Chapter 11 Atomic
Structure
• If atoms are too small to see, obviously
electrons are beyond our vision.
• We get clues from the atom on where
its electrons are located from energy it
absorbs and releases.
• This energy is in the form of waves of
various types from x-rays to radio
waves; all of which are classified as
electromagnetic waves.
Electromagnetic Waves
• A wave transmits energy from the
source to the receiver.
• Electromagnetic waves do this by
disturbing the electric and magnetic
fields surrounding the earth.
• The types of electromagnetic waves
are arranged on a spectrum from
greatest to smallest frequency.
Electromagnetic Spectrum
Color is Frequency
Media Frequencies
• AM radio - 535 kilohertz to 1.7 megahertz
• Short wave radio - bands from 5.9 megahertz
to 26.1 megahertz
• Citizens band (CB) radio - 26.96 megahertz to
27.41 megahertz
• Television stations - 54 to 88 megahertz for
channels 2 through 6
• FM radio - 88 megahertz to 108 megahertz
• Television stations - 174 to 220 megahertz for
channels 7 through 13
Household Gadgets
Frequencies
• Garage door openers, alarm systems, etc. Around 40 megahertz
• Standard cordless phones: Bands from 40 to
50 megahertz
• Baby monitors: 49 megahertz
• Radio controlled airplanes: Around 72
megahertz, which is different from...
• Radio controlled cars: Around 75 megahertz
• Wildlife tracking collars: 215 to 220
megahertz
• Cell Phones: 824 to 849 megahertz
• New cordless phones: around 900
megahertz!
Electromagnetic Spectrum
Wave Properties
• Wavelength (λ) is the distance
between two successive waves.
• Frequency (f) is the number of
waves passing a point per second.
Wave Properties
Wave Properties
• v = f * λ wave speed equation
• Wave speed (v) is the speed of
electromagnetic waves. In a vacuum
this is 3.0 x 108 m/s or 186,000 miles
per second!
• This is referred to as the “speed of
light” though all radiation travels at
this speed.
Energy of E/M Waves
• The energy of E/M radiation is
proportional to frequency.
• The more energetic waves are
higher in frequency like x-rays.
• E = h* f
• h is a constant in nature
• h = 6.626 x 10-34 J/Hz
Example
• What is the frequency and energy of
green light which has a wavelength of
510 nm?
• f = c/ λ = 3.0e8/510e-9 = 5.9e14 hz
• E = hf = 6.6e-34 * 5.9e14 = 3.9e-19 J
Bohr Model of the Atom
•
To explain the emission of light, Bohr
came up with a radical model of the
atom which had electrons orbiting
around a nucleus.
• This similarity between a planetary
model and the Bohr Model of the atom
ultimately arises because the attractive
gravitational force in a solar system and
the attractive Coulomb (electrical) force
between the positively charged nucleus
and the negatively charged electrons in
an atom are mathematically of the same
form.
Energy Levels of the Atom
• In the Bohr model,
electrons can
jump to higher
energy states by
the addition of
energy of certain
frequencies
• The excited electron is unstable and
falls down to ground state giving off
electromagnetic energy equal in
frequency to the energy change.
Hydrogen Emission Series
• The electrons can
move to only certain
locations from the
nucleus
• The result of this
restriction in
movement is
specific and distinct
wavelengths of EM
radiation is given off
as the electrons
transition.
Energy Quantified
• The energy transitions are quantified
and thus the transitions accept and
release only certain energies of E/M
radiation.
• The energy of each level for the
hydrogen atom is given by the
equation
• En = RH/n2
• RH is called the Rydberg constant
• RH = 2.180 x 10-18 J
Atomic Spectra
• White light emits all colors because a multitude of
electron transitions occur between energy levels.
• Spectral Lines give evidence of electron transitions.
• A specific species of gas emits only a limited amount
of spectral colors because of a more limited
availability of electron transitions.
Emission Spectra for H2 and He
More Emission Spectra Na, Ne, Hg