Chapter 2 Atoms, Molecules & Ions
Chapter 1: elements - substances which cannot
be decomposed into simpler substances by
normal chemical means
There are 118 elements – what makes one element
different from another?
How and why do elements combine to form compounds?
1. Dalton's atomic theory:
1. Each element is composed of small particles
called atoms.
2. All atoms of a given element are identical; atoms
of different elements have different properties
(including mass).
3. Atoms are neither created nor destroyed in the
course of chemical reactions; atoms of an
element are not changed into different types of
atoms by chemical reactions.
4. Compounds are formed when atoms of more than
one element combine; a given compound always
has the same the relative number and kinds of
atoms.
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How does Dalton's atomic theory help explain the
laws of constant composition, multiple
proportion, and conservation of mass?
So, elements differ from one another because
they are composed of different atoms......
How do the atoms of the various elements differ
from one another?
2.The Discovery of Atomic Structure
Cathode Rays and Electrons
Thomson, 1897: cathode rays
consist of (-) charged particles
with mass; same regardless of
source
Millikan, 1909: charge-to-mass
ratio of electron
Charge = 1.60 x 10-19 C
Mass = 9.10 x 10-28 g
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Note: the Coulomb (C) is the SI
unit of electrical charge
Radioactivity
Becquerel, Curie, Rutherford:
Three types of radiation
emitted:
 particles: positive charge
particles: negative charge
radiation: not charged
The Nuclear Atom
J. J. Thomson, 'plum pudding'
model
Rutherford, 1910
-particles (charge=+2)
scattered at 180 angle
by Au foil
o
implications?
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3. Modern view of atomic structure
Atoms are composed of:
Electrons: charge = -1; mass = 5.486x10-4 amu
Protons: charge = +1; mass = 1.0073 amu
Neutrons: uncharged; mass = 1.0087 amu
There are more fundamental particles in the nucleus, but
only these three particles have a bearing on chemical
behavior.
Note that charge is expressed as a multiple of electron
charge e, where 1e = 1.602 x 10-19C. An electron has a
charge of -1.602 x 10-19C or -1e or simply -1 (and a proton
has a charge of +1.602x10-19C or +1e or just +1)
Most atoms have diameters between 1x10-10 m and 5x10-10
m (we usually use the Angstrom Å, 1 Å = 1x10-10 m to
describe atomic dimensions)
Note that the diameter of an atomic nucleus is
approximately 10-4 Å, which is a small fraction of the total
diameter of an atom
Nuclei are incredibly dense
Atoms are mostly empty space
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The diameter of a tin atom is about 2.8 Å; how many tin
atoms would have to be lined up to span 1.0 cm?
Isotopes, atomic numbers, and mass numbers
What makes atoms of one element different from an atom
of another element?
All atoms of an element have the same number of protons
in the nucleus
number of protons is different for different elements
atoms are electrically neutral; # protons = # electrons
atomic number = # of protons
mass number = # protons + # neutrons
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Isotopes: atoms of a given element that differ in the
number of neutrons (and therefore in mass)
e.g., 12C, 14C
isotope symbols written as
element symbol
mass no.
or
element name – mass no.
E.g. How many neutrons and electrons are present in 12C?
What about carbon-14?
How many protons, neutrons, and electrons are in a 39K
atom?
4. Atomic Weights
Masses of individual atoms can be measured with a
high degree of accuracy, e.g.,
1
H has a mass of 1.6735 x 10-24 g
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O has a mass of 2.6560 x 10-23 g
How is this done?
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It proves to be much more convenient to use the atomic
mass unit (amu) when dealing with such small quantities:
1 amu = 1.66054 x 10-24 g
and
1 g = 6.022 x 1023 amu
(The amu is defined by assigning a mass of exactly 12
amu to the 12C isotope)
So, one 1H atom has a mass of 1.0078 amu and one 16O
atom has a mass of 15.9947 amu (nicer than gram units,
huh?)
There are three isotopes of C: 12C, 13C, 14C; on the periodic
table, the atomic weight of C is given as 12.011 amu.
Why?
Similarly, Cl has two isotopes, 35Cl and 37Cl. The atomic
weight of Cl is 35.453 amu…
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Average Atomic Masses
Most elements occur in nature as mixtures of isotopes;
the atomic weights of the elements are listed on the
periodic table as average atomic masses
The atomic weight of an element is also called the average
atomic mass, where
average atomic mass =
isotope mass)x(relative abundance)]
The sum is over all isotopes!
The relative abundance of an isotope is the fraction of an
element with a given mass, e.g., carbon is composed of
98.93% 12C and 1.07% 13C.
e.g., Ne has three isotopes of masses 19.99, 20.99, and
21.99 amu, with relative abundances of 90.92%, 0.25%, and
8.83%, respectively. Find the atomic weight of Ne.
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Problem du Jour
A rhodium atom has a diameter of about 2.5 x 10-8 cm.
(a) Express the diameter in angstroms (Å) and m.
(b) How many Rh atoms would have to be placed side-by-side
to span a distance of 6.0 m?
(c) If the atom is assumed to be a sphere, what is the volume
(m3) of a single Rh atom?
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Problem du Jour
Write the correct symbol, with both superscript and subscript,
for the following:
The isotope of argon with mass number 40
An particle
The isotope of hafnium that contains 107 neutrons
The isotope of silicon that has an equal number of neutrons
and protons
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