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Transcript
The Atom
The History of Atomic Models

Think with me…about sugar crystals,
you can see that they are small crystals
and every crystal is identical.
 You
may grind these particles into a very
fine powder, but each tiny piece is still
sugar.
 If you dissolve the sugar in water,
sugar particles become
invisible.
virtually
From Philosophy to Science
the
 You
could even look at the sugar solution
under a microscope and you’d still not
be able to see
the sugar.
 However, you know it is still there because
you can taste it.

These kind of observations and logic
patterns led ancient philosophers
to
ponder the design
of the
universe
From Philosophy to Science

There were two schools of thought of the
composition of the cosmos…
 is
everything in the universe continuous
and infinitely divisible
 Or, is there a limit to how small you can
get?

Particle theory was not the most popular
early opinion, but was supported as early
as Democritus
in ancient Greece.
From Philosophy to Science

Democritus proposed that all the matter
is composed of tiny particles called
“Atomos”
 These
“particles” were thought to be
indivisible

Aristotle did not accept Democritus’
atom, he was of the “matter is
continuous” philosophy
 Because
of Aristotle’s popularity
his
theory was adopted as the standard
From Philosophy to Science
By
the 1700’s nearly all chemists had
accepted the modern definition of an element
as a particle that is indivisible
It was also understood at that time that
elements combine to form compounds that
are different in their properties than the
elements that composed them
 However,
these understandings
were
based on observations
not
empirical evidence
From Philosophy to Science

There was controversy as to whether
elements always combine in the same
proportion when forming a particular
compound.
 In
the 1790’s, chemistry was
revolutionized by a new emphasis on
quantitative analysis because
of new
and improved balances

This new technology led
to the discovery of some new scientific
understandings
From Philosophy to Science

The Law of Conservation of Mass:
 Proposed
by Antoine Lavoisier
 States that mass is neither created nor
destroyed during ordinary chemical rxns
or physical changes.
 Which means the total mass of the
reactants must equal the total mass of
the products.
From Philosophy to Science

+
Carbon, C
Mass x
Oxygen, O
Mass y
Carbon Monoxide, CO
Mass x + Mass y

Carbon Monoxide, CO
Mass x + Mass y
+
Carbon, C
Mass x
Oxygen, O
Mass y

The Law of Definite Proportions:
 The
fact that a chemical compound
contains the same elements in exactly the
same proportions by mass regardless of
the size of the sample or the source of the
compound
 NaCl is NaCl no matter if it is table salt
(small crystals) or rock salt
crystals)
(large
From Philosophy to Science

The Law of Multiple Proportions:
 If
2 or more different compounds
are
composed of the same 2 elements, then the
ratio of the masses of the 2nd element
combined with a certain mass of the 1st
element is always
a ratio of small
whole numbers
From Philosophy to Science
+
Carbon
1
=
Oxygen
1
+
Carbon Monoxide,
1:1
=
Carbon
Oxygen
1
2
Carbon Dioxide,
1:2

In 1808, John Dalton proposed an
explanation for each of the proposed laws
 He
reasoned that elements were composed
of atoms & that only whole #’s of atoms
can combine to form
compounds
 His ideas are now called the Atomic
Theory of Matter
Atomic Theory
1.All matter is composed of
extremely small particles called
atoms
2. Atoms of a given element are
identical in size, mass, and
other properties; atoms of
different elements differ in size,
mass, & other properties
ELEMENT
2
ELEMENT
3
ELEMENT
4
Atomic Theory
3. Atoms cannot be
subdivided, created, or
destroyed
4. atoms of different
elements combine in
simple whole # ratios to
form chem compds
5. in chemical rxns, atoms
are combined,
separated, or rearranged
Atomic Theory
+
+

Through these statements, evidence could
be gathered to confirm or discount its
claims
 Not
all of Dalton’s claims held up
to
the scrutiny of experimentation
 Atoms CAN be divided into even smaller
particles
 Not every atom of an element
has an
identical mass
Atomic Theory
 Dalton’s
Atomic Theory of Matter has
been modified.
 What remains is…
1.All
matter is composed of atoms
2.Atoms of any one element differ in
properties from atoms of another element
 One
of the disputed statements of Dalton
was that atoms are
indivisible
Atomic Theory
 In
the 1800’s it was determined that
atoms are actually composed of several
basic types of smaller particles
 it’s the number and arrangement of these
particles that determine the atom’s
chemical properties.

