Chapter 4: The
Structure of the
Atom
Chapter Big Idea
Atoms are the fundamental
blocks of matter…
Section 1: Early Ideas About Matter
Section 1: Essential Questions & Vocabulary
• What are the similarities and differences of the atomic
models of Democritus, Aristotle, and Dalton?
• How was Dalton’s theory used to explain the conservation
of mass?
Vocabulary
• Dalton’s atomic theory
• Theory
Section 1: Big Idea
The ancient Greeks tried to explain
matter, but the scientific study of the
atom began with John Dalton in the
early 1800’s.
Roots of Atomic Theory
• How was science thousands of years
ago different from science now?
• Lacked controlled experimentation & tools for
scientific investigations
• Intellectual thought as truth
Roots of Atomic Theory
• Philosophers – scholarly thinkers
• Speculated about the nature of matter & formulated their own
explanations based on their own life experiences
• Common Conclusions:
• Matter composed of things such as earth, water, air, and fire.
• Matter could easily be divided into smaller and smaller pieces.
Democritus (460 BC -370 BC)
• Greek philosopher
• First to propose that matter was NOT
infinitely divisible.
• Atomos- Greek word meaning INDIVISIBLE.
• Matter is composed of atoms which move
through empty space
• Size, shape, and movement of atoms
determine the properties of matter
Democritus – Atomic Model & Analogy
~ 400 BC
Atomic Model
Analogy
Atoms are small, hard particles
that are all made of the same
material, but are formed into
different shapes and sizes
Legos
Democritus – Criticism
• What holds the atoms together?
THINK – PAIR - SHARE
• Why do you think it was hard for Democritus to
defend his ideas?
• Lack of experimentation
• Ahead of his time
Aristotle (384 – 322 BC)
• One of the most influential Greek
Philosophers
• Rejected Democritus’ notion of atoms
because it contradicted his own ideas about
nature.
• Because he was so influential, this led to
Democritus’ atomic theory to be rejected.
Aristotle – Atomic Model & Analogy
~ 300 BC – 1800’s
Atomic Model
Analogy
• All matter was made of only four
elements & four Properties
Death of Chemistry for 2000
• Fire, air, water, and earth
years!!!
• Hot, cold, dry, and wet
John Dalton (1766 – 1844)
• Marks the beginning of modern atomic
theory
• Revived Democritus’ idea of atoms
based on the results of his scientific
research
• Studied numerous chemical reactions
• Determined the mass ratios of the elements involved
in those reactions
Dalton’s Atomic Theory – 1803
• Matter is composed of extremely small particles called atoms which are
indivisible and indestructible.
• Atoms of a given element are identical in size, mass, and chemical
properties. Atoms of a specific element are different from those of another
element.
• Different atoms combine in simple whole-number ratios to form
compounds.
• In a chemical reaction, atoms are separated, combined, or rearranged.
Dalton – Atomic Model & Analogy
1803
•
•
•
•
Atomic Model
Analogy
All matter is made of atoms.
All atoms of a given element are alike, atoms of
different elements are different.
Atoms combine in whole-number ratios.
In chemical reactions, atoms are separated,
combined, or rearranged.
Billiard Ball
Conservation of Mass
• Law of conservation of Mass- mass is conserved in any
process.
• Dalton’s Atomic Theory – explains the conservation of mass in
chemical reactions as the result of SEPARATION,
COMBINATION, or REARRANGEMENT of atoms.
• Atoms are NOT created, destroyed or divided in the process.
Dalton’s Theory : Practice Question
• Which of these reactions show Dalton’s Theory?
Dalton’s Theory : Practice Question
Solution
• Which of these reactions show Dalton’s Theory?
Dalton’s Theory: Practice Question
• Six atoms of Element A combine with eight atoms of
Element B to produce six compound particles.
• How many atoms of Element A does each compound
particle contain?
• How many atoms of element B does each compound
particle contain?
• Are all of the atoms used to form compounds?
Dalton’s Theory Practice Solution
Element A (6)
Element B (8)
Compound (6 units)
• Only have 6 A Elements – can have up to 6 compound units
• Although you have 8 B Elements – not enough for 2 per
compound
• Must have 2 leftover B Elements
Section 2: Defining the Atom
Section 2: Essential Questions & Vocabulary
• What is an atom?
• How can the subatomic particles be distinguished in terms of
relative charge and mass?
• Where are the locations of the subatomic particles within the
structure of the atom?
Vocabulary
• Atom
• Nucleus
• Cathode ray
• Proton
• Electron
• Neutron
• Model
Section 2: Main Idea
An atom is made up of a
nucleus containing protons
and neutrons; electrons move
around the nucleus.
What is an Atom?
• Smallest particle of an element that retains the
properties of the element.
• How big is an atom?
