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Honors Chem Chapter 4The Tiny but Mighty Electron
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Electron: What do we already know?
• It has a negative charge (Thomson) • It is small in mass: approximately 0.0005 amu
or 9.11x10-28 grams (Milikan)• Previously, it was thought of traveling in orbits
around the nucleus (Bohr) • Currently, it is thought of as moving rapidly
outside of the nucleus in orbitals
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Orbital
• A probable space outside of the nucleus where an electron is likely to be found
• Electrons are organized in their orbitals according to relative energy
• Low energy: Closest to the nucleus• High energy: Farthest from the nucleus
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What do orbitals look like?
“In a science that we cannot see, a lot is left to the imagination…” –Correspondent from NOVA NOW
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Level 1
• In energy level 1, there is one orbital shape. Its called an s orbital and looks like a sphere.
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Level 2
• In energy level 2, there are two orbital shapes. First, an s orbital and then a p orbital. The p orbital is shaped like dumbbell. There are 3 of these shapes. Each one is a subshell.
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Level 3 • In energy level 3, there are three orbital
shapes. One s, Three p’s, and Five d’s.
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Level 4
• In energy level 4, there are four orbital shapes. One s, Three p’s, Five d’s, and Seven f’s
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Heisenberg Uncertainty Principle
• We cannot know the exact location and speed of an electron at the exact time. We can know one or the other precisely, but never both at the same time.
• WHY????
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Finding An Electron: Quantum Numbers
• We can assign an electron a series of 4 numbers to “guesstimate” where it can be found in an atom.
• 4 parts: n, l, m, s
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Quantum Numbers: Principle Energy Level
• Represented by the letter “n”
• Whole number: 1, 2, 3, 4, etc…• Represents the size of the orbital. The bigger
the number, the larger the orbital and also the further it is from the nucleus
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Quantum Numbers: Angular
• Represented by the letter “l” • Whole number: 0, 1, 2, and 3• Represents the shape of the orbital – For an “s” shape: l = 0– For a “p” shape: l = 1– For a “d” shape: l = 2– For an “f” shape: l = 3
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Quantum Numbers: Magnetic
• Represented by the letter “m” • Can be a range of numbers from the negative
integer to the positive integer• Represents the orientation (or number of
subshells) – For “s” shape: m = 0 (only one orientation)– For “p” shape: m= -1, 0, 1 (3 orientations)– For “d” shape: m = -2, -1, 0, 1, 2 ( 5 orientations)– For “f” shape: m = -3, -2, -1, 0, 1, 2, 3 (7 orientations)
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Quantum Numbers: Spin
• Represented by the letter “s”
• Can be either +1/2 or -1/2 • Represents the spin direction of the electron
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Why can’t electrons stay in one place?
• Electrons are “hit” by ambient radiation sources and when they are given more energy they are “promoted” to a higher energy level
• This is borrowed energy so what goes up must come down.
• This release of energy comes in the form of light and/or heat!
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Dual Nature: Particles and Waves
• Light is packets or bundles of energy, known as photons, that travel in waves (Einstein)
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Electrons and Energy
• Photoelectric effect: electrons are emitted from samples of matter when they are exposed to radiation energy—LIGHT!
• Electrons have energy, absorb energy, and release energy
• Electrons and Light exhibit properties of both particles and waves
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Why can’t electrons stay in one place?
• Electrons are “hit” by ambient radiation sources and when they are given more energy they are “promoted” to a higher energy level
• This is borrowed energy so what goes up must come down.
• This release of energy comes in the form of light and/or heat!
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Sources of Energy
• Light/Heat • Electricity • Chemical Reaction• Nuclear Reaction
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Ground State
• The ground state of an electron is its lowest-energy state.
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Excited State
• An excited state is when an electron has been promoted to a higher energy level.
• As the electron returns to its ground state, it releases the specific gained energy in the form of light
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Bright Line Spectrum
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Bright Line Spectrum
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Understanding Waves: Wavelength
• Wave = repetitive transfer of energy
• Wavelength (λ) = distance over which the wave’s shape repeats
• Generally measured in nanometers (1 x 109 nm = 1 m)
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Understanding Waves: Frequency
• Frequency (ν) = number of occurrences of a repeating event per unit time
• Measured in Hertz (Hz) OR “waves per second” (1/s = s-1)
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Frequency and Wavelength Related
• Frequency and Wavelength are inversely related.
• One variable increases while the other decreases.
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The Light Equation
•c = λν• The speed of light (c) is equal to 3.0 x108 m/s• The wavelength (λ) is in meters • The frequency (ν) is in Hertz
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Wave Practice
An orange light has a wavelength of 492nm. What is the frequency of this light?
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Energy and Frequency
• High frequency (short wavelength) is a high energy wave.
• Frequency and energy are directly related • Both variables increase together
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The Electromagnetic Spectrum
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The Visible Spectrum• ROY G BIV: Red, Orange, Yellow, Green, Blue,
Indigo, and Violet
Increasing energy
Increasing frequency
decreasing wavelength
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Bright Line Spectrum
• Light is a combination of multiple colors• Each atom has its own unique pattern of
electrons that absorb energy differently • No two atoms have the same bright line
spectrum: used for identification• Each line in the spectrum represents electrons
releasing a specific amount of visible energy—translating to a specific color.