The def. of an atom that emerged was, the
smallest particle of an element that
retains the chemical properties of that
original
element.
Atomic Theory

All atoms consist of 2 regions that
contain the subatomic particles
 The
 The

nucleus
electron cloud around the nucleus
The nucleus is a very small region located
near the center of the atom
 In
every atom the nucleus contains at
least 1 proton, which is positively charged
particle and usually contains 1 or more
neutral
particles called neutrons
Atomic Structure

The electron cloud is the region that
surrounds the nucleus
 This
region contains 1 or more elec-trons,
which are negatively charged subatomic
particles
 The volume of the
electron cloud is much
larger than the nucleus
Atomic Structure

The discovery of the first subatomic
particle took place in the late 1800’s.
A
power source was attached to two metal
ends of an evacuated glass tube, called a
cathode
ray tube.
 A beam of “light”
appears
between
the two
electrodes
called a cathode
ray.
Discovery of the Electron
Electric Current
Electric Current
Steering Coils
The electron beam is painting all 525 lines 30 times
per sec, it paints a total of 15,750 lines per sec.
 Investigators
began to study the ray and
they observed that…
An object placed in the path of the ray
cast a shadow on the glass
2. A paddle wheel placed the path
of the
cathode ray began to spin
3. Cathode rays were deflected
by
a magnetic field
4. The rays were deflected away
from a
negatively charged
object
1.
Discovery of the Electron


The first 2 observations support the idea
that the ray is composed of
tiny
individual particles traveling through the
vacuum tube
The second set of observations support the
evidence that the ray
is composed of
a substance that
is negatively
charged.
Discovery of the Electron

J.J. Thomson studied the rays and proved
that they were tiny negative particles
being emitted from the metal atoms.
 Dubbed

these tiny particles “electrons”
Robert Millikan then used an ingenious
investigation to calculate the mass to
charge ratio of an atom
 He
determined that the electrons were
not part of the mass
of the
atom.
Discovery of the Electron

What can their work help us conclude
about the atom?
 atoms
are composed of smaller particles,
and one of these compo-nents is negatively
charged
 atoms are neutral, so there must
be an
opposing (+) charge
 because E’s are essentially
mass-less,
an opposing substance that makes up the
mass
of the atom
First Atomic Model
Negative particles
embedded in a
sphere of positive
plasma-like matter.
THINK…
Chocolate Chip Cookie

In 1886, E. Goldstein observed in the
cathode-ray tube a new set of rays
traveling in the opposite direction than
the cathode rays
 The
new rays were called canal rays and
they proved to be positively charged
 And the particles mass were about
2000
X’s that of the electron
Discovery of the Proton
Discovery of the Proton

In 1932, the English physicist James
Chadwick discovered yet another
subatomic particle.
 the
neutron is electrically neutral
 It’s mass is nearly equal to the proton

Therefore the subatomic particles are the
electron, proton, and neutron.
Discovery of the Neutron
electron
e-
-1
0
9.11x10-28
proton
p+
+1
1
1.67x10-24
neutron
n0
0
1
1.67x10-24
Structure of the Atom

Scientists still didn’t really understand
how the particles were put together in an
atom.
 This
was a difficult question to resolve,
given how tiny atoms are.

Most thought it likely that the atom
resembled Thomson’s model
Atomic Structure


In 1911, Ernest Rutherford et al.
provided a more detailed picture
of
the internal structure of the atom
In his experiment, Rutherford directed a
narrow beam of alpha particles at a very
thin sheet of gold foil.
 Alpha particles (a) are He atoms
that
have been stripped
their electrons
Rutherford Model
of

According to Thomson’s model,
the
heavy, positive alpha particles should
pass easily through the gold, with only a
slight deflection
 And
mostly that’s how it happened.
 However, they found 1 in every
8000
particles had actually been deflected back
toward the source.
Rutherford Model

Rutherford suggested a new structural
model of the atom.
 He
stated that all the positive charge and
the mass is concentrated in a small core in
the center of the atom, AKA nucleus
 And that the atom is mostly empty space
with electrons surrounding
the
positively charged nucleus like planets
around the sun.
Rutherford Model
Rutherford Model

Rutherford’s planet system model was an
improvement over earlier models, but it
was still not complete.
 Physics
says that electrons can’t orbit the
nucleus without losing energy,
 Losing energy would cause the electron to
spiral into the nucleus.
 The attraction of the electron to the nucleus
would cause it to spiral into the nucleus as
well
Rutherford Model

