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Chapter 4 - Electrons
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Properties of Light
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What is light?• A form of electromagnetic radiation:
• energy that exhibits wavelike behavior as it travels through space
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Electromagnetic radiation is classified into two types:• Non-ionizing Radiation:
• Transfers energy causing vibrations, electron excitation, and heat
• Parts of the spectrum: radio, microwaves, infrared, and visible
• Ionizing Radiation:• High energy that ejects electrons and transforms
molecules into reactive unstable fragments• Parts of the spectrum: UV, X-ray, Gamma Ray
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Radiation is organized on the Electromagnetic Spectrum
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Wave PropertiesWavelength (λ) : • distance between
corresponding points on adjacent waves
• Unit: meters or nanometers
Frequency (ν):• number of waves that pass a
given point per unit time (usually 1 sec)
• Unit: 1/s = s-1 = Hertz (Hz)
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Properties of Light• How are frequency and wavelength related?
c = λνc : speed of light (m/s)
c = 3.00 x 108 m/sλ : wavelength (m)
ν : frequency of wave (s−1 = 1/s = Hz)
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Max Planck proposed …
Energy needed for electrons to move was quantized, or a quantum of energy is needed to move an electron.
E = hνE : energy of light emitted (J)
h : Planck’s constant (J·s) h = 6.626 × 10−34 J·s
ν : frequency of wave (s−1 = 1/s = Hz)
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E = hn
E = hn
Electrons can only move with a quantum of energy.
Every color represents a different amount of energy released.
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Louis de Broglie proposed …• that electrons be considered waves
confined to the space around an atomic nucleus.
• If electrons exist as waves, then they must have specific frequencies
• If they have specific frequencies, they also have specific energies
So energy is quantized!
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ELECTRON BEHAVIOR
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Who made this model of the atom?
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Electrons are always moving!
• The farther the electron is from the nucleus, the more energy it has.
• Electrons can change energy levels and emit a photon of light.
• Photon: packet of energy• Ephoton = hν
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Electron Movement Vocabulary
• Ground State: electrons in the lowest possible energy level
• Excited State: electrons absorb energy and move to a higher energy state
• Emission : when an electron falls to a lower energy level, a photon is released
• Absorption : energy must be added to an atom in order to move an electron from a lower energy level to a higher energy level
• How much energy must it absorb?• Quantum: amount of energy needed to move between energy
levels
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Electron Movement
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• When electrons move from the excited state to the ground state they emit a photon of light
• The light emitted by an element is viewed as a bright line emission spectrum
• Each band of light represents the energy released by an electron when it moves from higher to lower energy
Line Emission Spectrum of Hydrogen Atoms
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Because electrons move so much…
We cannot know an electrons exact speed and position at the same time
Heisenberg Uncertainty Principle:
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Schrodinger Wave Equation
• Equation that describes the wave & particle nature of an electron
• Don’t know the exact position…• Predict where electron will be most of the time
Where 90% of thee- density is foundfor the 1s orbital
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Electron Location
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DO NOW• How would you describe the location of your home?
• Like giving your house an address, each electron in an atom is assigned an address or location inside the atom.
• Remember: Do electrons ever stop moving?
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We do not know where the electrons are, but we can
describe where they might be
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• The address of all the electrons in an atom is called an Electron Configuration
Sodium #11
1s22s22p63s1
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First Location: ENERGY LEVEL
n = 1, 2, 3, 4 …
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ENERGY LEVEL
SUB-LEVEL
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Second Location: SUBLEVEL
These are the probability regions called: s , p , d , f
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SUB-LEVEL
ORBITAL
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Ene
rgy
1s
2s
3s
4s
5s
2p
3p
4p
3d
Third Location: ORBITALs = 1 p = 3 d = 5 f = 7
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Sublevels on the Periodic Table
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Examples:
• What is the electron configuration of Mg?
• What is the electron configuration of Cl?
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How to fill in an Orbital Diagram
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Ene
rgy
Step 1 – Aufbau Principle:Electrons are represented as arrows and drawn one at a time.Electrons start at the lowest possible energy.
1s
2s
3s
4s
5s
2p
3p
4p
3d
Element: HydrogenNumber of Electrons:
Electron Configuration:
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Ene
rgy
Paired electrons
Step 2 – Pauli Exclusion Principle:Electrons have opposite spins.Draw one up and one down.
1s
2s
3s
4s
5s
2p
3p
4p
3d
Element: HeliumNumber of Electrons:
Electron Configuration:
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Ene
rgy
Step 3 – Hund’s Rule:Electrons occupy equal energy orbitals so that a maximum number of unpaired electrons resultsDraw electrons in each orbital first then pair them.
Unpaired electrons
1s
2s
3s
4s
5s
2p
3p
4p
3d Element: OxygenNumber of Electrons:
Electron Configuration:
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What is the electron configuration of S?E
nerg
y
1s
2s
3s
4s
5s
2p
3p
4p
3d
Element: SulfurNumber of Electrons:
Electron Configuration:
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What is the electron configuration of Sr?E
nerg
y
1s
2s
3s
4s
5s
2p
3p
4p
3d
Element: StrontiumNumber of Electrons:
Electron Configuration:
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Valence Electrons
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Why is location important?
location influences behavior
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Valence Electrons• valence electrons:
• electrons in the outermost energy level• can be lost, gained, or shared• electrons that are responsible for reactions
• Elements in a group have similar properties because they have the same number of valence electrons
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How many valence electrons in each group?
Only the main block elements!
1
2 3 4 5 6 7
8
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• All electrons under the outer most level
Inner Core Electrons