A.Amplitude: Height of the wave. The higher the wave the greater the intensity. B.Wavelength: ( λ, “lambda”) in nanometers (1 x 10 -9 m) measured from crest to crest or from trough to trough. C. Frequency: ( ν, “nu”) The number of times a wave passes a fixed point. Measured in cycles/second (1/s) 1 cycle/second = Hertz (Hz) ex) Radio FM 93.3 megahertz (MHz) is 93.3 x 10 6 cycles/sec.
Visible Light Radar Microwaves Infrared Radio/TV Ultraviolet X-Rays Gamma Rays Low Long High Short RedOrangeYellowGreenBlueIndigoViolet Energy Low High
Speed of Light D. (c) 3.0 x 10 8 m/s or 186,000 miles/sec. The relationship between wavelength and frequency can be shown with the following equations: or this is an indirect relationship. If λ then ν .
Quantum Theory A.Planck’s Hypothesis: (Max Planck 1900) 1. An object absorbs or emits light in little packet or bundle called a quantum (quanta –plural). 2. Energies are quantized. (Think steps not a ramp.) 3. Equation relating energy (E) to frequency ( ν, nu) (h= Planck’s Constant ) E = h This is a direct relationship, as ν , E e-e- e-e- X e-e- e-e-
Light (Electromagnetic Radiation) 1. A quantum of light is called a photon. 2. Light travels through space in waves. 3. Light acts like a particle when it interacts with matter. 4. This shows the Dual Nature of Light.
3. The Atom A. Atomic Emission Line Spectra: contains only certain colors or wavelengths ( ) of light. 1. Every element has its own line spectrum (fingerprint).
Double Slit Experiment Double Slit Experiment Video This Experiment Proves the Dual Nature of Light – Photons of Light travel through spaces in waves, but act like particles when they interact with matter.
Continuous Spectrum – White Light Line Spectrum – Excited Elements Line Emission Spectra of Selected Elements
Gas Discharge Tubes Electricity is added to the gas which causes the electrons to jump to a higher or excited state. They immediately fall back to the ground state and give off particular wavelengths of light. We see a blending of wavelengths without the spectroscopes.
White Light gives off a Continuous Spectrum a blending of every possible wavelength
Spectroscope Uses a diffraction grating to diffract the light into particular wavelengths of light.
A Line Spectrum results from excited elements - as electrons of an element gain energy and rise to an excited state they then fall back to their ground state in the same pattern producing the same energy drop each time which we see as individual wavelengths of light.
Atomic Spectra Hydrogen Lithium Mercury Helium Although Bohr’s atomic model explained the line spectra of hydrogen, it failed for heavier elements.
B. Bohr Model of Hydrogen: (1911) 1. Bohr said the energy of an electron was quantized (only certain orbits that represented different amounts of energy.) 2. Bohr labeled each energy level with a quantum number (n). 3. n=1 lowest level or ground state. 4. When electrons absorb energy they jump to a higher ( excited ) state. n=2 n=3 n=4 n=5 n=6 n=7 5. Radiation ( light ) is emitted when an electron falls back from a higher level to a lower level.
Path of an excited electron as it “falls” back to the Ground State When electrons gain energy, they jump to a higher energy level (excited state). Electrons are not stable at the excited state and will immediately fall back to a lower level or ground state. As they fall, they emit electromagnetic radiation. Depending on how far they fall determines the type of radiation (light) released.
electrons fall to n = 1 and give off ultraviolet light. electrons fall to n = 2 and give off visible light. electrons fall to n = 3 and give off infrared light. Hydrogen Atom
Werner Heisenberg: (1927) 1. Heisenberg’s Uncertainty Principle : states the position and momentum (speed, direction & mass) of a moving object cannot simultaneously be measured and known exactly. 2. You cannot predict future locations of particles. 3. He found a problem with the Bohr Atom - no way to observe or measure the orbit of an electron.
4. Quantum Mechanics A. Quantum Mechanical Model of the atom combines previous ideas and treats the electron like a wave that has quantized energy. Impossible to state the exact position or momentum of an electron, but you can state a probability of where the electron is located.
B. Electron Density Where the density of an electron cloud is high there is a high probability that is where the electron is located. If the electron density is low then there is a low probability.
C. Atomic Orbitals - region around the nucleus where an electron with a given energy is likely to be found (not the same as Bohr’s orbits!) 1.Orbitals have characteristic shapes, sizes, & energies. 2. Orbitals do not describe how the electron moves. 3. The drawing of an orbital represents the 3-dimentional surface within which the electron is found 90% of the time.
4. Sublevels can have 4 different shapes s – orbital spherical
1s, 2s & 3s orbitals Superimposed on one another
Quantum Numbers - Finding an address for each electron: 1) “state” Principle Quantum Number (n) or the energy level ranges from n=1 to n=7 2) “city” Sublevel shape either s, p, d,or f. 3) “street” Orbital The orientation in space ex) x,y,z axis 4) “house” Spin The cw or ccw motion of electrons.
1s, 2s,2p and 3s orbitals superimposed on each other
2. The number of Sublevels in an energy level equals the Principle Quantum Number (n). 3. Orbital: There are always an odd number of orbitals. s-sublevel has 1 orbital p-sublevel has 3 orbitals d-sublevel has 5 orbitals f -sublevel has 7 orbitals Orbitals in higher principle levels get larger. A max of 2 electrons fit in each orbital.
Electron Spin a.Two electrons in each orbital have opposite spins. ( clockwise and counterclockwise ) b. The opposite spin is written: or ___
c. Pauli Exclusion Principle: 1. Each orbital can only hold 2 electrons. 2. The electrons must have opposite spins. s-sublevel = max 2 electrons p-sublevel = max 6 electrons d-sublevel = max 10 electrons f-sublevel = max 14 electrons incorrect: ↑↑↑ incorrect: ↑↑ correct: ↑↓
d. Hund’s Rule: Electrons will spread out among the orbitals before they pair up. incorrect ↑↓ ↑ __ correct ↑ ↑ ↑
E. Electron Configurations: 1. Shows the distribution of electrons among the orbitals. Describes where the electrons are found & what energy they possess. 2. The Aufbau Principle: electrons are added up one at a time to the lowest energy orbital.
Electron Configuration Examples: Ex) electron configuration for Na: 1s 2 2s 2 2p 6 3s 1 Ex) orbital filling box diagram for Na:
3. Electron Dot Diagrams: Write the symbol for the element. Place dots around the symbol to represent the s & p valence s & p electrons only. Do NOT include d & f orbitals in diagram. p orbital electronss orbital electrons
Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Pauli Exclusion Principle: No more than 2 e- are put in each orbital and they must have opposite spin. Hund’s Rule: electrons spread out among equal energy orbitals in a sublevel (like charges repel) Aufbau Principle: Electrons fill lowest energy levels first (n=1)
In 1926, Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom probability Equation for the probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger