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Chapter 0
A Very Brief History of
Chemistry
Chemistry: The Molecular Nature
of Matter, 7E
Jespersen/Hyslop
Jespersen/Hyslop, Chemistry7E, Copyright © 2015 John Wiley & Sons, Inc. All Rights Reserved
Chapter in Context
 Scope and purpose of the chemical sciences –
Chemistry’s big ideas
 Formation of the elements – supernovas
 Distribution of substances around the world –
elements and the Earth
 Atomic theory – explanatory and predictive
power
 The structure of the atom – key experiments
Jespersen/Hyslop, Chemistry7E, Copyright © 2015 John Wiley & Sons, Inc. All Rights Reserved
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The Four Big Ideas
1. Atomic theory – John Dalton, 1813
Described atoms and how they interact with one
another
2. Careful laboratory observation
Can lead to understanding the atomic world
3. Energy changes and probability
Lead to predictions about chemical interactions
4. Geometric shapes of molecules are important
Affect properties, reactivity, and function (e.g.
DNA, RNA, and proteins)
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Supernovas and the Elements
The Big Bang:
 14 billion years ago
 Explosion of energy and subatomic particles
 Extreme temperature, pressure, and density
 As the Big Bang cooled
 Initially only quarks exist
 After 1 second, quarks form protons and neutrons
 After 3 minutes, nucleosynthesis begins of light
nuclei (e.g., helium, lithium)
 After further cooling, electrons join nuclei to form
atoms.
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Supernovas and the Elements
 Universe was 91% hydrogen, 8% helium, 1%
other light atoms
 Uneven distribution of matter resulted in star
formation
 Formation of elements occurred in the stars
 Small atoms combined, due to high pressure at
the center, to create slightly heavier elements
 New elements concentrated in the stars’ centers
 Heavier elements then combined into new, even
heavier elements
 Cycle kept repeating
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Supernovas and the Elements
 Iron is the heaviest element created in stars
 Causes the nuclear reactions to stop and the
star to cool and collapse in on itself
 allows for even heavier elements to form
 Eventually, the star disintegrates





Called a supernova
Spews its content into space
Remnants rejoin to form new stars
Cycle begins again
Some of the debris combines to form moons,
planets, and asteroids
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Distribution of the Elements
 Earth formed 4.5 billion years ago
 Result of gravitational forces
 Earth heated up





Iron and nickel melted
Migrated to the core
Outer core is superheated lava
Mantel is superheated rock
Crust is the surface
 10 miles thick
 Contains the familiar elements (gold, silicon,
carbon, etc.)
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Atomic Theory
 Most significant theoretical model of nature
Atoms




Tiny submicroscopic particles
Make up all chemical substances
Make up everything in macroscopic world
Smallest particle that has all properties of given
element
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Atomic Theory
Three important ideas
1. Law of Definite Proportions


In a given compound, the elements are always
combined in the same proportion by mass
Always find 1 g H to 8 g O in water
2. Law of Conservation of Mass
 No detectable gain or loss of mass occurs in
chemical reactions. Mass is conserved.
 A closed vessel with 16 g O and 2 g H will weigh
18 g after water is formed from them
3. Dalton’s atomic theory
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Atomic Theory
Law of definite proportions and law of
conservation of mass
 Based on laboratory observations of mass and
volume
 Discussed in detail in a later chapter
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Dalton’s Atomic Theory
John Dalton
 Developed underlying theory to explain
 Law of Conservation of Mass
 Law of Definite Proportions
 Reasoned that if atoms exist, they have
certain properties
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Dalton’s Atomic Theory (cont.)
1. Matter consists of tiny particles called atoms
2. Atoms are indestructible
 In chemical reactions, atoms rearrange but
do not break apart
3. In any sample of a pure element, all atoms
are identical in mass and other properties
4. Atoms of different elements differ in mass
and other properties
5. In a given compound, constituent atoms
are always present in same fixed
numerical ratio
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Proof of Atoms
 Early 1980’s, use
Scanning Tunneling
Microscope (STM)
 Surface can be
scanned for
topographical
information
 Image for all matter
shows spherical
regions of matter
STM of palladium
 Proof of atoms
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Discovery of Subatomic Particles
 Late 1800s and early 1900s
 Cathode ray tube experiments showed that
atoms are made up of subatomic particles
 Discovered negatively charged particles
moving from the cathode to the anode
 Cathode – negative electrode
 Anode – positive electrode
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Discovery of Electron
JJ Thomson (1897)
 Modified cathode ray tube
 Made quantitative
measurements on
cathode rays
 Discovered negatively
charged particles
 Electrons (e –)
 Determined charge to mass ratio (e/m) of these
particles
 e/m = –1.76 x 108 coulombs/gram
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Millikan Oil Drop Experiment
 Determining charge on Electron
 Calculated charge on electron
 e – = –1.60 × 10–19 Coulombs
 Combined with Thomson’s experiment to get
mass of electron
 m = 9.09 × 10–28 g
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Discovery of Atomic Nucleus
Rutherford’s Alpha Scattering Experiment




Most alpha () rays passed right through gold
A few were deflected off at an angle
1 in 8000 bounced back towards alpha ray source
Gave us current model of nuclear atom
Jespersen/Hyslop, Chemistry7E, Copyright © 2015 John Wiley & Sons, Inc. All Rights Reserved
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Discovery of Proton
 Discovered in 1918 in Ernest Rutherford’s lab
 Detected using a mass spectrometer
 Hydrogen had mass 1800 times the electron mass
 Masses of other gases whole number multiples of
mass of hydrogen
Proton
 Smallest
positively
charged particle
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Rutherford’s Nuclear Atom
 Demonstrated that nucleus:
 has almost all of mass in atom
 has all of positive charge
 is located in very small volume at center of atom
 Very tiny, extremely dense core of atom
1
 Where protons (1 p ) and
1
neutrons (0 n ) are
located
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Discovery of Neutron
 First postulated by Rutherford and coworkers
 Estimated number of positive charges on nucleus
based on experimental data
 Nuclear mass based on this number of
protons always far short of actual mass
 About ½ actual mass
 Therefore, must be another type of particle
 Has mass about same as proton
 Electrically neutral
 Discovered in 1932 by James Chadwick
Jespersen/Hyslop, Chemistry7E, Copyright © 2015 John Wiley & Sons, Inc. All Rights Reserved
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Atomic Structure
0
(1 e ,
 Electrons
or e –)
 Very low mass
 Occupy most of atom’s space
 Balance of attractive and repulsive forces controls
atom size
1
 Attraction between protons ( 1 p ) and
0
electrons ( 1 e ) holds electrons around nucleus
 Repulsion between electrons helps them spread out
over volume of atom
 In neutral atom
 Number of electrons must equal number of protons
 Diameter of atom ~10,000 × diameter of nucleus
Jespersen/Hyslop, Chemistry7E, Copyright © 2015 John Wiley & Sons, Inc. All Rights Reserved
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Properties of Subatomic Particles
 Three kinds of subatomic particles of principal
interest to chemists
Particle
Mass (g)
Electrical
Charge
Symbol
10–28
–1
Proton
1.673  10–24
+1
0
1 e
1
1
H
,
1
1p
Neutron
1.675  10–24
0
1
0n
Electron
9.109 
Jespersen/Hyslop, Chemistry7E, Copyright © 2015 John Wiley & Sons, Inc. All Rights Reserved
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Atomic Notation
Atomic number (Z)
 Number of protons that atom has in nucleus
 Unique to each type of element
 Element is substance whose atoms all contain
identical number of protons
 Z = number of protons
Isotopes
 Atoms of same element with different masses
 Same number of protons ( 11 p )
 Different number of neutrons ( 01n )
Jespersen/Hyslop, Chemistry7E, Copyright © 2015 John Wiley & Sons, Inc. All Rights Reserved
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Atomic Notation
Isotope Mass number (A)
 A = (number of protons) + (number of neutrons)
 A=Z+N
 For charge neutrality, number of electrons and
protons must be equal
Atomic Symbols
 Summarize information about subatomic particles
 Every isotope defined by two numbers Z and A
A
 Symbolized by Z X
Ex. What is the atomic symbol for helium?
He has 2 e–, 2 n and 2 p Z = 2, A = 4
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2 He
24
Isotopes
 Most elements are mixtures of two or more
stable isotopes
 Each isotope has slightly different mass
 Chemically, isotopes have virtually identical
chemical properties
 Relative proportions of different isotopes are
essentially constant
 Isotopes distinguished by mass number (A):
e.g.,
 Three isotopes of hydrogen (H)
 Four isotopes of iron (Fe)
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Example:
What is the isotopic symbol for Uranium235?
 Number of protons ( 11 p ) = 92
= number of electrons in neutral atom
 Number of neutrons ( 01n ) = 143
 Atomic number (Z ) = 92
 Mass number (A) = 92 + 143 = 235
 Chemical symbol = U
 Summary for uranium-235: 235
92
U
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Learning Check:
 Fill in the blanks:
symbol
neutrons protons
electrons
131I
78
53
81Br
46
35
35
36
29
29
65
29 Cu
53
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Carbon-12 Atomic Mass Scale
 Need uniform mass scale for atoms
Atomic mass units (symbol u)
 Based on carbon:
 1 atom of carbon-12 = 12 u (exactly)
 1 u = 1/12 mass 1 atom of carbon-12 (exactly)
Why was 12C selected?
 Common
 Most abundant isotope of carbon
 All atomic masses of all other elements ~ whole
numbers
 Lightest element, H, has mass ~1 u
Jespersen/Hyslop, Chemistry7E, Copyright © 2015 John Wiley & Sons, Inc. All Rights Reserved
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Calculating Atomic Mass
 Generally, elements are mixtures of isotopes
e.g. Hydrogen
Isotope
Mass
% Abundance
1H
1.007825 u
99.985
2H
2.0140 u
0.015
How do we define atomic mass?
 Average of masses of all stable isotopes of given
element
How do we calculate average atomic mass?
 Weighted average
 Use isotopic abundances and isotopic masses
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Learning Check
Naturally occurring magnesium is a mixture of 3
isotopes; 78.99% of the atoms are 24Mg (atomic
mass, 23.9850 u), 10.00% of 25Mg (atomic mass,
24.9858 u), and 11.01% of 26Mg (atomic mass,
25.9826 u). From these data calculate the average
atomic mass of magnesium.
24Mg
0.7899 x 23.9850 u = 18.946 u
0.1000 x 24.9858 u = 2.4986 u 25Mg
0.1101 x 25.9826 u = 2.8607 u 26Mg
Total mass of average atom =
24.3053 u rounds up to 24.31 u
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