Download Atomic Structure

The University of Trinidad and Tobago
Course CCCH110B
Course Instructors:
Balram Mahabir
Foundations Unit
Point Lisas Campus.
Ph: 642-8UTT (8888) ext.: 25096 or 374-8104
Email: [email protected]
Atomic Structure
An atom is composed of three types of subatomic particles: the proton, neutron, and
electron.
Particle Mass
Proton
Charge (multiple of 1.602 x10 -19 coulombs)
1.6727 x10 -24 g +1
Neutron 1.6750 x10 -24 g 0
Electron 9.110 x10 -28 g -1
Protrons and neutrons have similar masses and electrons are much lighter (over 1,000
times lighter).
Protons and electrons have equal and opposite charges while neutrons have no charge.
We have the following simple picture of the atom.
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The atom is comprised of a positively charged nucleus composed of protons and neutrons.
This small nucleus is surrounded by orbiting electrons. Because the protons and neutrons are
so much more massive than the electrons, virtually all the mass of the atom is located in the
nucleus. The light negatively charged electrons move around in an orbit in the space around
the nucleus.
We use the following symbol to describe the atom:
A= Z + N, where N is the number of neutrons.
If you add or subtract a proton from the nucleus, you create a new element.
If you add or subtract a neutron from the nucleus, you create a new isotope of the same
element you started with.
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In a neutral atom, the number of positively charged protons in the nucleus is equal to the
number of orbiting electrons.
The Hydrogen Atom
Let's look at the simplest example of an atom, the hydrogen atom.
The atom consists of a proton and an electron held together by the electromagnetic force
between the positively charged proton and the negatively charged electron.
The electron orbits around the proton because it is the lighter particle, sort of like the earth
orbits around the sun, There are, however, big differences in the picture of the earth going
around the sun and the electron going around the nucleus. This is because protons, neutrons
and electrons exist on a length scale so small that quantum mechanics is required to
understand the electron's orbit around the nucleus. We will learn more about the quantum
theory of the atom later.
When we add neutrons to the nucleus of
1
1H
we can make the isotopes of hydrogen. Here
are three common isotopes of hydrogen.
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If we add a proton to the hydrogen nucleus we would get helium (a different element). Here
are two common isotopes of helium.
Another example is carbon.
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Because the element symbol and atomic number are redundant, you will often see isotopes
written without the atomic number. For example, you would see 12C only.
Now you might think that an atomic nucleus with lots of protons (like 12C ) would fly apart from
the electrical repulsions between positively charged protons. It turns out that these forces of
electrical repulsions are overcome by an attractive force between protons and neutrons
called the strong nuclear force. At small distances inside a nucleus, this force is stronger than
the electromagnetic forces of repulsion, but at larger distances it becomes much weaker.
Example Problem
Let's work out an example problem involving elemental symbols. How many electrons,
protons and neutrons are contained in the isotope 3517Cl?
The number of protons is given by the atomic number, the bottom number, so the number of
protons is 17. This is a neutral atom, so there will be an equal amount of negatively charged
electrons to balance out the positively charged protons, thus making the number of electrons
17 also.
We know that the atomic mass is
A=Z+N
where N is the number of neutrons. Rearranging the equation we get
N=A-Z
Plugging in the numbers we already know, we get
N = 35 - 17 = 18 = the number of neutrons
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Atomic Mass
Grams is not a very convenient unit for atomic masses, so a new unit called the atomic
mass unit (amu) is defined.
1 amu = 1.660551 x 10 -24 g
Particle Mass(g)
proton
Mass(amu)
1.6727 x 10 -24 g 1.007316
neutron 1.6750 x 10 -24 g 1.008701
electron 9.110 x 10 -28 g 0.000549
Using an instrument called a mass spectrometer we can very accurately measure the mass
of atoms and molecules. Here are some measured isotope masses using a mass
spectrometer.
Isotope Mass(amu)
That
12
2
H
2.0140
4
He
4.00260
8
Be
8.005305
12
C
12.000000
16
O
15.994915
24
Mg
23.985042
C has a mass of exactly 12.000000 amu is not a coincidence. A mass spectrometer
can only measure mass differences accurately. To solve this problem the
12
C isotope is
defined to have a mass of exactly 12.000000 amu's. Then everything else is measured
relative to 12C.
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As you might expect different isotopes of the same element will have different masses. If you
look at the periodic table, however, you'll notice that there is only one number listed for the
mass of each element. How can you only have one mass if there is more than one isotope of
each element?
The answer is that the mass under each element is the weighted average of all of the
isotope masses for that element. In this weighted average, the weights are the percent
abundance that each isotope occurs in nature.
For example, if you analyzed a lump of pure carbon from the planet Earth, you would find that
98.89 % of all carbon atoms on earth are 12C atoms, and
1.11 % of all carbon atoms on earth are 13C atoms.
So the weighted average mass of carbon is
(.9889)(12.000000 amu 12C ) + (.0111)(13.0039 amu 13C ) = 12.011 amu
It is possible that on a planet in a galaxy far away, the natural abundances of carbon isotopes
may be different, and therefore they would have slightly different numbers under carbon in
their periodic table. The masses of the isotopes, however, are the same everywhere.
Let's look at another example. The natural abundance of
30.19 %. If the atomic weight of
63
Cu is 62.93 amu and
63
Cu is 69.09 % and for
65
(.6909)(62.93 amu 63Cu ) + (.3091)(64.93 amu 65Cu ) = 63.55 amu
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Cu is
Cu is 64.93 amu, what is the
average atomic weight for natural copper?
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