Transcript File - Mrs. Hille`s FunZone

The Development
of Atomic Models:
A Historical Perspective
Model of an Atom
An IDEA of what it looks like
(a working representation)
Atomic Models
Democritus’ Idea
• An object CANNOT
be divided indefinitely.
• There is a smallest
particle.
atom: (Gk. atomos—indivisible)
Atomic Models
Democritus’ Idea
was CORRECT!
• There is a basic unit
of matter—the atom.
• Chemists found this
out by looking at the mass ratio
of substances.
• But the atom is NOT indivisible.
Atomic Models
John Dalton
• Framed the first
experimental model
• Experimented by
weighing substances
• Assigned relative
masses to elements
Atomic Models
John Dalton
• Believed that:
 Atoms of different
elements form
compounds.
 Atoms of different
elements have
different masses.
law of definite
composition
Every compound has a definite
composition by weight.
Example:
Water’s composition by weight
is always an 8:1 ratio.
Atomic Models
Joseph Thomson
• Created a vacuum
in a tube
 Light came from the
“cathode” end.
• Noted that cathode rays are
affected by magnets and
electricity, but not by gravity
Atomic Models
Joseph Thomson
• Decided that the cathode rays
were negatively (−) charged
and calculated their charge-tomass ratio
• Found that every element he
tested emitted these cathode
rays, later called electrons
Atomic Models
Thomson’s Model
• The electrons (e−) are
in a positive jell-like
substance.
• The electrons and the
positive substance cancel out.
• Electrons can be removed from
the atom.
Atomic Models
Thomson’s Model
“Plum-pudding” model
(It looks more like a
chocolate chip cookie.)
Atomic Models
Thomson’s Model
• Shows inaccuracies
with Dalton’s model:
 Atoms aren’t solid.
 Atoms aren’t
indestructible.
Question
What do we now know is part
of the atom, but is missing
from Thomson’s model?
1. Cathode rays
2. Electrons
3. Nucleus
Atomic Models
2 Theories
1. Continuous: Matter can be
subdivided forever.
2. Particulate: A smallest
particle exists.
Atomic Models
Ernest Rutherford
Discovery of the Proton
• 1910 Alpha Particle Experiment:
He aimed alpha particles at a
thin gold foil.
 Alpha particles are relatively
heavy and positively (+)
charged.
Atomic Models
Ernest Rutherford
Exper. Fact
Conclusion
A few alpha
particles
bounced back
completely.
Atoms have a
dense positive
charge—the
proton (p+).
Atomic Models
Ernest Rutherford
Exper. Fact
Conclusion
Most alpha
particles
passed through
the gold foil.
Atoms are
mostly empty
space.
The nucleus is
only 1/100,000th
the size of the
atom.
Atomic Models
Rutherford’s Experiment
alpha particle slightly deflected
zinc sulfide
gold foil
screen
beam of alpha
particles
alpha particle greatly deflected
Atomic Models
Rutherford’s Model
• Protons form a
small, central
nucleus.
• A proton is 1800
times as massive
as an electron.
• Electrons are somewhere away
from the nucleus.
Question
What do we now know is in
the atom, but was missing
from Rutherford’s model?
1.
2.
3.
4.
Neutrons
Protons
Electrons
Nucleus
Rutherford’s Model
This is NOT a planetary model.
atomic number (Z)
the number of protons in the
nucleus of an atom
Atomic Models
James Chadwick
Discovery of the Neutron
• According to Rutherford’s
model, there is not enough
mass with just a proton and
an electron.
Atomic Models
James Chadwick
Discovery of the Neutron
• In 1932, Chadwick discovered
a neutral particle with almost
the same mass as a proton.
• He called it a neutron (N).
Atomic Models
Niels Bohr
Energy Levels for Electrons
• Where are the electrons?
• Are they moving?
• If so, in what manner?
• Why don’t the electrons fall
into the nucleus? Don’t protons
attract them?
