Light and Electrons

John Dalton

atoms —> solid spheres

J J thompson

plum pudding model (electrons in a positive mass)

Rutherford

nucleus discovered (majority empty space)

Bohr

electrons in fixed energy levels

Quantum model

electron cloud

Bohr’s Model

shows that the electrons ar ein fixed energy levels and explains the bright line spectra

only works well for hydrogen and doesn’t explain behavior fully

Electromagnetic spectrum

radio, microwave, infrared, visible UV, X ray Gamma

low —> high energy

Relationship between energy frequency and wavelength

if energy increases, frequency increases, but wavelength decreases

speed of light equations

c=λν

• c = 3.00×1083.00×108 m/s

• λ = wavelength

• ν = frequency

Planck’s equation

E=hν

• h = 6.626×10−346.626×10−34 J·s

• ν = frequency

Photon

packet of energy

Bright line spectrum

specific wavelengths are emitted

happens when electrons fall to lower energy levels

flame tests = electrons get excited and release colored light

Aufbau principle

fill the lowest energy first

hund’s rule

fill orbitals singly first

Pauli’s exclusion

max 2 electrons per orbital

core electrons

inner electrons

valence electrons

on the outer shell (show reactivity)

ground state

normal configuration

excited state

electrons move up

electron behavior

atoms absorb energy, electrons jump up

atoms releases energy, electrons fall and emit light

Heisenburg’s uncertainty principle

you can’t know the exact poision and speed of an electron

Electron cloud model

electrons exist in probability regions and not fixed paths

more accurate than bohr model