Transcript Matter is Made up of Atoms

The Structure of the
Atom
Chapter 4
The Beginning…
 Democritus- Greek Philosopher (460-370 BC)
 Came up with the term “atomos”- now called “atoms”
 Believed that atoms could not be created, destroyed,
or further divided
 Matter is composed of empty space through which atoms
move
 Atoms are solid, homogenous, indestructible, and
indivisible
 Different kinds of atoms have different sizes and shapes
 The different properties of matter are due to size, shape
and the movement of atoms
 Aristotle- Greek Philosopher (384-322 B.C.)
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One of the most influential philosophers
Wrote extensively on many subjects
Rejected Democritus’ idea that the “nothingness” of
empty space could exist His
denial went largely unchallenged for two thousand
years!
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Dalton (Schoolteacher- England) (1766-1844)
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Atomic Theory
1.
2.
3.
4.
5.
All matter is made up of atoms.
Atoms are indestructible and cannot be divided into
smaller particles.
All atoms of one element are exactly alike, but they
are different from atoms of other elements.
Different atoms combine in simple whole number
ratios to form compounds
In a chemical reaction, atoms are separated,
combined, or rearranged.
 Dalton’s atomic theory explains the
conservation of mass when a compound
forms from its component elements.
Subatomic Particles
and the Nuclear Atom
Discovering the Electron
 Sir William Crookes- English
Physicist
Saw a flash of light within
one of the vacuum tubes
with metal electrodes at
opposite ends
 Accidental discovery of the
cathode ray (ray of
radiation originating from
the cathode end of the
tube)
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ELECTRONS!!
 Scientists continued their research using
cathode ray tubes, and by the end of the
1800’s were fairly convinced of the following:
1.
2.
Cathode rays were actually a stream of charged
particles
The particles carried a negative charge. (exact
value of the negative charge was not known)
JJ Thompson
 Series of experiments to determine the
ratio of its charge to mass
 He concluded that the mass of the
charged particles (electrons) was much
less than that of a hydrogen atom
(lightest known atom)
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SHOCKING!- this meant Dalton was
WRONG and atoms were divisible into
smaller subatomic particles
 Proposed a model of the atom that came
known as the plum pudding model.
JJ Thomson
Robert Millikan
 Determined the charge of an electron
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His experimental set up and technique was
accurate to within 1% of the currently accepted
value.
Oil Drop Experiment
Ernest Rutherford
 Studied how positively charged alpha
particles interacted with solid matter
Thompson’s
“Plum Pudding ”
Model
Rutherford’s
results when
alpha particles
where beamed
through foil
Rutherford
 He expected most of the fast moving and
relatively massive particles to pass
straight through the gold atoms
 He concluded:
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an atom consisted mostly of empty space
through which atoms move
There was a tiny, dense region centrally
located within the atom that contained all
of an atom’s positive charge and virtually
all of its mass (aka NUCLEUS)
Eight years later…. (1920)
 Rutherford refined his concept of the
nucleus
 It contained positively charged
particles called PROTONS
Few years later…
 James Chadwick (Rutherford’s
coworker)
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Showed that the nucleus also contained
another subatomic particle- a neutral
particle called a NEUTRON
The neutron has a mass equal to that of a
proton
How Atoms Differ
Section 4.3
Atomic Number and Masses
 Atomic Number: the number of
PROTONS in the nucleus of an
atom of an element
 Also tells you the number of
ELECTRONS since elements
have no over all charge.
 The number of PROTONS
determines the IDENTITY of
the element
Isotopes
Have different mass numbers because they
have a different amount of NEUTRONS
*If they had a different amount of PROTONS, they
would be a different element completely*
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Because isotopes have a different number of
NEUTRONS, they also have a different mass
number
Mass Number = Protons + Neutrons
Sodium Isotopes
Example: NEON
Atomic Number:10
Mass Number: 20.180
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# of Protons: 10
# of Electrons: 10
# of Neutrons: 10
Mass of Individual Atoms
 Atomic Mass Unit (amu) is defined as
1/12 the mass of a carbon-12 atom.
 The atomic mass of an element is the
weighted average mass of the isotopes
of that element.
75%
25%
Average Atomic Mass
 To calculate the atomic mass of an
element
1. Multiply the mass of each isotope by
its natural abundance % (expressed
as a decimal)
2. Add the products
Unstable Nuclei and Radioactive Decay
Radioactivity
 Nuclear Reactions- involve a change in
an atom’s nucleus (unlike most chemical
reactions)
 Radioactivity- a process in which
substances spontaneously emit radiation
 Radiation- the rays and particles emitted
by the radioactive material
 Radioactive atoms undergo significant
changes that can alter their identities
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One element can change into atoms of
another element
 Unstable nuclei lose energy by emitting
radiation in a spontaneous process (aka
radioactive decay)
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They continue to do this until they form a
stable nonradioactive atoms
Types of Radiation
 Alpha Radiation
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Deflects towards a negatively charged
plate
Alpha Particles- 2 protons and 2 neutrons
(+2 charge)
 Beta Radiation
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Radiation that is deflected towards a
positively charged plate
Consists of fast moving electrons called
beta particles
 Gamma Radiation
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Gamma Rays are high-energy radiation
that possess no mass
Because gamma rays are massless, the
emission of gamma rays by themselves
cannot result in the formation of a new
atom