Bohr Model of

the Atom

Bohr

• To explain spectral lines, Bohr came up with a radical model of the atom which had electrons orbiting around a nucleus:

• Electrons can only be in "special" orbits.

• All other orbits just were not possible

• Electrons could "jump" between these special orbits

• Each orbit represented a certain amount

of energy and that the electron would

move in the orbit without radiating energy

(although this violated classical physics)

• The allowed orbits were called stationary

states

• The lowest possible energy state was

called the ground state

Spectrum Formation

• a photon would be absorbed when

the electron jumped from one state to

a higher state

• a photon would be emitted if the

electron dropped to a lower state

• Each line in a spectrum corresponds to an

electron moving from 1 stationary state to

another

Electron Energy

• The total energy of the electron is

negative and proportional to the distance

from the nucleus.

• This means that it takes energy to pull the

orbiting electron away from the nucleus.

Example

• Determine the

wavelength of

light emitted

when an electron

drops from

energy level E to

level B in the

atom shown in

the diagram

Solution

• E = Ef – Ei

• E = -21. 76 x 10-19 J – (-3.49 x 10-19 J)

• E = -18.27 x 10-19 J

• (atom lost, photon gained it)

• =1.09 x 10-7 m

• Different lines in

the spectrum

would be caused

by different

transitions

• High frequency

photons would

be emitted by

large transitions

Summary

• The Bohr model accurately explained atomic

spectra (emission and absorption)

• The energy of each orbit is quantized

• Explained why atoms are stable

• Explained the chemical and physical properties

of the elements

Problems

• Could not explain why electrons could only

be found in certain orbits

• Could not explain why some lines in the

emission spectra were brighter than others

• Worked only for hydrogen and atoms with

a single electron (for example He+)