km

Thomas Young

• examined one light source sent through two openings

• observed constructive and destructive interference patterns
  

LIGHT IS A WAVE

Wavelength

• (lambda) is the distance between identical points on successive waves. (m or nm) (1m = 1 x 10^9nm)

Frequency

• (n) is the number of waves that pass through a particular point in 1 second (Hz = 1 cycle/s).

Electromagnetic radiation

lambda x v = c

• lambda = wavelength

• v = frequency

• c = speed of light = 3.00 × 10^8 m/s

Electromagnetic Spectrum

• Red Martians Invade Venus Using X-ray Guns

  • Visible Light: 400 nm - 700 nm

Planck

• Heated Solids Problem

  • Solids emit electromagnetic radiation over a wide range of wavelengths when solids are heated. Radiant energy emitted by an object at a certain temperature corresponds to their wavelength and frequency

  • Energy(light) is emitted or absorbed in discrete units(quantum)

  • E = h x v

    • E: energy of a photon of light (J)

    • h = Planck’s constant = 6.63 × 10^-34 Jxs

    • v = frequency of light (1/s)

Einstein

• Photoelectric Effect(1905)

  • Light has both wave and particle nature

  • Photon is a “particle” of light

  • hv = KE + W

    • hv = energy of the photons of light hitting metal

    • W = work function

    • KE = energy of ejected electron

Relationship between frequency and wavelength

Inversely proportional

Relationship between frequency and energy

Directly proportional

Hz in a GHZ

10^9

nm in a m

1 × 10^9

Bohrs Model of the Atom

• Have quantized energy values(n=1, n=2, n=3…)

• Able to jump to higher orbitals when they absorb enough energy

• When electrons move towards nucleus, light is emitted and the energy is dependent on the energy difference

Energy of jump

En = -Rh (1/n²)

E = - Rh (1/nf² - 1/2i²)

• n = 1,2,3, …

• Rh (Rydberg constant) = 2.18 × 10^-18 J

Ion

• An atom/group of atoms that has a net positive or negative charge due to the loss or gain of electrons

Shell*

• Area around the nucleus where electrons reside (n)

Orbital

• Specific area in the shell when the electrons reside

Valence electrons

• Electrons in the outermost shell

Octet Rule

• To be “happy” atoms want 8 valence electrons – LOW

  ENERGY (**Exception – the first shell only wants 2 electrons)

Cation

• Ion with a positive charge

• Neutral atom loses one or more electrons

Anion

• Ion with a negative charge

• Neutral atom gains one or more electrons

• ide

Work function

• How strong/tight electron is bound to the metal

Emission Spectra

• Charging gas in a tube which separates light into different components with a prism

J in a Kj

1000J

Shortcoming of Bohr

• Didn’t account for emission spectra of atoms that have more than one electron

• Conflict with wavelike properties

Heisenberg Uncertainty Principle

• Impossible to know the momentum p (m x v) and the position of a particle with certainty

Schrodinger

• Equation describes particle and wave nature of the e-

• Hydrogen atom

• Wave function(psy):

  • Energy of e- with a given psy

  • psy ² = probability of finding electron in a volume of space

Electron density

• Probability that an electron will be found in a particular region of an atom

Atomic orbital

• Probability of locating an electron in space

Principal quantum number n

• Distance of electron from the nucleus and average energy

• n = 1, 2, 3, 4, …

Angular momentum quantum number /

• Shape of the atomic orbital

• Orbitals

  • l = 0

  • l = 1

  • l = 2

  • l = 3

Magnetic Quantum Number ml

• Orientation of the orbital in space

Spin quantum number ms

• ms = -1/2 or +1/2

Pauli exclusion principle

• No two electrons in an atom can have the same four quantum numbers

Hunds Rule

• The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins

Aufbaus Principle

• When in the ground state, electrons must fill orbitals from the lowest energy to the highest energy

Shell

• Electrons with the same value of n

Subshell

• Electrons with the same values of n and l

Orbital

• Electrons with the same values of n, l, and ml

Total orbitals

n²

Total electrons

2n²

Paramagnetic

• When there are unpaired electrons that are attracted to external magnetic fields

Diamagnetic

• When electrons are paired

Isoelectronic

• Have the same number of electrons and the same ground-state electron configuration

Constructive

• Crest to crest - > Constructive Interference

Destructive

• Crest to trough → Destructive Interference