electrons and the em spectrum. let’s light stuff on fire!
TRANSCRIPT
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Electrons and the
EM Spectrum
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Let’s light stuff on fire!
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Models of the Atom So far, the model of the
atom consists of protons and neutrons making up a nucleus surrounded by electrons. After performing the gold foil experiment, Rutherford hypothesized a model of the atom that looked much like the one below.
electron
neutron
proton
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What about the electrons??• Rutherford didn’t know
exactly where the electrons were located in the atom, just that they surrounded it, or why chemical bonding occurred.
• Bohr figured out that electrons orbit in energy levels around the atom.
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The Bohr Model
• Niels Bohr (1852-1962) was a student of Rutherford and believed the model needed improvement.
• Bohr proposed that an
electron is found only in
specific circular paths, or
orbits, around the nucleus.
(write this in before “However, this model….”)
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Models of the Atom
However, this model could not explain the chemical and physical properties of the elements.
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Models of the Atom
For example it could not explain:
1. why metals give off certain colors when heated in a flame
-or- 2 .why objects heated to high temperatures first glow dull red, then yellow, then white
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So what is light?
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Light is all about the Electrons!
• Light is a form of electromagnetic radiation that travels like a wave through space as PHOTONS.
• When electrons get excited, they jump up to higher energy levels and then fall back down
• Depending on how high they jump, they will give off a different color of LIGHT
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Atomic Emission Spectra
ground state
excited state
ENERGY IN PHOTON OUT
GA
IN e
nerg
y
LOS
E energy
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Energy of Electrons
When atoms are heated, bright lines
appear called line spectra
An electron absorbs energy to “jump”
to a higher energy level. (excited)
When an electron falls to a lower
energy level, energy is emitted. (ground
level state)
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Emission Spectrum
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Atomic Emission Spectra
• Every element has a UNIQUE emission spectrum. The colors that you see represent the element’s electrons jumping through the energy levels!
• LONG JUMPS are represented by HIGH energy colors (violet & blue)
• SHORT JUMPS are represented by LOW energy colors (red).
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C. Johannesson
EM Spectrum
LOW
ENERGY
HIGH
ENERGY
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EM Spectrum
LOW
ENERGY
HIGH
ENERGY
R O Y G. B I V
red orange yellow green blue indigo violet
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Properties of Light
• Movement of excited electrons to lower energy levels, and the subsequent release of energy, is seen as light! Before 1900, scientists thought light behaved solely as a wave. This belief changed when it was later discovered that light also has particle-like characteristics. This is called the wave-particle duality of light.
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First, Let’s look at the WAVE nature of light
Light as a WAVE
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Wave Properties
A
Lesser frequency
greater frequency
crest
trough
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Properties of Light
• The significant feature of wave motion is its repetitive nature, which can be characterized by the measurable properties of wavelength and frequency.
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Waves
• Wavelength () - length of one complete wave
• Frequency (f) - # of waves that pass a point during a certain time period– hertz (Hz) = 1/s
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Properties of Light
• Electromagnetic radiation is a form of energy that exhibits wavelike behavior (wavelength, frequency, ect) as it travels through space. Together all forms of electromagnetic radiation form the electromagnetic spectrum.
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On the electromagnetic spectrum, the lowest energy waves (longest wavelength and lowest frequency) are radio waves. The highest energy waves (shortest wavelength and highest frequency) are gamma rays.
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• FREQUENCY and WAVELENGTH are INVERSELY proportional. (f ↑ ↓)
• ENERGY and FREQUENCY are DIRECTLY proportional. (E↑ f ↑)
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Now, let’s look at the PARTICLE nature of light
Light as a PARTICLE
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Properties of Light
• In the early 1900’s, scientists conducted experiments involving interactions of light and matter that could not be explained by the wave theory of light.
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Properties of Light
• One experiment involved a phenomenon known as the photoelectric effect. The discovery of the photoelectric effect led to the description of light as having both wave and particle properties.
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• A Quantum of light energy is called a PHOTON.
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EM Spectrum• Frequency & wavelength are inversely
proportional
c = f c: speed of light (3.00 108 m/s): wavelength (m, nm, etc.)f: frequency (Hz)
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Example 1
If the frequency of a wave is 500 hz, what is the wavelength?
C = λf
C = 3.0 x 108 m/s
f = 500 hz
λ = ?
3.0 x 108 = 500 x λ λ = 600,000 m
One sig fig or 6 x 105
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Quantum Theory
E: energy (J, joules)h: Planck’s constant (6.626 10-34 J·s)f: frequency (Hz)
E = hf
The energy of a photon is proportional to its frequency.
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Example 2
What is the energy of a wave if the frequency is 300. hz?
E = hf
f= 300. hz
h = 6.626 x 10-34
E = 300. x 6.626 x 10-34
E = 1.99 x 10-31 3 sig figs
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Example 3
• If the energy of a wave is 9.00 x 10-19 J, find frequency and wavelength
E = hf
E= 9.00 x 10-19 hz
h = 6.626 x 10-34
9.00x 10-19 = 6.626 x 10-34 x f
f = 1.36 x 1015 hz 3 sig figs
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If the energy of a wave is 9.00 x 10-19 J, find frequency and
wavelength C = λf
If f = 1.36 x 1015 hz
C = 3.00 x 108 m/s
3.00 x 108 m/s = λ x 1.36 x 1015
λ = 2.21 x 10 -7m 3 sig figs.