Download Development of the Atomic Model

Development of the
Atomic Model
Remember Rutherford?
• Many scientists in the early 20th century found
Rutherford’s model to be incomplete.
•It did not explain how the electrons
occupied the space around the
nucleus.
•Nor did it explain why the negatively
charged electrons were not pulled into
the atom’s positively charged nucleus.
•Nor did it account for the differences in
chemical behavior among the various
elements.
Light: is it a wave? Or is it
a particle?
•Before 1900, light was thought to be a
wave.
•But then it was found to have some
particle-like character.
•Still, many properties can be
described in terms of waves, and thus
an understanding of the wave nature
of light is needed.
Wave Description of Light
• Electromagnetic Radiation- a form of energy
that travels thru space like a wave
•Examples: visible light, microwaves that
warm and cook your food, X rays that
doctors and dentists use to examine
bones and teeth, and waves that carry
radio and television programs to your
home.
•Moves at a constant speed (3.0 x
108m/s)
Picture of the Electromagnetic
Spectrum
Waves are measured by:
• Wavelength (l)- distance between
corresponding points of adjacent waves
(range from m to nm)
• Frequency (n)- number of waves that
passes a given point in a specific amount
of time
• Units: 1 wave/1 second= Hz (hertz)
•A longer wavelength has a lower
frequency and lower energy.
•A shorter wavelength has a
higher frequency and higher
energy.
•UV light is so dangerous
because it is high in energy.
A Typical Wave
• Include
•Wavelength
•Trough
•Crest
•Amplitude
Amplitude
Wavelength, frequency and the
speed of light are all related…
c=ln
• l and n are inversely proportional
• c is the speed of light (3.0 x 108 m/sec)
Limitations of the Wave Model
•Does not describe important aspects
of light’s interactions with matter.
•Cannot explain why heated
objects emit only certain
frequencies of light at a given
temperature.
•Does not explain why some metals
emit electrons when colored light
of a specific frequency shines on
them.
The Photoelectric Effect
• Photoelectric Effect- emission of electron
from a metal when light shines on the metal.
• This stream of electrons creates an electric
current.
• For a given metal, no electrons will be
emitted if the light’s frequency is below a
certain minimum—no matter how long the
light was shined on it.
Max Planck
•Suggested that hot objects do not
emit electromagnetic energy (light)
continuously (as a wave would) but
in small, specific amounts he called
quanta.
•Quanta: minimum amount of
energy that can be gained/lost by
an atom
Planck summed up the relationship
between quantum of energy (E) and
frequency (n) of radiation in an
equation…
E= hn
Energy= Planck’s constant  frequency of
radiation
(Joules)
(6.626x10-34Js)
Hz
Practice Problems
•What is the energy of a quantum
of light with a frequency of 4.31 x
1014 Hz?
•A certain violet light has a
wavelength of 413 nm. What is
the frequency of the light?
Albert Einstein
• Einstein expanded on Planck’s theory by
introducing wave-particle duality of light.
• That is, while light has many wave like
characteristics, it can also be thought of as a
stream of particles or bundles of energy.
(each of which carries a quantum of
energy).
• Einstein called these particles photons.
• Photon: a particle of electromagnetic
radiation having no mass and carrying a
quantum of energy.
• The energy of a particular photon depends on
the frequency of the radiation.
• Einstein explained the photoelectric effect by
proposing that electromagnetic radiation is
absorbed by matter only in whole numbers of
photons.
• In order for an electron to be emitted from a
metal surface, the electron must be struck by
a single photon possessing at least the
minimum energy needed to knock the
electron loose.
•According to E=hn, minimum energy
corresponds to frequency; so, if the
photon’s frequency is below
minimum, the electron won’t be
released.
•Electrons of different metals require
different minimum frequencies
because they are bound more or
less tightly.
So what does this have to
do with atoms?
Have you ever wondered how light is
produced in the glowing tubes of neon
signs?
• The light of the neon sign is produced
by passing electricity through a tube
filled with neon gas.
• Neon atoms in the tube become
excited.
• These excited and unstable atoms
then release energy by emitting light.
Line-Emission Spectrum
•Ground state- lowest energy state of an
atom.
•Excited state- a state in which an atom
has a higher potential energy than in its
ground state.
•When an excited atom returns to its
ground state, the energy is given off in
the form of electromagnetic radiation.
