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CHAPTER 2
Atom and the Periodic Table
Elements
There are millions of different pure chemical substances that are
observed in the world. Early scientists suggested that all of these
substances were composed of a much smaller number of substances,
called elements.
Element - A substance that cannot be separated into more simple
substances by chemical means.
Compound - A substance composed of two or more elements.
Example - calcium oxide
calcium carbonate  calcium oxide + carbon dioxide
calcium oxide  calcium + oxygen
carbon dioxide  carbon + oxygen
Early Theories of Matter
The ancient Greeks discussed two possibilities for the essential
property of matter.
Matter is continuous (no “particles” of matter) - Plato, Aristotle,
and a majority of Greek philosophers.
Matter is discrete (composed of particles) - Democritus,
Leucippus, and a small number of Greek philosophers.
Does this process have an end?
yes – particles of matter exist
no – matter is continuous
Dalton’s Atomic Theory
A comprehensive theory that accounted for the above observations was proposed by John Dalton, an English chemist, in 1808.
There were three parts to the theory.
1) Elements are composed of particles, called atoms.
a) All atoms of the same element are identical in size, mass,
and chemical properties.
b) Atoms of different elements differ in their size, mass, and
chemical properties.
Dalton’s Atomic Theory (continued)
2) Chemical substances (compounds) are composed of atoms of
more than one element.
a) In any particular pure chemical compound the same kinds of
atoms are present in the same relative numbers.
3) Chemical reactions can rearrange atoms, but atoms cannot be
created, destroyed, or converted from atoms of one element to atoms of a
different element.
We now know that some of the hypotheses in Dalton’s atomic
theory are not completely correct; however, the theory represents a good
starting point in understanding the composition of matter.
Atomic Structure
In Dalton’s atomic theory the smallest particles (atoms) could not
be further broken down. However, a series of experiments, beginning in
the mid-19th century, demonstrated that atoms themselves were
composed of smaller particles.
Radioactivity
In 1895, Antoine Becquerel discovered that some substances
(such as radium and uranium) spontaneously emit “radiation”, a process
called radioactivity.
Three types of radioactivity were found:
alpha () radiation: positively charged particles, now known to
be He2+ nuclei (2 protons + 2 neutrons)
beta () radiation: negatively charged particles, now known to be
electrons
gamma () radiation:
uncharged, now known to
be high energy photons
(particles) of light
Electrons
J. J. Thompson (1897) found that when a high voltage was applied
across two electrodes at low pressure a beam of particles moved from the
negative to the positive electrode. The particles, named electrons, were
negatively charged and the same regardless of the gas between the
electrodes or the metal used in the electrodes.
Charge and Mass of the Electrons
Thompson’s experiment made it possible to find the value for the
charge to mass ratio for an electron.
Charge to mass ratio = charge of an electron = 1.76 x 108 C/g
mass of an electron
The charge of an electron (and therefore the mass of the electron) was
first determined experimentally by Millikan (1909).
Charge of electron = - 1.6022 x 10-19 C (C = Coulomb)
Mass of electron = 9.11 x 10-28 g
Consequences of the Discovery of
the Electron
Based on the work of Thompson and Millikan, it was known that
atoms contained electrons. This suggested the following:
1) Atoms must also contain a positively charged particle to
balance the negative charge of the electron.
2) The mass of the positively charged particle must be much
larger than the mass of the electron.
The “Plum Pudding” Model
To account for the predictions about the positive particles in an
atom, Thompson suggested that most of the space within an atom
consisted of a positively charged substance, with electrons embedded
within, the “plum pudding” model.
Rutherford and the Nuclear Atom
To test Thompson’s plum pudding model, Ernest Rutherford
(1909) carried out an experiment where a beam of positively charged
particles (alpha particles) were directed at a thin sheet of gold metal.
The results of this experiment were inconsistent with the plum
pudding model. Rutherford proposed a new model, called the nuclear
model of the atom, that did account for the experimental results.
Subatomic Particles
particle
charge
Coulombs
mass
elementary
kg
amu
proton (p+)
+ 1.60 x 10-19
+1
1.673 x 10-27
1
neutron (n)
0
0
1.675 x 10-27
1
electron (e-)
- 1.60 x 10-19
-1
9.11 x 10-31
0
________
1 amu = 1.6605 x 10-27 kg
mp/me = 1836.
Atomic Structure
nucleus
electron charge cloud
1) The protons and neutrons of the atom are found in a small region in
the center of the atom, called the nucleus. This region contains most of
the mass of the atom, and all of the positive charge.
2) Electrons in the atom form a diffuse cloud of negative charge centered
on the nucleus and occupying most of the volume of the atom.
3) The size of the charge for the proton and electron is the same. The
charge for the proton is positive, and the charge for the electron is
negative. Neutrons have no charge.
4) The type of element for an atom is determined by the number of
protons in the atomic nucleus.
Element (new definition) - An element is a pure chemical
substance composed of atoms, each of which has the same number of
protons in the nucleus.
Hydrogen - one proton per atom
Helium - two protons per atom
Lithium - three protons per atom
.
.
.
.
Similarly, we can now define a compound (new definition) as a
pure chemical substance composed of two or more different kinds of
atoms.
The Periodic Table
Atomic Number and Mass Number
1) The atomic number (Z) is equal to the number of protons in the atom.
2) Since atoms are electrically neutral, the number of electrons in an
atom is also equal to Z, the atomic number.
3) The mass number (A) is equal to the number of protons + neutrons in
the atom.
a) Because protons and neutrons have a mass of approximately 1
(in amu) and electrons have a mass of approximately 0 (in amu) the mass
number is equal to the approximate mass of the atom in amu.
b) Based on the above, the number of neutrons in an atom is
equal to A - Z. So for an atom:
# protons = Z
# electrons = Z
# neutrons = A - Z
Notation For Atoms
We use the following general notation to represent isotopes of
atoms.
mass number
symbol for element
atomic number
Since we can use the symbol for the element and the periodic
table to determine Z, the atomic number, we often omit Z in giving the
symbol for the atom.
Example: 3416S = 34S We can omit the subscript because all
sulfur atoms contain 16 protons.
Isotopes
The atomic number determines the number of protons and
electrons in an atom. This does not place any restrictions on the number
of neutrons in the atom.
It is possible for atoms of the same element to have different
numbers of neutrons. These different types of atoms are called isotopes.
Isotopes of Hydrogen
normal hydrogen
deuterium
tritium
1H
2H
3H
Note that to a very good approximation isotopes of a
particular element are chemically identical to one another.
Example of Notation for Isotopes
As an example of using the above notation, consider the following naturally occurring isotopes of oxygen (Z = 8).
protons
neutrons
electrons
mass number symbol
8
8
8
16
16O
8
9
8
17
17O
8
10
8
18
18O
Example: How many protons, neutrons, and electrons are there in
one atom of 56Fe? What is the approximate mass of one atom of 56Fe in
amu and in kg?
Atomic Mass Units (amu)
The mass of a single atom of an element, expressed in SI units, is
an extremely small number. For example, the mass of a single atom of
16O is 2.6560 x 10-26 kg. For convenience, we often express values for
atomic mass in terms of atomic mass units (amu).
Atomic mass units are defined as follows
12.00 amu = mass of one atom of 12C (exact)
From this we get 1. amu = 1.6605 x 10-27 kg (approximate)
The mass of any other atom (or particle) is found relative to the
ratio of its mass to the mass of a 12C atom, which can be measured
experimentally.
Mass of particle (amu) =
mass particle • (12.00 amu)
mass 12C atom
Example: A mass spectrometer is a device for determining values
for mass for atoms (and molecules). In a particular experiment, the ratio
(mass X/mass 12C) is measured and found to be equal to 2.581. What is
the mass of the atom X (in amu)?
Atomic Mass in the Periodic Table
Because different isotopes of an element have different masses,
the question arises as to which mass should be given in the periodic
table.
For short lived radioactive elements the mass number of the most
stable isotope of the element is listed.
Element
Z
A
technetium (Tc)
43
98
radon (Rn)
86
222
plutonium (Pu)
94
244
Average Atomic Mass
For naturally occurring elements, the value for mass given in the
periodic table is the average atomic mass, based on the natural
abundance of the isotopes that is observed.
In general, we find the average atomic mass as follows:
Mave = f1 M1 + f2 M2 + f3 M3 + … = i=1n fi Mi
where f1, f2,...are the fractions of each isotope observed in nature
M1, M2,…are the corresponding masses for each isotope (in amu)
Note the following
f1 + f2 + f3 + …= 1
fx = % X
100 %
Non-chemical Example
A person has a box of sandwiches. Half of the sandwiches are
6.0 ounces, and half of the sandwiches are 10.0 ounces. What is the
average weight of a sandwich?
Average weight = (0.50)(6.0 oz) + (0.50)(10.0 oz)
= 8.0 ounces
We use the same procedure in finding the average mass of an atom. We
multiply the fraction of each isotope by the mass of that isotope, and then
add the results to find the average mass.
Chemical Example
There are three naturally occurring isotopes of the element
magnesium. Based on the information below, find the average atomic
mass of a magnesium atom.
