http://mrhlanc.tripod.com/hhl3/page13.html
Electron Configuration
I Radiant Energy
A. study of atomic structure uses light
B. have properties of waves and particles.
C. waves
1. light is electromagnetic waves
a. electric field
b. magnetic field
c. oriented 90 degrees to each other
d. wave travels perpendicular to fields.
e. form of electromagnetic radiation.
f. ex. x-rays, gamma rays, radio waves
2. all waves have four characteristics
a. amplitude
1. wave height - origin to crest
2. provides intensity / brightness
b. wavelength
1. distance between crest
2. visible light 400 - 750 nm
c. frequency
1. number of times a wave completes a cycle in 1 second.
2. unit - s-1 or hertz (Hz)
d. speed
1. 3.00 x 108 m/s
2. = c/v relates frequency, wavelength, and light speed
3. Electromagnetic Spectrum
a. gamma to radio 10-11 to 10-1 m
b. visible spectrum 4.00x10-7 to 7.50x10-7 m
1. violet shortest wavelength highest frequency
2. red longest wavelength lowest frequency
3. ROY G BIV
II Quantum Theory
A. Planck's Theory
1. problem: why does the emitted spectrum change with temperature
2. explained by Max Planck
3. proposed revolutionary idea
a. energy emitted or absorbed in specific amounts
b. called a quantum
c. contradicted classical physics: energy continuous
d. quantum: basis for today's modern model of the atom
4. relationship between frequency and energy
a. E = hv
b. h: Planck's constant 6.6262 x 10 -34 j-s
5. energy absorbed or emitted by atoms is quantized
6. detectable with only extremely small particles (subatomic)
B. the photoelectric effect
1. electrons ejected from metal surface when light hits it.
2. each metal requires light of a specific frequency (energy).
3. Einstein thought light acts like particles of energy called photons.
4. carries specific amount of energy - Planck's equation.
5. electrons of metals can absorb only specific amounts of energy.
6. electrons absorb all the energy or none of it.
7. too little or too much will not eject the electron.
8. demonstrates the relationship between frequency and energy - why x-rays are dangerous
and radio waves are not.
C. dual nature of radiant energy
1. Compton proved the dual nature of light
a. light is like a particle
b. has a frequency and wavelength
III Another Look at the Atom
A. line spectrum
1. spectrum that contains only certain colors, wavelengths
2. elements absorb energy and then emit unique line spectrums - atomic emission spectrum
3. Na - yellow
N - orange
Hg - blue
B. the Bohr model of hydrogen atom
1. related Rutherford planetary model (nucleus with circling electrons) to Planck's quantum theory
2. explained the unique spectral lines of elements
3. electrons are quantized
4. electrons possess specific amounts of energy
5. electrons restricted to specific orbits (distance from the nucleus)
6. electrons - lowest energy - closest to nucleus - first energy level
7. Bohr labeled each level using quantum numbers (n )
8. first energy level n =1
9. electrons absorb energy jump to a higher level (higher energy) n = 2, 3, or 4 farther from the nucleus
10. higher energy called excited state
11. energy emitted when electron returns to the lower energy level
12. amount emitted = to the difference between the energy level involved
13. Bohr used Planck's equation - calculated the frequencies for the emissions
14. matched the observed frequencies for hydrogen
15. did not work well with multi electron atoms
16. suggests more to understand about electron arrangement
C. matter wave
1. wave property of matter
2. proposed by De Broglie (1924)
3. quantized (particle) and wavelength (energy)
4. developed equation for the wavelength of matter
a.
= h
mv
b. needed experimental evidence
5. Davisson and Germer did it
a. electrons reflected from metal surfaces
b. same reflection pattern as x-rays
c. demonstrated electrons reflected as a wave
6. not observable with large object - wavelength too small to be detected
D. Heisenberg's uncertainty principle
1. position and momentum of a moving object can not be determined simultaneously
2. the act of measuring disturbs the position and behavior of the electron
3. no way to measure the orbits of Bohr's atom
4. instead, use probability to describe the location of electrons in an atom
IV A New Approach to the Atom
A. quantum-mechanical model
1. quantized
2. wavelike
3. can not know exact location
B. probability and orbital
1. areas of greatest probability (electron density) described as a cloud
2. orbitals - certain regions where electrons with given energy are likely to be found
a. have shape
b. size
c. energy
d. does not imply how electrons move
C. electron cloud includes the area with a 90% probability of finding an electron
D. several orbital shapes
1. s spherical
2. p dumbbell
3. d and f more complex
E. electron energy determines the orbital it occupies
F. orbitals and energy
1. principal quantum number (n) - relative distance from the nucleus
a. more energy higher principal value
b. farther from the nucleus
2. sublevel quantum number (l) - shape of orbital
a. 4 shapes
b. divisions of an energy level
c. energy values s<p<d<f
d. number of sublevels in an energy level = to principal Q.N.
1. 1 s
2. 2 s, p
3. 3 s, p, d
4. 4 s, p, d, f
3. magnetic quantum number (m) - orientation of orbital
a. position of orbital in relation to the nucleus
b. number of position
1. s -1
2. p - 3
3. d - 5
4. f - 7
4. spin quantum number (s) - spin
a. clockwise
b. Counterclockwise
G. electron spin
1. spinning charge creates a magnetic field
2. clockwise spin north pole up
3. counterclockwise spin north pole down
4. Pauli exclusion principle
a. orbitals can hold only 2 electrons
b. must have opposite spins
c. limits the number of electrons in each energy level and sublevels.
5. maximum number of electrons
a. s - 2
b. p - 6
c. d - 10
d. f - 14
V Electron Configuration
A. distribution of electrons in an atom
B. used by chemists to understand how atoms interact
C. determined by distributing the atom's electrons among the levels, sublevels, and orbitals
D. determining electron configuration
1. electrons normally at ground state - lowest energy
2. based on three principles
a. aufbau - electron added one at a time to lowest energy orbitals
b. Pauli exclusion - orbitals hold a maximum of 2 electrons with opposite spins
c. Hund's rule - each orbital orientation of a sublevel must contain 1 electron before being paired up
E. exceptions to the aufbau principle
1. 4s13d5 rather than 4s2 3d4
2. little chemical consequence
TOP