Download Chapter 3 - EZWebSite

Chapter 3
Atoms:
The Building Blocks
Of Matter
Section 3-1
The Atom:
From Philosophical Idea
To Scientific Theory
Two Theories on the Nature of
Matter



It is continuous and can be
infinitely broken down into
smaller and smaller parts.
It has some fundamental particle
that can be reached that cannot
be broken down any further.
The idea of the fundamental
particle is the one accepted
today.
The study of matter lead to two
basic laws of chemistry

The Law of Conservation of
Matter – During ordinary
chemical and physical reactions,
matter is neither created nor
destroyed, it is conserved.
The study of matter lead to two
basic laws of chemistry

The Law of Definite Proportions
– 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.
John Dalton – (1766-1844)



English
Schoolteacher
Proposed an
explanation for the
Law of
Conservation of
Mass and the Law
of Definite
Proportions in
1808.
This was called
Dalton’s Atomic
Theory. It has five
statements.
Dalton’s Atomic Theory


All matter is
composed of
submicroscopic,
indivisible
particles called
atoms.
Atoms of the
same element
are the same,
atoms of
different
elements are
Dalton’s Atomic Theory


Atoms cannot
be broken
down, created,
destroyed, or
changed into
other atoms.
Atoms of
different
elements
combine in
whole number
ratios to form
Dalton’s Atomic Theory

In chemical
reactions, atoms
are combined,
separated, or
rearranged.
Dalton’s Atomic Theory is not
Completely Correct




Atoms are
divisible.
Atoms of the
same element
are not always
the same.
Atoms are not
submicroscopic.
Atoms can be
changed into
other atoms.
Section 3-2
The Structure of the Atom
Structure of the Atom


An atom is the
electron
smallest particle
of an element
that retains the
properties of
that element.
It contains two
regions, the
nucleus and the
electron cloud.
neutron
proton
The Regions of the Atom




The Nucleus
Small, Dense
region in the
center of an
atom.
Holds at least
one proton,
most hold at
least one
neutron.
The nucleus is


The Electron
Cloud
Region around
the nucleus that
contains
electrons.
The Three Subatomic
Particles



Electrons – Negatively charged,
found outside of the nucleus.
Smallest subatomic particle.
Protons – Positively charged,
found in the nucleus. Middle
sized subatomic particle.
Neutron – No charge, found in
the nucleus. Largest subatomic
particle.
Discovery of the Electron


In the late 1800’s many experiments
were done by passing current
through gases at low pressure.
These were carried out in sealed
glass tubes called cathode ray tubes.
Discovery of the Electron


When a current was passed through
the tube, the anodes (+) glowed.
The theory was that the particles that
made up the cathode ray were
negatively charged, and moved from
cathode to anode.
Joseph John Thompson (1856 –
1940)




Conducted
experiments to
support this idea in
1897.
Called these
particles
corpuscles.
Determined it had
a very large
negative charge
with a very small
mass.
Conclusion was
The Thompson Model of the Atom


Thompson’s model
of the atom is
different from
Dalton’s because it
shows negative
electrons stuck in
positive ball of
matter.
There are two
problems. First, it
does not show how
easy it is to
remove electrons,
and second, it
does not show
Assumptions from Thompson’s
Experiments



Atoms must be divisible. This
proves part of Dalton’s theory
wrong.
Since atoms are neutral, there
must be a positively charged
particle to counteract the
negatively charged electron.
Because electrons are so small,
there must be another particle
that gives the atom its mass.
Lord Ernest Rutherford (1871 –
1937)



Thompson’s
graduate student.
He, Ernest
Marsden and Hans
Geiger worked
together firing
alpha particles at a
sheet of gold foil.
Observed that
most of the
particles passed
through the foil, but
some were
deflected back at
Discovery of the Nucleus



To the scientists,
this was like firing
a gun at a piece of
tissue paper and
having it bounce
off back at them.
Rutherford
theorized that the
particles were
bouncing off a
small, dense,
positively charged
region in the atom.
He called this the
nucleus.
The Rutherford Model of the Atom



Rutherford
developed a new
atomic model in
1913.
In it, the electrons
move in space
around the
positively charged
nucleus.
The problem with
his model is that
the nature of
electrons is not
accurately
described by this
The Atomic Nucleus




The nucleus contains protons and
neutrons.
It is the number of protons that
determines the identity of an
element.
Atoms with more than one proton
should not exist.
Nuclear forces hold the nucleus
together. These are strong
attractions between protons and
protons, neutrons and neutrons, and
Atoms are very small and
Numerous.



Neon for example is a gas that
makes up only 0.002% of our
air.
In each breath you take, you
breath in 5 * 1017 atoms of neon.
If the nucleus of an atom were
increased in size to 1cm3, it
would have a mass of 2 * 108
tons.
Section 3-3
Counting Atoms
There are three numbers that
identify an atom



Atomic Number
Mass Number
Atomic Mass
Atomic Number



The number of protons in the
nucleus.
This is the number that identifies
an element.
This is the whole number on the
periodic table.
Mass Number







Most elements have isotopes.
Isotopes – Atoms of the same elements
with different numbers of neutrons.
Mass Number – The number of protons
plus the number of neutrons.
There are two ways to write the isotope of
an element.
Write the name, then a hyphen, then the
mass number – Carbon – 14
Write the symbol, mass number as a
superscript, atomic number as a subscript.
An isotope is also called a nuclide.
Atomic Mass



Do not use grams to measure
mass of an atom, instead we
use atomic mass units (amu).
Average atomic mass is the
weighted average of the masses
of all of the isotopes of an
element.
This is the decimal number on
the periodic table.
There are three ways to compare
numbers of atoms.



Moles
Avogadro’s Number
Molar Mass
The Mole

The amount of
substance that
contains as
many particles
as there are in
12 grams of
carbon – 12.
Avogadro’s Number



The number of
particles in one
mole of a pure
substance.
Named after
Italian scientist
Amadeo
Avogadro (1776
– 1856).
6.02 * 1023
particles.
Molar Mass



The mass of one mole of a pure
substance.
Expressed in grams.
It is equal to the atomic mass
(decimal number on the periodic
table) in grams.