Transcript DEVELOPMENT OF THE ATOMIC MODEL

DEVELOPMENT OF THE
ATOMIC MODEL
From Democritus to Rutherford
c.400 B.C.
Ancient Greek Philosophy
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Everything in the
universe is made of
one or more of the
basic “elements:”
Earth, Fire, Water, Air
What makes up the elements?
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ARISTOTLE
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Matter is infinitely
divisible; no matter
how small a piece is, it
can always be divided
into smaller pieces
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DEMOCRITUS
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There exists a
“smallest piece” of
matter, which cannot
be divided any further.
These pieces are
called “ατομοσ,” or
“atoms”
Whose Argument Prevails?
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Aristotle’s viewpoint enjoyed the support of
most of the world because he was more
well-known and because Democritus had
no evidence to back up his claim since
these “atoms” would be too small to see.
Antoine Lavoisier 1743-1794
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Father of Modern Chemistry
Chemical Revolution – 1770-1790
Stated the Law of Conservation of Mass
Oxygen for combustion
Decomposed Water into two gases, hydrogen
and oxygen and then reformed the exact same
amount of water.
1803: John Dalton
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Dalton was an
English schoolteacher
Began teaching
mathematics and
chemistry at the age
of 12
Revived the idea of
Democritus’ “smallest
piece” of matter
Dalton’s Atomic Theory
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All matter is made of tiny particles called “atoms”
Atoms are indivisible and indestructible
Atoms of the same element are identical
Atoms of different elements differ in some
fundamental way
Atoms combine in simple whole number ratios to
form compounds
Atoms are rearranged in chemical reactions but
cannot be created or destroyed
Three Laws Explained by the Atomic Theory
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Law of Conservation of Mass (Antoine Lavoisier)
Law of Definite Proportions - compounds
always contained the same mass ratio of one
element to another. (Joseph Proust)
Law of Multiple Proportions – When elements
combine in different ratios, each new ratio is a
unique compound. (John Dalton)
So what?
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These results could only be explained by
assuming that matter was made of atoms
– tiny building blocks – and that these
atoms only came in certain sizes.
Dalton’s View of an atom
1897: J.J. Thomson
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English physicist
Worked with
Cathode-Ray Tubes
(CRTs)
Credited with the
discovery of the
electron
CRTs
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Mysterious particles
emanated from the
cathode end
These particles were
deflected by magnetic
and electric fields
They were very small
and negatively
charged
These particles were
called “electrons” and
were assumed to be a
part of all matter
Thomson’s Atomic Model
Electrons
Positively charged “goo”
A.K.A. the “Plum-Pudding Model”
1909-11: Robert Millikan
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Set out to discover
the charge of a single
electron
Famous experiment
called the “oil-drop
experiment”
Using his results and
the charge-to-mass
ratio from Thomson,
the mass of an
electron was found
The Oil Drop Experiment
To view an animation of this experiment click below
OIL DROP EXPERIMENT
1910: Ernest Rutherford
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Expert in radiation
Famous “Gold-Foil
Experiment”
Discovered the
presence of the
nucleus by firing
alpha particles at a
sheet of gold foil
The Gold Foil Experiment
To view an animation of this experiment click below
GOLD FOIL EXPERIMENT
Rutherford’s Atom
Rutherford’s Atomic Model
Electrons
Empty Space
Nucleus
Positively charged
Made of “protons”
Gold Foil Conclusions
The atom has a nucleus
 The nucleus has a positive charge
 The nucleus is very small and very dense
 Most of the atom is empty space
 The electron resides in the region outside
the nucleus
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1932: James Chadwick
A fellow researcher with Rutherford,
Chadwick discovered years later that the
nucleus was not made of only one particle
– the proton – but of two particles.
 This second particle was called the
“neutron” because it had no electrical
charge
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Subatomic Particles
Electron (e-) – mass 1/1840 amu, charge -1,
found in space around the nucleus
Proton (p+) – mass of 1 amu, charge +1,
found in the nucleus
Neutron (n) – mass of 1 amu, no charge,
found in the nucleus
The Atom
Made up of 3 fundamental subatomic
particles: protons, neutrons, and electrons
 Very small and very dense nucleus
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(nucleus make up over 99% of atom’s mass)
Nucleus contains protons and neutrons
 Electrons occupy the empty space outside
the nucleus
 # of protons = # of electrons
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What’s in the atom?
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Nucleons – particles in the nucleus (protons and
neutrons)
The combined total of the protons and neutrons
is called the mass number
The number of protons is called the atomic
number.