Atomic Models
Niels Bohr
Energy Levels for Electrons
• 1913 spectroscopy
experiment
Spectroscopy
Heated atoms absorb
electromagnetic radiation.
They emit energy as light,
which is part of the
spectrum.
Electromagnetic
Spectrum
• It contains different
wavelengths of energy—
electricity, radio, microwave,
visible, cosmic rays.
• As the wavelength decreases,
there is more energy in each
wave.
Electromagnetic
Spectrum
• A continuous spectrum
comes from a source that
emits all energies (colors)
visible to our eyes.
• A line spectrum comes from
a source that emits only
exact (quantized) energy.
Electromagnetic
Spectrum
a continuous spectrum
a bright-line spectrum of hydrogen
Electron Motion
• How do the e− make the line
spectra?
 By only giving off exact
amounts of energy.
• Bohr said they do this by moving
only in restricted steps—from
level to level (orbit to orbit).
• These levels are therefore said
to be quantized.
Question
What causes the line spectra?
1. e− orbiting
2. e− moving between levels
3. p+ orbiting
4. p+ moving between levels
Electron Levels
• Electrons can move from
orbit to orbit.
• A more distant orbit is
higher in energy.
Electron Levels
• So, an electron absorbs
energy, “jumps” to a higher
orbit, and then “falls” back
down, emitting energy.
• The lowest level is called
the ground state.
Electron Orbits
• Since the electron falls back
down from one exact orbit
to another, it releases an
exact amount of energy
(quanta).
• This exact energy packet
produces a line of the line
spectra.
Question
Why do the e− give off line spectra
and not continuous spectra?
1. They move too fast.
2. They are too far from the
nucleus.
3. They move only from one
exact level to another.
4. They can only give off energy
bursts, not continuously.
Bohr’s Model
It is a planetary model.
Bohr’s Model
• Bohr’s orbits (energy levels)
are called principal energy
levels.
• Seven levels can
be measured.
Spectroscopy
• Bunsen and Kirchoff
invented the spectroscope.
• Different elements, when
heated, produce different
sets of wavelengths.
Spectroscopy
• It acts like a fingerprint to
identify elements.
• It gives evidence that
different elements have
different numbers of
electrons in different
locations.
Principal Energy Levels
• There are 7.
• Each can hold more than
1 electron.
• Bohr’s model works well
only for atoms with one
electron.
Level (n)
1
2
3
4
5
6
7
Total # e− (2n2)
2
8
18
32
50* (32)
72* (18)
98* (8)
Where is the Electron?
• In Bohr’s model:
The electron is in an exact
location, orbiting like a planet.
Where is the Electron?
• Problems with Bohr’s model:
 It only works for atoms with
1 electron.
 Electrons are not planets
made up of many atoms;
they are a part of an atom.
 Forces between atoms are
not the same as forces
within atoms.
Where is the Electron?
• Conclusion:
 Electrons are NOT in exact
orbits.
Light
• Has properties of a wave
 Refraction
 Reflection
 Diffraction
Light
• Einstein suggested that
light consisted of massless
particles called photons.
• De Broglie suggested the
opposite: that particles
could act like waves.
de Broglie’s
Hypothesis
• All particles act like waves.
• The bigger the particle, the
smaller the wave.
• The waves are only
significant for very small
particles.
de Broglie’s
Hypothesis
• An exact number (from 1–7)
of de Broglie wavelengths
can fit in a Bohr orbit
around an atom, depending
on the size of the orbit.
Heisenberg Uncertainty
Principle
• We can know where an
electron is or where it is
going; but we cannot know
both at the same time.
Heisenberg Uncertainty
Principle
• When an electron is hit by
photons (light), it
disappears and reappears
somewhere else, but it
does not travel between.
Heisenberg Uncertainty
Principle
• Bohr’s precise, planet-like
orbits were replaced by
orbitals.
Heisenberg
& Electrons
• Electrons are in orbitals
(areas where electrons are
most likely found).
• Orbitals look like fuzzy,
shaded clouds.
• Darker shading indicates a
higher probability of an
electron existing there.