson
1897 discovery of the electron. A high voltage electric current was passed through a gas at low pressure, negatively charged particles were observed to travel between the electrodes. The so called cathode rays were the same no matter what gas was used, we now know them as electrons. Electric and magnetic fields were applied to the beam of electrons and deviations from the straight line were used to calculate the charge to mass ratio. Plum pudding model suggested.
Millikan
1910 measurement of electronic charge (-1.60x10^-19) by observing the rate of fall of charged oil droplets, since the mass to charge ratio was known, electronic mass could also be determined.
Rutherford
1909 discovery of the nucleus
Chadwick
1932 discovery of the neutron as the particles emitted when bombarding beryllium and boron with alpha particles (+ve charge)
Gold foil experiment
Alpha particles directed at a very thin gold foil, only a very atoms thick. Most alpha particles pass through the electron clouds of the atoms and are undeflected however a very small proportion (1/20,000) approach close to a nucleus and are deflected significantly (>90) due to repulsive interactions with a positively charged nucleus. This disproved the plum pudding model (negatively charged electrons in a positively charged sphere). The small number of deviations suggested a small dense centre of positive charge, the nucleus surrounded by a large volume of mainly empty space containing the electrons.
EM radiation
A form of energy consisting of oscillating electric and magnetic fields that travel through space at the speed of light, c (2.998x10^8ms^-1).
Speed of light (c) =
frequency (nu) x wavelength (lambda)
Electromagnetic spectrum
Radiowave, microwave, infrared, visible, ultraviolet, x-ray, gamma-ray.
Plank 1900
EM radiation can only be emitted and absorbed in packets or quanta of radiation, later called photons. The energy of a photon is proportional to its frequency.
Energy (E) =
Planks constant (h, 6.626x10^-34Js) x frequency (nu)
Photoelectric effect
When UV radiation strikes metal surface, electrons are ejected only when the frequency is above a certain threshold which is specific to the metal, regardless the intensity of the radiation. Einstein reasoned that photons must transfer a minimum value of energy to atoms on the metals surface. Once above the threshold the excess energy of the photon is converted to the kinetic energy of the ejected electron. hv=Φ+Eke. Work function (Φ) is the minimum energy. This illustrates have EM radiation acts as a particle.
Wave particle duality of light
The observation of diffraction shows EM radiation behaves as a wave. Young demonstrated that when light passes through 2 closely spaced slits, each slit gives rise to a circular wave which have constructive/destructive interference as the amplitudes of the waves add/subtract together generating areas of light and dark. Light has wave particle duality.
Rydberg equation
frequency (nu) = Rh[(1/n(1)^2)-(1/n(2)^2); Rh = 3.29x10^15 Hz
n(2)>n(1) (emission spectrum where electrons fall from higher energy levels to lower)
Atom
The smallest unit quantity of an element that may exist either alone or in chemical combination with another element
Atomic symbol notation
Atomic mass - top left
Atomic number - bottom right
Bohrs model
Atomic spectroscopy demonstrates that the energy of an electron in an atom is quantised. Electrons travel around in quantised circular orbits around the nucleus, electrons are held in these orbits by attractive electrostatic forces with the nucleus. Orbit frequencies are given by:
Frequency = -Rh(1/n^2)
The Rydberg equation gives the difference in frequencies between orbits.
Failure of the Rutherford model
Predicts unstable atoms! Electrons loose energy as they travel meaning they should eventually collapse into the nucleus.
Successes of the Bohr model
Worked well for hydrogen
Explained the Rydberg formula for the emission lines of the hydrogen spectra
Quantum numbers and quantisation introduced
Failures of the Bohr model
Useless for anything except hydrogen
Didn’t explain why only certain orbits were allowed
Radial unit
0.529A, the mathematics behind the Bohr model led to the calculation of the radius of hydrogen
de Broglie equation
Proposed all matter had wave like properties
λ = h/mv
mv = p (momentum)
de Broglies vision of the atom
Electrons move in a wave like motion around the nucleus, each electron orbit was considered to be a fixed number (integer) times the wavelength.
Davisson & Germer
Shon an electron beam at a nickel crystal, the resultant diffraction pattern had areas of light and dark (constructive/destructive interference) which is only relevant for waves.
Wave-particle duality
replacing the classical model where everything is known with probability.
Heisenberg uncertainty principle
When objects exhibit wavelike properties, you cannot measure their position and momentum simultaneously.
ΔpΔq > h/4π (lower limit to uncertainty)
Δp = uncertainty in momentum Δq = uncertainty in position p = mv
Failures of de Broglie
Based on electron particles moving in a wave like motion, a new wave mechanical treatment was needed
Schrodinger wave equation (SWE)
Used to determine orbital energies for electrons, using the idea of an electron acting as a wave that alters with position. Can only be solved for 1 electron systems, where it gives all energy levels possible for that electron. Treats electrons as spread out/delocalised.
Wavefunction (Ψ)
A mathematical function that varies with position, E(total) varies depending on where you are (x,y,z). Each wave function describes a different orbital type wth a characteristic energy. It is not measurable.
