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Chapter 3 Atoms
Chapter 3
Section 1 The Atom: From
Philosophical Idea to Scientific Theory
Objectives
•
Explain the law of conservation of mass, the
law of definite proportions, and the law of
multiple proportions.
•
Summarize the five essential points of
Dalton’s atomic theory.
•
Explain the relationship between Dalton’s
atomic theory and the law of conservation of
mass, the law of definite proportions, and the
law of multiple proportions.
Section 3.1 Defining the Atom
Around 460 B.C. Greek philosopher,
Democritus, developed the idea of atoms.
He asked this question:
If you break a piece of matter in
half, and then break it in half
again, how many breaks will
you have to make before you
can break it no further?
Democritus thought that it ended at some point, a smallest
possible bit of matter. He called these basic matter particles,
atoms.
Democritus’s Atomic Philosophy

Atoms were indivisible and
indestructible
Atom comes from the
Greek word atomos, meaning
“indivisible”
His ideas:
 were based on philosophy
 didn’t explain chemical behavior
 lacked experimental support
Three Atomic Laws
• Rules that explain how matter interacts in chemical reactions:
1790’s- improved balances revolutionized quantitative
analysis of chemical reactions.. allowed accurate
measure of masses of elements and compounds
1. Law of conservation of mass
mass is neither created nor destroyed during
ordinary chemical reactions or physical changes
Three Atomic Laws
• Rules that explain how matter interacts in chemical reactions:
2. 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 source
3. Law of multiple proportions
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
Chapter 3
Section 1 The Atom: From
Philosophical Idea to Scientific Theory
Law of
Conservation of
Mass
Chapter 3
Section 1 The Atom: From
Philosophical Idea to Scientific Theory
Law of Multiple Proportions
John Dalton (1766-1844)
Dalton transformed Democritus’s ideas
on atoms into a scientific theory.
 Dalton used experimental methods not
just philosophy
 Studied ratio in which elements
combine in chemical reactions
Developed an atomic theory based on results of
experiments


Note: Dalton was always very smart and at the age of twelve
he took over a Quaker school in Cumberland, England.
Designated much of his life to researching color blindness
because both he and his brother suffered from it.
Dalton’s Elements and Compounds
Dalton’s Atomic Theory (p. 64)
1. All elements are composed of
extremely small particles called
atoms.
2. Atoms of an element are
identical in size, mass and
other properties; atoms of
different elements differ in
size, mass, and other properties.
Dalton’s Atomic Theory (p. 64)
3. Atoms cannot be subdivided, created, or
destroyed (law of conservation of mass)
4. Atoms of different elements combine in
simple whole-number ratios to form
chemical compounds. (law of definite
proportions and multiple proportions)
5. In chemical reactions, atoms are combined,
separated, or rearranged. (law of
conservation of mass)
Scanning Tunneling Microscope (STM)
Gerd
Binnig
Dr. Heinrich
Rohrer
• Invented in 1981 by Gerd Binnig and Heinrich Rohrer
working for IBM in Zurich, Switzerland.
• Won 1986 Nobel Prize in physics for their discovery
• open door to field of Nanotechnology
Scanning Tunneling Microscopy
Use of Scanning Tunneling Microscopy
by IBM

Moving Atoms: Making the World's
Smallest Movie
Scanning Tunneling Microscope (STM)
 STM provides ability to look at individual atoms
 detects the electrons of surface atoms
 maps the position and spacing of atoms
Modern Atomic Theory
Not all aspects of Dalton’s atomic theory have
proven to be correct. We now know that:
Atoms are divisible into even smaller particles.
A given element can have atoms with different
masses (isotopes).
Some important concepts remain unchanged:
All matter is composed of atoms.
Atoms of any one element differ in properties from
atoms of another element.
Dalton’s Atomic Theory has not been discarded.. only modified!
Chapter 3
Atoms
Chapter 3
Section 2 The Structure of the Atom
Objectives
• Summarize the observed properties of cathode rays that
led to the discovery of the electron.
• Summarize the experiment carried out by Rutherford and
his co-workers that led to the discovery of the nucleus.
• List the properties of protons, neutrons, and electrons.
• Define atom.
• Describe nuclear forces
Section 3.2 Structure of the Atom
 One change to Dalton’s atomic theory - atoms
ARE divisible
 Into subatomic particles:
Electrons, protons, and neutrons
Atom – smallest particle of an element that retains
the chemical properties of that element
Discovery of Electron- J.J. Thomson p. 69
1897 - J.J. Thomson used cathode ray tube to show
presence of negatively charged particles...electrons
Note: Called the negatively charged particles “corpuscles”
Thomson’s Atomic Model
J. J. Thomson