• ~1.3 x 10-10 m
If you could increase
the size of an atom to
make it as big as an
orange. In this new
scale, the orange
would be as big as
Earth!!
WHAT IS AN ATOM LIKE?
Now that scientists were convinced on the
existence of atoms, a new set of questions
arises!!!
Cathode Ray Tube
• When an electric charge is applied, a ray of radiation travels from the
cathode to the anode, called a cathode ray.
• Cathode rays are a stream of particles carrying a negative charge.
Link to video on Cathode Ray Tubes by clicking on the image
J.J Thomson (1856-1940)
• Completed a series of cathode-ray tube
experiments at Cambridge University.
• Studied mathematics and physics
• Won the Nobel Prize in 1906
JJ Thomson: Discovery of the Electron
1897

Measured the effects of both magnetic and electric
fields on the cathode ray to determine the charge to
mass ratio.
 Particles that compose cathode rays are negatively charged
 Charge to mass ratio of the cathode-ray is always the same.
 Concluded that all cathode rays are composed of identical
negatively charged particles called ELECTRONS
 Experiments revealed the electron has a very large charge for
its tiny mass
PROOF –
THERE MUST
BE A
PARTICLE
SMALLER
THAN THE
ATOM!!
JJ Thomson – Atomic Model & Analogy
1897
Atomic Model
•
•
Discovered the presence of a negative
particles in the atom.
Atoms are made from a positively charged
substance with negatively charged electrons
scattered about, like raisins in a pudding or
chocolate chips in a cookie.
Analogy
Plum Pudding
Or
Chocolate Chip Cookie
Robert Millikan (1868 – 1953)
• American Physicist
• Determined the charge of an electron
• Nobel Prize Winner in 1923
Oil Drop Experiment – Charge of an Electron
~1910
• Oil Droplets are dropped
• X-Ray knocks electrons from air which
then attach to the oil droplets.
• Millikan could vary the electric field
strength to makes the oil droplets move
more slowly, rise, or become suspended.
• Could calculate the charge on the
droplets based on their rate of fall.
• The magnitude of charge on the drops
always changed by a discrete amount
• Smallest common denominator = 1.602 x 10 -19 C
The Electron – Mass???
• JJ Thomson - charge to mass ratio of the electron.
• Millikan - charge of the electron.
• CAN THOMSON’S INFORMATION ALONG WITH
MILLIKAN’S BE USED TO GET THE MASS OF THE
ELECTRON?
Mass of electron = 9.1 x 10 -28 g
Plum Pudding Model
New questions arise
• Recap of what we know:
• Matter is neutral!
• All matter is made up of atoms which have electrons (- charge)
• Electrons are much lighter than the lightest atom known!
• Questions:
• If electrons are part of all matter and they possess a
negative charge, how can matter be neutral?
• If the mass of an electron is so small, what accounts
for the mass of a typical atom?
Ernest Rutherford (1871-1937)
• JJ Thomson’s student
• Early work – discovered radioactive half-life
• Nobel Prize in Chemistry – 1908
• Father of nuclear physics
• Studied how alpha particles interacted with matter
• Alpha particles – positively charged particles
Rutherford’s Gold Foil Experiment
• Aimed a narrow beam of alpha particles at a thin sheet of gold foil.
• A flash of light is produced when the particle strikes the gold foil
• WHAT DID HE EXPECT?
• Positive charge is evenly distributed
• Path of α – particle should not be altered
• α – should continue in a straight path
Rutherford’s Gold Foil Experiment
What actually happened!!
Rutherford’s Gold Foil Experiment
Conclusions
• Atoms consist of mostly empty space
• Almost all of the atom’s positive charge
and mass is contained in a tiny dense
region at the center of the atom
• He called this dense region – the NUCLEUS
• Electrons are held within the atom by their
attraction to the positively charged
nucleus
• Opposite charges attract
Rutherford – Atomic Model & Analogy
1911
Atomic Model
•
•
Atoms have a small, dense, positively
charged center that he called NUCLEUS
Nucleus is tiny compared to the atom
because the atom is mostly empty space
Analogy
Cherry with a Pit
James Chadwick (1891-1974)
1932 - Neutron
• Worked with Rutherford (after WWI)
• Proved the existence of neutrons
• Elementary charges devoid of any electrical
charge.
• Nobel Prize in Physics – 1935
• Manhattan Project – Development of
Atomic Bomb
Bohr – Atomic Model & Analogy
1913
•
•
Atomic Model
Analogy
Theorized electrons move in definite orbits
around the nucleus like planets circle the
sun.
Energy levels are located at certain
distances from the nucleus.