•This is the light we see.
•When that light is passed thru a prism,
it separates into a series of specific
frequencies of visible light called the
line-emission spectrum.
Hydrogen
Helium
Carbon
•Scientist had expected to see
a continuous range of
frequency…so why only
specific intervals???
•This lead to a new theory of
the atom…
Quantum Theory
•In order for an atom to fall from
its excited state back to its
ground state, it must emit the
energy as a photon.
•That energy is equal to the
difference in energy between the
atom’s initial state and its final
state.
Bohr Model
•Linked the electron in his hydrogen
atom with its photon emission.
•The electron can only circle the
nucleus in a certain path (called an
orbit) which has a fixed amount of
energy.
•When the electron is in that orbit,
the atom has a definite, fixed
energy.
•The electron stays in lowest energy
orbit (close to the nucleus) unless it
gains enough energy to move to the
next orbit; then it has to release
the energy as a photon to return to
lower energy orbit.
•This lowest energy orbit is
separated from the nucleus by a
large empty space where the
electron cannot exist.
•This is comparable to being on a
ladder, you must be on one of the
rungs, not in between because you
can’t stand in midair.
•In the same way, an electron
can be in one orbit or another,
but not in between.
•But how does this explain the
observed spectral lines?????
•While in orbit, the electron can
neither gain nor lose energy. It
can however, move to a higher
energy orbit by gaining an
amount of energy equal to the
difference in energy between the
higher energy orbit and the
initial lower energy orbit.
•When an atom is in an excited state,
its electron is in a higher energy orbit.
•When the atom falls back from the
excited state, the electron drops to a
lower energy orbit.
•A photon is emitted equal in energy to
the energy difference between the
levels.
•The energy of each emitted photon
corresponds to a particular frequency
of emitted radiation and therefore a
particular color.
Limits of Bohr’s model:
•Didn’t explain elements with more
than one electron nor the chemical
behavior of atoms.
•To most scientists of the time, the
Bohr model contradicted common
sense. They didn’t see why
electrons couldn’t exist in a limitless
number of orbitals with slightly
different energies.
•Why were electrons allowed
only in certain orbitals with
definite energy?
•In order to understand the
answer, scientists had to
change their view of the
electron.
Electron Configuration
the arrangement of electrons
in an atom
•The quantum model of the atom
improves on the Bohr model because it
describes the arrangements of electrons
in atoms other than hydrogen.
•Because atoms of different elements
have different numbers of electrons, a
distinct electron configuration exists
for the atoms of each element.
But why do we care???
•We care because the behavior
of electrons is what Chemistry
is all about.
•All the interactions of
elements and characteristic
properties of elements deal
with the electrons.
Electrons line up around the nucleus
according to 3 basic rules:
•Aufbau Principle
•Pauli Exclusion Principle
•Hund’s Rule
Aufbau principal
•Electrons occupy the lowest-energy
orbital that can receive it.
•Orbital with lowest energy is 1s.
•Sometimes sublevels of one energy
level might be higher than sublevels of
the next energy level.
•(Beginning with the 3rd energy level,
energies of the sublevels in different
main energy levels begin to overlap.)
•Represent orbitals with ____
Pauli Exclusion Principal
•No two electrons in the same atom
can have the same set of 4 quantum
numbers.
•What that means is: in order for
two negative electrons to occupy the
same orbital, they must have
opposite spins.
•Written like:
_____
Hund’s Rule
•Orbitals of equal energy are each
occupied by one electron before any
orbital is occupied by a second
electron, and all electrons in singly
occupied orbitals must have the same
spin.
•___ ___ ___
Electron configuration notation
•Eliminates the lines and arrows of
orbital notation. Instead, the
number of electrons in a sublevel is
shown as a superscript of the
sublevel letter
•Hydrogen’s orbital notation:_____
1 s
•Hydrogen’s configuration: 1s1
Let’s put it all together…
•Take a look at the periodic
table.
•Give the electron configuration
and orbital (arrow) diagram for
the following elements.
•Na, N, Y, Br, Xe, Ca, Mn
Nobel Gas Notation
•Shortcut for electron configuration.
•Each noble gas ends with the period
number and 6 electrons.
•Use { } and the closest noble gas.
•Examples: Ga, Zr, Cl
•Nobel gas configuration shows an
outer main energy level fully occupied
by 8 electrons.