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Atoms
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2
Chapter
Atoms
2.1 The Atomic Theory of Matter
Chemists make their observations in the macroscopic world and seek to understand the
fundamental properties of matter at the level of the microscopic world (i.e. molecules and
atoms). The reason why certain chemicals react the way they do is a direct consequence of
their atomic structure.
The word "atom" is derived from the Greek word "atomos", meaning indivisible. The
philosopher Democritus (460-370 B.C.) believed that matter was composed of fundamentally
indivisible particles, called "atomos".
Dalton's atomic theory of 1803:
1. Each element is composed of extremely small indivisible particles called atoms.
2. All atoms of a given element are identical; the atoms of different elements are different
and have different properties (including different masses).
3. Atoms of an element are not changed into different types of atoms by chemical
reactions; atoms are neither created nor destroyed in chemical reactions, they simply
combine, separate or rearrange.
4. Compounds are formed when atoms of more than one element combine in a specific
simple ratio; a given compound always has the same relative number and kind of
atoms.
Atoms are the basic building blocks of matter; they are the smallest units of an element:
Atoms are the smallest particle of an element which retains the chemical properties of
that element
Simple "laws" (i.e. theories) of chemical combination which were known at the time of Dalton:
1. The law of constant composition (in a given compound the relative number and kind
of atoms are constant). For example, water is made up of hydrogen and oxygen in the
fixed ratio of H:O 2:1, or a fixed mass ratio of 11% to 89%.
2. The law of conservation of mass (the total mass of materials present after a
chemical reaction is the same as the total mass before the reaction).
In a chemical reaction mass of reactants = mass of products (Lavoisier’s law)
Atoms
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2.2 The Discovery of Atomic Structure



1803 Dalton - the atom is a indivisible, indestructible, tiny ball
1850 Evidence is accumulating that the atom is itself composed of smaller particles
The current model...
Cathode rays and electrons
Electrical discharge through partially evacuated tubes produced radiation. This radiation
originated from the negative electrode, known as the cathode (thus, these rays were termed
cathode rays).






The "rays" traveled towards, or were attracted to the positive electrode (anode)
Not directly visible but could be detected by their ability to cause other materials to
glow, or fluoresce
Traveled in a straight line
Their path could be "bent" by the influence of magnetic or electrical fields
A metal plate in the path of the "cathode rays" acquired a negative charge
The "cathode rays" produced by cathodes of different materials appeared to have the
same properties
These observations indicated that the cathode ray radiation was composed of negatively
charged particles (now known as electrons).
J.J. Thompson (1897) used results from
cathode ray tube (commonly abbreviated
CRT) experiments to discover the electron.
He was able to measure the charge to
mass ratio for a stream of electrons (using
a cathode ray tube apparatus) at 1.76 x
108 coulombs/gram.
The side image is of J.J. Thomson and a
cathode ray tube from around 1897, the
year he announced the discovery of the
electron. Only the end of the CRT can be
seen to the right-hand side of the picture.
The image below is of a CRT used by Thomson in his experiments. It is about one meter in
length and was made entirely by hand.
Atoms
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The amount the cathode
ray bent from the
straight line using either
the electric field or the
magnetic field allowed
Thomson to calculate
the e/m ratio, but he was
not able to determine
the mass of the electron.
However, from his data,
if the charge of a single
electron
could
be
determined, then the
mass of
a
single
electron
could
be
determined.
Robert Millikan (1909) was able to successfully measure the charge on a single electron
(the "Millikan oil drop experiment"). This value was determined to be 1.60 x 10-19 coulombs.
Thus, the mass of a single electron was determined to be:
(1 gram/1.76 x 108 coulombs)*(1.60 x 10-19 coulombs) = 9.10 x 10-28 grams
Note: the currently accepted value for the mass of the electron is 9.10939 x 10-28 grams.
The nuclear atom
J.J. Thompson model of the atom (1900)


The atom consists of a sphere of
positive charge within which was
buried negatively charged electrons
Also known as the "plum pudding"
model of the atom
Rutherford model of the atom (1911)
(the gold-foil experiment...image aside)
Observations




Most of the alpha particles pass straight through the gold foil.
Some of the alpha particles get deflected by very small amounts.
A very few get deflected greatly.
Even fewer get bounced of the foil and back to the left.
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Conclusions



