1. Planetary Model At first, Bohr thought the atom was much like the sun (nucleus) with the planets (e-) orbiting around it.

2. States of the Atom Ground state – lowest state (energy) of an atom (e- closest to the nucleus) Excited state - the atom has more or higher energy than it has in its ground state. When it returns to ground state it gives off the energy it gained in the form of electromagnetic radiation.

3. Photon Emission and Absorption -the amount of energy needed to excite an e- is equal to the amount of energy given off by an e- when it returns to its next lower level: E 2 – E 1 = E photon -energy levels are like the rungs of a ladder – you must be on one or the other Floor (nucleus) Ground state (lowest level of energy) Increasing energy E1E1 E2E2 E2E2 E1E1

4. Continuous spectrum – the emission of a continuous range of frequencies of electromagnetic radiation. This is what light was thought to be until the discovery of the line emission spectrum for hydrogen. Line emission spectrum – series of frequencies of light emitted by an element. Every element has its own line emission spectrum.

5. Niels Bohr Danish physicist Linked hydrogen’s electron with photon emission Electrons can circle the nucleus only in allowed paths, or orbits Orbits are separated from one another and the nucleus by a large empty space where electrons cannot exist

6. Electrons as Waves Electrons were found to have wave-like properties: -diffraction – bending of a wave as it passes by the edge of an object -interference – a reduction or increase in energy when waves overlap Heisenberg Uncertainty Principle: it is impossible to determine simultaneously both the position and velocity of an electron or any other particle

7. Quantum Theory -describes mathematically the wave properties of electrons and other very small particles