quantum mechanics
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Unit 3, Part 2Quantum Mechanics
Chemistry Notes
Light & Wave Motion
• Waves carry energy from 1 place to another
• Light = Electromagnetic Radiation
–carries energy through space
–Behaves as a wave or as a stream of particles of energy
(wave-particle nature of light)
A Typical Wave
http://www.pas.rochester.edu/~afrank/A105/LectureV/FG03_003_PCT.gif
Wave Descriptions
• Wavelength: distance from crest to crest (or trough to trough); unit = m
• Wave speed: how fast the wave travels over time; unit = m/s
• Frequency: the # of waves that pass an area in a given amount of time; unit = 1/s = Hz
The Nature of Light
• The Photon – a bundle of a specific amount of electromagnetic energy
–Generally considered particle-like but also demonstrates wave-like properties
The Electromagnetic Spectrum
Gamma (shortest)
X-ray
Ultraviolet/UV
Visible
Infrared/IR
Microwave
Radio (longest)
The Electromagnetic Spectrum
http://www.yorku.ca/eye/spectrum.gif
Looking At Wavelength
Frequency
Short Wavelength
High Frequency
Energy
Short Wavelength
High Energy
Or
High Amplitude
High Energy
http://www2.glenbrook.k12.il.us/gbssci/phys/Class/sound/u11l2a2.gif
Visible Light Spectrum
Red
Orange
Yellow
Green
Blue
IndigoViolet http://cache.eb.com/eb/image?
id=73584&rendTypeId=35
Long Short
Absorption & Emission Spectra
Absorption
something is being “taken in”
Emission
something is being “given off”
--Atoms will release energy as photons of light; the energy of the photons corresponds to the energy changes experienced by the atom
Absorption & Emission Spectra
The “fingerprint” of an element; unique Absorption – when an atom absorbs
energy; electrons jump to higher energy levels; dark line spectra
Emission – when an atom releases photons of energy; electrons return to lower energy levels; bright line spectra
Absorption & Emission Spectra
http://www.astro.columbia.edu/~archung/labs/fall2001/images/spectrum_line.gif
Hydrogen Atom Example
Notice the spectral lines in the same positions? These represent specific energy level changes of the electrons in the atom.
http://www.solarobserving.com/pics/hydrogen-spectra.jpg
Dark Line
Bright Line
The Bohr Model
http://library.thinkquest.org/19662/images/eng/pages/model-bohr-2.jpg
Light Energy Released
Light Energy Taken In
Bohr Model of Atom
• Small nucleus with e- orbiting like planets around the sun
• Electrons occupy specific energy levels (orbits)
• Currently, atomic models suggest that e- occupy 3D regions called orbitals within an electron cloud
Ground State vs. Excited State
http://imagine.gsfc.nasa.gov/docs/teachers/lessons/xray_spectra/images/emit.gif
http://content.answers.com/main/content/wp/en/thumb/c/c4/180px-Energylevels.png
Ground State = lowest energy level
Excited State = a higher energy level
Higher energy levels are closer together…
Energy Levels Are Quantized
1. Atom absorbs energy, electrons jump from ground state to higher energy level
2. Electrons return to ground state, atom gives off photon of energy
The photon is quantized = carries a specific amount of energy that depends on the size of the “jump” from 1 energy level to the next
Chemical Behavior of Atom
• Determined by electron structure
• How the electrons are arranged determines how they will behave in a reaction
Locating Electrons in an Atom
• Impossible to predict the exact location
• The most probable location is described as the orbital within electron cloud
• Describe the locations of electrons (or its address) as the electron configuration
The Electron Configuration
• Pauli Exclusion Principle
–No 2 electrons can have the same electron figuration
–No 2 electrons can occupy the same location at the same time
–Remember - electrons are always moving!
Electrons & Energy Levels
• Energy levels are quantized = specific amounts of energy
• Energy levels are called the principle energy levels
• Each principle energy level can be divided into specific sublevels
The Hydrogen Example
• Atomic Number = 1
–1 proton
–1 electron
• Hydrogen has only 1 electron but more than 1 orbital. The electron can only occupy one orbital at a time.
