2.  The Bohr model was proposed:  1913  by Neils Bohr  After observing the H line emission spectrum

3.  Why did Bohr study hydrogen?  Simplest atom – only has one electron  The electrons of the gas are easily excited by a current

4.  1. How are the electrons moving?  As particles with a definite circular path 2. Where are they located? - In rings, or orbitals, around the nucleus 3. How much energy do they have? - The energy of an electron can be calculated

5.  E = -2.178 x 10 -18 (Z 2 /n 2 )  Z= atomic number  N= ring number or energy level  Calculations of energy change for transitions:   E = E final – E initial   E = -2.178x10 -18 J (Z 2 /n f 2 – Z 2 /n i 2 )

6.  Electrons closer to the nucleus have LOWER energy values. At higher energy levels, the energy is HIGHER

8.  Each orbit is an energy level designated by the variable n.  n = 1 is the ground state.  n = ∞ is when the electron has been removed from the atom.  Energy is quantized, meaning there are no levels in between the levels designated by n = 1, 2, 3, 4, etc.

9.  When atoms are excited by an outside energy source heat, flame, or electric current, the electrons can be promoted to higher energy states.  However, this situation is highly unstable and the electron will eventually return to a lower energy state. When the electron returns to a lower energy level, energy is given off in the form of quanta (light), also known as a photon of light.  How do we represent this electron transition in our model?

10.  Using a spectroscope, Bohr noticed that hydrogen gave off a spectrum rather than a continuous spectrum (rainbow) when excited.  He calculated the energy associated with each transition based on the wavelength of light given off. Each energy value corresponded to a certain electron transition from an excited state back to the n=2 energy level.

11. Wavelength of light observed (nm) Frequency of light (/ sec) Energy of light (Joule) N initial (excited state) N final (lower energy level) 410. nm7.32 x10 16 -4.84 x10 -19 62 434 nm6.91 x10 16 -4.57 X10 -19 52 486 nm6.17 x10 16 -4.08 X10 -19 42 656 nm4.57 x10 16 -3.03 X10 -19 32

12.  Hydrogen produces lines in the visible, ultraviolet, and infrared regions of the electromagnetic spectrum  Each set of spectral lines has a name:  Balmer series  Lyman series  Paschen series