Bohr Model of the Atom & Light Monday, October 31st, 2016

Dalton’s Atomic Model

Plum Pudding Model (Thomson)

The Development of a New Atomic Model Rutherford’s model was incomplete: How were electrons distributed? What prevents negative electrons from crashing into positive nucleus? 1900s-New Atomic Model (revolutionary) Relationship between light & electrons

Niels Bohr (Born in Denmark 1885-1962) Student of Rutherford

Niels Bohr’s Model (1913) Electrons orbit the nucleus in circular paths of fixed energy (energy levels).

How many Electrons Can Fit in Each Energy Level? Energy level is also called “principal quantum number” Energy Level = n n max. # electrons 1 2 2 8 3 18 4 32 5 50 6 72 Max # e = 2n2 Big Question: Why?

Bohr Model of an Atom Fixed energy related to the orbit Electrons cannot exist between orbits The higher the energy level, the further it is away from the nucleus An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive)

Sketch the Bohr Model for the following elements – which is unreactive? Nitrogen Argon

Sketch the Bohr Model for the following elements – which is unreactive? Nitrogen Argon

} Further away from the nucleus means more energy. Increasing energy Fifth Further away from the nucleus means more energy. There is no “in between” energy Energy Levels Fourth Increasing energy Third Second First

The closer an electron is to a proton, the more stable the atom! The most stable location for an electron is as close to the nucleus as it can get… that is its ground state configuration (the lowest possible energy level the electron can be at) But electrons can jump from energy level to energy level! An atom is said to be in an excited state when its electrons are at an energy level higher than the ground state.

Electrons absorb or emit energy when they jump from one energy level to another. A quantum of energy is the amount of energy required to move an electron from one energy level to another.

How do we determine (calculate) energy? A scientist (Max Planck) determined hot objects emit energy in packets called “quanta” “quantum” - minimum energy that can be gained or lost by an atom E = h Where E = energy, h = Planck’s constant (6.626 x 10-34 Js), and  = frequency (in s-1)

Photons are bundles of light energy that is emitted by electrons as they go from higher energy levels to lower levels.

High Freq Electromagnetic Spectrum Low Freq High E Low E Light emitted produces a unique emission spectrum.

Wavelength & Frequency Speed of light (c) = 3.00 x 108 m/s Wavelength ( ) = distance between waves Frequency () = # waves that pass point in given time c =   as  increases,  decreases as  increases,  decreases (slinky analogy)

Emission Spectra White light - continuous spectrum Hydrogen atoms - line-emission spectrum Big Question - Why did hydrogen atoms only give off specific frequency (colors) of light?

The “Fingerprints” of Atoms

The atom is quantized, i.e. only certain energies are allowed. Atomic Spectra The atom is quantized, i.e. only certain energies are allowed.

Neils Bohr Solved the Mystery Electron circled nucleus in “orbit” of fixed energy Absorption-electrons can “hop” from ground state to excited state Emission-when electrons “fall” from excited state to ground state Energy difference corresponds to hydrogen’s spectral lines