Transcript document

History of the Atomic Model
Once upon a time…
Some info from http://mooni.fccj.org/~ethall/period/period.htm
Dmitri Mendeleev organizes our elements
In 1869, a Russian scientist named Dmitri Mendeleev,
drew up our first periodic table of elements arranged
according to increasing
atomic masses. (this was before
e
knowing of protons, neutrons and electrons, he just
intuitively believed that chemical properties were related
to their masses)
The table revealed a periodic relationship among the
elements’ properties. This means there were repeating
patterns of related chemical properties.
He also left holes in it where he felt new elements would
be discovered later on.
Mendeleev’s Original 1869 Table
Mendeleev’s table
as published in
1869, with many
gaps and
uncertainties.
A later 1898 Version
The table, rotated
ninety degrees, as
shown in textbooks
in 1898 when Marie
Curie discovered
Radium. Mendeleev
promptly added Ra
to his table also.
Henry Mosely fine-tunes the periodic table
In 1913, Henry Mosely (an Englishmen)
discovered a method to calculate the atomic
number (the positive charge of the nucleus)
through the bombardment of electrons
creating a unique X-ray spectrum of every
atom. This process was called X-ray
diffraction.
e
In 1914, Henry Mosley arranged the periodic
table today according to an element’s atomic
number…and we still do it this way today!
This method lead to better periodicity and
explained why argon, with an atomic mass of
39.95 amu (atomic number of 18), can come
before potassium with the atomic mass of
39.10 amu (atomic number of 19).
Democrotis with first mention of atoms
(460B.C. - 370B.C.)
The greek philospher Democrotis was among the first to
suggest the existence of atoms…a discontinuous theory
of matter.
The word atom comes from “atomos” meaning
indivisible or indestructible.
Atoms were thought of as the building blocks (the
Legos) of everything.
Aristotle (another greek philospher around 350BC)
believed in a thought that all matter was made of only 4
elements: earth, wind(air), fire and water.
Aristotle
John Dalton’s Solid Sphere Model
In 1803, John Dalton was the first to show laboratory evidence of
the existence of atoms.
Dalton’s model of the atom was a solid sphere with no subatomic
particles…the atom was indivisible!
Dalton also proposed the first atomic theory which stated that
1. all things are made up of atoms and
2. all atoms of the same element are identical (later determined
this not to be true…isotopes)
3. these atoms can combine in simple whole number ratios to
form compounds.
Thomson’s Discovery of the Electron,
shatters Solid Sphere Model
He used a cathode ray tube to discover the
presence of electrons…which were even smaller
than atoms!
JJ Thomson’s Plum Pudding Model
In 1898 JJ Thomson proposed
The Plum Pudding Model for
the structure of an atom,
shattering the solid sphere
model.
The model pictured a sphere of
positive electricity (“pudding”)
studded with negatively charged
electrons (the “plums”).
Later this model proved to be
wrong by his student E.
Rutherford.
Earnest Rutherford’s Gold Foil Experiment
His expected results.
The actual results…Shocking!
He shot alpha particles (+ charged) through the atoms of gold in a
thin sheet of foil, and studied how they were scattered by the
atoms. Oddly, some of the particles “bounced” back off the foil but
most when straight through.
The circles are atoms of gold, the dots are the nucleus and the
arrows are the alpha particles (positively charged electromagnetic
radiation).
Rutherford Nuclear Atom
Rutherford (a student of
Thomson) set out to prove
Thomson’s model correct,
however ended up proving it
wrong.
By 1911, Rutherford had
theorized a new model with a
positively charged dense
central nucleus surround by
a cloud of negative electrons.
Although very, very tiny the
nucleus contains 99.9% of an
atoms mass surrounded
mostly by empty space.
Rutherford’s Model
Millikan’s Oil Drop Experiment
When he sprayed oil droplets into a chamber and
bombarded them with X-rays to place a negative charge
on them, the charged droplets were attracted to the
positive plate. Changing the strength of the electrical
field offset the attraction and allowed Millikan to
determine the charge of the electrons.
He measured the charge of the electrons (1.60 X 10-19
coulombs) and later used this to calculate their
extremely small mass (9.11 x 10-28 grams).
Bohr Model of the Atom
The Bohr model has protons and
neutrons in the nucleus surrounded
by orbitals with definite electron
paths, kind of like a planetary
model.
Atoms in their normal, energetically
stable state are said to be in their
ground state.
The greater the distance of the
orbitals from the nucleus, the
greater is the energy of the
electrons in that shell (because
they are overcoming the
electrostatic attraction towards the
nucleus).
Bohr Planetary Model
Neils Bohr
Neils Bohr studied light emitted from
excited (electrically stimulated) gas
atoms called atomic emission
spectrums (bright-line spectrums). It
turns out that every element emits a
unique spectrum (combination) of
wavelengths when excited. (or brightThis emission spectrum can be used as a type
fingerprint for determining
lineofspectrum).
what elements are present. Atomic emission spectrums are used today to
determine what elements are present in far away stars.
