Transcript Atomic Spectra Bohr Model Notes

Warm Up
Draw the Bohr Model for Aluminum and
Neon.
Electrons in Atoms
Unit 2- Continued
Everything you ever wanted to know about
where the electrons hang out!
Section 1: Early 1900’s
Scientists started doing a lot of experiments
looking at the absorption and emission of
light by matter.
Found that there is a relationship between
light and an atom’s electrons.
Light behaves as a wave
Transfer of energy
Draw the Wave!
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Amplitude: height of the wave from the origin to the crest
Wavelength ( ) : the distance between the crests (m, cm,
nm)
Frequency (v): number of waves to pass a given point per
unit of time (waves/second = Hz)
An Important Relationship
The frequency and wavelength of all waves,
including light, are inversely related.
As the wavelength of light increases, the
frequency decreases.
Electromagnetic Radiation
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Includes radio waves, radar, microwaves, visible
light, infrared light, ultraviolet light, X-rays, and
gamma rays
Wave Particle Duality
http://www.youtube.com/watch?v=DfPeprQ7oGc
The Photon
Photon- a particle of
electromagnetic
radiation having no
mass, carrying a
quantum of energy.
Photoelectric Effect
Looks at the emission of
electrons from a metal
when light shines on
the metal.
Light causes electrons to
be ejected from the
metal.
So, what happens when photons
hit an atom and eject an electron?
The electron goes from
the ground state to an
excited state.
As the electron returns to
the ground state, it
gives off the energy
that it gained- LIGHT
Energy Levels
Energy levels are
not evenly spaced
• Energy levels
become more closely
spaced the greater
the distance from the
nucleus
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Another Cool Illustration
Color
The energy given off has a certain wavelength.
Wavelength determines the colors that we see.
Flame Test
Look at the color produced by the flame…
Determine the wavelength by comparing the
color to those in the visible range on the
Electromagnetic Spectrum.
Warm Up
You have two different samples… sample A.
glows red and sample B. glows violet.
a. Draw what the waves might look like?
b. Which has the longer wavelength?
b. Which has the smaller frequency.
Atomic Spectra
White light is a combination of all the
wavelengths in the visible range of the
Electromagnetic Spectrum.
Spectral Tubes
Each element has a
unique line-emission spectra
Atomic Line Spectrum
Interpretation of Atomic Spectra
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The line spectrum is related to energy transitions in
the atom.
Absorption = atom gaining energy
Emission = atom releasing energy
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All samples of an element give the exact same pattern
of lines.
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Every atom of that element must have certain,
identical energy states
Atomic Spectrum Activity
Using Atomic Spectral Data
Bohr Model
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Electrons orbit around a
nucleus
Each orbit has a fixed
energy and because of this
cannot lose energy and fall
into the nucleus
Energy Level of an
electron is the region
around the nucleus where
the electron is likely to be
moving
This helped explain the
spectral lines
Absorption- the electron
gains energy and
moves to a higher
energy level.
Emission- when the
electron falls to a
lower energy level.
Schrodinger Wave Equation
Developed an equation that treated
electrons as waves and described the
location of electrons.
Helped lay the foundation for modern
quantum theory (atomic model).
The Quantum Model
Finally– the truth (as we know it!)
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Electrons can behave as both waves and particles.
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Electrons can be considered waves with specific
frequencies confined to the space around the
nucleus.
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Electrons can also be considered negatively
charged particles.
Quantum Theory
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Estimates the probability of finding an
electron in a certain position
We denote the position of the electron as a
“fuzzy” cloud
This volume of space where an electron is
most likely to be found is called an orbital.
The atomic orbitals have distinct shapes
Work on Wave WS- 15 min
Go to shape ppt