Discovery of electron charge: Millikan

• Oil drop experiment
  

  Atomized oil falling between two charged plates, charge required to suspend the atoms used, in conjunction with mass to determine charge

Discovery of the electron: Thomson

• Cathode ray tube
  

  Tube with very little air, light rays going from cathode to anode, makes few gas molecules fluoresce, the rays were deflected by magnetic/electric fields, therefore they carried a charge

Plum pudding: Thomson

• Positive batter
• Negative "berries" throughout

Nuclear model: Rutherford

• Discovery of the proton: Gold foil experiment
  

  Positively charged alpha radiation shot at thin gold foil, some particles deflected at large angles, this would require a dense positive charge (positive nucleus)

The neutron: Chadwick

• Mass of electrons + protons ≠ mass of most elements
  

  Presence of neutral particles

Light's wave nature:

• Wavelength:

  
  

  Distance between peaks

  
  

  Spd. of Light = Wve. Lngth. X Freq.

  
  

• Frequency:

  
  

  Nb of complete cycles passing a given point

  
  

• Amplitude:

  
  

  Distance from peak to midline

  
  

• Velocity:

  
  

  Speed of wave

Problem 1: Blackbody radiation

• All objects above 0 K emit radiation
  

  This would mean humans (at 310 K) would emit X-Rays

Answer 1: Planck

• Atoms emit/absorb radiation in discrete quantities (quanta)
• Quanta: Smallest amount of absorbable/emittable energy (electromagnetic radiation)
• E of a single quantum is proportional to the freq. of the radiation
  

  E = hv

• Since E is always emitted in multiples of hv, it acts like a particle

Problem 2: Photoelectric effect

• Metals emit electrons/current when hit by light
• The speed of the emission should vary by wavelength and intensity
• Would produce lag when turning smt on
  

  Experiments proved otherwise (min. freq. needed regardless of current, no lag time)

Answer 2: Einstein

• Light should be thought of as a stream of photon particles rather than a wave
• Using Planck's quantum theory: electrons must absorb one photon of enough energy to jump (min. freq. needed)
• Electron makes jump as soon as it absorbs photon

Problem 3: Atomic Spectroscopy

• When atoms absorb energy, they emit light
• Spectroscopy: single gas emits light when given electrical energy to rid the gained energy
• Different gas = different absorption
  

  More electrons means more light

Answer 3: Rydberg

• Finds a general formula for the Hydrogen line spectrum
• Certain wavelengths required for electrons to jump orbits
  
  

Answer 3.5: Bohr's model

• 1) Single electron
• 2) In orbit, it does not emit energy
• 3) If it changes orbit, it's energy must change by 1 photon
  
  

Each jump represents a freq.

• Initial state: fundamental level (resting
• Excited: moving up/down
• Ionized: free
• Level determines light emission
  

Uncertainty principle: Heisenberg

• Problem with wave-like particles: impossible to know both momentum and position simultaneously
• The electron does not orbit nucleus predictably
  
  

Schrodinger: Orbital model

• We can predict where the electron should be
  
  

Quantum mechanics:

• Probability that the electron will be in a certain place

n

• Values: 1 to ∞
• Describes a shell
• As n increases, orbital gets larger, electrons a further from nucleus and less stable

l

• Values: 0 to n-1
• Describes shape of orbital
• Defines a subshell
  

  • 0: S
  • 1: P
  • 2: D
  • 3: F

mL

• Values: -l to l
• Describes orbital orientation

ms

• Values: 1/2 or -1/2
• Defines electron spin

1s

• Sphere
  
  

2s & 3s

• Sphere
• Larger in size
• Electron density is 0 at nodes
  

  Nb of nodes: n-1

  
  
  

2p

• Tear drop lobes
• Differ in orientation (x,y,z)

Schrodinger:

• His equation can be modified to account for electron repulsion
• It's only approximate

Coulomb's law:

• The potential energy between two same charges decreases as they get further away (repulsion)

  
  

• The potential energy between two different charges is negative and becomes more negative as they get closer

Shielding:

• Positive charge in nucleus (attracts electrons)
• Core electrons shield this attraction
• Easier to remove electrons

Penetration:

• Some electrons penetrate closer to nucleus
• This lowers it's energy, adding stability

Alkali metals:

• Low melting point
• Gas and water produced when contact w/ water
• Oxidizes in air

Transition metals:

• Not easy to predict

Periodic table:

• Group (column):

  
  

  Down: + orbitals, + electrons, bigger radii

  
  

• Period (row):

  
  

  Right to left: Noble to alkali = bigger radii

  
  

  Alkali = less valence = less stable = less attraction = bigger

  
  

  Attraction on e = nb. protons - nb. core e

Cations and anions:

• Cations (+): lose e, energy shared less, more attraction, so smaller
• Anions (-): gain e, energy shared more, less attraction so bigger

Ionization energy:

• Down a group: + orbitals, - IE, easier to ionize
• Exceptions:
  

  Elements prioritize symmetry

  
  

  B resists less than Be

  
  

  O resists less than N

Electron affinity:

• Increases: Left to right, bottom to top
• Radius gets smaller, attraction from nucleus gets bigger