Download Chem12-Unit 1-1-Development of Atomic Theory

1
No.
Atomic Theory / Model
Development of Atomic Theory
Phenomena that
Diagram illustrating the
Can’t Be Explained by Old
Atomic Theory / Model
Theory / Model
1
(i) 400 B.C.
Democritus
(A Greek
Philosopher
(ii) 1807 –
John
Dalton
Atomic
Theory
Solid-Particle Model (1805) Democritus coined the word “Atom”.
Dalton’s Atomic Theory:
1. Atoms are the basic particles of matter.
2. They are indivisible and cannot be created
or destroyed.
3. Atoms of a given element are identical,
having the same mass and the same
chemical properties.
4. Atoms of different elements combine with
one another in simple whole number
ratios to form molecules of compounds.
5. In chemical reactions, atoms are
combined, separated, or re-arranged.
Phenomena or Laws
Explained by Current
Theory / Model
Using deductive logic, the
Greeks hypothesized that
matter cut into smaller and
smaller
pieces
would
eventually reach what they
called the “atom”, the word
coined by Democritus literally
meaning indivisible.
(i). Law of
Conservation
of Matter
(Antoine Lavoisier)
(ii). Law of Constant
Composition
Static electricity was observed.
Electron
was
discovered,
providing the crucial clue that
atom is not indivisible, but have a
sub-structure. (Read page 7 of
this notes for the details of J.J.
Thomson’s and Robert Milliken’s
Experiments and their discovery
of the e/m ratio and the mass of
an electron
(i). Charge-to-Mass
(1799-Joseph Proust)
(iii). Law of Multiple
Proportions
(John Dalton)
Billiard Ball Model
2 1800’s (i).1897-J.J
Thomson:
Experiments
with Cathode
Ray tubes
(ii). 1909Robert
Milliken:
Oil Drop
Experiment
Thomson Atomic Theory (1897):
Plum Pudding Model – Negatively
charged electrons are imbedded in a large
positively charged atom, just like plums
embedded in the famous English dessert,
pudding. Atoms may gain or lose electrons
to form ions in solution. Particular atoms and
ions gain or lose a specific number of
electrons. Electricity is composed of
negatively charged particles. Electrons are a
component of all matter. Electrons have a
specific fixed electric charge.
ratio was found by
experiment:
e / m = −176
. ×108 C / g
(ii). Charge
& Mass of
Electron were
found by experiment:
(a) e = −160
. × 10 −19 C
(b) m = 911
. × 10 −28 g
Raison Cookie Model
3 1910 Ernest
Rutherford:
Gold Foil
Experiment
(See page
172 of this
notes for
details of
this
experiment)
4 1913 Niels Bohr:
Theorized
Atomic
Spectra of
Hydrogen
(Bohr
Atomic
Theory)
5 1927 Erwin
Schrödinger
:
Rutherford Atomic Theory (1911):
Nuclear Model – An atom is made of
an equal number of negatively charged
electrons and positively charged
protons. Electrons move in orbits
around a very dense and positively
charged nucleus. The proton in the
nucleus carries the positive charge. A
very strong nuclear force holds the
positive charge within the nucleus. Most
of the atom is empty space
Planetary / Nuclear Model
Radioactivity
was
first
observed and proton was
discovered: While most of the
+vely charged alpha particles
went through the foil with no
visible effect, a few were
deflected from their path and
some actually bounced back in
the direction from which they
came, suggesting a dense
nucleus with a +ve charge.
Atom is mostly empty
space with an extremely
dense, +vely charged
nucleus. The +ve charge
is due to the proton
located in the nucleus.
(The neutral particle,
neutron, also in the
nucleus was found by
James Chadwick later in
1937)
Solar System Model – The nucleus
According to the Rutherford Model,
an
orbiting
electron
should
continuously emit electromagnetic
radiation, lose energy, and collapse
the atom. But it didn’t. Atomic
Spectra of discrete wavelengths and
frequencies were observed, that
couldn’t be explained by Thomson’s
Atomic Theory. In 1885, J. Balmer
found an empirical mathematical
relationship between the wavelengths of
the lines in the visible region of the
spectrum. Johannes Rydberg extended
Electrons do not radiate
energy as they orbit the
nucleus.
Each
orbit
corresponds to a state of
constant energy based on
Planck’s
Theory
of
quantized energy. The
Model was able to derive
Rydberg’s
constant
in
calculating
the
wave
numbers of the spectral lines
and to successfully predict
the infrared and ultraviolet
spectra for hydrogen. It
offered
a
reasonable
explanation for Mendeleev’s
Periodic Law.
lies in the center of the atom and the
electrons move in certain “allowed” orbits
around it. These circular orbits represent
the quantized energy levels of the
electrons. The orbit closest to the nucleus
represents the ground state of the energy
level designated by a quantum number n
=1. As the electron absorbs a quantum of
energy, it jumps to a higher level, excited
state. When the electron jumps back down
to the n =1 level, it is accompanied by an
emission of radiation. There are a
maximum number of electrons allowed in
each orbit.
Wave-Particle Model–An electron is
regarded as a wave that has quantized its
energy. Electrons have no precise orbits.
Instead, their motion can only be described
by the probability of finding them in certain
regions surrounding the nucleus. These
regions are called orbitals. These orbitals
form an electron cloud around the nucleus.
The cloud is most dense where the
probability of finding the electron is highest:
electron density. According to the
Uncertainty Principle, the position and the
momentum of an electron cannot both be
known exactly at the same time.
Solar System/Orbit Model
his Equation so that all of the wavelengths
could be predicted. It couldn't explain
why the protons in the nucleus don't fly
apart and it couldn't totally account for
the total mass of the nucleus.
In Bohr’s model, there is no way
to observe or to measure the
orbit of an electron in an atom.
The Bohr’s model was unable to
explain completely
(i) the Zeeman Effect
(ii) the spectral details of the
atoms that have several
electrons.
Electron Cloud / Quantum
Mechanical Model
(i) Heisenberg’s
Uncertainty
Principle (1927):
( ∆x )( ∆mv) ≥
h
4π
(ii) In 1924 Louis De
Broglie suggested
electron could
behave as a wave
as well as a particle:
hf = mc 2
2
3
Figure 2 The solid lines show the intensity of the colours of light emitted
Figure 3 Scientists used to think that as the intensity or
by a red-hot wire and a white-hot wire. Notice how the curve becomes
brightness of light changes, the total energy increases
higher and shifts toward the higher-energy UV as the temperature increases. continuously, like going up the slope of a smooth hill.
The dotted line represents the predicted curve for a white-hot object,
As a consequence of Planck’s work, Einstein suggested
according to the existing classical theory before Planck.
that the slope is actually a staircase with tiny steps,
where each step is a quantum of energy.
4
Figure 6 In the photoelectric effect, light shining on a metal liberates electrons from the metal surface. The ammeter (A) records
the electric current (the number of electrons per second) in the circuit.
5
Figure 2 The original Bohr orbit for n = 2 is circular (red line). Sommerfeld’s revision is three, slightly elliptical orbits for n = 2 (blue
lines). [The secondary quantum number relates primarily to the shape of the electron orbit. The number of values for l equals the
volume of the principal quantum number.]