Transcript Chapter 4-Student notes

Chapter 4
The Hook
Burning money demo—
Draw a picture of what happened.
What color did you see? Any idea why?
Why didn’t the money burn?
Chapter 4
Everything you ever wanted to know about
where the electrons hang out!
Building on the Atomic Theory
What did Thompson determine?
What did Rutherford’s gold
foil experiment prove?
Just write the words… we will talk in class!
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 can behave as a wave
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.
C = v
Where:
C= speed of light 3.00 x 108 m/sec
 = wavelength (m, cm, nm…)
v = frequency (1/sec or sec-1)
Electromagnetic Radiation
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Includes radio waves, radar, microwaves, visible
light, infrared light, ultraviolet light, X-rays, and
gamma rays
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.
Photoelectric Effect
Only occurs at certain frequencies!
Wave Particle Duality
Explained by Dr. Quantum
Leave some space here to write a reflection on
the video clip.
Sometimes Light Acts Like Particles!
What would happen if the frequency of the
wave increased so much that you could
hardly tell where one wave ended and
another began?
Light would start acting more like a particle
than a wave.
Max Plank
Objects emit small packets of energy- Quanta
Quantum- the minimum quantity of energy that can
be lost or gained by an atom.
E = hv
E = Energy
h = 6.626 x 10 -34 Js (Joule x sec)
V = frequency (1/sec)
Take a look at the WS
Let the units be your guide!!!!!
The Photon
Photon- a particle of
electromagnetic
radiation having no
mass, carrying a
quantum of energy.
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
The Visible Spectrum
From about 400nm to 700nm in wavelength.
Blue (400nm) has a shorter wavelength than
red (700nm).
Spectral Analysis of Emitted Light
from Excited Atoms
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When emitted light from excited atoms is passed
through a prism a spectrum of discrete lines of
separate colors (separate energies) is observed
rather than a continuous spectrum of ROY G
BIV.
Each element has a
unique line-emission spectra
Interpretation of Line Spectrum of
Elements
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The light atoms give off contain very specific
wavelengths called a line spectrum
light given off = emission spectrum
Continuous Spectrum
Atomic Line Spectrum
Atomic Spectrum Activity
Interpretation of Atomic Spectra
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The line spectrum is related to energy transitions of
electrons 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 because every atom of that element
must have certain, identical energy states
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Energy of an atom is quantized – limited to
discrete values
• If the atom could have all possible energies, then the
result would be a continuous spectrum instead of lines
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.
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.
Where are the electrons?
Heisenberg Uncertainty
Principle:
It is impossible to know
the position and the
velocity of an electron
at the same time.
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).
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