Chapter 7.1

# Chapter 7.1 - 1 THE NATURE OF LIGHT - Light is also known...

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1 THE NATURE OF LIGHT - Light is also known as electromagnetic radiation. - Light is emitted by oscillating charges. - Oscillating charges create oscillating electric and magnetic fields. - Light is very peculiar in that it is a particle and a wave at the same time. - Our physical intuition tells us that this is impossible. - Waves are spread out as in ocean waves - Particles are in one place (localized) as in a bowling ball. Q: How can something be spread out and localized at the same time? A: I have no idea. It is a mystery of nature. THE WAVE NATURE OF LIGHT - Light moves at a constant speed through a particle medium. - c – speed of light in a vacuum ( air) - c = 2.997 x 10 8 m/s (670,000,000 mi/hr) - Waves have two components. - wavelength - λ (Greek lambda) - distance between wave crests - frequency - ν (Greek nu) - how often wave crest moves up and down at a single point - number of beats per second: Hertz – Hz - 1 Hz = 1 /s = 1 s -1 - think of a boat bouncing up and down on waves 1 1 sin x ( ) 18.85 0 x 0 5 10 15 20 1 0.5 0 0.5 1 λ + -

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2 - Frequency and wavelength are related ν λ = c - if we know ν , we can calculate λ Example: If light has a frequency of 1.5 x 10 15 Hz, what is the wavelength of the light? Recall 1 Å = 10 -10 m 7 10 2. 0 1 0 m 2000Å 1 0m - - × ×= λ ν If light has a wavelength of 721 nm, what is the frequency of the light? THE PARTICLE NATURE OF LIGHT - light comes as particles called photons - energy of a photon is proportional to frequency E = h ν - h = Planck’s constant h = 6.626 x 10 34 J s a) How much energy does a photon with a frequency of 1.90 x 10 14 Hz have? ν = 6.626 x 10 J s × 1.90 x 10 s -1 = 1.26 x 10 -19 J = 1.26 x 10 J/photon b) How many photons of frequency 1.90 x 10 Hz comprise 54 kJ of energy? We can reconsider the units for the answer of part a) as J/photon.
3 ELECTROMAGNETIC SPECTRUM Name Wavelength Frequency (Hz) Radio 300 km to 0.3 m 10 3 – 10 9 Microwave 30 cm to 1 mm 10 9 – 3 x 10 11 Infrared 1.0 mm to 780 nm 3 x 10 – 4 x 10 14 Visible 780 nm to 390 nm 4 x 10 – 8 x 10 Ultraviolet 390 nm to 1 nm 8 x 10 – 3 x 10 17 X-ray 10 Å to 0.06 Å 3 x 10 – 5 x 10 19 Gamma 1.5 Å to 0.3 ym 2 x 10 18 - 10 33 LINE SPECTRA OF THE ELEMENTS The light emitted by pure elements has specific energies. - Therefore light of only specific wavelengths can be seen. (i. e., different colors can be seen) - This emitted light is called a line spectrum . (pl. spectra) Line spectra tell us that atoms can only have certain energy levels. - The atoms can not have any arbitrary value of energy. BOHR’S MODEL OF THE HYDROGEN ATOM History - Scientists before Bohr knew atom was made of nucleus and electrons. - They didn’t know where the electrons were or how they behaved. - They also knew each element had a distinct line spectrum. Model - Bohr assumed electrons traveled in orbits around nucleus. - Bohr also assumed that electrons could only have specific orbits. - Specific orbits were labeled with a quantum number - Energies of orbits are E R n n n H = - = 1 12 34 2 , , , , K R H = Rydberg constant (for hydrogen) = 2.18 x 10 -18 J

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4 SCHEMATIC OF BOHR MODEL Using the Bohr Model to Calculate Spectra Spectra result from light being emitted or absorbed when atom changes energy, i. e. when electron goes to different orbit.
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## This note was uploaded on 09/27/2010 for the course CHEM 220 taught by Professor Bates during the Spring '10 term at Skyline College.

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Chapter 7.1 - 1 THE NATURE OF LIGHT - Light is also known...

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