The Atom
Christopher G. Hamaker, Illinois State University, Normal IL
© 2005, Prentice Hall
Dalton Model of the Atom
• John Dalton proposed that all matter is made up
of tiny particles.
• These particles are molecules or atoms.
• Molecules can be broken down into atoms by
chemical processes.
• Atoms cannot be broken down by chemical or
physical processes.
Dalton’s Model
• According to the law of definite composition, the
mass ratio of carbon to oxygen in carbon dioxide
was always the same. Carbon dioxide was
composed 1 carbon atom and 2 oxygen atoms.
• Similarly, 2 atoms of hydrogen and 1 atom of
oxygen combine to give water.
• Dalton proposed that 2 hydrogen atoms could
substitute for each oxygen atom in carbon dioxide
to make methane with 1 carbon atom and 4
hydrogen atoms. Indeed, methane is CH4!
Dalton’s Theory
• A Summary of Dalton’s Atomic Theory:
1) An element is composed of tiny, indivisible,
indestructible particles called atoms.
2) All atoms of an element are identical and have the
same properties.
3) Atoms of different elements combine to form
compounds.
4) Compounds contain atoms in small whole number
ratios.
5) Atoms can combine in more than one ratio to form
different compounds.
Dalton’s Atomic Theory
• The first two parts of Dalton’s theory were later
proven incorrect.
• Proposals 3, 4, and 5 are still accepted today.
• Dalton’s theory was an important step in the
further development of atomic theory.
Subatomic Particles
• About 50 years after Dalton’s proposal, evidence
was seen that atoms were divisible.
• Two subatomic particles were discovered.
– Negatively charged electrons, e–.
– Positively charge protons, p+.
• An electron has a relative charge of -1 and a
proton has a relative charge of +1.
Thomson Model of the Atom
• J.J. Thomson proposed a subatomic model of the
atom in 1903.
• Thomson proposed that
the electrons were
distributed evenly
throughout a homogeneous
sphere of positive charge.
• This was called the “Plum
Pudding” Model of the
atom.
Mass of Subatomic Particles
• Originally, Thomson could only calculate the
mass-to-charge ratio of a proton and an electron.
• Robert Millikan determined the charge of an
electron in 1911.
• Thomson calculated the masses of a proton and
electron:
– An electron has a mass of 9.11 × 10-28 g
– A proton has a mass of 1.67 × 10-24 g
Rutherford’s Gold Foil Experiment
• Rutherford fired alpha particles at thin gold foils.
If the “plum pudding” model of the atom was
correct, most a-particles should pass through
undeflected.
• However, some of the alpha
particles were deflected
backwards.
Explanation of Scattering
• Most of the alpha particles passed through the foil
because an atom is largely empty space.
• At the center of an atom is the atomic nucleus
which contains the atom’s protons.
• The a-particles that
bounced backwards
did so after striking
the dense nucleus.
Rutherford’s Model of the Atom
• Rutherford proposed a new model of the atom:
– The negatively charges electrons are distributed
around a positively charged nucleus.
• An atom has a diameter of
about 1 × 10-8 cm and the
nucleus has a diameter of
about1 × 10-13 cm.
• If an atom were the size
of the Astrodome, the
nucleus would be a
marble.
Subatomic Particles Revisited
• Based on the heaviness of the nucleus, Rutherford
predicted that it must contain neutral particles in
addition to protons.
• Neutrons, n0, were discovered about 30 years
later. A neutron is about the size of a proton
without any charge.
Atomic Notation
• Each element has a characteristic number of
protons in the nucleus. This is the atomic
number, Z.
• The total number of protons and neutrons in the
nucleus of an atom is the mass number, A.
• We use atomic notation to display the number of
protons and neutrons in the nucleus of an atom:
Using Atomic Notation
• An example:
• The element is sodium (symbol Na).
• The atomic number is 11 – sodium has 11 protons.
• The mass number is 23 – the atom of sodium has
23 protons + neutrons.
• The number is neutrons is: A – Z = 23 – 11 =
12 neutrons.
Isotopes
• All atoms of the same element have the same
number of protons.
• Most elements occur naturally with varying
numbers of neutrons.
• Atoms of the same element that have a different
number of neutrons in the nucleus are called
isotopes.
• Isotopes have the same atomic number but
different mass numbers.
Isotopes Continued
• We often refer to an isotope by stating the name of
the element followed by the mass number.
– Cobalt-60 is
60
37
– Carbon-14 is
14
6
Co
C
• How many protons and neutrons does an atom of
mercury-202 have?
– The atomic number of Hg is 80, so it has 80 protons
– Hg-202 has 202 – 80 = 122 neutrons
Simple & Weighted Averages
• A simple average assumes the same number of
each object.
• A weighted average takes into account the fact
that we do not have equal numbers of all the
objects.
• A weighted average is calculated by multiplying
the percentage of the object (as a decimal number)
by its mass for each object and adding the
numbers together.
Average Atomic Mass
• Since not all isotopes of an atom are present in
equal proportions, we must use the weighted
average.
• Copper has two isotopes:
– 63Cu with a mass of 62.930 amu and 69.09% abundance
– 65Cu with a mass of 64.928 amu and 30.91% abundance
• The average atomic mass of copper is:
– (62.930 amu)(0.6909) + (64.928 amu)(0.3091)
= 63.55 amu
Periodic Table
• We can use the periodic table to obtain the atomic
number and atomic mass of an element.
• The periodic table shows the atomic number,
symbol, and atomic mass for each element.
Conclusions
• Atoms are composed of protons, neutrons, and
electrons.
• The protons and neutrons are located in the
nucleus and the electrons are outside the nucleus.
• Atoms are mostly empty space.
• The number of protons is referred to as the atomic
number for the atom.
Conclusions Continued
• All atoms of the same element have the same
number of protons.
• Isotopes are atoms with the same number of
protons but differing numbers of neutrons.
• The mass number for an isotope is the total
number of protons plus neutrons.
• The atomic mass of an element is the weighted
average of the masses of all the naturally
occurring isotopes.
Complete the following table:
Isotope
Symbol
Atomic
Number
# of
protons
# of
# of
Neutrons electrons
C -12
6
Ca2+
6
20
Br-80
36
Te-128
7
10
1. Copper has two isotopes:
Cu-63
Cu-65
Relative abundances
69.2%
30.8%
Atomic masses (amu)
62.93
64.93
Calculate the average atomic mass of copper.
2. Element X has two natural isotopes.
Isotope1
Isotope 2
Relative abundances
19.91 %
80.09 %
Atomic masses (amu)
10.012
11.009
Calculate the atomic mass of this element and name
it.