Download Atoms and Atomic Theory

ACP Chemistry 2011/12
Atoms and Atomic Theory
 The History of Atomic Theory
Circa. 400-5 BC. Greek philosopher Democritus proposes the idea of matter being made up of small, indivisible
particles (atomos).
Late 18th Century. Lavoisier proposes the Law of conservation of mass and Proust proposes the Law of
constant composition.
Early 19th Century. Using the previously unconnected ideas above, John Dalton formulates his Atomic
Theory.
 Dalton’s Atomic Theory
(i)
Elements are made from tiny particles called atoms.
(ii)
All atoms of a given element are identical. (N.B. see isotopes below).
(iii)
The atoms of a given element are different to those of any other element.
(iv)
Atoms of different elements combine to form compounds. A given compound always has
the same relative numbers and types of atoms. (Law of constant composition).
(v)
Atoms cannot be created or destroyed in a chemical reaction they are simply rearranged
to form new compounds. (Law of conservation of mass).
 Structure of the Atom and The Periodic Table
Several experiments were being carried out in the 19th and 20th centuries that began to identify the subatomic particles that make up the atom. A summary of those experiments is given below.
Scientist
Experiment
Knowledge gained
Relating to
Crookes
Cathode Ray Tube
Electron
J.J. Thomson
Cathode Ray
Deflection
Negative particles of
some kind exist
Mass/charge ratio of
the electron determined
Milikan
Oil Drop Experiment
Charge on the electron
Electron
Rutherford, Marsden,
and Geiger
Gold Foil Experiment
Nucleus present in
atom
The nucleus of an atom
and the proton
Electron
In the first part of the 20th Century, following Chadwick’s discovery of the neutron, Bohr proposed the
idea that the atom was made up of the nucleus containing protons and neutrons that was being orbited
by electrons in specific, allowed orbits. This particle model of the electron and atom was expanded a
few years after Bohr’s original ideas to incorporate the wave nature of the electrons.
PARTICLE
CHARGE
MASS
POSITION in atom
PROTON
+1
1 amu*
Nucleus
NEUTRON
0
1 amu*
Nucleus
ELECTRON
-1
1/1836 amu*
Outside of the nucleus
*atomic mass units
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ACP Chemistry 2011/12
The atomic numbers (in the table below printed above the symbol and sometimes referred to as Z) and
mass numbers (in the table below printed below the symbol and sometimes referred to as A) that
appear on the periodic table have specific meanings.
Atomic number the number of protons in the nucleus of one atom of the element
Since all atoms are neutral it also tells us the number of electrons surrounding the nucleus. (N.B.
When atoms lose or gain electrons they become charged and form ions).
Mass number the number of protons the number of neutrons in one atom of the element
Group
1A
2A
THE PERIODIC TABLE
1
2
Period
3
Period
4
Period
5
Period
6
Period
7
4A
Group
5A
6A
5
B
11
13
Al
27
31
Ga
70
49
In
115
81
Tl
204
6
C
12
14
Si
28
32
Ge
73
50
Sn
119
82
Pb
207
7
N
14
15
P
31
33
As
75
51
Sb
122
83
Bi
209
7A
1
H
1
Period
Period
3A
3
Li
7
11
Na
23
19
K
39
37
Rb
86
55
Cs
133
87
Fr
223
4
Be
9
12
Mg
24
20
Ca
40
38
Sr
88
56
Ba
137
88
Ra
226
T R A
21
22
Sc
Ti
45
48
39
40
Y
Zr
89
91
57
72
La*
Hf
139 178
89
104
Ac** Rf
226
58
Ce
140
90
Th
232
*Lanthanides
**Actinides
N S
23
V
51
41
Nb
93
73
Ta
181
105
Db
59
Pr
141
91
Pa
231
I T
24
Cr
52
42
Mo
96
74
W
184
106
Sg
60
Nd
144
92
U
238
I O
25
Mn
55
43
Tc
99
75
Re
186
107
Bh
61
Pm
147
93
Np
237
N
26
Fe
56
44
Ru
101
76
Os
190
108
Hs
62
Sm
150
94
Pu
242
M E T A L S
27
28
29
30
Co
Ni
Cu
Zn
59
59
64
65
45
46
47
48
Rh
Pd
Ag
Cd
103 106 108 112
77
78
79
80
Ir
Pt
Au
Hg
192 195 197 201
109 110 111
Mt
Ds
Rg
63
Eu
152
95
Am
243
64
Gd
157
96
Cm
247
65
Tb
159
97
Bk
251
66
Tb
163
98
Cf
251
67
Ho
165
99
Es
254
68
Er
167
100
Fm
253
69
Tm
169
101
Md
256
8
O
16
16
S
32
34
Se
79
52
Te
128
84
Po
210
70
Yb
173
102
No
254
9
F
19
17
Cl
35.5
35
Br
80
53
I
127
85
At
210
71
Lu
175
103
Lr
257
KEY:
Metal
13
Al
27
14
Si
28
Semi Metal
Non-metal
15
P
31
In this example Al is a metal, Si is a semi-metal (metalloid) and P is a non-metal.
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8A
2
He
4
10
Ne
20
18
Ar
40
36
Kr
84
54
Xe
131
86
Rn
222
ACP Chemistry 2011/12