Quantum Numbers
• 4 quantum numbers make up the electron configuration–(n) = principle energy level
–(l) = energy sublevel
–(m) = actual orbital
–(s) = spin
Quantum Number Analogy
• Hospital Complex Analogy–(n) = building
–(l) = floor
–(m) = room
–(s) = bed
(assuming 2 beds per room)
Electron Configurations Written
• Example = 2p3
2 = principle energy level
p = sublevel
3 = # of electrons in sublevel
Quantum Number (n)
• Principle energy level
• n = 1, 2, 3, 4, 5, 6, or 7
• n = corresponds to the row (period) on the periodic table
• As n increases, energy increases
• n refers to orbital size; as n increases, the orbital size increases
• 1 = close to nucleus, 7 = far away
Quantum Number (l)
• Energy sublevel
• Not all e- in an energy level have the exact energy
• l= s, p, d, or f
• Sublevels may overlap
• A filled sublevel is spherical in shape
• Orbital shapes can be complex
Quantum Number (m)
• The actual orbital of the electron
• Each orbital holds 2 electrons
• s sublevel – 1 orbital, spherical shape
• p sublevel – 3 orbitals, dumbell shape
• d sublevel – 5 orbitals, weird shape
• f sublevel – 7 orbitals, complex shape
Orbital Shapes
http://lincoln.pps.k12.or.us/lscheffler/OrbitalShapesOverheads_files/image002.jpg
Quantum Number (s)
• Spin of the electron
• Clockwise (+½) or Counterclockwise (-½)
• Electrons in the same orbital have opposite spins
3d ___ ___ ___ ___ ___
4s ___
3p ___ ___ ___
Principle Energy Level (n)
Energy Sublevel (l)
Orbital (m) holds 2 electrons total
Example Orbital Diagrams
1s 1s2s 1s2s2p
http://www.uwgb.edu/dutchs/PETROLGY/WhatElmsLookLike.HTM
Electron Configuration Guidelines
• Elec. configuration lists ALL e- in the atom, giving the principle energy level, sublevel, and e- per sublevel
• An atom is neutral: the # of protons equals the # of electrons, described by the atomic number
• The period (row) describes the number of energy levels present
Electron Configuration Guidelines
• 2n2 = the maximum number of electrons in an energy level
–n = the energy level
n = 1, then 2 electrons
n = 2, then 8 electrons
n = 3, then 18 electrons
Electron Configuration Guidelines
• The nth energy level has n sublevels present
–n = 1, then 1 sublevel
–n – 2, then 2 sublevels• Any element is “put into” the energy
level (period) of its highest energy level electrons, even if these electrons are not filled in last
Electron Configuration Guidelines
• Lower energy levels fill in first; then higher energy electrons are filled in
• Hund’s Rule: in any sublevel, one electron is placed in each orbital before pairing begins– Some exceptions (discussed later)
• Degenerate: orbitals have the same energy & appear at same “level” on an orbital filling diagram
Electron Configuration Guidelines
Place one electron in each
orbital before pairing!
Degenerate Orbitals: All 3 have the same energy and appear at the same level
Hund’s Rule
Electron Configuration Guidelines
• Order of Orbital Filling
You will use the Periodic Table to help remember this!
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d…
Orbital Filling Clues
http://pegasus.cc.ucf.edu/~jparadis/chem2045/chapter06.html
Electron Configuration Guidelines
Reminder!
Each Orbital Holds 2 Electrons
s 1 orbital 2 electrons
p 3 orbitals 6 electrons
d 5 orbitals 10 electrons
f 7 orbitals 14 electrons
Orbital Filling Diagrams
• See notes packet for diagram template
• Aufbau Principle (to build up)–Electrons fill lowest energy states
first–If multiple orbitals, electrons fill all
unoccupied orbitals first before pairing electrons with opposite spins (Hund’s Rule)
Exceptions You Need to Know
• Copper (Cu) & Chromium (Cr)–Electrons are found in unexpected
places; there is extra stability in having half-filled d orbitals
–The completion of a d sublevel is desired, so an electron will fill the d sublevel before the next s
–1 electron in each d5
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