The atomic number identifies the element.
Electrons found in the space outside the nucleus
Lots of empty space
Isotopes
Particles with the same number of protons
and electrons but different numbers of
neutrons.
 Have different mass numbers.
 Have different masses.
 React the same chemically.
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Rutherford’s Dilemma
If the electron is in orbit around the nucleus,
it should be emitting radiation, but it is not.
What prevents the electron from being
pulled into the nucleus?
Electromagnetic Properties
Wavelength – distance between
consecutive waves.
 Frequency – the number of waves that
pass a point in a given amount of time,
usually 1 s.
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Quantum Theory
Proposed by Max Planck
 Two main ideas:
1. Energy changes are not continuous but
rather occur in small increments called
“Quantums”.
2. The energy of a quantum is directly
proportional to the frequency of the
radiation.
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….. Quantum theory continued.
E = hν
E = energy of a quantum
h= Planck’s constant
(6.63 x 10-34 J.s)
V = frequency
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What is the energy of a quantum with a
wavelength of 550 nm?
Photoelectric Effect
Smart Guy
In explaining the
photoelectric effect
Albert Einstein
showed that radiant
energy, such as light,
can posses particlelike properties.
Neils Bohr
Hydrogen’s Spectrum
Questions?
Why do the lines always occur at the same
place and why only 4 lines?
 Ans: Only specific energy changes are
possible in an atom. The lines are
representations of those energy changes.
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Bohr reasoned that if an electron could
occupy an infinite number of possible
orbits, its jumps from these orbits should
give rise to an infinite number of different
energy radiations……….THUS
A continuous spectrum
Bohr’s Interpretation
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Electrons in atoms can only occupy certain fixed
orbits or “energy levels”.
These energy positions are quantized, meaning
only certain values are possible within an atom.
To move from one energy orbit to another one,
an electron must absorb or emit a quantum of
energy exactly equal to the energy difference
between the two positions.
Electronic Transitions
Since Bohr . . . The Modern Model
The Quantum Mechanical Model
 Has a nucleus
 Electrons are in a “cloud” of negative
charge.
 An electron “orbit” is an area where the
electron is most “likely” to be.
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Quantum Numbers
Describe the properties of orbitals and
electrons in those orbitals.
1.
Principal q.n. (n) – designates the main
energy level or shell. (Bohr)
Values: 1, 2, 3, 4, 5, ……
n=1; means the electron is located in the
first energy level, which has the lowest
energy.
2. Angular momentum q.n. – (l) designates
the shape of the atomic orbital.
values: 0, 1, 2, 3, . . . n-1
so if n=3, l can be 0, 1, or 2.
if l = 0, (s) then it is spherical
if l = 1, (p) then it is dumbbell
if l = 2, (d) complex
if l = 3, (f) complex
Orbital Shapes
3. Magnetic q. n. (ml) – designates the
orbital’s orientation in space.
values – (from –l through 0 to +l)
ex. If l = 1 (p) then Ml can = -1, 0, +1
-1 corresponds to px
0 corresponds to py
+1 corresponds to pz
4. Spin q. n. (ms) – describes the spin of the
electron on its axis; clockwise or counter
clockwise.
values - +1/2 or – 1/2
Electron Configurations
Describes the arrangement of electrons in
an atom.
 Each electron in an atom has a set of 4 q.
n.
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Rules that govern orbital filling
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Aufbau principle – electrons enter orbitals of
lowest energy first.
Hund’s rule – when electrons enter orbitals of
equal energy, degenerate orbitals, each orbital
receives one electron, with parallel spins before
any receive two.
Pauli exclusion principle – no two electrons in
an atom can have the same set of 4 q.n. They
must have opposite spins.
Order of orbital filling
1869: Dmitri Mendeleev
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Russian chemist
Arranged elements in
tabular form so that
elements with similar
properties were in the
same column
When listed in order
by mass, elements
generally repeat
properties in groups
of 8 (Law of Octaves)
The First Periodic Table
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Most tables at the time listed elements by mass
Mendeleev also arranged elements by mass, but
left several “holes” in his table and occasionally
reversed the order of elements to fit the
properties of others in that column
The “holes” were later filled in with newly
discovered elements that had the properties
predicted by Mendeleev’s table.
The reason for the reversal of elements was
explained later by Henry Moseley, who noted
that the elements were in order by atomic
number (number of protons) rather than by mass
Introducing the Elements
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The Element Song
Periods- Horizontal Rows (7 periods)
Groups/families – Verticle columns (18
groups)