Chapter 3 – Atoms: The building
Blocks of Matter
I. Philosophical Ideas
• A. Democritus
• 1. Greek philosopher (thinker)
• 2. Supported the particle theory of matter in
400 B.C.
• 3. Called the basic particle the “atom”,
which in the Greek language meant
“indivisible”
• 4. Believed matter could be divided in half
over and over again until it reached the point
where it could no longer be divided, and this
point would be the atom
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B. Aristotle
1. Greek thinker shortly after Democritus
2. Did not believe in atoms
3. Thought all matter was continuous
4. The world believed Aristotle for nearly 2000
years, rather than Democritus
• C. Neither of these Greek thinkers had any
experimental evidence to support their ideas
• 1. Their beliefs were philosophical ideas – not
scientific theories
• 2. It was the 18th century before scientists
began to gather evidence for the atomic theory
II. Atomic Theory
• A. Accepted foundations of the 1700’s
• 1. An element cannot be broken down
by ordinary chemical means
• 2. Elements combine to form
compounds
• 3. Properties of compounds are
different from properties of its elements
• 4. Transformation of one substance into
a new substance is a chemical reaction
• B. Improved scientific instruments brought
about the discovery of basic scientific laws
• 1. Law of Conservation of Mass
• a. Matter is neither created or destroyed
• b. Total mass before a chemical reaction is
always equal to total mass after the reaction
• 2. Law of Definite Proportions
• a. A certain compound always contains the
same elements
• b. These elements are in exactly the same
proportions by mass, regardless of the size of
the sample
• 3. Law of Multiple Proportions
• a. Two or more different compounds can
be composed of the same elements
• b. The mass of the second element in the
compound compared to the mass of the
first will always be a ratio of small whole
numbers
Carbon dioxide
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C. Dalton’s Theory
1. John Dalton – English school teacher in the early
1800’s
2. Explained the Law of Conservation of Mass, the
Law of Definite Proportions, and the Law of
Multiple Proportions
a. All matter is composed of atoms
b. Atoms in an element are identical to each other
c. Atoms in one element are different from atoms of
any other element
One element
Different
elements
• d. Atoms cannot be subdivided, created or
destroyed
• e. Atoms combine in whole number ratios to
form compounds
• f. Atoms combine, separate, or rearrange in
chemical reactions
• g. By experimenting with measurements of
mass, Dalton changed Democritus’ idea into a
theory
Chapter 3 – Section II
• I. The Structure of the Atom
electrons
, protons
and neutrons
• A. The smallest particle of an element that still
contains all the properties of the element is an
atom
• B. All atoms consist of two regions
• 1. The very small, positively charged center of the
atom is the nucleus
• a. Positively charged particles in the nucleus are
protons
• b. Neutral particles in the nucleus are neutrons
• 2. Surrounding the nucleus is a very large space
containing negatively charged particles called
electrons
• 3. Protons, neutrons, and electrons are called
subatomic particles
• II. Discovery of the Electron
• A. In the 1800’s investigations were done to show
the relationship between electricity and matter
• 1. Electric charges were passed through gases at
low pressure
• 2. Gases at atmospheric pressure do not conduct
electricity very well
• 3. Gases at low pressure were contained in glass
tubes called cathode-ray tubes
• B. When current passes through a cathode-ray
tube, the surface opposite the cathode will glow
• 1. Scientists believed the glow was caused by a
stream of particles called a cathode-ray
• 2. The ray traveled from the cathode (-) to the
anode (+)
• 3. Experiments were devised to test this hypothesis
• a. The ray was deflected by a magnet just like
electricity in a wire
• b. Rays were deflected away from negative objects
• 4. From these experiments, the ray was believed
to be negative
• 5. JJ. Thomson measured the ratio of the charge
of the particles in the ray to their mass
• a. The ratio was always the same regardless of the
material used
• b. All cathode rays must, then, be made of
identical negative particles, which he called
electrons
• c. Electrons must be present in the atoms of all
elements
• d. This gave evidence that atoms are divisible
• e. The electron, then, is a subatomic particle
• C. The charge of the electron was measured by
Robert Millikan in his famous oil drop experiment
• 1. The simplest atom is the hydrogen atom
• 2. The electron was found to be ____1____
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2000
• the mass of the simplest atom
• 3. Since atoms are neutral, they must have a
positive charge to balance the negative electron
• 4. The mass of electrons is so small, that other
particles must account for the atom’s mass
• D. Thomson’s model of the atom was known as
the plum pudding model
• 1. He believed the electrons were spread out
evenly among the positive particles of the atom
• 2. His model is like the seeds in a watermelon,
where the seeds would be like the electrons
• III. Discovery of the Atomic Nucleus
• A. Ernest Rutherford, with the help of Hans