Born interpretation
Ψ^2, the probability of an electron being in a certain volume of space (probability per unit volume). Electron density per unit volume. dt(tau) how electron density varies as one moves through space in small volume amounts.
n
PQN - determines the energy (shell); 0=s, 1=p, 2=d
l
SQN - types of orbital for a given n and their shapes; (n-1), (n-2) … 0
m(l)
MQN - orbital orientations for each l value; ±l each holds 2 electrons
m(s)
spin
Polar coordinates
Atoms are approx. spherical (r,θ,Φ)
r = distance from nucleus θΦ = angular deviations
Can be factorised into R(r) x Y(θ,Φ)
Radial (measured in radial units/0.529A) and angular wave functions
1s
Never = 0
2s
≠0 at r=0, =0 at r=0
2p
=0 at r=0, ≠0 elsewhere
3s
≠0 at r=0, =0 at 2 places
3p
=0 at r=0, =0 in one other place
3d
=0 at r=0, ≠0 elsewhere
Simplifications of R(r)
Z=1, a(0) = 1 radial unit
Nodes
where Ψ=0, can be a radial node or nodal plane.
n-(l+1)
R(r)^2
The electron density at a specific point in space as a function of distance (r) from the nucleus only. Does not take into account the amount of space available for the electron meaning the highest probability of finding an electron is said to be at the nucleus in an s orbital.
Radial distribution function (rdf)
The probability of finding an electron in a spherical shell of thickness(dr) at distance (r) from around the nucleus. The ref max tells us the most probable distance (r) from the nucleus of finding an electron. 4π^2R(r)^2
Angular wavefunctions
S orbitals have a fixed value, and do not vary with θ,Φ hence s orbitals are spherical. P(z) only depends on θ hence no variation in x,y plane. P(x) and P(y) depend on both θ,Φ.
Boundary surface
Represents 95% probability of where electron density is located for a given orbital, R(r) never touches the x axis hence is only equal to 0 at ∞. Always indicate the sign of the wave function even though they are based on (r,θ,Φ)^2.
Angular nodes
Either planes or cones and do not depend on r only θ,Φ
D orbitals
dxy, dyz, dxz (3 between the axises) dx^2-y^2 (on the axis), dz^2
Orbital approximation
For multi electron systems the orbitals are based on the solutions to the SWE for a 1e- system. Hence each energy level (orbital) will be described by 3 quantum numbers (r,θ,Φ). The wave functions of systems with many electrons can be viewed as the product of the wav functions from each of the single electrons. The energies of the orbitals are no longer degenerate, in H orbital energies do not depend on l.
Multiple electron SWE solution
Rdf max for He2+ closer to nucleus due to greater PQN
Aufbau principle
Electrons eneter and fill lower energy orbitals before filling higher energy orbitals
Paulis exclusion principle
No two electrons in the same atom can be in the same quantum state. This means that no two electrons can have the same set of 4 quantum numbers; i.e. they must have different spins if n,l&m are the same.
Hund’s Rule of maximum multiplicity
When there are degenerate orbitals available electrons will enter the orbitals singly and only when all the orbitals are half filled will pairing up occur to avoid repulsion and maximise exchange energy.
Relative exchange energies, K
Calculated considering the number of pairs of parallel spins that exist for electrons of equal energy in a given electronic configuration. More exchanges = Higher value of K = More stable the arrangement. Largest k values arise for half filled and fully filled orbitals.
Electron population of transition metals
transition metals only populate their outermost s-orbitals when they are in their elemental forms. Electrons in the outer s orbitals will always move to the d orbitals as this gives a lower energy arrangement.
Screening
In all atoms there are attractive forces between the electrons and nucleus. In atoms with more than 1 electron, there is electron-electron repulsion to consider, these screen an electron from the full nuclear charge.
Z(eff) - Effective nuclear charge
Z(eff) = Z - S. As Z increases, Zeff increases for a given PQN because an additional e- does not fully screen an additional proton in the nucleus.
Covalent radius of a non-metallic element
1/2 the internuclear separation of neighbouring atoms of the same element in a molecule
Metallic radius
1/2 the experimentally determined distance between nuclei of nearest neighbour atoms in the solid state
Ionic radius
Measure of an ion size
Atomic radii own a group …
- Get larger
- Increase in PQN means an increase in radial nodes
- Further from the nucleus you will find the rdf max
Atomic radii across a period
Zeff increases pulling e- cloud closer to the nucleus giving rise to the d/f block contraction. F orbitals are very diffuse and do not screen nuclear charge very well, thus e-‘s with higher PQNs than that of an f orbital will experience an increased Zeff and be held tightly. 4d&5d elements in the same group have a similar radii
Van der Waals radius
1/D internuclear distance of closest approach between two atoms of the same type in different molecules
Ionization energy I(e)
Energy change on removing an electron from an atom to infinite distance in the gas phase
Electron affinity A(e)
Reverse of ionisation potential, also quoted for the gas phase, should be negative as energy is released.
Cation size trends
Radii are smaller than those for free atoms because fewer nodes, ref max closer to nucleus.
Anion size trends
E- being added p orbitals thats already 1/2 filled causing repulsion not offset by increase in PQN hence are bigger.
Electronegativity
the power of an atom to attract electrons to itself when it is part of a compound