Niels Bohr proposed a new model that
would allow the electrons to be outside
the nucleus and in orbit around the
nucleus.
 His
model coupled Rutherford’s model
with a new concept of energy in Physics
called quantum mechanics,
 Bohr proposed that the electrons aren’t on
any random orbit around the nucleus,
they are on
“special” orbits
Bohr Model

Bohr’s Model restricts the orbits on which
an electron can be
 The
bases for what orbit an electron is
allowed is entirely based on how much
energy the electron has
 If it has any more energy or any less energy
it would be forced to be on a different path
of different energy
 The
energy of the electron is quantized,
which means it is
of a very
specific quantity
Bohr Model

Each path or level of energy that the electron is on is given a
label of “n”
 Such that n=1 is the closest energy level to the nucleus
 n=2 is higher in energy and outside of, but adjacent to n=1,
and so on…

Each energy level c

an only hold




a certain number of electrons (2n2)
n=1 can hold 2 electrons
n=2 can only hold 8 electrons
n=3 can hold 18 electrons
Etc.
Bohr Model
Bohr Model of the Atom
Model describes the paths
of electrons as energy levels.
The electrons are only allowed
to have a certain amount of energy
which restricts their path around
the nucleus.
Bohr Model

With the exception of Hydrogen, every
nucleus contains 2 kinds of particles
protons and neutrons
 they
make up the mass of the atom (Mass
Number = Protons + Neutrons)

Proton has a charge equal to but opposite
of the charge of an elec.
 Atoms
are neutral because they
equal #’s of protons
electrons
Atomic Structure
contain
&

The atoms of different elements differ in
the # of protons in their nuclei and
therefore in their positive charge
 The
# of protons the atom contains
determines the atom’s identity
 Only Oxygen contains 8 protons
 Only Fluorine contains 9 protons
 Only Neon contains 10 protons
Structure of the Atom

The nucleus is composed of a densely
packed cluster of protons, which are all
electrically positive
 Don’t
like charges repel?
 Why don’t they fly apart?

When 2 protons are in very close
proximity, there is a strong force
attraction between them.
 similar
attraction exists
when neutrons are close
Structure of the Atom
of

These short-range p+-n0, p+-p+, &
n0n0 forces hold the nuclear particles
together, A.K.A strong nuclear forces.
 When
these nuclear forces are strong
enough the atom is stable
 If the forces are not strong enough
the
atom (heavier atoms) the atom
is
unstable and becomes radioactive.
Structure of the Atom

Basic Truth: All atoms contain the same
basic parts, but atoms of different
elements have different numbers of
protons.
 The
PT lists atoms in consecutive order by
their Atomic Number (Z)
 The atomic number is directly
related to
the number of protons
in the nucleus
of each atom
of that element
Counting Atoms

The atomic number is found above the
elemental symbol on the PT and it
defines the type of element
 Atomic
#47 can only be Ag and it also
can only have 47 protons in each nucleus
 Because atoms are neutral,
we
know from the atomic number
the atom must also contain
47 electrons.
Counting Atoms

The total number of protons & neutrons
determines the mass of the atom
 Called
the Mass Number
 A Carbon atom, has 6 protons and 6
neutrons, so its mass number is 12

If you know the atomic number & mass
number of an atom of any element, you
can determine the atom’s composition
Counting Atoms
ATOMS OF THE 1ST TEN ATOMS
NAME
SYMBOL
ATOMIC #
p+
n0
MASS #
e-
Hydrogen
H
He
Li
Be
B
C
N
O
F
Ne
1
2
3
4
5
6
7
8
9
10
1
2
3
4
5
6
7
8
9
10
0
2
4
5
6
6
7
8
10
10
1
4
7
9
11
12
14
16
19
20
1
2
3
4
5
6
7
8
9
10
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon

Every Cl atom has 17 protons, w/o
exception, but not every Cl atom has 18
neutrons.
 Atoms
with the same # of protons but
contain different #s of neutrons are called
isotopes.
 Since isotopes of an element have different
#s of neutrons they have different masses
Counting Atoms

Isotopes are chemically alike because they
have identical numbers of protons and
electrons
 It’s
the electrons and protons that are
responsible for chemical behavior

Isotopes can be noted using hyphen
notation (Cl-35 vs Cl-37)
 elemental
symbol hyphen mass number
Counting Atoms
or Na-23
or Na-24
Isotope: one of two or more atoms having
the same number of protons but
different numbers of neutrons