Solar System
Modern – Atomic Model & Analogy
Late 20th Century – 21st Century
Atomic Model
•
•
•
•
Schrodinger, Heisenberg, Einstein & many other scientists
Electrons move at high speeds in an electron could around
the nucleus
In the ELECTRON CLOUD, electrons orbit around the nucleus
billions of times in one second
Electron’s motion is dependent on the AMOUNT of ENERGY
they contain
Analogy
Cotton Balls
Atomic Theory Ted Ed talk !!
Completing the Model of the Atom
• All atoms have 3 fundamental subatomic particles
• Protons, neutrons, electrons
• Atoms are spherically shaped, with small, dense positively charged nucleus
surrounded by negatively charged electrons.
• Most of an atom consists of fast moving electrons that are held within the atom
by their attraction to the positively charged nucleus.
• Nucleus is composed of neutrons (neutral charge) and protons (positive charge)
• Scientists have determined that protons and neutrons are composed of “quarks”
• Scientists unsure if and how “quarks” affect chemical behavior.
• Chemical behavior can be explained by considering only the atom’s electrons.
Section 3: How Atoms Differ
Section 3: Essential Questions & Vocabulary
• His the atomic number used to determine the identity of an atom?
• What is an isotope?
• Why are atomic masses not whole numbers?
• Given the mass number and the atomic number, how are the number of
electrons, protons, and neutrons in an atom calculated?
Vocabulary
• Atomic Number
• Atomic Mass Unit (AMU)
• Isotope
• Atomic Mass
• Mass Number
Atomic Number
• ATOMIC NUMBER – the number of protons in the nucleus of an atom
• the number of protons in an atom identifies it as an atom of a particular
element.
• All atoms are neutral.
• So the number of protons
must be equal to the
number of electrons.
Atomic Number Practice
Atomic Number Practice
Atomic Number Practice SOLUTIONS
6
Oxygen
Sodium
8
11
17
Chlorine
Uranium
6
92
Mass Number
• Sum of the number of protons and neutrons in the nucleus
• Always a whole number.
• Not on the periodic table.
© Addison-Wesley Publishing Company, Inc.
Mass Number : Practice
Isotopes
• Atoms of the same element with different mass numbers.
 Nuclear symbol:
Mass #
Atomic #
12
6
 Hyphen notation: carbon-12
C
Isotope Question?
• If isotopes are the same element but have different mass
number, what is different in each isotope?
• Number of Neutrons
© Addison-Wesley Publishing Company, Inc.
Natural Abundance of Isotopes
• In nature, most elements are found as mixtures of isotopes.
• Usually, the relative abundance of each isotope is constant
regardless of where the element is obtained.
Relative Atomic Mass
• 12C atom = 1.992 × 10-23 g
 atomic mass unit (amu)
 1 amu = 1/12 the mass of a
12C
atom
1p
= 1.007276 amu
1 n = 1.008665 amu
1 e- = 0.0005486 amu
© Addison-Wesley Publishing Company, Inc.
C. Johannesson
Atomic Mass
• Weighted AVERAGE of all isotopes of that element
• Atomic mass found on the Periodic Table
• Round to 2 decimal places
Write the relative
abundance
percentage as a
decimal.
Average Atomic Mass : Practice
• EX: Calculate the avg. atomic mass of oxygen if its abundance in
nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.
Average Atomic Mass Practice
Solution
• EX: Calculate the avg. atomic mass of oxygen if its abundance in
nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.
+
16O
= 16 x 0.9976 = 15.9616
17O
= 17 x 0.0004 = 0.0068
18O
= 18 x 0.0020 = 0.0360
Average Atomic Mass for O = 16. 00 amu
Add all
isotope
masses
together!!
Average Atomic Mass : Practice
• EX: Find chlorine’s average atomic mass if approximately 8 of every 10
atoms are chlorine-35 and 2 are chlorine-37.
Average Atomic Mass : Practice
Solution
• EX: Find chlorine’s average atomic mass if approximately 8 of every 10
atoms are chlorine-35 and 2 are chlorine-37.
1. Find the relative percentage of each isotope
• Chlorine – 35 : 8 out of 10 atoms
% 𝑨𝒃𝒖𝒏𝒅𝒂𝒏𝒄𝒆 =
•
Chlorine – 37 : 2 out of 10 atoms
% 𝑨𝒃𝒖𝒏𝒅𝒂𝒏𝒄𝒆 =
𝟖 𝒂𝒕𝒐𝒎𝒔
𝟏𝟎 𝒂𝒕𝒐𝒎𝒔
𝒙 𝟏𝟎𝟎% = 80 %
𝟐 𝒂𝒕𝒐𝒎𝒔
𝟏𝟎 𝒂𝒕𝒐𝒎𝒔
𝒙 𝟏𝟎𝟎% = 20 %
2. Calculate the average atomic mass
35Cl
= 0.80 x 35 = 28.0
37Cl = 0.20 x 37 = 7.4
+
Average Atomic Mass = 35.4 amu