Download Investigating Atoms and Atomic Theory

Investigating Atoms and Atomic Theory

Students should be able to:
Describe the particle theory of matter.
 Use the Bohr model to differentiate among the
three basic particles in the atom (proton,
neutron, and electron) and their charges,
relative masses, and locations.
 Compare the Bohr atomic model to the
electron cloud model with respect to their
ability to represent accurately the structure of
the atom.

Atomos: Not to Be Cut
The History of Atomic Theory
Atomic Models

This model of the
atom may look
familiar to you. This is
the Bohr model. In
this model, the
nucleus is orbited by
electrons, which are
in different energy
levels.

A model uses familiar ideas to
explain unfamiliar facts
observed in nature.

A model can be changed as
new information is collected.
 The
atomic
model has
changed
throughout the
centuries,
starting in 400
BC, when it
looked like a
billiard ball →
Who are these men?
In this lesson, we’ll learn
about the men whose quests
for knowledge about the
fundamental nature of the
universe helped define our
views.
Democritus

This is the Greek
philosopher Democritus who
began the search for a
description of matter more
than 2400 years ago.
 He asked: Could matter
be divided into smaller
and smaller pieces
forever, or was there a
limit to the number of
times a piece of matter
could be divided?
400 BC
Atomos



His theory: Matter could
not be divided into
smaller and smaller
pieces forever, eventually
the smallest possible
piece would be obtained.
This piece would be
indivisible.
He named the smallest
piece of matter “atomos,”
meaning “not to be cut.”
Atomos


To Democritus, atoms
were small, hard
particles that were all
made of the same
material but were
different shapes and
sizes.
Atoms were infinite in
number, always
moving and capable
of joining together.
This theory was ignored and
forgotten for more than 2000
years!
Why?

The eminent
philosophers of the time,
Aristotle and Plato, had a
more respected, (and
ultimately wrong)
theory.
Aristotle and Plato favored the earth, fire,
air and water approach to the nature of
matter. Their ideas held sway because of
their eminence as philosophers. The
atomos idea was buried for approximately
2000 years.
Foundations of Atomic Theory

This was until the
18th century when
scientists began to
gather evidence
favoring the atomic
theory of matter.

All chemists in the late
1700’s accepted that an
element cannot be further
broken down by ordinary
chemical means,
however, there was
controversy as to whether
elements always combine
in the same ratio when
forming a particular
compound.

Aided by new balances and methods, scientist began to
accurately measure the masses of the elements and
compounds—this led to the discovery of various laws:

Law of Conservation of Mass (Lavoisier, 1777) :
mass is neither created nor destroyed during ordinary
chemical reactions or physical changes.
Also found out that pure chemical compound is
composed of a fixed proportion of elements.


Law of Definite Proportion: (Proust, 1797) A
chemical compound contains the same elements
in the same proportions by mass regardless of the
size of the sample or source of the compound.

Law of Multiple Proportions: If two or more
different compounds are composed of the same
elements, then the ratio of the masses of the
second element combined with a certain mass of
the first element is always a small whole number
ratio.
Law of Definite Proportion
Water
Hydrogen + Oxygen
18.0 g
2.0 g
16.0 g
Mass ratio = 16.0 g of O = 8 gives 8:1 ratio
2.0 g of H
8 Oxygen to 1 Hydrogen ratio
Always results from the same combinations of
atoms.

Law of Multiple Proportions
Also known as Dalton’s Law
CO and CO2

CO = C to O ratio is 1:1 O= 16g C = 12g
CO2 = C to O ratio is 1:2 O = 2(16g) = 32g
C = 12g
If CO and CO2 were each produced by the
reaction of 12 g of carbon with oxygen,
the resulting compounds would contain 16
g and 32 g of oxygen. The ratio of 16 g of
O to 32g of O is ½ . This indicates that in
 CO there are half as many oxygen atoms
for every carbon atom in CO2.

Dalton’s Model

In the early 1800s,
the English
Chemist John
Dalton performed a
number of
experiments that
eventually led to
the acceptance of
the idea of atoms.
Dalton’s Theory



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He deduced that all
elements are composed of
atoms. Atoms are
indivisible and
indestructible particles.
Atoms of the same element
are exactly alike.
Atoms of different elements
are different.
Compounds are formed by
the joining of atoms of two
or more elements.
.
 This
theory
became one
of the
foundations
of modern
chemistry.
Dalton’s Theory

Not all of Dalton’s Atomic Theory have
proven to be correct…we know that atoms
are divisible into smaller particles, and know
that for a given element can have atoms
with different masses…the theory has been
modified to explain the new observations.
Important concept is:
 All matter is composed of atoms
 Atoms of any one element differ in
properties from atoms of another element
remain unchanged.

