ELECTRONS IN ATOMS
A. The Evolution of Atomic Models
1. Dalton's Theory
a. Invisible, indivisible particles
2. J.J. Thomson Model
a. "Plum Pudding Model"
b. electrons are stuck in a lump of positively charged matter
c. Baed on discovery of the electron
3. Ernest Rutherford Model
a. "Nuclear model"
b. electrons surround a dense positive core
1. atoms are mostly empty space
c. Based on discovery of the nucleus
4. Niels Bohr Model
a. "Planetary" or "Orbital" model
b. elecrons move around the nucleus in fixed paths
1. Paths are called Energy Levels
2. electrons did not collapse into nucleus (unexplained phenomenon)
c. Energy levels are like a ladder
1. Need exact amount of energy to change levels, can't go half way
d. Quantum -- amount of energy needed to jump one energy level
B. Quantum Mechanical Model
1. erwin Schrodinger, 1926
2. Mathematical formula to determine the location of an electron in a Hydrogen atom
a. Sort of like the Bohr model --> electrons have fixed energy levels
b. Different --> probability of finding the electron in certain positions in the electron cloud (energy levels are more like a region of space)
1. Represented as a fuzzy cloud
2. High probability --> more dense cloud
3. Low probability --> less dense cloud
4. Cloud only shows electron position 90% of the time
C. Atomic Orbitals
1. Energy levels = principal quantum number (n)
a. 1, 2, 3, ..., n
2. sublevels within each principal energy level
a. # of sublevels = principal quantum number
1. 1st level has 1 sublevel, 2nd level has 2 sublevels, etc.
3. Atomic orbitals - regions in which electrons are likely to be found
a. Represented by letters (s, p, d, f)
b. s = spherical, p = 3 dunbells on x, y, z axis, d = 4 clover leaves and 1 dumbell, f = too complex to visualize
c. Node - area where probability of finding electrons is low
Principal # of sublevels type of sublevel # of atomic Max. # of electrons
energy level (n) orbitals (n^2) (2n^2)
n = 1 1 1s 1 * 2
n = 2 2 2s, 2p 4 * 8
n = 3 3 3s, 3p, 3d 9 * 18
n = 4 4 4s, 4p, 4d, 4f 16 * 32
* - s = 1 orbital, p = 3, d = 5, f = 7
II. Electron Arrangement in Atoms
A. Electron Configurations
1. Def. --> the arrangement of electrons around the nucleus
2. Three Rules
a. Aufbau principle: electrons enter orbitals of lowest energy first
b. Pauli Exclusion Principle: an atomic orbital may hold two electrons at most
c. Hund's Rule: when electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spin
3. Shorthand version involves energy leel, sublevel, & # of electrons
a. Ex. --> 1s^2
B. Exceptional electron configurations
1. Use these rules up to Vanadium (#23)
2. filled energy sublevels are more stable than half-filled sublevels
3. Half-filled energy sublevels are more stable than other configurations
III. Physics and the Quantum Mechanical Model
A. Light and Atomic Spectra
1. elements emit light when heated (Flame test lab!)
a. Atomic emission spectrum (colors of light emitted by that element, remember ROY G BIV)
b. each element has it's own atomic spectrum
B. The Quantum Concept and the Photoelectric Effect
1. Photoelectric effect --> metals eject electrons when light shines on them
a. Photon --> light quanta (bundle of energy)
b. Wave-Particel Duality of Light --> light exibits wave and particle properties
C. An Explanation of atomic spectra
1. Ground State --> lowest energy level on an electron
2. When atoms are heated , the electrons jump to a higher energy level (excited state), when they cool again and return to their ground state, they release a photon, the energy of the photon dictates the color that is observed
a. Red -- highest amount of energy, Violet -- lowest amount of energy
D. Quantum Mechanics
1. New branch of chemistry
2. classical mechanics deals with motion of large bodies
3. quantum mechanics deals with motion of subatomic particles and atoms as waves
Outline based upon:
Matta, M. S., Staley, D. D., Waterman, E. L., & Wilbraham, A. C. (2000). Chemistry, Addison-Wesley. (5th ed.). Menlo Park, CA: Prentice Hall, pp. 361-384.
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