Isotope
percent
f
M(amu)
24Mg
78.70 %
23.98504
25Mg
10.03%
24.98584
26Mg
11.17%
25.98259
Periodic Table
The periodic table is an arrangement of the chemical elements
based on similarities in their physical and chemical properties
The periodic table contains a large amount of useful information
about the chemical elements.
Organization
There are several ways in which the elements in the periodic table
may be classified.
Rows = Periods
Columns = Groups
This is the more important classification.
Elements in the same group usually have similar physical and chemical
properties.
Simplified Periodic Table
1A
2A
3A
4A
5A
6A 7A
8A
1A
2A
3A
4A
5A
6A 7A
8A
You are responsible for knowing the names/symbols for elements 1-57,
72-86, and 92.
Major Groups in the Periodic Table
1A
2A
3A
4A
5A
6A 7A
8A
Metals, Nonmetals, and Metalloids
Metals:
Usually solid at room temperature (exceptions Cs, Fr, Hg)
Shiny metallic luster
Good conductors of electricity and heat
Malleable (can be hammered into thin sheets)
Ductile (can be drawn into thin wires)
Nonmetals:
Can be solid, liquid, or gas at room temperature
Dull colored (as solids)
Poor conductors of electricity and heat
Not malleable, not ductile
Metalloids (semimetals): Intermediate between metals and nonmetals
Metals, Nonmetals, Metalloids in the Periodic Table
1A
2A
3A
4A
5A
6A 7A
8A
Examples of Elements (as found in nature)
nickel
germanium
(metal)
(metalloid)
sulfur
(nonmetal)
The Mole
It is nearly impossible to work with individual atoms in the
laboratory because of their small size and mass. For example, one 12C
atom has a mass of 1.993 x 10-26 kg, far to small to measure directly with
an analytical balance..
It is convenient to have a unit representing a number of atoms
that can easily be measured by the usual techniques in the laboratory.
This unit is called the mole.
Avogardo’s Number
The unit mole is simply a number. By definition, one mole of
anything is equal to Avogadro’s number of things. Avogadro’s number,
an experimentally determined quantity is NA = 6.022 x 1023.
Note that there is no difference between the concept of moles and
more common terms used for specific numbers of things.
1 dozen = 12
1 score = 20
1 thousand = 1000
1 mole = 6.022 x 1023
Example (Avogadro’s Number)
If I have three dozen eggs, then how many eggs do I have?
1 dozen eggs = 12 eggs (a conversion factor)
number of eggs = 3 dozen eggs 12 eggs = 36 eggs
1 dozen
If I have three moles of carbon atoms, then how many carbon
atoms do I have
1 mole carbon atoms = 6.022 x 1023 carbon atoms
number of carbon atoms = 3 moles C 6.022 x 1023 C atoms
1 mole
= 1.81 x 1024 carbon atoms
Moles is the fundamental SI unit for quantity of substance. The
symbol for moles is n, and the abbreviation for the unit moles is mol.
Significance of Avogadro’s Number
There must be something special about the number 6.022 x 1023
(Avogadro’s number). The significance is as follows. Consider a
collection of identical objects. The following relationship will apply.
If one object has a mass of X amu…
…then one mole of objects has a mass of X g.
This is a subtle point. Consider the unit “thousand”. We could
make the following statement
If one object has a mass of X g…
…then one thousand objects has a mass of X kg.
Example: The mass of one penny is 2.50 g. The mass of one
thousand pennies is
mass one thousand pennies = 1000 (2.50 g) = 2500. g 1 kg
= 2.50 kg
The same principle applies to the concept of moles
mass of one 19F atom = 19.00 amu,
so mass of one mole of 19F atoms = 19.00 g.
Check (Note that 1 amu = 1.6605 x 10-27 kg)
mass one mole 19F = (6.022 x 1023) (19.00 amu)
= 1.1442 x 1025 amu 1.6605 x 10-27 kg
1 amu
= 0.01900 kg = 19.00 g
Use of the Mole Concept
We may use the concept of moles to reinterpret atomic mass.
Atomic mass of S = 32.06 amu = 32.06 g/mol
We interpret the atomic mass of sulfur (or any other element in
the periodic table) in two equivalent ways:
The average mass of one sulfur atom is 32.06 amu.
The mass of one mole of sulfur atoms is 32.06 g.
Use of the Mole Concept (Example)
A chemist has a 14.38 g sample of copper (Cu).
a) How many moles of copper does she have?
b) How many atoms of copper does she have?
End of Chapter 2
“…the ultimate particles of all homogeneous bodies are perfectly
alike in weight, figure, and so forth.”
- John Dalton, A New System of Chemical Philosophy (1808)
“Elements arranged according to the size of their atomic weights
show clear periodic properties.” - D. I. Mendeleev (1869)
“I don’t believe that atoms exist!” - Ernst Mach (1897)