Thomson’s Plum Pudding
Model
 In
1897, the
English scientist
J.J. Thomson
provided the first
hint that an atom
is made of even
smaller particles.
Thomson Model
He proposed a
model of the atom
that is sometimes
called the “Plum
Pudding” model.
 Atoms were made
from a positively
charged substance
with negatively
charged electrons
scattered about,
like raisins in a
pudding.

Thomson Model
 Thomson
studied
the passage of an
electric current
through a gas.
 As the current
passed through
the gas, it gave
off rays of
negatively
charged particles.

A cathode ray can also be deflected by a
magnet.
Thomson Model
 This
surprised
Thomson,
because the
atoms of the gas
were uncharged.
Where had the
negative charges
come from?
Where did
they come
from?
Thomson concluded that the
negative charges came from within
the atom.
A particle smaller than an atom had
to exist.
The atom was divisible!
Thomson called the negatively
charged “corpuscles,” today known
as electrons.
Since the gas was known to be
neutral, having no charge, he
reasoned that there must be
positively charged particles in the
atom.
But he could never find them.
The U.S. physicist Robert A. Millikan
(1868–1953) carried out experiments to
find the quantity of an electron’s charge.
From his data, he found that the
charge on each oil droplet was a
multiple of 1.60  10–19 coulomb,
meaning this must be the charge of
an electron.
•Using this charge and Thomson’s
charge-to-mass ratio of an electron,
Millikan calculated an electron’s mass.
•Millikan’s values for electron charge and
mass are similar to those accepted today.
•An electron has one unit of negative
charge, and its mass is 1/1840 the mass
of a hydrogen atom.
Protons and Neutrons
If cathode rays are electrons given off by
atoms, what remains of the atoms that
have lost the electrons?
 For example, after a hydrogen atom (the
lightest kind of atom) loses an electron,
what is left?

You can think through this problem using four
simple ideas about matter and electric charges.
1.
Atoms have no net electric charge; they are electrically
neutral.
2.
Electric charges are carried by particles of matter.
3.
Electric charges always exist in whole-number multiples
of a single basic unit; that is, there are no fractions of
charges.
4.
When a given number of negatively charged particles
combines with an equal number of positively charged
particles, an electrically neutral particle is formed.
It follows that a particle with one unit of
positive charge should remain when a
typical hydrogen atom loses an electron.
 In 1886, Eugen Goldstein (1850–1930)
observed a cathode-ray tube and found
rays traveling in the direction opposite to
that of the cathode rays.

•
He concluded that they were composed of
positive particles.
•
Such positively charged subatomic
particles are called protons.
The Neutron!!!!
In 1932, the English physicist James
Chadwick (1891–1974) confirmed the
existence of yet another subatomic
particle: the neutron.
 Neutrons are subatomic particles with no
charge but with a mass nearly equal to
that of a proton.

Rutherford’s Gold Foil
Experiment

In 1908, the
English physicist
Ernest Rutherford
was hard at work
on an experiment
that seemed to
have little to do
with unraveling the
mysteries of the
atomic structure.
 Rutherford’s
experiment Involved
firing a stream of tiny positively
charged particles at a thin sheet of
gold foil (2000 atoms thick)
 Gold
Foil Experiment


Most of the positively
charged “bullets” passed
right through the gold
atoms in the sheet of
gold foil without changing
course at all.
Some of the positively
charged “bullets,”
however, did bounce
away from the gold sheet
as if they had hit
something solid. He
knew that positive
charges repel positive
charges.




