Download Chapter 2 Atoms, Ions, and Molecules

Chapter 2
Atoms, Ions, and Molecules
What is matter made of? Continuous or particulate?
Ancient view: Four elements (fire, air, earth, water)
Democritus (460–370 BC): “father of atomism”.
Aristotle (384–322 BC) held that Democritus’ views were impossible. Atomic concept suppressed for
~2000 years. But by the 1700s, things had started to change.
Observations Leading to the Atomic View of Matter
Law of Conservation of Mass: total mass of substances does not change during a chemical reaction.
[Matter can not be created or destroyed.]
–Lavoisier, 1789
Law of Definite Composition: no matter its source, a substance is composed of the same elements in
the same fractions (or ratio) by mass.
–Proust, 1799
mass of element
 fraction by mass =
total mass of substance
 percent by mass =
mass of element
total mass of substance
100%
Example: A 15.00 gram sample of water contains 13.32 grams of oxygen. Calculate the % by
mass of oxygen and hydrogen in water. (Water contains only hydrogen and oxygen.)
Law of Multiple Proportions: If elements A, B form a series of substances, the different masses of B
that combine with a fixed mass of A will be in a ratio of small whole numbers.
–Dalton, 1803
Example: Two compounds of copper and bromine with different properties
Substance
mass of bromine that
combines with 1g copper
I
1.26 g
II
2.52 g
properties
yellowish green solid
MP = 504 °C, BP = 1345 °C
light sensitive, water insoluble
black solid
MP = 498 °C, BP = 900 °C
not light sensitive, is water soluble
bromine = element B, copper = element A, ratio of B w/fixed mass of A = 2.52:1.26 or 2:1
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Dalton’s Atomic Theory (1808)
1. matter is made up of microscopic, indivisible
particles (atoms) that cannot be created or destroyed
Our Understanding Today
2. all atoms of an element are identical in mass, and
have identical physical and chemical properties
3. atoms of different elements have different masses,
physical properties, and chemical properties
4. atoms of different elements combine in simple
whole numbers to form compounds
5. atoms of an element cannot be converted into atoms of other
elements; chemical reactions involve reorganization of the atoms
But Dalton’s Atomic Theory couldn’t explain it all—
 Why do atoms combine the ways they do?
 What about the electrically charged particles that were being observed?
Observations Leading to the Nuclear Model of the Atom
 Cathode Ray Tube Experiments (Figure 2.2): J.J. Thomson, 1897
The particles discovered were called electrons. Thomson determined the mass/charge ratio of the
electron in a related experiment. Another experiment was needed to determine the mass or charge.
 Millikan’s Oil Drop Experiment ( Figure 2.3), 1909
Control of oil drop by electric field allowed for calculation of charge of 1 electron:
e– charge = –1.602 10–19 C
Since the mass/charge ratio was already known, the mass could also be calculated:
e– mass = 9.109 10–28 g
So, where’s the positive charge? Speculation was that it was spread throughout the atom.
 Rutherford’s Alpha Particle Scattering Experiment ( Figure 2.6), 1910
To test this speculation, Rutherford shot α particles (positively charged and relatively massive) at
gold foil. Most of the α particles passed through the foil relatively unaffected, but a few bounced
back nearly right back at the source. Conclusion: Atoms have a dense, positively charged core
called the nucleus.
By 1932, Chadwick discovered the neutron (neutral, not electrically charged).
Figure 2.6
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Describing Atoms
 atomic number, Z: # of protons in the nucleus of an atom
 mass number, A: sum of # of protons + # of neutrons in the nucleus of an atom

