1. Niels Bohr
In the Bohr Model (1913) the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun.
They are not confined to a planar orbit like the planets are.
Niels Bohr ( ) received the Nobel Prize, for his theory of the hydrogen atom, in 1922.
Niels Bohr was born in Copenhagen. He is best known for his ground breaking work in atomic theory, which earned him the Nobel Prize in Physics in 1922.
He was forced to flee Denmark in Bohr spent the remaining war years in the United States, where he participated in the Manhattan Project.
In 1955 he organized the first Atoms for Peace Conference. He died on November 18, 1962 in Copenhagen.
Worked on the atomic bomb project in WW II, but after the war, became a strong proponent of peaceful uses of atomic energy.
   Niels Bohr was born in Copenhagen in Denmark in His father was a professor of physiology at the University of Copenhagen. Niels attended the same university and was a distinguished soccer player as well as a brilliant student.    Bohr studied at J. J. Thomson´s Cavendish Laboratory and at Rutherford´s laboratory. At the young age of 28, while working with Rutherford, he invented the first effective model and theory of the structure of the atom. His work ranks as one of the truly great examples of an imaginative mind at work. He was awarded the 1922 Nobel Prize for physics for his study of the structure of atoms.    During World War 2, Bohr and his family escaped from occupied Denmark to the United States. He and his son, Aage, acted as advisers at the Los Alomos Atomic Laboratories, where the atom bomb was developed. Thereafter, Bohr concerned himself with developing peaceful uses of nuclear energy. Aage Bohr, Neil´s son was awarded the Nobel Prize for physics in 1975.

2. Bohr’s Model
Bohr proposed that electrons traveled in orbits around the nucleus of the atom.
Each orbit has a definite amount of energy-the further away from the nucleus, the more energy it has
Electrons “jump” from one orbit to another- they never hang out in between. Like rungs on a ladder .

3. Bohr’s Model
Electrons traveling within their orbit do so without radiating any energy
Electrons absorb or radiate energy when they move between orbits (energy levels)
The wavelength of the energy absorbed or emitted depends on orbital involved
This energy can be seen on the visible spectrum.

4. Bohr Model of Atom e- e- e-
Increasing energy
of orbits
n = 3 e-
n = 2
n = 1 e- e-
In 1913, Niels Bohr proposed a theoretical model for the hydrogen atom that explained its emission spectrum.
– His model required only one assumption: The electron moves around the nucleus in circular orbits that can have only certain allowed radii.
– Bohr proposed that the electron could occupy only certain regions of space
– Bohr showed that the energy of an electron in a particular orbit is
En = – hc n2
where  is the Rydberg constant, h is the Planck’s constant, c is the speed of light, and n is a positive integer corresponding to the number assigned to the orbit.
n = 1 corresponds to the orbit closest to the nucleus and is the lowest in energy.
A hydrogen atom in this orbit is called the ground state, the most stable arrangement for a hydrogen atom.
As n increases, the radii of the orbit increases and the energy of that orbit becomes less negative.
A hydrogen atom with an electron in an orbit with n >1 is in an excited state — energy is higher than the energy of the ground state.
Decay is when an atom in an excited state undergoes a transition to the ground state — loses energy by emitting a photon whose energy corresponds to the difference in energy between the two states.
A photon is emitted
with energy E = hf
The Bohr model of the atom, like many ideas in
the history of science, was at first prompted by
and later partially disproved by experimentation.

5. Ground state – lowest energy state of atom
Excited state – atom has higher potential energy (has absorbed right amount of energy to move to a higher energy level)

6. Bohr’s contributions to the understanding of atomic structure:
Electrons can occupy only certain regions of space, called orbits.
Orbits closer to the nucleus are more stable — they are at lower energy levels.
Electrons can move from one orbit to another by absorbing or emitting energy, giving rise to characteristic
spectra.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

7. Emission Spectrum of Hydrogen
1 nm = 1 x 10-9 m = “a billionth of a meter”
410 nm
434 nm
486 nm
656 nm

8. HYDROGEN

9. Bohr’s limitations…
Only used Hydrogen - Couldn’t explain the spectra of more than one electron
Couldn’t explain the chemical behavior of atoms
Didn’t explain why or how the electron’s jumped.

10. Properties of Light
Visible light is a kind of electromagnetic radiation
has wavelike behavior as it travels
All electromagnetic radiation travels at 3.00 x 108 m/s through vacuum (empty space) – a little slower through air

11. Energy travels in waves…
Wavelength- distance between peaks (crests) of a wave (m,cm)
Frequency- number of waves that pass a given point in one second. Measured in Hertz - Hz
Wavelength - lambda distance between corresponding points on adjacent waves – measured in meters
Frequency – number of waves that pass a given point in a specific time – usually one second
Measured in waves/sec (waves per second) . It is called Hertz Hz named after man Heinrich Hertz
Wavelength and frequency are related mathematically ( speed of light m/sec equals wavelength times frequency. I
The greater the energy, the greater the frequency!

13. Visible Spectrum C = λ f 400nm-700nm
Speed of light= c= (wavelength) (frequency)
2.99 x 10 8 Sometimes rounded to 3.0 x 108
C = λ f

15. Particle description of Light
Photoelectric effect – emission of electrons from a metal when light shines on it
Mystery – no electrons were emitted by certain metals if frequency of light was below a certain minimum
Wave description of light says any frequency of light should be enough…..

16. Planck 1901 and Einstein 1905
Plank suggested hot objects emit energy in small specific packets called quanta
A quantum of energy is minimum amount of energy that can be lost or gained by an atom

17. Einstein 1905
Light has wavelike properties but also is a stream of particles
Each particle carries a quantum of energy
Photon – indivisible packet of electromagnetic radiation with zero mass carrying a quantum of energy

18. How to find the frequency of a wave if you know the wavelength…
C = λ f
Speed of light = wavelength X frequency
c= speed of light = 2.99 x 10 8 m/s
If λ = x m …then
f = (2.99 x 10 8 m/s) / (656.5 x m/s)
f = x cycles/s or Hz (Hertz)

19. You can calculate the energy of a photon of light
What is a photon of light?
A tiny, indivisible packet of light
Planck discovered an equation to calculate the energy in a photon……what a great guy, even got a Nobel Prize in Physics in 1918!
The units for energy is Joules, J

20. Planck’s equation…
How to find the energy of a photon at a particular frequency
E (photon) = h f
Where h = Planck’s constant 6.63 x
f = frequency

21. How to find the energy of a photon at a particular frequency
E (energy) = h f
h = planck’s constant
E = (4.567 x ) (6.63 x )
E = 3.03 x Joules