This could only mean that the gold atoms in the
sheet were mostly open space. Atoms were not
a pudding filled with a positively charged
material.
Rutherford concluded that an atom had a small,
dense, positively charged center that repelled
his positively charged “bullets.”
He called the center of the atom the “nucleus”
The nucleus is tiny compared to the atom as a
whole.
Rutherford

Rutherford reasoned
that all of an atom’s
positively charged
particles were
contained in the
nucleus. The
negatively charged
particles were
scattered outside the
nucleus around the
atom’s edge.
Bohr Model
 In
1913, the
Danish scientist
Niels Bohr
proposed an
improvement. In
his model, he
placed each
electron in a
specific energy
level.
Bohr Model

According to
Bohr’s atomic
model, electrons
move in definite
orbits around the
nucleus, much like
planets circle the
sun. These orbits,
or energy levels,
are located at
certain distances
from the nucleus.
Wave Model
The Wave Model
Today’s atomic
model is based on
the principles of
wave mechanics.
 According to the
theory of wave
mechanics,
electrons do not
move about an
atom in a definite
path, like the
planets around the
sun.

The Wave Model


In fact, it is impossible to determine the exact
location of an electron. The probable location of
an electron is based on how much energy the
electron has.
According to the modern atomic model, at atom
has a small positively charged nucleus
surrounded by a large region in which there are
enough electrons to make an atom neutral.
Electron Cloud:




A space in which
electrons are likely to be
found.
Electrons whirl about the
nucleus billions of times
in one second
They are not moving
around in random
patterns.
Location of electrons
depends upon how much
energy the electron has.
Electron Cloud:



Depending on their energy they are locked into a
certain area in the cloud.
Electrons with the lowest energy are found in
the energy level closest to the nucleus
Electrons with the highest energy are found
in the outermost energy levels, farther from
the nucleus.
Indivisible Electron
Greek
X
Dalton
X
Nucleus
Thomson
X
Rutherford
X
X
Bohr
X
X
Wave
X
X
Orbit
Electron
Cloud
X
X
The table below summarizes the properties
of these subatomic particles.
Properties of Subatomic Particles
Particle
Symbol
Relative
charge
Relative mass (mass
of proton = 1)
Actual mass
(g)
Electron
e–
1–
1/1840
9.11  10–28
Proton
p+
1+
1
1.67  10–24
Neutron
n0
0
1
1.67  10–24
Size of atoms

An atom is small, but the particles
associated with the atom are much
smaller!
Atomic Numbers and Mass
Numbers
The atomic number of an element is the
number of protons in the nucleus of each
of the elements. In a neutral atom, the
atomic number also designates how many
electrons the atom contains.
 The atomic number identifies an element.


For each element listed in the table below, the number of
protons equals the number of electrons
Mass Number



The total number of protons and neutrons in an
atom is called the mass number.
If you know the atomic number and mass
number of an atom of any element, you can
determine the atom’s composition.
The number of neutrons in an atom is the
difference between the mass number and
atomic number.
Mass Number

The composition of any atom can be
represented in shorthand notation using atomic
number and mass number.
•
The atomic number is the subscript.
•
The mass number is the superscript.

Au is the chemical symbol for gold.
ISOTOPES

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

Isotopes- Atoms having different mass numbers
but the same atomic numbers.
Nuclide is the term for any isotope of any
element.
EX; Hydrogen has 3 isotopesHydrogen, Deuterium, Tritium
There are three different kinds of neon
atoms.
 How do these atoms differ?

•
All have the same number of protons (10).
•
All have the same number of electrons (10).
•
But they each have different numbers of neutrons.


Because isotopes of an element have different
numbers of neutrons, they also have different
mass numbers.
Despite these differences, isotopes are
chemically alike because they have identical
numbers of protons and electrons, which are the
subatomic particles responsible for chemical
behavior.
Analogy to Dogs…
•
Their color or size doesn’t change the fact
that they are all dogs.
•
Similarly, the number of neutrons in
isotopes of an element does not change
which element it is because the atomic
number does not change.
Relative Atomic Masses

Masses of atoms measured in grams are
very small. In order to set up a relative
scale of atomic mass, one atom is
arbitrarily chosen as the standard and
assigned a relative mass value. The
masses of all the other atoms are then
expressed in relation to this defined
standard.

The average atomic masses are important
to chemist because they indicate relative
mass relationships in chemical reactions.
•
It is more useful to compare the relative
masses of atoms using a reference
isotope as a standard.
•
The reference isotope chosen is carbon12.
This isotope of carbon has been assigned a
mass of exactly 12 atomic mass units.
 An atomic mass unit (amu) is defined as onetwelfth of the mass of a carbon-12 atom.


