Chapter 5 Electrons in Atoms
Atomic Structure Part II
Electrons in Atoms Chapter 5
I. Light and Quantized Energy - Properties of Waves 5.1 pgs. 117-126
1. Definition:
2. Wavelength (λ
λ) distance from peak to peak
3. Frequency (ν)
a. number of peaks that pass at a given
point each sec
b. can be called cycles per second (peak/sec)
c. cps now called 1 Hertz (Hz)
4. Velocity (c = speed of light)
a. distance a given peak moves in a unit of time
b. velocity (m/s) = frequency (ν) x wavelength(λ
c=ν x λ
II. Behavior of Light
A. Newton (1600) thought light consisted of particles (beam of light is a stream of particles)
B. Maxwell (1864) thought light was a wave phenomenon.
1. some say light is like waves, some say its like particles
2. modern theory says that it behaves as both "wave/particle duality"
3. Max Planck (early 1900's) said:
a. light is made up of bundles of energy called photons (or quanta)
b. the energy of each photon is proportional to the frequency of the light
(Quantum Theory)
A quantum is
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Chapter 5 Electrons in Atoms
example: CONTINUOUS SPECTRUM
*** when white light is passed through
a prism, it is separated into a band of colors
from red violet. It's called a continuous
spectrum
E =
c. the work of Planck & Einstein led to
h x ν
E = Energy in Joules, ν = frequency and h = Planck’s constant 6.6262x10-34J
J is the symbol for
Energy of a quantum is related to
According to Planck’s theory, for a given frequency
d. The photoelectric effect
In the photoelectric effect, electrons, called photoelectrons, are emitted from a metal’s
surface when light of a certain frequency shines on the surface. (example: solar calculator)
Einstein said light can both wavelike and particlelike natures. That is, while a beam of light
has many wavelike characteristics, it also can be thought of as a stream of tiny particles, or
bundles or energy, called photos.
A Photon is
III. Bright line spectrum
A. a spectrum that shows
separate bright lines, each with a
specific wavelength
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Chapter 5 Electrons in Atoms
B. bright-line spectra occur when an element is heated and the colored light given off is
viewed through spectroscope. Each element has a unique set of lines, characteristic of that
element (like a fingerprint)
IV. Electromagnetic Spectrum
A. visible light (like the continuous
spectrum) is only one type of radiation.
All other types are not visible to the
human eye.
B. all forms of electromagnetic
radiation travels at the speed of light.
1. speed of light = 3.00 x 108 meters/sec
2. use formula:
C= V x λ
3. each line spectrum has a particular frequency (V). If know wavelength (λ), we can find V
using c as a constant.
C. The energy in a photon of light is directly proportional to the frequency
of the light.
frequency,
energy
1.
2. can find the energy of a single quantum (photon) of radiation at any given frequency.
3. proportionality constant that relates the two is called Planck's constant (h).
4. formula:
E = h x ν
ν
example: a spectral line has frequency of 3.5x10 12 hertz. What is the energy of a photon of
radiation of this frequency?
V. Electron energy levels in Bohr's Model 5.2 pgs. 127-128
A. There are certain different orbits in which an electron can travel around a nucleus.
1. each circular orbit (or shell) is at a fixed distance from the nucleus
2. the greater the radius of that shell, the greater the energy of the
electron in that shell.
3. these electron orbits are known as energy levels
B. When electrons absorb energy from an outside source, they jump from lower to higher
energy levels.
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Chapter 5 Electrons in Atoms
when they fall back to their original levels , energy is emitted (light); the same
amount as was absorbed.
C. In every atom in its normal state, all electrons are in the lowest energy levels available
(energetically stable)
VI. Atoms and Radiation
A. When all of the lowest energy levels
are occupied, the atom is in the ground
state (unexcited)
B. When electron moves to higher
energy level, atom is in the excited
state, and is energetically unstable.
C. Bright line spectrum of an element
represents the energy levels in its
atoms
.problems with Bohr's Model:
.only explained some of the lines in the bright line spectrum
.really only worked for hydrogen
.need sublevels and electron cloud model to account for all of the lines.
VII. The Quantum Mechanical Model of the Atom 5.2 pgs. 129-134
A. Mechanics
1. Classical Mechanics - Newton's Laws of Motion (Newtonian Mechanics)
2. Quantum Mechanics a. Louis de Broglie - particles could have properties of waves
b. Schrodinger - described the behavior of electrons in terms of quantized
energy changes "quantum mechanics"
c. Heisenberg - uncertainty principle
B. Principal Energy Levels
1. Energy Levels
Bohr -
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Chapter 5 Electrons in Atoms
Principal Quantum Numbers (N) Number of electrons
1
2
3
4
2. Sublevels
Principal Quantum
Numbers (N)
Sublevel Present
1
2
3
4
Orbital:
1s
2s2p
3s3p3d
4s4p4d4f
Rules for filling orbitals
1. Pauli Exclusion Principle
– Each orbital can hold TWO electrons with opposite spins.
– No two electrons in an atom can have the same 4 quantum numbers.
– Each e- has a unique “address”:
2. Aufbau Principle - Electrons fill the lowest energy orbitals first. Electrons to be added must be
placed in unfilled orbitals of lowest energy for stable configuration.
3. Hund’s Rule - In sublevel the second electron can't be added until each orbital in sublevel
contains one electron
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Chapter 5 Electrons in Atoms
3. Shapes
C. Electron Configurations pgs. 135-141
1s2 = Helium
1s22s1 = Lithium
1s22s22p63s23p64s2 = Calcium
Try W = ?
Try a few! Mg, Fe, Ru, Ir, Ca+2 = Cl-1 =
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