Geiger, is given credit for discovering the nucleus
of the atom
• 1. He bombarded a piece of gold foil with alpha
particles (these particles re the same thing as a
Helium nucleus)
• 2. He assumed the mass and
• charge were uniformly distri• buted throughout the atom of
• the gold foil (like Thomson’s
• plum pudding model indicated)
• 3. He expected the alpha particles to pass through
the foil with only slight deflections
• 4. Many went straight through as if going through
empty space
• 5. Some were slight deflected
• 6. He was shocked to see that some bounced
straight back, as if hitting a brick wall or being
sharply repelled
• 7. He described this phenomenon
• as if a 15-inch artillery shell were
• shot at a piece of tissue paper,
• and it bounced straight back at you
• 8. He concluded that a very small area of the atom
was composed of a very densely packed bundle of
positively charged matter
• 9. This positive matter became known as the
nucleus of the atom
• 10. The volume of the nucleus was found to be
very small compared to the area of the electron
cloud {marble (nucleus) in the middle of a football
field (electron cloud)}
• B. Neils Bohr discovered the energy levels of the
atom (locations of the electrons), and his model is
known as the planetary (or solar system) model of
the atom
• IV. Composition of Atomic Nucleus
• A. Protons are postively charged, have a mass of 1
atomic mass unit, and are located in the nucleus
• B. Neutrons are neutral, have a mass of 1 atomic
mass unit, and are located in the nucleus
• C. Electrons are negatively charged, have no mass
(1/2000 mass of proton or neutron), and are located
in the electron cloud surrounding the nucleus
• D. Atoms of different elements
• have different numbers of
• protons, and atoms of the same
• element all have the same number of protons
• E. The short-range forces that hold protons
together, neutrons together, and protons and
neutrons together are called nuclear forces
• F. The force that holds the nucleus together is
often referred to as the strong force
• G. The distance from the center of the nucleus to
the outer edge of the electron cloud is known as
the atom’s radius
• H. The unit for measuring the
• size of an atom is the picometer
Strong force
Chapter 3 – Section III
• Counting Atoms
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I. Atomic Number
A. The number of protons in an
atom is its atomic number or Z number
B. The atomic number identifies the element
C. Since atoms are neutral, the number of protons
and electrons are always the same in any atom
II. Isotopes
A. The simplest atom is hydrogen
B. All hydrogen atoms have 1 proton
C. Hydrogen atoms can have different numbers of
neutrons
• 1. When atoms of the same element have different
numbers of neutrons, they are isotopes of each other
• 2. Protium is the most common form of hydrogen
(1 proton, 0 neutrons; 99.985% of all hydrogen is
this type)
• 3. Deuterium is a hydrogen atom with 1 proton and
1 neutron (0.015% of all hydrogen is this type)
• 4. Tritium is the radioactive hydrogen atom with 1
proton and 2 neutrons
• (only trace amounts of it
• are found in nature, but
• it can be made artificially
• in the laboratory)
• III. Designating Isotopes
• A. There are two methods of indicating when an
element is an isotope
• 1. Hyphen notation – the name or symbol of the
element is written with a hyphen after it and the
mass number (U-235, U-238, C-12, C-14, B-10,
B-11 are some examples)
• 2. Nuclear symbol – the symbol of the element is
written with a superscript showing the mass
number and a subscript showing the atomic
number (U235, U238)
92
92
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• B. The term nuclide is sometimes used for the
term isotope when referring to a specific isotope
(Example: One radioactive nuclide of uranium is
U238)
• 92
• IV. Relative Atomic Masses
• A. Since the mass of an atom is so small, it is
usually more convenient to use relative atomic
masses
• B. The atomic mass unit (amu) was developed by
scientists for comparing atomic mass
• 1. The standard for comparison is the Carbon-12
atom
• 2. It is assigned a mass of exactly 12 amu
• 3. 1 amu is 1/12 the mass of a Carbon-12 atom
• 4. Since Carbon-12’s mass
• is made up of 6 protons
• and 6 neutrons, each proton
• and each neutron is a mass
• of 1 amu
• C. Isotopes of an element, while having different
masses, do not behave differently from each other
chemically
• V. Average Atomic Masses of Elements
• A. The average of the masses of all the naturally
occurring isotopes of an element is the element’s
average atomic mass
• B. Average atomic mass is rounded to 2 decimal
places when used in calculations
• VI. Relating Mass to Numbers of Atoms
• A. The amount of a substance that contains as
many particles as there are atoms in 12 grams of
Carbon-12 is called a mole
• B. The number of particles in one mole of a pure
substance is called Avogadro’s number (which is
6.022 x 1023)
• C. Another definition for mole is the amount of a
substance that contains Avogadro’s number of
particles
• D. The mass of one mole of a pure substance is its
molar mass
• 1. Units for molar mass are g/mol
• 2. Molar mass is the atomic mass of the element
from the periodic table
• How many
• moles are
• in a mole
• of moles?