Transcript Atomic Emission Spectra – Copy

Let’s Practice
1.Calculate the speed of a wave whose
wavelength is 0.25m and whose frequency
is 760 Hz.
2.Calculate the frequency of a blue light
wave whose wavelength is 5.22 x 10-7m.
3. Calculate the energy of the blue light in
#2.
1.Calculate the speed of a wave whose wavelength is 0.25m and whose
frequency is 760 Hz.
C = λν = (0.25m)(760Hz) = 19 m/s
2.Calculate the frequency of a blue light wave whose wavelength is 5.22 x 107m.
Ν = c/λ = (3.00 x 10-8m/s)/ 5.22 x 10-7m
= 5.75 x 1014Hz
3. Calculate the energy of the blue light in #2.
E = hν = (6.626 x 10-34Js)(5.75 x 1014Hz)
= 3.81 x 1019J
GLOW IN THE DARK STARS
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•
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•
laser light - 670 nm wavelength
stars glow? _______
flashlight - 500 nm wavelength
stars glow? _______
UV light - 254 nm wavelength
stars glow? _______
glowing star light - 520 nm wavelength
• On a separate piece of paper, do the following:
A. Convert wavelength in nm to meters.
B. Using Planck’s energy formula and the speed of light formula,
determine the energy for each.
C. Answer the questions:
1. Would infrared light cause the stars to phosphoresce?
2. Would microwave light cause the star to phosphoresce?
3. Give another example of minimum energy photons.
ATOMIC EMISSION SPECTRA
AND QUANTUM THEORY
Atomic Emission Spectra
• We talked before about flame tests and
why the elements produce a color when
heated. These elements produce a
continuous emission of light.
• The electrons in the atoms absorb energy
and become excited. Excited and
unstable electrons then drop back to their
stable level, releasing energy by emitting
light.
Atomic Emission Spectra
• A neon light works the same way.
• We can turn a continuous emission spectra into
a discontinuous one by refracting the light.
• An Atomic Emission Spectrum is a set of
frequencies of electromagnetic waves emitted by
atoms of the element. They are usually distinct
color lines. Each element’s atomic emission
spectrum is unique, can be used to identify the
element, and can be used to determine if the
element is part of an unknown compound.
Atomic Emission Spectra Hydrogen
• Page 144 shows the emission spectrum of
hydrogen’s one electron. Notice that it is
discontinuous – it is made up of only
certain frequencies of light – Figure (b).
Spectra of Four Elements
1. Hydrogen: which emits 4 colors of light that's in the
visible light range. Note that other frequencies, such as
UV light might be emitted, but we can't see them.
2. Helium: It has 2 electrons and we see 7 colors.
3. Mercury: spectra shows 8 colors. Mercury also
produces a lot of UV light which in fluorescent bulbs is
normally converted to visible light by the use of certain
minerals that capture UV light and emit visible light.
4. Uranium: Uranium emits many frequencies of colors. It
appears that the elements that have more electrons
emit more colors. So there seems to be a connection.
Atomic Emission Spectra
Atomic Emission Spectra
• Acts as a fingerprint for elements.
• Is the same everywhere in the universe for
a particular element.
• Allows us to determine the elements in a
star far away.
QUANTUM THEORY
• 1913: Neils Bohr comes up with the quantum model of the
hydrogen atom. He also correctly predicts the frequencies of
the spectral lines in the hydrogen atomic emission spectrum.
• He theorized that the quantized energy that Max Planck
suggested, Einstein proposed, and Rydberg calculated could
be the reason that certain frequencies of light were seen
coming from hydrogen. He thought hydrogen’s electron could
only occupy certain energy levels within the atom. Light that
equaled the difference in the levels could cause the electron
to jump to the higher level - this is called the EXCITED
STATE. Go to the website
(http://www.chemistryland.com/CHM130W/10ModernAtom/Spectra/ModernAtom.html) to see this. Go to the topic
“Bohr’s Atom” about halfway down.
•
When the electron fell back into the original orbit that same
frequency of light would be emitted. This would appear as a
single line on the atomic emission spectra.
BOHR’S THEORY
• Electrons were now viewed to be in orbit around the
nucleus. The electrons could only orbit at certain
distances which represent distinct energy (quantum)
levels. These energy levels were labeled n=1, n=2, and
so forth. They were called the principle quantum
numbers.
• The smaller the orbit, the lower the atom’s energy state
or energy level. Each orbit is assigned a quantum
number, n.
• Bohr said only certain atomic energies are possible for
each atom’s electron and so only certain frequencies of
electromagnetic radiation can be emitted.
• See page 147, figure 5-11.
Hydrogen Atom Emissions
Balmer Series –
visible light
Lyman Series –
ultraviolet light
Paschen Series –
Infrared light
Bracket Series –
Infrared light
QUANTUM THEORY HISTORY
• The spectra of elements showed that light
waves also behaved like particles. Who
was brave enough to ask, "If waves could
behave like particles, can particles behave
like waves?"
Quantum History - 1928
Louis de Broglie stated that Bohr’s quantized electron
orbits had characteristics similar to waves. Particles
could have wavelike behaviors. The energy had
wavelike characteristics. (You do not need to know the
equation.)
