Lesson 3 : The Bohr Model

Bohr Model of an Atom  Electrons orbit the nucleus in fixed energy ranges called orbits (energy levels)  An electron can move from one energy level to another by gaining or losing discrete amounts of energy  Electrons cannot be found between energy levels (think of energy levels like rungs on a ladder... )  The lowest energy level is closest to the nucleus, the highest is farthest away  The electron energy levels are quantized. Think of it like potential energy, the higher you are the more you have

Absorption Vs Emission  When an electron (e-) absorbs (gains) energy (in whole photons or “quanta”) it “jumps” to a higher energy level  This is called the EXCITED STATE  When an e- emits (loses) energy it falls to a lower energy level and the energy emission is given of as photons (light)  This is called the GROUND STATE  The return to ground state is what we see as color in the flame test

So how was the “color” made in the flame test?  Scientists use the Bohr model to explain this phenomenon Hydrogen Spectrum Flame Test

There is NO net change in energy  Energy absorbed = energy released = energy of light produced  Sometimes (like the flame test) this light is in the small section of wavelengths called the visible spectrum and we can see it  Most of the time the human eye cannot

Bohr’s Hydrogen Model  Turn to page 8 in your ref. packet  When an electron falls from n=6 to n=3 what wavelength of light will be emitted?  1094 nanometers  What region of the spectrum does that wavelength correspond to?  Infrared  Would we see it?  Not with our naked eye

Hydrogen’s Line Spectrum  Hydrogen emits four visible wavelengths of light  Visible light is emitted when an excited electron “falls” from n= 3, 4, 5, or 6 back to n=2 410 nm; violet 434 nm; blue 486 nm; cyan 656 nm; red Faint lines are UV

Practice  What color of light will be emitted if an e- goes from:  n=6 to n=2?  Violet (410nm)  Red/orange (656 nm)  Blue (434nm)  n=5 to n=2?  n= 3 to n=2?

9 Evidence for Energy Levels  Bohr realized that this was the evidence he needed to prove his theory.  The electric charge temporarily excites an electron to a higher orbit. When the electron drops back down, a photon is given off.  The red line is the least energetic and corresponds to an electron dropping from energy level 3 to energy level 2.

Chapter 510 Radiant Energy Spectrum  The complete radiant energy spectrum is an uninterrupted band, or continuous spectrum.  The radiant energy spectrum includes most types of radiation, most of which are invisible to the human eye.  The visible spectrum is the range of wavelengths between 400 and 700 nm.

Chapter 511 The Wave/Particle Nature of Light  In 1900, Max Planck proposed that radiant energy is not continuous, but is emitted in small bundles. This is the quantum concept.  Radiant energy has both a wave nature and a particle nature.  An individual unit of light energy is a photon.

Electromagnetic Spectrum (EM) 12 EM is the complete range of electromagnetic radiation Wavelength increases Frequency increases Energy increases

13 Wave Nature of Light  Light travels through space as a wave, similar to an ocean wave.  Wavelength is the distance light travels in one cycle.  Frequency is the number of wave cycles completed each second.  As frequency increases, energy increases  Small Wavelength = Large Frequency = Big Energy

 Wavelength ( λ ) – the shortest distance between equivalent points on a continuous wave. Wavelength is measured is units of length - m, mm, µm, nm  Amplitude – the height from the origin to the crest (or trough)  Frequency ( ν ) – the number of waves that pass a given point in one second 14 Anatomy of a Wave

Inverse Relationship Between Wavelength(λ) and Frequency (ν)  When λ increases, ν decreases  When λ decreases, ν increases  When ν decreases, λ increases  When ν increases, λ decreases  The longer the wavelength of light, the lower the frequency. The shorter the wavelength of light, the higher the frequency. 15