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Chapter 3: The Atom
A History of Atomic Theory & Basic Atomic Structure
Big Idea: Physical, chemical and nuclear changes are explained
using the location and properties of subatomic particles.
The Atom: From Philosophical
Idea to Scientific Theory
Section 3.1
Democritus – 450 BC
 All matter is made up of tiny, indivisible particles
called atomos
 Today, we define atom as the smallest particle
of an element that retains the chemical identity
of that element
 Aristotle asked:
What holds the
tiny particles
together?
Democritus: ???
Aristotle – 384 BC
 Aristotle rejected Democritus reasoning and proposed
that matter was a continuum composed of mass and
form
Marble (mass)  Statue (form)
 Later the simplest forms of matter were proposed to be:
Earth,
Water,
Fire,
Wind
Foundations of Atomic Theory
 The transformation of a substance(s) into one or
more new substances is known as a chemical
reaction.
 Law of Definite Composition: a chemical
compound contains the same elements in exactly
the same proportions by mass (regardless sample
size or source)
Sugar: 42.1 % Carbon
51.4 % Oxygen
6.5 % Hydrogen
Whether you have a teaspoon or a truckload!
Foundations of Atomic Theory
 Law of Conservation of Mass: mass is neither
created nor destroyed during ordinary chemical
reactions or physical changes
HgO  Hg + O
433.2 g
401.2g + 32g
 Law of Multiple Proportions: if two or more
different compounds are composed of the same
two elements, then the ratio of the masses of the
second element combined with a certain mass of
the first element is always a ratio of small whole
numbers
H2O
Water
H2O2
Peroxide
2g H
16g O
2g H
32g O
1:2 Ratio
2
Dalton - 1808
Dalton’s Atomic Theory
All matter is composed of extremely small particles
called atoms.
Atoms of a given element are identical in size,
mass, and other properties; atoms of different
elements differ in size, mass, and other properties.
Atoms cannot be subdivided, created, or destroyed.
satisfies Law of Conservation of Mass
Atoms of different elements combine in simple
whole-number ratios to form chemical compounds.
satisfies Law of Multiple Proportions
In chemical reactions, atoms are combined,
separated, or rearranged.
Modern Atomic Theory
 Not all aspects of Dalton’s atomic theory have
proven to be correct.
 Atoms can be split into even smaller particles.
 A given element can have atoms with different masses
(called isotopes)
 Some important concepts remain unchanged
 All matter is composed of atoms
 Atoms are rearranged in chemical reactions
 Atoms of any one element are never identical to
atoms of another element
The Structure of the Atom
Section 3.2
The Discovery of the Electron
 Atom is the smallest particle of an element that retains
the chemical properties of that element.
 1897 - Joseph John Thomson’s cathode-ray tube (CRT)
The Discovery of the Electron
Joseph John Thomson’s cathode-ray tube (CRT)
Negatively Charged
Electrode
Positively Charged
Electrode
Cathode Ray Tube
 Scientists studied the flow of electric current in a glass
vacuum tube with electrodes at each end.
 When connected to electric current the remaining gas
glowed forming a BEAM OF LIGHT.
 The beam always originated at the NEGATIVE electrode
and toward the POSITIVE electrode.
 The electrode is named by what type of particle it attracts
 Cathode: Negative (-)
 Anode: Positive (+)
 JJ Thomson used magnets to deflect the beam
proving that particles had a negative charge.
 These negatively charges particles were called
electrons.
 Major contribution to the atom:
 Electrons are in all atoms!
CRT Video
 Cathode Ray Tube Experiment
Robert A. Millikan - 1909
 Continued Thomson’s work –
 performed the Oil Drop Experiment
 confirmed the negative charge of an electron and
measured the mass of an electron
 The electron has mass, though 1836 x less than that of
soon to be discovered proton.
J.J. Thomson’s Plum Pudding Model
 Thomson proposed that the electrons of an atom
were spread evenly throughout a positively charged
ball of matter.
 Known as Plum-pudding model
Plum Pudding Video
The Discovery of the Atomic Nucleus
 Earnest Rutherford’s Gold Foil Experiment - 1909
Gold Foil Experiment
 Set up Gold Foil with a detection sheet around it.
 Set up radioactive source- emitted alpha particles.
 ALPHA PARTICLES shot at gold foil.
 MOST particles went through the gold foil
 But SOME particles BOUNCED back
Gold Foil Conclusions
1.
2.
3.

The atom is made up of mostly EMPTY SPACE
The center of the atom contains a POSITIVE
CHARGE
Rutherford called this positive bundle of matter
the NUCLEUS
Rutherford’s major contribution to the atom was
the discovery of the nucleus. The volume of this is
very small compared with the total volume of an
atom.
Rutherford’s Model of
the Atom
atomic radius ~ 100 pm = 1 x 10-10 m
nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
Composition of the Nucleus
 When two protons are extremely close to each
other, there is a strong attraction between them.
 A similar attraction exists when neutrons are very
close to each other or protons and neutrons.
 The short-range proton-neutron, proton-proton,
and neutron-neutron forces that hold the nuclear
particles together are referred to as nuclear forces.
Structure of the Atom
 The nucleus is a very small region located at the
center of an atom.
 The nucleus is made up of at least one positively
charged particle called a proton and usually one or
more neutral particles called neutrons.
 Surrounding the nucleus is a region occupied by
negatively charged particles called electrons.
 P, N, E are often referred to as subatomic particles.
Particle
Electron
Proton
Neutron
Symbol
ep+
no
Charge Mass
Actual Mass
Number (kg)
-1
0
9.109 x 10-31
+1
1
1.673 x 10-27
0
1
1.675 x 10-27
Ions and Isotopes
Section 3.3
Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei
Mass Number
A
Z
Atomic Number
1
1H
2
1H
235
92
U
X
Element Symbol
3
1H
(D)
238
92
U
(T)
An ion is an atom, or group of atoms, that has a net positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
Cl-
17 protons
18 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
Relative Atomic Mass
 One atom is the standard – Carbon
 Mass of other elements are based off of the standard
 Carbon: 6 p and 6 n = 12 amu
Atomic Mass Unit
 1/12 mass of Carbon atom
 Periodic table lists weighted average atomic masses of elements
(like a GPA calculation)