Transcript early_Atomic Theory notes_academic - wths

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What is matter ?
To describe the
Democritus model
of the atom
Democritus ((& Leucippus)) ~ 400 BC
Pg 87-91
• Early philosophers thought that the material world
must be made up of tiny indivisible particles called
atomos, (meaning indivisible or uncuttable) that
move through space
– Later Democritus could not answer questions from
Aristotle regarding his notion of empty space and the
“atomic” view of matter faded away for centuries
Scientists studying gases…you can feel
the wind, thus air must be composed of
invisible particles of air
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To describe the
Dalton model ofPg38-39
the
atom
John Dalton - Devised an
atomic theory based on the
Greek idea of “atomos”
Atom-smallest particles of an
element that retain the
chemical identity of the element
Dalton’s Atomic Theory
1) Each element is composed of extremely small
indivisible particles called atoms
2) All atoms of a given element are identical to one
another in mass and other properties, but the
atoms of one element are different from the atoms
of all other elements.
3) Atoms of an element are not changed into atoms of
a different element by chemical reactions; atoms are
neither created nor destroyed in chemical
reactions.
4) Compounds are formed when atoms of more than
one element combine; a given compound always has
the same relative number and kinds of atoms.
Dalton’s Theory Explained several simple
laws of chemical combination known at the time
• Law of constant composition (definite
proportions)– in a compound the relative
numbers and kinds of atoms are constant
(pure H2O is 11% H and 89% O by mass)
• Law of conservation of matter – matter is
neither created nor destroyed
total massbefore = total massafter
Law of Multiple Proportions
if 2 elements A&B combine to form more than one
compound, different masses of B combine with the same
mass of A in the ratio of small whole numbers
There is twice as much
oxygen by mass in H2O2
as there is in H2O
Pg38-39
A problem with Dalton’s Theory
Are atoms “really” the smallest particles we
know?
Are atoms “really” indivisible?
What are atoms made up of?
Are atoms of the same elements “really”
identical?
Cathode Rays an Electrons
JToJdescribe
Thompson
the
Thomson model of the
atom
To state relative charge
& mass of electron and
proton
Cathode Rays
Pg 92-94
• Partially evacuated tubes produced radiation under
high voltage
• Called cathode rays = initiated from cathode side
• Rays could not be seen but made materials
fluoresce
Cathode Rays
• Experiments showed cathode rays are deflected by electric or
magnetic fields
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• Thomson found that cathode rays are the same
regardless of what material the cathode is made of
• How do these findings help us answer the
question…are cathode rays radiation (energy) or
particles?
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Electrons
• Cathode rays deflected by “field” = must be
electrical particle not pure energy
• Deflected TOWARD positive field = must
be negative particle
• Rays were the same no matter what substance
cathode was made of = must be a fundamental
particle (all the same)
“So how do all these particles that make up matter go
together?
Pg42
JJ Thomson’s “Plum Pudding” – Atomic Model
• Small negatively charged electrons embedded in a
positively charged atom
• Like seeds (negative electrons) embedded in a
watermelon (positive atom)
Finding the Center
“Ernest Rutherford
and
Gold
Foil”
To describe the
Rutherford model of the
atom
To state relative charge
& mass of electron,
proton and neutron
Radioactivity & Rutherford
Pg41-42
Revealed three types of radiation:
alpha (α) beta (β) gamma (γ) –
How does each type respond to electric/magnetic
field? And what does that imply about it?
Rutherford’s Gold Foil Experiment
Pg42
• almost all α particles passed directly through the foil
without deflection
• a small % slightly deflected, on the order of 1 degree,
consistent with Thomson’s plum-pudding model
Rutherford’s Gold Foil Experiment
Pg42
•BUT eventually … a small amount of
scattering was observed at large angles & some
particles were even scattered back in the
direction from which they had come
What implications
does this finding
have about the atom?
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Rutherford’s Atom
Pg42
• most of the mass of each atom and all of its positive
charge reside in a very small, extremely dense region
• Rutherford called this region the nucleus
LiNkeD
Protons
LiNkeD
Pg42-43
Subsequent experimental studies led to the
discovery of positive particles (called protons) in
the nucleus.
• Protons were discovered in 1919 by Rutherford.
• HOWEVER, when the nuclear mass was
computed based on the number of protons, the