cathode ray experiment showed presence of electrons
“plum pudding” model - believed e- like plums in +
charged “pudding,”
charge-to-mass ratio of an electron
Joseph John “J.J.” Thomson
1856-1940
Famous for discovery of the first subatomic particle, the electron, and for
work on the atomic model. It was soon superseded by his student Ernest
Rutherford’s nuclear model.
1906 received Nobel Prize for physics for his research into the electrical
conductivity of gases. Thomson was a great teacher and an outstanding
scientist. Seven of his students and assistants also received Nobel Prizes
for their work.
Cathode Ray Tube (CRT)


Cathode ray tube (CRT) is a vacuum tube containing an
electron gun (source of electrons) and fluorescent screen used
to view images. It accelerates and deflects the electron beam
onto fluorescent screen to create images. Can be found in
everyday devises such as televisions, video games,
computers, video cameras, monitors, automated teller
machines, oscilloscopes, and radar displays.
CRTs have been superseded by more modern display
technologies such as LCD, plasma display, and OLED, which
as of 2012 offer lower manufacturing and distribution costs.
CRT - Cathode Ray Tubes
Television
Computer Monitor
CRT’s pass electricity thru gas
contained vacuum tube
Charge and Mass of Electron- Robert Millikan
Charge of the
electron is
1.592 × 10−19 coulomb
The oil drop apparatus

determined the actual e- charge: 1.592 × 10−19 coulomb)
(Somewhat lower than the modern of 1.602 × 10 −19 coulomb)

determined actual e- mass: 1/1837 the mass of H atom
(9.11 x 10-28 g)
Robert Millikan's Oil drop experiment
Robert Millikan (1868-1953)
 Doctorate in Physics from Columbia University
 1923 Nobel Prize for Physics
 Determined charge on the electron as well as
Planck’s constant
Millikan set up his experiment by spraying oil drops, which
fell between positively charged plate and negatively
charged plate. He balanced the electric charge with
electron's weight, as gravity pulled down on the oil drop,
and the electric charges attracted and repelled it. Using
this method, he discovered that the charge of an electron
is negative. The weight of each drop was determined by
observing it's rate of free fall through the air. Originally, he
sprayed droplets of water through an electrical field that
was intended to suspend them in midair, but the water
droplets evaporated too fast, so Millikan switched to using
oil.
Robert Millikan
Starting in 1908, while a professor at the University of Chicago, Millikan
worked on an oil-drop experiment in which he measured the charge on a
single electron. J.J. Thomason had already discovered the charge-to-mass
ratio of the electron. However, the actual charge and mass values were
unknown. Therefore, if one of these two values were to be discovered, the
other could easily be calculated. Millikan and his then graduate student
Harvey Fletcher used the oil-drop experiment to measure the charge of the
electron (as well as the electron mass, and Avogadro’s number, since their
relation to the electron charge was known). Millikan showed that the results
could be explained as integer multiples of a common value (1.592 × 10−19
coulomb), the charge on a single electron. That this is somewhat lower than
the modern value of 1.602 176 53(14) x 10−19 coulomb is probably due to
Millikan's use of an inaccurate value for the viscosity of air.
Conclusions from Electron study:
 Because atoms are electrically neutral,
they must contain a positive charge to
balance the negative e-
 Because electrons have so much less
mass than atoms, atoms must contain
other particles that account for most of
their mass.
Discovery of Protons: Eugen Goldstein
 1886 - Eugen Goldstein observed +
particles (now called “proton”)
 particles with +charge & relative mass
of 1 (or 1840x e-)
 called them “Canal rays” (anode rays)
Eugen Goldstein (1850-1930)