The atom is 99.99%
empty space.
The nucleus contains a
positive charge and most
of the mass of the atom.
The
nucleus
is
approximately 100,000
times smaller than the
atom.
Rutherford (1919) discovers protons - positively charged particles in the nucleus
Chadwick (1932) discovers neutron - neutral charge particles in the nucleus.
We will discuss Bohr’s model and the quantum model later
2.3 Modern view of atomic structure
Physicists have identified a long list of particles which make up the atomic nucleus.
Chemists, however, are primarily concerned with the following sub-atomic particles:



electron
proton
neutron
Electron
The electron is negatively charged, with a charge of -1.602 x 10-19 Coulombs (C). For
convenience, the charge of atomic and sub-atomic particles are usually described as a
multiple of this value (also known as the electronic charge). Thus, the charge of the
electron is usually simply referred to as -1.
Proton
The proton has a charge of +1 electron charge (or, +1.602 x 10-19 C)
Atoms
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Neutron
Neutrons have no charge, they are electrically neutral.
Note: Because atoms have an equal number of electrons and protons, they have no
net electrical charge.
Protons and neutrons are located in the nucleus (center) of the atom. The nucleus is small
compared to the overall size of the atom. The majority of the space of an atom is the space in
which the electrons move around.
Electrons are attracted to the protons in the nucleus by the force of attraction between
particles of opposite charge.
Note: The strength of attraction
between electrons and protons in
the nuclei for different atoms is the
basis of many of the unique
properties of different atoms. The
electrons play a major role in
chemical
reactions.
In
atomic
models,
the
electrons
are
represented as a diffuse electron
cloud
The mass of an atom is extremely small. The units of mass used to describe atomic particles
is the atomic mass unit (or amu).
An atomic mass unit is equal to 1.66054 x 10-24 grams.
How do the different sub-atomic particles compare as far as their mass?
Proton = 1.0073 amu
Neutron = 1.0087 amu
Electron = 5.486 x 10-4 amu
From this comparison we can see that:



The mass of the proton and neutron are nearly identical
The nucleus (protons plus neutrons) contains virtually all of the mass of the atom
The electrons, while equal and opposite in charge to the protons, have only 0.05% the
mass
The size of an atom is quite small also, the typical range for atomic diameters is between 1 x
10-10 and 5 x 10-10 meters.
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Note: A convenient unit of measurement for atomic distances is the angstrom (Å). The
angstrom is equal to 1 x 10-10 meters. Thus, most atoms are between 1 and 5
angstroms in diameter.
Parts of an Atom
Symbol
Mass in
amu
Relative Electric
Charge
Proton
p
1.0073
+1
Nucleus
Neutron
n
1.0087
0
Nucleus
Electron
e
0.0005
−1
Orbital/Electron Cloud
Subatomic Particle
Atomic Location
Isotopes, Atomic Numbers and Mass Numbers
What characteristic feature of sub-atomic particles distinguishes one element from another?



All atoms of an element have the same number of protons in the nucleus
Since the net charge on an atom is 0, the atom must have an equal number of
electrons.
What about the neutrons? Although usually equal to the number of protons, the
number of neutrons can vary somewhat. Atoms which differ only in the number of
neutrons are called isotopes. Since the neutron is about 1.0087 amu (the proton is
1.0073), different isotopes have different masses.
Carbon
All atoms of the element Carbon (C) have 6 protons and 6 electrons. The number of protons
in the carbon atom are denoted by a subscript on the left of the atomic symbol:
This is called the atomic number, Z, and since it is always 6 for carbon, it is somewhat
redundant and usually omitted. Another number, the "Mass Number", A, is a superscript
on the left of the atomic symbol. It denotes the sum of the number of protons and neutrons in
the particular isotope being described. For example:
Refers to an isotope of carbon which has (as expected for the element carbon) six protons,
and six neutrons. The following isotope of carbon:
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Has 6 protons (atomic number) and 8 neutrons (8=14-6 or A-Z). This isotope is also known
simply as "carbon 14". Carbon 12 is the most common form of carbon (~99% of all carbon).
An atom of a specific isotope is called a nuclide.
Hydrogen, for example has three isotopes:
number of
protons
number of
neutrons
Mass in
amu
Hydrogen (protium)
Heavy hydrogen (deuterium)
1
1
0
1
1
2
Super-heavy hydrogen (tritium)
1
2
3
Isotope
Tritium, however, is unstable and the nucleus decays (breaks apart). This is an example of
radioactivity.
Since all atoms are composed of protons, electrons and neutrons, all chemical and physical
differences between elements are due to the differences in the number of these sub-atomic
particles. Therefore, an atom is the smallest sample of an element, because dividing an atom
further (into sub-atomic particles) destroys the element's unique identity.
Exercises
Complete the table by filling in the empty boxes.
Element
Calcium (Ca)
Nickel (Ni)
Gold (Au)
Atomic Number
79
6
Protons
Neutrons
20
28
20
Mass number
59
118
14
Write the symbol or name for each of the following elements, as appropriate.
a.
b.
c.
d.
e.
f.
g.
h.
i.
gold
Hg
uranium
sodium
Mn
Zn
fluorine
Fe
lead
Isotope
Symbol
Atoms
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2.4 The Periodic Table
As more and more elements were discovered and characterized, efforts were made to see
whether they could be grouped, or classified, according to their chemical behavior. This
effort resulted, in 1869, in the development of the Periodic Table.
Certain elements show similar characteristics:


Lithium (Li), Sodium (Na) and Potassium (K) are all soft, very reactive metals
Helium (He), Neon (Ne) and Argon (Ar) are very non-reactive gasses
If the elements are arranged in order of increasing atomic number, their chemical and
physical properties are found to show a repeating, or periodic pattern..
As an example of the periodic nature of the atoms (when arranged by atomic number), each
of the soft reactive metals comes immediately after one of the nonreactive gasses.
The elements in a column of the periodic table are known as a family or group. The labeling
of the families are somewhat arbitrary, but are usually divided into the general groups of:



Metals (everything on the left and middle region)
Non-metals (upper diagonal on the right hand side - green, salmon and red)
Metalloids (atoms in the boundary between the metals and metalloids: Boron(B),
Silicon(Si), Germanium(Ge), Arsenic(As), Antimony(Sb), Tellurium(Te), and maybe
Astatine(At)). These are some of the more useful materials for semi-conductors.
Another convention is the 'A' and 'B' designators with column number labels (either in Roman
or Arabic numerals). These columns have different types of classifications:
Group
1A or I
2A or II
6A or VI
7A or VII
8A or VIII
Name
Alkali metals
Alkaline earth metals
Chalcogens ("chalk formers")
Halogens ("salt formers")
Noble gases (or inert, or rare gases)
Elements
Li, Na, K, Rb, Cs, Fr
Be, Mg, Ca, Sr, Ba, Ra
O, S, Se, Te, Po
F, Cl, Br, I, At
He, Ne, Ar, Kr, Xe, Rn
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Atoms
Group
I
1
2
3
4
5
6
7
II
III
IV
V
VI
VII
VIII
1
2
H
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
87
88
89
104
105
106
107
108
109
110
Fr
Ra
Ac
Rf
Db
Sg
Bh
Hs
Mt
Ds
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Legend
Solid
Liquid
Gas
Synthetic
Alkali metals
Alkali earth metals
Transition metals
Rare earth metals
Other metals
Noble gases
Halogens
Other nonmetals
The elements in a family of the periodic table have similar properties because they have the
same type of arrangement of electrons at the periphery of their atoms.
The majority of elements are metals. Some properties of metals are:


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


high luster
high electrical conductivity
high heat conductivity
solid at room temperature (except Mercury [Hg])
ductile
malleable
Note: hydrogen is a non-metal (at left hand side of the periodic table)
Non-metals can be solid, liquid or gas at room temperature. Their properties are opposite
those of metals
Atoms
EXERCISES
Q.1 Give the correct terms that correspond to the following statements:
a. Six elements falling on the stair-step line in the Periodic Table.
b. The total number of periods in the current Periodic Table.
c. Name given to elements of group II.
d. Name given to elements that are assigned the columns 3 up to 12.
e. Very reactive elements that lose one electron easily.
f. Nonmetals with eight valence electrons.
g. The person accredited with the discovery of the atomic number.
Q.2 Answer with True/False and correct the false statements properly:
a. Most metals are solids at room temperature.
b. Nonmetals have the general property of being lustrous.
c. All the halogens are nonmetals.
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