Atomic Theories

Planck’s Quantum Theory

  • Radiant energy emitted by atoms occurred in small bundles called quantum. (Quanta plural)
  • Atoms could only absorb or emit radiant energy in one or multiples of one quantum (discrete quantities only)
  • Quantum Theory doesn’t apply to macroscopic objects. Energy at atomic level occur in quantum

Photoelectric effect

  • Light is an electromagnetic wave consisting of a continuous series of wavelengths
  • Light consisted of a stream of energy packets/quanta - later called photons
  • A photon of red light contains less energy than a photon of UV light
  • An electron cannot break free from an atom unless a certain minimum quantity of energy is absorbed

Bohr’s model

Postulates:

1: Atoms and electrons can only exist in specific energy states. When moving within the allowed energy state (stationary state) electron does NOT emit/lose energy

2: Each energy level corresponds to an orbit

3: Electron can travel in an orbit without radiating energy

4: Electron may only change its energy by jumping from one allowed energy state to another

Success & Failures of the Bohr Model

  • %%Worked well for hydrogen in predicting the spectra%%
  • %%Explained Mendeleev’s periodic law%%
  • ==Only worked well for Hydrogen and got worse as atomic number increased==

Spectroscopy

  • Analysis of the way matter absorbs/releases radiation. Can identify elements, obtain info on bonding, determine concentration
  • Atomic spectra: Absorption spectra & Emission spectra
    • Absorption Spectra (dark line spectrum): Excitation (photons absorbed)
    • Emission Spectra (bright line spectrum): Relaxation (photons released)
    • Continuous Spectrum: contains all frequencies of visible light (white light)
  • Used in astronomy, forensics and can distinguish substances with similar physical and chemical properties
  • We only see the visible range (Balmer series), even though many other transitions are happening
    • UV - Lyman series. IR - Paschen series
  • Excitation: electron absorbs energy (heat, light, electricity) and moves to a higher energy level
  • Relaxation: electron releases energy (electromagnetic radiation) and moves to a lower energy level
  • Ground state: the lowest energy for that electron (closest to the nucleus)
  • Quantized energy: atoms can only absorb or release the exact amount of energy required for electrons to move from one allowed energy state to another

De Broglie’s Matter Waves

  • Matter has both wave and particle characteristics
  • Electrons are paired with a standing wave
    • zero vibration/amplitude (destructive interference)
  • Wavelength of a matter wave: λ = h / mv
    • heavy objects have short wavelengths, tiny particles with low masses have longer wave lengths

Wave particle duality of electrons

  • Waves are continuous travelling disturbances, Particles are discrete bundles. Distinctions break down on subatomic level
  • Diffraction patterns
    • Bands of light and dark produced on a screen by waves.
    • Bands of light occur when the waves reinforced each other (in phase) and underwent constructive interference.
    • Bands of dark occur when the waves cancel each other (out of phase) and underwent deconstructive interference

Heisenberg’s Uncertainty Principle

  • It is impossible to determine simultaneously the exact position and momentum of a particle
    • we cannot know the exact location of the electron, only where it is most likely to be
  • Δ(mv) Δ(x) ≥ h

Schrodinger’s wave equation

  • Gives 3d probability of where an electron is most likely to be found (orbitals)

Superposition

  • An electron may exist in many possible states at the same time, interacting with each other
  • Superposition disappears when we look at a particle (open box)

Probability plots + Radical distribution

  • Antinode: 90 - 95% high probability of finding an electron - constructive interference

  • Node: 0% probability of finding an electron - deconstructive interference

Quantum Numbers

Principal Quantum Number

  • main energy level of an electron. n

Secondary Quantum Number

  • shape or type of orbital. l. for each number of n, there are the same number of subshells

Magnetic Quantum Number

  • how many orientations in 3d space. ml

Spin Quantum Number

  • an electron has an up spin and down spin

Challenges of Quantum Mechanical Model

  • Basic concepts that are applicable in macroscopic world become inapplicable at subatomic level
  • Superconductivity (conduct electricity with zero resistance) is explained using quantum mechanics. The model sets limits for temperatures at which superconductivity takes place

Energy Level Diagram

  • Energy of each subshell increases s p d f

Pauli’s Exclusion Principle

  • no two electrons can have the same set of quantum numbers

Aufbau Principle

  • electrons must be added by starting with lowest available energy level + subshell

Hund’s Rule

  • electrons must be spread out within a subshell with the same spin before doubling up

Magnetism

  • when spins of an electron aren’t paired
  • Domains: groups of atoms magnetically influencing each other
  • Permanent Magnet = domains aligned

Ferromagnetism

  • Fe, Co, Ni - small enough to infuence eachother
  • Strongly attracted to a magnetic field
  • Forms domains, permanent magnet if domains align
  • Unpaired electron spins

Paramagnetism

  • Weakly attracted to a magnetic field
  • No domains form, no permanent magnet possible
  • Unpaired electron spins

Diamagnetism

  • Weak repulsion to a magnetic field
  • Only paired electron spins (Noble gases + Alkaline Earth Metals)

Anomalous Configuration

  • half-filled and filled subshells are more stable than unfilled subshells
  • Electron promotion