Credited by many for the discovery of
the existence of protons
discoverer of anode rays (canal rays)
Worked in Berlin and Potsdam
Observatories
Ernest Rutherford’s
1911- Gold Foil Experiment
 Alpha particles fired at gold foil
 Particles hitting detecting screen were recorded
Discovery of Nucleus: Ernest Rutherford
 most particles passed thru
“It was as if you fired a 15-in
 few deflected
[artillery] shell at a piece of
 VERY FEW greatly deflected tissue paper and it came
back and hit you.”
Conclusions:
a) small, dense nucleus
b) + charged nucleus
Rutherford’s Atomic Model
Experimental evidence:

atom mostly empty space
 almost all mass in nucleus
 e- distributed around nucleus…occupy most of the
volume
 called “nuclear model”
Rutherford’s Atom 3:08
Discovery of Neutrons: James Chadwick
1932 – James Chadwick confirmed existence of
the “neutron”
 particle with no charge
 n0 mass = p+ mass
1935 - Nobel Prize for Physics
Prepared the way for the fission of Uranium-235 which led to developing
atomic bomb (Manhattan Project)
Atomic Scientists
• Democritus named the atom
• John Dalton proposed the Atomic Theory
• J.J. Thomson’s cathode-ray tube experiments measured
the charge-to-mass ratio of an electron. (plum-pudding
model)
• Robert A. Millikan’s oil drop experiment measured the
charge of an electron
With this information, scientists were able to determine
the mass of an electron.
Atomic Scientists
• Rutherford’s gold foil experiment led to the discovery
of the nucleus.. a very densely packed bundle of
matter with a positive electric charge.
• James Chadwick proved existence of neutrons.. no
charge
Subatomic Particles
Particle
Electron
(e-)
Proton
(p+)
Neutron
(no)
Charge
Mass
Number
Actual Mass
(g)
Location
Outside
Nucleus
1-
0
9.11 x 10-28
1+
1
1.67 x 10-24
Nucleus
0
1
1.67 x 10-24
Nucleus
(Electron
cloud)
Chapter 3
Section 2 The Structure of the Atom
The Structure of the Atom

An atom is the smallest particle of an element that
retains the chemical properties of that element.

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.
Chapter 3
Section 2 The Structure of the Atom
The Structure of the Atom

