Transcript Electrons in Atoms

Electrons in Atoms
Chemistry
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Electrons in Atoms Objectives
1. Compare the wave and particle models of light.
2. Define a quantum of energy and explain how it is related to an
energy change in matter.
3. Contrast continuous electromagnetic spectra and atomic
emission spectra.
4. Compare the Bohr and quantum mechanical models of the
atom.
5. Explain the impact of de Broglie’s wave particle duality and the
Heisenburg uncertainty principle on the modern view of
electrons in atoms.
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Light and Quantized Energy
Even though Rutherford’s nuclear model was a major
discovery, it did not explain how electrons actually
behaved.
-did not:
a. state how electrons occupy the space around the
nucleus
b. account for the differences in chemical behavior
of the various elements
This chapter describes how electrons are arranged and
how it plays a role in chemical behavior.
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Light and Quantized Energy
By the early 1900’s, the puzzle of chemical behavior
was beginning to be understood.
-observed certain elements emit a visible light when
heated
To understand this relationship, you must understand
the way light behaves.
electromagnetic radiation: form of energy that
exhibits wavelike behavior as it travels through space
-visible light, microwaves, x rays, radio waves,
are a few examples
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Light and Quantized Energy
Waves are characterized by:
-wavelength (l): shortest distance between equivalent
points on a continuous wave, usually crest to crest or
trough to trough
♦measured in m, cm, or nm
-frequency (n): number of waves that pass a given
point per second
♦measured in hertz (Hz), 1/s or s-1
-amplitude: waves height from the origin to a crest or
from origin to the trough
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-speed of light (c): all electromagnetic waves travel at
a speed of 3.0 x 108 m/s in a vacuum
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Light and Quantized Energy
Speed is represented by:
- c=ln
♦ wavelength & frequency are inversely proportional
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Light is actually a continuous range of wavelengths and
frequencies.
-can be separated into colors by a prism, which is
called the visible spectrum, though the visible
spectrum makes up a tiny portion of the complete
electromagnetic spectrum
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Pictures came from http://www2.chemistry.msu.edu/faculty/reusch/VirtTxtJml/Spectrpy/UV-Vis/spectrum.htm
electromagnetic spectrum: includes all forms of
electromagnetic radiation, with the only differences in
the types of radiation being their wavelengths and
frequencies.
-p 120, Fig 5-5
Pictures came from http://www.lbl.gov/MicroWorlds/ALSTool/EMSpec/EMSpec2.html
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Particle Nature of Light
Though light behaves as a wave, it does not
adequately describe important aspects of light’s
interactions with matter.
-why heated objects emit only certain
frequencies at given temperatures
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Max Planck began searching for an explanation for
these phenomena.
-he concluded that matter can gain or lose energy
only in small, specific amounts called quanta, the
minimum amount of energy that can be lost or gained
by an atom
-proposed the emitted light was quantized and
showed a mathematical relationship between quanta
and frequency
Equantum = hn
where h = Planck’s constant = 6.626 x 10-34 J
and Equantum and n are proportional
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Photoelectric Effect
the wave model of light could not explain a phenomenon
called 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.
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While a beam of light has many wavelike
characteristics, it also can be thought of as a stream
of tiny particles, or bundles of energy, called photons
-photon: a particle of electromagnetic radiation with
no mass that carries a quantum of energy.
Extending Planck’s idea of quantized energy, Einstein
calculated that a photon’s energy depends on its
frequency.
For the photoelectric effect to occur, a photon must
possess, at a minimum, the energy required to free
an electron from an atom of the metal.
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Atomic Emission Spectra
atomic emission spectrum: set of frequencies of the
electromagnetic waves emitted by atoms of that
element
-not a continuous range, but several distinct lines of
color
-unique to each element (fingerprint)
-can be used to identify that element
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Quantum Theory of the Atom
Neils Bohr proposed a quantum model to explain why
hydrogen’s atomic emission spectrum was
discontinuous.
-hydrogen has only certain allowable energy states
♦ground state: lowest allowable energy level
♦excited state: when an atom gains energy and
goes into a higher energy state than the ground
state
-the larger the orbital, the higher the energy state
♦he assigned a quantum number, n, to each orbit
(n = 1 is the ground state, or first orbit)
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♦there are 7 orbits, or energy levels
Electrons can only circle the nucleus only in allowed
paths, called orbits, though they can jump from one
orbit to another, if
1) space allows and
2) there is enough energy for the electron to make
the jump
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Bohr stated that in the ground state, an atom does not
radiate energy, though it will in a higher energy state.