mass value was much less than the actual mass…
• What does this then imply about the atom???
The Neutron
Pg42-43
…and subsequent experimental studies led
to neutral particles (called neutrons) in the
nucleus.
• Neutrons were discovered in 1932 by the
British scientist James Chadwick (1891–
1972).
“Planetary” Model
Pg43
• Every atom has an
equal number of
electrons and protons,
and so atoms have no
net electrical charge.
• The electrons are
attracted to the
• The vast majority of an
protons in the nucleus
atom’s volume is the
by the force that exists
space in which the
between particles of
electrons reside.
LinKed opposite electrical
charge.
Rutherford’s Paradox
• You cannot simply explain electron as
“orbiting” the nucleus…because (as per
classical physics)
the e- will loose
energy and be
more attracted
toward nucleus
Pg 228-229
Properties of Particles
Link
Pg43-44
• charge of electron is −1.602 × 10−19 C and that of
a proton is +1.602 × 10 −19 C (charges of particles
are expressed as multiples of this charge)
• mass of atoms very small so instead of grams we use
amu (“atomic mass units”) (1 amu =1.66054 × 10−24 g )
• Atoms are extremely small (SI unit angstrom; Å is
used) (1 Å = 0.0000000001m)
Atomic Notation
To draw a diagram of an
atom given its atomic
notation
To explain & illustrate
the concept of isotopes
Pg45
Practice
Pg45
How many protons , neutrons and electrons
are there in each?
p+:
n0:
e-:
p+:
n0:
e-:
p+:
n0:
e-:
Another Problem with Dalton’s Theory
Pg45-46
• Most elements are uniform mixtures of 2 or more unique
substances called isotopes.
• Isotopes of an element have very similar chemical
properties but their atoms have slightly different masses
What about the different atoms, of the same element,
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could make their masses different???
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Atomic Mass Scale
Pg46-48
• Masses of atoms are measured in atomic mass
units (amu)
• 1 amu = 1.66054x10-24g
• 1g = 6.02214x1023amu
• Defined by assigning a mass of exactly 12amu
to an atom of the 12C isotope of carbon.
Average Atomic Weight
Pg47-48
• Every sample of an element has the same isotopic
composition (the average mass per atom is the same
from sample to sample)
• So…we can use the masses of each isotope and its
relative abundance to determine the average atomic
weight for the element!
Average Atomic Weight
Pg47-48
Naturally occurring chlorine is 75.78% 35Cl which has an atomic
calculate
the37atomic
mass of 34.969 amu,To
and
24.22%
Cl, which has an atomic mass of
weight
an element
36.966 amu. Calculate
thefor
average
atomic mass (that is, the atomic
weight) of chlorine. given the mass and
abundance the naturally
occurring isotopes.
Solution The average atomic mass is found by multiplying the
abundance of each isotope by its atomic mass and summing these
products. Because 75.78% = 0.7578 and 24.22% = 0.2422 we have
Average atomic mass = (0.7578)(34.969 amu) + (0.2422)(36.966 amu)
= 26.50 amu + 8.953 amu
= 35.45 amu
This answer makes sense: The average atomic mass of Cl is
between the masses of the two isotopes and is closer to the value of
35Cl which is the more abundant isotope
Practice
Example: Copper occurs naturally as two isotopes Cu-63
and Cu-65. Given the atomic mass and the % natural
abundance of each isotope below, calculate atomic weight of
Cu (as on the periodic table).
Isotope
Cu-63
Cu-65
mass
62.930 amu
64.928 amu
natural abundance
69.09%
30.91%
Radioactivity – spontaneous emission of radiation
(high energy & particles)
Pg41-42
To state the properties
of alpha, beta and gamma
radiation
• In a nuclear reaction, protons and
neutrons are rearranged
• In a chemical reaction, electrons
rearranged
Often, at least one isotope is unstable – the strong
force of the nucleus is overcome by positive-positive
repulsion of the+ protons
Radioactive Decay - breakdown of atom
– Nucleus becomes unstable
– Radiation is emitted until nucleus becomes stable
– Stable, non-radioactive atom is formed
– Can occur naturally, or be caused by
bombarding the atom with energy
LinKeD
3 Main Types of Radiation
Particles Charge Mass
Alpha
Stopped By
+2
4
Paper or light clothing
-1
0
0 1
1
2 p+
2 no
no e1 eno p+
no no
Heavy Clothing,
Sunscreen or Lead
(1cm thick)
Gamma
None
None
0
Concrete or Lead
(10cm thick)

4
2
2
Beta
e
0
0

Penetrating Ability
• HALF-LIFE is the is the time that it takes for
1/2 a sample to decompose
To relate the amount of
• The rate of a nuclear
transformation depends only on
radioactive sample, or its
the “reactant” (original
compound)
concentration.
radioactivity,
to a given
half-life.
• For each duration (half-life),
For each duration (halflife), one half of the substance decomposes
For example
Ra-234 has a half-life of 3.6 days; if you
start with 50 grams of Ra-234, how much do
you have …
After 3.6 days :
After 7.2 days :
After 10.8 days :