He determined that this must mean that each electron must jump exact,
discrete amounts within set orbitals within an atom. This contributed to his
discovery that there are quantum energy levels (or orbital locations) for
electrons. Therefore, an electron in a given orbital has a definite amount of
energy.
Mechanism of Bright-line Sprectra
1st Step atoms absorb
energy….Radiant energy, heat
energy, or electrical energy
absorbed by the atom causes
the electron to move (“jump”)
from its lower-energy orbit
ground state to a higher-energy
orbit (an excited state).
2nd Step Immediately after, the
atoms release energy… radiant
energy (or photons) are emitted
in the form of visible light or Xray or UV rays when the electron
moves from a higher-energy
orbit back to its lower-energy
orbit.
1st
2nd
This lead to Quantum Theory of
Atomic Structure
Atoms and molecules could emit (or absorb) energy
only in discrete quantities called “quanta”, like small
packages or bundles, there were not unlimited energy
contents for particles.
Observations of electron behavior in the 1920’s were
inconsistent with the planetary model because electrons
moved like a particle and yet also like a wave.
Researches such as Albert Einstein (1905), Louis de
Broglie (1924) and Erwin Schrodinger (1926) devised a
mathematical equation to describe the wave-particle
behaviors of electrons.
Quantum-mechanical Model
of the Atom
It is also known as the Shrodinger Model or ElectronCloud Model.
Like the Bohr model, this model restricts the energy of
electrons to certain levels. Unlike the Bohr model, The
Q-M model does not show the definite paths of
electrons; it shows the most probable location of an
electron as a cloud (region of space)
An electron’s location at any time can not be pinpointed
exactly, yet you have a good idea of where it should be.
Also, each energy level contains subshells called
orbitals (these are the s, p, d and f orbitals).
Quantum mechanical Model
It is very hard to draw. Here are 2 attempts.
Photoelectric Effect
This is the theory that light radiation is both a particle
and a wave (Albert Einstein worked on this).
In other words it is made of photons (tiny bundles of
energy) that behave like waves in their movement.
A high-intensity, low-frequency light will not emit
electrons from a metal, but a low-intensity, highfrequency beam will. Depends on the frequency of the
light not its intensity.
Electronic eye in a store is an example of this effect.
Einstein’s main conclusion was that all matter exhibits
wave properties.
Wave-Particle Duality Cartoon
Wave-Particle Dualtity
Wave/particle duality is the possession by physical entities (such as light and
electrons) of both wavelike and particle-like characteristics. On the basis of
experimental evidence, the German physicist Albert Einstein first showed
(1905) that light, which had been considered a form of electromagnetic
waves, must also be thought of as particle-like, or localized in packets of
discrete energy (see the photoelectric effect).
The French physicist Louis de Broglie proposed (1924) that electrons and
other discrete bits of matter, which until then had been conceived only as
material particles, also have wave properties such as wavelength and
frequency. Later (1927) the wave nature of electrons was experimentally
established. An understanding of the complementary relation between the
wave aspects and the particle aspects of the same phenomenon was
announced in 1928.
Max Planck’s Eqtn  E= h· f
E is the energy emitted or absorbed
h is Planck’s constant 6.63 x 10-34 J-s or J/Hz
f is the frequency of radiation in Hz or 1/s
Energy is always emitted in multiples of h f (h f, 2h f, 3
h f ) from an atom.
This is what tells us there are discrete energy levels
within electron orbitals.
Wavelength and frequency of light waves when
multiplied by each other will always equal a constant, c,
which is the speed of light 3.0 x 108 m/s2. So c =
wavelength x frequency.
SPECTRAL LINES
Spectral lines (emission spectra) are a result of
transition of electrons between energy levels.
Their frequency is related to the energy spacing
between levels using Planck’s relationship (E=hf )
The frequencies present give off an atomic fingerprint
called a bright-line spectra.
We can use the color (determined by the frequency) of
light emitted to identify the metal atoms present.
Flame Test
We will be doing a Flame Test Lab to identify the presence
specific metals in our compounds.
We will be adding heat energy to metal ionic compounds. The
metal ions will absorb energy and electron(s) will “jump” up to a
higher energy level. This is called an “excited state”.
This excited state is highly unstable, so instantly (within a
millisecond) the electron(s) will return to the lower energy level
and emit a photon of light energy in return.
Each metal ion will emit a different frequency of light (dif. color)
related to the energy levels that the electron moved within.
Other emission spectra uses
An emission spectrum (or bright-line spectrum)
is a more quantitative way to determine the
identity of atoms. This method looks through a
spectroscope at the specific multiple
wavelengths that are released.
This method can be used to determine the
identities of atoms that are in stars in the sky.