Transcript Discovery of Atomic Structure

Democritus (460-370 B.C.)
Democritus was one of a few Greek philosophers who believed
that all matter in the world was made of of indivisible parts which
he called “atomos,” which means “indivisible.”
Although this theory was to be discovered to later be the truth,
Democritus’ ideas faded until the seventeenth century in Europe.
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John Dalton (1766-1844)
Dalton is known for his atomic theory
Theory states that atoms are the smallest chemical
building block of nature
His four postulates:
http://en.wikipedia.org/w
iki/John_Dalton#Atomic
_theory
Each element is composed of extremely small particles called
atoms
All atoms of a given element are identical; the atoms of
different elements are different and have different properties
(including different masses)
Atoms of an element are not changed into different types of
atoms by chemical reactions; atoms are neither created nor
destroyed in chemical reactions
Compounds are formed when atoms of more than one
element combine; a given compound always has the same
ratio and kind of atoms
ATOMS ARE SMALL AND CANNOT BE DIVIDED (later
changed!)
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Discovery of Atomic Structure
By 1850 scientists knew that atoms
were composed of charged particles.
Electrostatic attraction:
Like charges repel
Opposites attract
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J.J. Thomson
Experimented with
cathode rays and
found that the
properties didn’t
change
Distinguished charges
within atoms, positive
and negative charges
Plum-pudding model
Found electron
charge ratio
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Cathode Rays and Electrons
C.R. 1st discovered in mid-1880s from studies of
electrical discharge thru partially evacuated tubes
(CRTs)
Cathode rays = radiation produced when high voltage
is applied across the tube.
The voltage causes negative particles to move from
the negative electrode (cathode) to the positive
electrode (anode).
The path of electrons can be altered by the presence
of a magnetic field.
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Consider cathode rays leaving the positive electrode
through a small hole….
If they interact with a magnetic field perpendicular to
an applied electric field, then the cathode rays can be
deflected by different amounts.
Amount of deflection depends on applied magnetic
and electric fields.
Deflection also depends on the charge-to-mass ratio
of an electron.
Thomson determined the charge-to-mass ratio
of an electron in 1897.
Charge-to-mass ratio = 1.76 x 108 C/g
C: Coulomb, SI unit of electric charge
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Robert Millikan
Millikan built on J.J.
Thomson’s work on
electrons by measuring
their charge with his
famous oil-drop
experiment.
Using Thomson’s ratio,
Millikan calculated
electron’s mass of
9.10 x 10-28 g which
proved how much
smaller electrons are
than nucleus particles.
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Millikan Oil-Drop Experiment
Sprayed oil drops over the hole in a positively
charged plate and measured the electrostatic force
of attraction.
Found the charge on the electron to determine its
mass
Concluded the charge on the electron must be 1.60
x 10-19 C
Mass of electron = 1.60 x 10-19 C = 9.10 x 10 -28 g
1.76 x 108 C/g
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Radioactivity
(Spontaneous emission of radiation)
Type
Alpha
particle
Beta
particle
Positron
Gamma ray
Symbol
Charge
Mass (amu)
2+
4.002 60
0 e
-1
1-
0.000 548 6
0 e
+1
1+
0.000 548 6
0
0
He
β or
β or
γ
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-Conducted an experiment using alpha
particles to bombard gold foil to conclude:
1. Nucleus of an atom contains positive
particles that we now call protons.
2. The nucleus is a small dense area in the
atom.
-Studied three types of radioactive emissions:
alpha, beta & gamma
-Concluded that
alpha
particles were
He nuclei
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A Positively Charged Nucleus
Rutherford shot alpha particles though a thin
piece of gold foil.
Some of these particles were deflected
instead of passing straight through
Recall “like repels like.”
When a + alpha particle encountered a
nucleus of a gold atom, it was deflected by
the dense positively charged nucleus.
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James Chadwick suspected there were neutral particles when
he saw in experiments like Rutherford’s that some of the
particles were deflected backwards, meaning that they had no
charge. Chadwick had discovered the neutron!
neutron
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Niels Bohr
Bohr proposed that an electron circles
the nucleus in allowed orbits at specific
energy levels.
orbital
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Scientist Contributions
Thomson:
“Discovered” electron (1897)
Cathode ray experiments
“Plum pudding” atomic model
Millikan:
Mass of electron
Oil-drop experiment (1909)
Rutherford:
Positively charged nucleus (1911)
Gold foil experiments
Discovered proton (1919)
Chadwick: Discovered neutron (1932)
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Small Numbers
Electronic Charge: 1.609 x 10-19 C
Charge on an electron: -1.609 x 10-19 C
Charge on a proton: +1.609 x 10-19 C
Atomic Mass Unit (amu): 1.66054 x 10-24 g
Proton mass: 1.0073 amu
Neutron mass: 1.0087 amu
Electron mass: 5.486 x 10-4 amu
Unit of length used to note atomic dimensions:
1 Angstrom(Å) = 1x10-10 m
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Atomic Number
Number of protons or electrons in an element
Identifies the element
Atomic Mass
Nucleus contains most of the mass of an atom.
Protons and neutrons are each ~ 1.67 x 10-24 g.
Electrons are each ~ 9.11 x 10-28 g.
Use atomic mass unit (amu) instead of gram.
The mass of one proton is ~ 1 amu.
Mass Number
The sum of the number of protons and number of
neutrons in the nucleus
Is approximately equal to the average atomic mass
shown on periodic table.
Number of neutrons = mass number – atomic number
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Isotopes
Atoms of the same element with different
numbers of neutrons
Have the same number of protons
Example: Carbon-12 and Carbon-14
Radioactive Isotopes
Unstable in nature
Can be used to date fossils and rocks
The time it takes for half of the radioactive atoms
in a piece of the fossil to change to another
element is its half-life.
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Isotopes
AX
Z
Isotopes have the same Z, but different A.
Isotopes have different numbers of neutrons.
An atom of a specific isotope is called a
nuclide.
Nuclides of hydrogen include:
1H
= hydrogen (protium)
2H = deuterium (heavy hydrogen)
3H = tritium (3H is radioactive.)
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