Surrounding the nucleus is a region occupied by
negatively charged particles called electrons.
• Protons, neutrons, and electrons are often referred to
as subatomic particles.
Chapter 3
Section 2 The Structure of the Atom
The Structure of the Atom
• Atoms are electrically neutral because they contain
equal numbers of protons and electrons
• The nuclei of atoms of different elements differ in their
number of protons and therefore in the amount of
positive charge they possess.
• Thus, the number of protons determines that atom’s
identity.
Chapter 3
Section 2 The Structure of the Atom
Properties of Subatomic Particles
Chapter 3
Section 2 The Structure of the Atom
Forces in 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 when protons and neutrons are very
close together.
The short-range proton-neutron, proton-proton, and
neutron-neutron forces that hold the nuclear particles
together are referred to as nuclear forces.
Chapter 3
Section 2 The Structure of the Atom
The Sizes of Atoms
• The radius of an atom is the distance from the center
of the nucleus to the outer portion of its electron
cloud.
• Because atomic radii are so small, they are expressed
using a unit that is more convenient for the sizes of
atoms.
• This unit is the picometer, pm.
Objectives
• Explain what isotopes are.
• Define atomic number and mass number, and
describe how they apply to isotopes.
• Given the identity of a nuclide, determine its number
of protons, neutrons, and electrons.
• Define mole, Avogadro’s number, and molar mass,
and state how all three are related.
• Solve problems involving mass in grams, amount
in moles, and number of atoms of an element.
Atomic Number
• Atoms of different elements have different numbers of
protons.
• Atoms of the same element all have the same number
of protons.
• The atomic number (Z) of an element is the number
of protons of each atom of that element.
Atomic Number
Atomic Number
Atomic number (Z) of element = # of p+ in
nucleus of each atom of that element.
Element
Carbon
# of protons Atomic # (Z)
# of electrons
6
6
Phosphorus
15
15
Gold
79
79
electrically neutral atoms: # of p+ = # of e+
Isotopes and Mass Number
• Isotopes are atoms of the same element that have
different masses.
• The isotopes of a particular element all have the same
number of protons and electrons but different numbers
of neutrons.
• Most of the elements consist of mixtures of isotopes.
• The mass number is the total number of protons
and neutrons that make up the nucleus of an
isotope.
Mass Number
Designating Isotopes
• Hyphen notation: The mass number is written with a
hyphen after the name of the element.
• uranium-235
• Nuclear symbol: The superscript indicates the mass
number and the subscript indicates the atomic
number.
235
92
U
• Nuclide is a general term for a specific isotope of an
element.
Mass Number
Mass number is # of p+ and n0 in nucleus of an isotope:
Mass # = p+ + n0
p+
n0
e- Mass #
Oxygen - 18
8
10
8
18
Arsenic - 75
33
42
33
75
Phosphorus - 31
15
16
15
31
Isotope
Sample Problem A
How many protons, electrons, and neutrons are there in an
atom of chlorine-37?
Atomic number = 17
made up of 17 protons, electrons, 17.
Mass number = 37
mass number − atomic number = 37 (p+ plus n0) − 17 p+ = 20 n0
An atom of chlorine-37 is made up of 17 e-, 17 p+, and 20 n0.
Complete Symbols
Superscript →
Mass
number
Subscript →
Atomic
number
Complete Symbols
Find each of these:
a)
b)
c)
d)
e)
number of protons 35
number of neutrons 45
number of electrons 35
Atomic number 35
Mass Number 80
80
35
Br
Complete Symbols
If an element has an atomic number of 34 and a
mass number of 79, what is the:
a)
b)
c)
d)
number of protons 34
number of neutrons 45
number of electrons 34
complete symbol
79
34
Se
Symbols
 If an element has 91
protons and 140 neutrons
what is the
a) Atomic number
b) Mass number
c) number of electrons
d) complete symbol
Symbols
 If an element has 78
electrons and 117 neutrons
what is the
a) Atomic number
b) Mass number
c) number of protons
d) complete symbol
Isotopes
 Dalton
wrong about elements of
same type being identical…
 Element’s
atoms can have
different # neutrons.
different
mass #
Isotopes
 (flavors)
Isotopes
 Frederick
Soddy (1877-1956)
proposed idea of isotopes in 1912

Isotopes - atoms of same element with
different masses, b/c varying #s of
neutrons
 1921
Nobel Prize in Chemistry
 crater w/ his name on far side of Moon
Naming Isotopes
We
can also put mass
number after name of
element:
carbon-12
carbon-14
uranium-235
Isotope
Protons Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
1
1
2
Hydrogen-3
(tritium)
Nucleus
Isotopes
Elements
occur in
nature as
mixtures of
isotopes.
Atomic Mass



How heavy is an oxygen atom?
 Depends, b/c different kinds of oxygen
atoms exist.
We’re more concerned with average
atomic mass.
Based on abundance (%) of each variety
of that element in nature.

Don’t use grams - #’s tooooo small
Measuring Atomic Mass
 Atomic
Mass Unit (amu)
 one-twelfth mass of C-12 atom
 Carbon-12
 Each
chosen b/c of its isotope purity
isotope has own atomic mass
 we determine average from %
abundance.
Calculate the average:
Multiply
atomic mass of each
isotope by abundance
(decimal), then add results.
 If
not told otherwise, mass of isotope
expressed in atomic mass units
(amu)
Atomic Masses
Atomic mass is average of all naturally
occurring isotopes of that element.
Isotope
Symbol
Carbon-12
12C
Carbon-13
13C
Carbon-14
14C
Composition
of the nucleus
6 protons
6 neutrons
6 protons
7 neutrons
6 protons
8 neutrons
% in
nature
98.89%
Atomic
mass
(amu)
12.00
1.11%
13.00
<0.01%
14.00
What is the average atomic mass of Carbon?
12.01