Isotopes
Atoms with the same number of protons and electrons, but different numbers of neutrons are called
isotopes.
All Atoms of the same element contain the same number of protons and electrons but may have
different numbers of neutrons.
Since it is the electrons in atoms that affect the chemical properties of a substance, isotopes of the
same element have the same chemical properties.
Practice:
1. Consider the following pairs; does either pair represent a pair of isotopes?
(a)
40
19
K and
40
18
(b)
90
38
Sr and
2.
Determine the number of protons, electrons and neutrons in,
(a)
210
82
(b)
34
16
Ar
94
38
Sr


Pb
S


On some periodic tables you will find atomic mass numbers that are not integers. What does this
mean? A good starting point is to analyze what it does not mean.
For example, the atomic mass of Cl is often quoted on a periodic table as 35.5 and can be represented
by the following symbol;
35.5
17

Cl
This does not mean that there are 17 protons, 17 electrons and 18.5 neutrons in an atom of chlorine. It
is not possible to have a fraction of a neutron, there can only be a whole number of neutrons in an
atom.
So what does it mean, and where does the 0.5 come from? Here is the explanation.
The non integer values mean that there is more than one isotope of chlorine that exists in nature, in
this case 35Cl and 37Cl. A quick calculation will tell you that these two species have the same number
of protons and electrons, but different (whole) numbers of neutrons (18 and 20 respectively). That is,
they are isotopes of one another. These isotopes happen to exist naturally in the following abundance;
35
Cl 75% and 37Cl 25%.
A simple calculation can be applied to work out the average atomic mass when considering all the
isotopes present in a natural sample.
Average atomic mass 
 (% of each isotope)(atomic mass of each isotope)
100

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ACP Chemistry 2011/12
In this case
Average atomic mass =
((35)(75)  ((37)(25))
 35.5 (A non-integer in this case)
100
Another example is provided by Boron. Boron has two isotopes
18.7% and 81.3% respectively.

Average atomic mass =
10
B and 11B. They have the abundance
((10)(18.7)  ((11)(81.3))
 10.8 (Again a non-ineger)
100
Practice:
1.
The Noble gas, Neon has three isotopes of masses 22, 21 and 20. If the isotopes have the
abundance 8.01%,
1.99% and 90.00% respectively, what is the average atomic mass of neon
atoms?
2.
A naturally occurring sample of an element consists of two isotopes, one of mass 85 and
one of mass 87. The abundance of these isotopes is 71% and 29%. Calculate the average
atomic mass of an atom of this element.
3.
If 69Ga and 71Ga occur in the %’s 62.1 and 37.9, calculate the average atomic mass of
gallium atoms.
4.
Naturally occurring Chlorine molecules, Cl2 have masses of 70, 72 and 74. They occur in
the percentages 56.25%, 37.50% and 6.250%. What is the average atomic mass of chlorine
atoms? What is the relative abundance of 35Cl and 37Cl isotopes?
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