A carbon-12 atom has six protons and six
neutrons in its nucleus, and its mass is set at 12
amu.
• The six protons and six neutrons account for
nearly all of this mass.
• Therefore, the mass of a single proton or a
single neutron is about one-twelfth of 12 amu,
or about 1 amu.
Average Atomic Masses

Most elements in nature occur naturally as
mixtures of isotopes. The percentage of
each isotope in the naturally occurring
element is nearly always the same, no
matter where the element is found. Atomic
masses for the elements are averages for
these naturally occurring mixtures of the
isotopes.
In nature, most elements occur as a
mixture of two or more isotopes.
 Each isotope of an element has a fixed
mass and a natural percent abundance.



Chlorine occurs as two isotopes: chlorine-35 and
chlorine-37.
If you calculate the arithmetic mean of these two
masses

(34.968 amu + 36.966 amu)/2), you get an
average atomic mass of 35.986.
•
However, this value is higher than the actual
value of 35.453.

To explain this difference, you need to
know the natural percent abundance of the
isotopes of chlorine.
•
Chlorine-35 accounts for 75 percent of the
naturally occurring chlorine atoms;
chlorine-37 accounts for only 24 percent.

Because there is more chlorine-35 than
chlorine-37 in nature, the atomic mass of
chlorine, 35.453 amu, is closer to 35 than
to 37.
The atomic mass of an element is a
weighted average mass of the atoms in a
naturally occurring sample of the element.
 A weighted average mass reflects both the
mass and the relative abundance of the
isotopes as they occur in nature.

Understanding Relative Abundance of Isotopes

The atomic mass of
copper is 63.546 amu.
Which of copper’s two
isotopes is more
abundant: copper-63 or
copper-65?
Compare the atomic mass to the mass of each
isotope.
 The atomic mass of 63.546 amu is closer to 63
than it is to 65.
 Determine the most abundant isotope based on
which isotope’s mass is closest to the atomic
mass.
 Because the atomic mass is a weighted average
of the isotopes, copper-63 must be more
abundant than copper-65.

You can determine atomic mass based on
relative abundance.
 To do this, you must know three things:
the number of stable isotopes of the
element, the mass of each isotope, and
the natural percent abundance of each
isotope.


To calculate the atomic mass of an
element, multiply the mass of each isotope
by its natural abundance, expressed as a
decimal, and then add the products.

Carbon has two stable isotopes: carbon-12,
which has a natural abundance of 98.89
percent, and carbon-13, which has a natural
abundance of 1.11 percent.
•
The mass of carbon-12 is 12.000 amu; the mass
of carbon-13 is 13.003 amu.
•
The atomic mass of carbon is calculated as
follows:

Atomic mass of carbon = (12.000 amu x 0.9889) +
13.003 amu x 0.0111) = (11.867 amu) + (0.144 amu) =
12.011 amu
Let’s try one!

Element X has two naturally occurring
isotopes. The isotope with a mass of
10.012 amu (10X) has a relative
abundance of 19.91 percent. The isotope
with a mass of 11.009 amu (11X) has a
relative abundance of 80.09 percent.
Calculate the atomic mass of element X.
The mass each isotope contributes to the
element’s atomic mass can be calculated
by multiplying the isotope’s mass by its
relative abundance. The atomic mass of
the element is the sum of these products.
• Isotope 10X:

mass = 10.012 amu

relative abundance = 19.91% = 0.1991
• Isotope 11X:

mass = 11.009 amu

relative abundance = 80.09% = 0.8009


Use the atomic mass and the decimal form
of the percent abundance to find the mass
contributed by each isotope.
for 10X: 10.012 amu x 0.1991 = 1.993 amu
for 11X: 11.009 amu x 0.8009 = 8.817 amu
Add the atomic mass contributions for all
the isotopes.
For element X, atomic mass = 1.953 amu +
8.817 amu = 10.810 amu

Why is the atomic mass of an element
usually not a whole number?
 The atomic mass of an element is usually
not a whole number because it is a
weighted average of the masses of the
naturally occurring isotopes of the
element.

Subatomic Particles of the atom

Each basic element has a certain number
of electrons and protons, which
distinguishes each element from all other
basic elements. In most elements, the
number of electrons is equal to the
number of protons. This maintains an
electrical balance in the structure of atoms
since protons and electrons have equal,
but opposite electrostatic fields.