How can something be a
particle and a wave?
• Go to the website: (http://www.chemistryland.com/CHM130W/10ModernAtom/Spectra/ModernAtom.html) to see this. Go to the topic
“Electron: Wave and Particle” about halfway down.
• This wave motion would be n=1. The electron starts out
as a particle orbiting but also defining a wave-like
oscillation. Next, the oscillation is shown without the
electron particle. Finally the wave is changed into an
electron cloud representing the probability of the
electron’s position. In textbooks you might see the
electron represented by a particle in orbit or as a cloud.
That's because of electron's dual nature of being both
wave and particle.
Quantum History
• Werner Heisenberg came along and concluded that it is
impossible to make any measurement on an object
without disturbing the object. It is fundamentally
impossible to know precisely both the velocity and
position of a particle at the same time.
• In other words, the light used to measure the particle
changes it.
• Now watch the Captain Quantum video. You can find it
on my website on the assignments page.
Quantum History
• Erwin Schödinger, in 1926, developed an atomic model
in which the electrons are treated as waves. This model
is called the quantum mechanical or wave mechanical
model.
Quantum Mechanical Model
• This model does not describe the path of the electron
around the nucleus.
• The three dimensional region around the nucleus called
an atomic orbital describes the electron’s probable
location.
• It is a fuzzy cloud in which the density of the cloud at a
given point is proportional to the probability of finding the
electron at that point.
• The electron cloud has no definite boundary – it is
arbitrarily drawn at 90%. See page 152, figure 5-15.
Quantum Mechanical Model
Quantum Mechanical Model
1. The atom has a dense, centrally located,
positively charged nucleus full of protons
and neutrons surrounded by mostly
empty space where the electrons are.
2. The energy of electrons is quantized.
3. Electrons exhibit both wave and particle
behaviors.
Quantum Mechanical Model
4. The absolute location of an electron is
impossible to determine – its location and
velocity cannot be determined at the
same time.
5. The electrons travel in orbitals that have
characteristic sizes, shapes, and energies,
but do not describe how the electrons
move.
DO NOW
• Pick up two handouts
• Turn in corrected paperwork – if you .
Make sure you have the old paperwork,
the corrected pages, and the yellow sheet
all paper clipped together.
• Get out problem set homework and Atomic
Emission Spectra notes.
DO NOW
• Pick up two handouts
• Turn in “Glow in the Dark Stars”
homework.
• Get out problem set homework and Atomic
Emission Spectra notes.
Quantum Mechanical Model
So how do electrons position themselves
outside the nucleus?
• As the progression of elements were built by adding one
proton and one electron at a time, the position of the
protons was always in the center of the atom in the
nucleus.
• However, electrons repelled each other, so as each
electron got added for each new element, they would
find a position and shape that maximized their distance
from each other. Amazingly, the way they positioned
themselves followed a fairly basic pattern.
Quantum Mechanical Model
• Quantum Numbers are used to describe
the most probable location of an electron.
• The electrons fill the orbitals in a pattern.
• No matter what element, the pattern is the
same.
Quantum Numbers
• There are four Quantum numbers.
• They each tell something different about
the electron’s location
Principle quantum number, n
- Is the energy level number.
- Gives information about the relative size
and energy of the atomic orbitals.
- It can have values of 1, 2, 3, 4….
- The greatest number of electrons possible
in any one level is 2n2.
- Example: The maximum number of electrons
that can occupy the first level is 2(1)2 = 2; the
fourth level is 2(4)2 = 32.
Angular Momentum Quantum Number, l
• Is the energy sublevel number.
• It gives information as to the shape of the
orbitals.
• The first four levels are s, p, d, and f.
• Example: The first energy level has only an s
sublevel. The second energy level has an s and
p sublevel. This third energy level has s, p, and
d sublevel.
Angular Momentum Quantum Number, l
Angular Momentum Quantum Number, l
- The quality of the spectroscopic lines were
labeled sharp, principle, diffuse, and
fundamental.
- It was believed that the different orbitals were
responsible for the quality of lines; for example,
orbitals that created “sharp” lines were given the
name “s”. “p” orbitals made the principle lines,
etc.
- This turned out not to be true, but the names
stuck.
Magnetic Quantum Number, m
• Gives information about the orientation in
space of an orbital.
• The s sublevel has one orbital, the p
sublevel has three orbitals, the d sublevel
has 5 orbitals, and the f sublevel has 7
orbitals.
• This number determines which p, d, or f
orbital the electrons are in.
Electron Spin Quantum number, s
• Indicates the direction of the electron spin.
• Spin is either clockwise or
counterclockwise.
• Is designated with a +1/2 or a –1/2 .
Quantum Numbers
For example, what are the four quantum
numbers for the last electron in an oxygen
atom?
The last electron is located in the 2p orbital
with a down spin.
n=2 l=1
m = -1
s = -1/2
What’s Next?
• Thankfully, we will not learn how to write
the quantum numbers for each electron in
an atom.
• We will learn electron configuration and
orbital diagrams to depict the probable
location of electrons.
• Stay tuned…..