-as an excited atom falls back to its ground state
(lower energy state), it emits a photon of radiation
Ephoton = E2 - E1 = hn
where E2 = higher energy state and
E1 = lower energy state
-these photons emitted occur only at certain
frequencies, giving off distinct colors
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This shows the relative
energies of the electron
transitions responsible
for hydrogen’s four
spectral lines.
-called the Balmer series
(visible series)
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Though Bohr’s model explained hydrogen’s spectral
lines well, it did not
-explain the spectrum of any other element
-fully account for the chemical behavior of
atoms
Why was his theory so important?
It laid the groundwork for later atomic models
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Quantum Mechanical Model
Scientists (mid 1920’s) were convinced that Bohr’s
model was incorrect.
-Louis de Broglie proposed an idea to explain the
fixed energy levels
♦Bohr’s quantized orbits have wave-like
characteristics
♦light has both wave and particle characteristics:
‘If waves have particle-like behavior, can particles of
matter (including electrons) behave like waves?’
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de Broglie knew that if an electron has wave-like motion
and is restricted to circular orbits, the electron is
allowed only certain wavelengths, frequencies and
energies
-developed a wave equation, the de Broglie equation,
which predicts all moving particles have wave
characteristics
l= h
mv
where v = velocity (speed)
and m = mass
We now have begun making great strides in
understanding the atom.
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Heisenburg Uncertainty Principle
Werner Heisenburg used the work of Rutherford, Bohr
and de Broglie to make his own conclusion, which had
profound implications for the atomic model.
-Heisenburg Uncertainty Principle: states it is
fundamentally impossible to know precisely both the
velocity and position of a particle at the same time.
-even though scientists at the time had trouble
understanding this concept, it was able to describe the
fundamental limitations on what can be observed
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Schrodinger Wave Equation
Erwin Schrodinger used de Broglie’s ideas and the
Heisenburg Uncertainty Princple to further the waveparticle theory
-derived an equation treating hydrogen’s electron as a
wave
-worked with all atoms, not just hydrogen’s
-electrons do not travel around the nucleus in neat
orbits, but exist in regions called atomic orbitals, a
three dimensional region around the nucleus that
indicates the probable location of an electron
We now have a working quantum mechanical model
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of the atom.
Quantum Mechanical Model
Did: limits an electron’s energy to certain
values
Did Not: makes no attempt to describe the
electron’s path around the nucleus.
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Intro to Spectroscopy Lab
A white light source can be separated into the individual
wavelengths in the visible region of the spectrum (400
nm to 650 nm.) This separation of the different
wavelengths of visible light can be done with a
diffraction grating (which works like a prism.) The
diffraction grating can be adjusted so that only the
energy of a specific wavelength of light is passed
through a sample, and a detector can then determine
how much of that energy was absorbed by the
sample.
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Intro to Spectroscopy Lab
This is how the Spectronic 20 works. Refer to the
diagram below to follow the path of the light through
the Spectronic 20. The diffraction grating used to
separate the wavelengths is turned when you turn the
knob on top of the instrument. In doing so, you
determine the specific wavelength of light that passes
through your sample, and the Spectronic 20 gives you
a readout of the light that was absorbed by your
sample.
detector
cuvette
diffraction
grating
light
source
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Intro to Spectroscopy Lab
Purpose
This lab will familiarize you with the operation of the
Spectronic 20. At the conclusion of the lab, you
should understand the function and know the names
of each of the controls on the spectrophotometer.
You should also be able to set the spectrophotometer
to a given wavelength and zero it using a blank.
Hypothesis
I will demonstrate the proper way to use a
spectrophotometer.
Materials/Equipment
spectrophotometer
chalk distilled water
2 cuvettes
Kimwipes
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Terminology:
1. A cuvette looks like a test tube. It fits inside the
spectrophotometer to hold your sample. It is different
from a test tube because the glass it is made of is more
uniform in width than a test tube. This is important
because the light must pass through both walls of the
cuvette before it reaches the detector, and the thickness
of the glass should not affect the readings you get from
the spectrophotometer.
2. Kimwipes are small tissues like Kleenex. They are used
to wipe fingerprints and dirt off the cuvette before
inserting it into the spectrophotometer.
Safety Considerations:
1. Always wear goggles and an apron in the lab.
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2. Please be careful with the cuvettes; they are expensive.