A
write as Z X or AX or “atomic name–A” (where X = atomic symbol)
e.g.,
 isotopes: atoms with same Z but different A. (atoms of same element but with different
masses)
e.g.,
Atomic Mass
standard: 12C atom mass = 12 atomic mass units (amu) exactly
1
1 amu = 12
the mass of a 12C atom [= 1 dalton (Da), often used in biochemistry and biology]
Using mass spectrometry, we can measure the masses of other atoms relative to 12C.
mass of 20 Ne atom
e.g.,
= 1.66604
mass of 12 C atom
 mass of 20Ne atom = 1.66604
12 amu = 19.9924 amu
Mass spectrometry also gives information about the relative
abundances (fractions) of each isotope of the element.
The weighted average of the masses of the isotopes is called the
atomic mass.
e.g.,
atomic mass of Ne = 0.904838 (19.9924 amu)
+ 0.002696 (20.9940 amu)
+ 0.092465 (21.9914 amu)
20.180 amu
10
atomic symbol
Ne
atomic mass
Neon
20.180
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The atomic mass is
closest to the mass of
the(Z)
most abundant
atomic number
isotope of the element.
atomic name
The Periodic Table
Mendeleev proposed the first periodic table in 1871, similar to today’s table
1A
(1)
1
H
8A
(18)
2
He
3
Li
2A
(2)
4
Be
3A
(13)
5
B
4A
(14)
6
C
5A
(15)
7
N
6A
(16)
8
O
7A
(17)
9
F
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
22.99
24.31
19
K
1.008
6.941
9.012
10.81
55.85
43
Tc
44
Ru
16.00
19.00
20.18
20
Ca
4B
(4)
22
Ti
5B
(5)
23
V
6B
(6)
24
Cr
7B
(7)
25
Mn
39.10
40.08
44.96
47.88
50.94
52.00
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
(98)
101.07
102.91
106.42
107.87
112.41
114.82
118.71
121.76
127.60
126.90
131.29
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
(10)
28
Ni
1B
(11)
29
Cu
2B
(12)
30
Zn
26.98
28.09
30.97
32.07
35.45
39.95
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
55
Cs
56
Ba
71
Lu
72
Hf
132.91
137.33
174.97
178.49
180.95
183.84
186.21
190.23
192.22
195.08
197.00
200.59
204.38
207.2
208.98
(209)
(210)
(222)
87
Fr
88
Ra
103
Lr
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Cn
113
Nh
114
Fl
115
Mc
116
Lv
117
Ts
118
Og
(267)
(268)
(271)
(270)
(269)
(278)
(281)
(281)
(285)
(286)
(289)
(289)
(293)
(294)
(294)
57
La
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
138.91
140.12
140.91
144.24
(145)
150.36
151.96
157.25
158.93
162.50
164.93
167.26
168.93
173.04
89
Ac
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
[227]
232.04
231.04
238.03
[237]
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(223)
[226]



(262)
95.94
54.94
14.01
10
Ne
3B
(3)
21
Sc
92.91
(8)
26
Fe
8B
(9)
27
Co
12.01
4.003
METALS, METALLOIDS, NONMETALS
Columns: “Groups”; Rows: “Periods”
Group Names to LEARN
ELEMENTS IN THE
O 1A(1): Alkali Metals
SAME GROUP
O 2A(2): Alkaline Earth Metals
ARE USUALLY
O 6A(16): Chalcogens
CHEMICALLY SIMILAR
O 7A(17): Halogens
O 8A(18): Noble Gases (somewhat old-fashioned, but still in use)
Know/be aware of the A and B group designations as shown above. While now not official, they are
too popular and enduring to not be aware of. (They also have uses that we will take advantage of.)
 “A” group elements are “main group” or “representative” elements.
 “B” group elements are “transition metals.”
The 14-column-wide section at the bottom would fit in between Groups 2A and 3B if there were space.
Two numbering conventions are in use.
 American convention: “A” and “B” groups, as shown above
[Europeans reverse the “A” and “B” designations for 3 to 8!]
New International convention: Consecutive numbers 1 to 18 (future: 1 to 32?)
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LEARN the names/symbols of the following elements now. (You’ll pick up others as we go along.)
Al
aluminum
Au
gold
P
phosphorus
Ar
argon
He
helium
Pt
platinum
As
arsenic
H
hydrogen
Pu plutonium
Ba
barium
I
iodine
K
potassium
Bi
bismuth
Kr
krypton
Rn radon
B
boron
Fe
iron
Si
silicon
Br
bromine
Pb
lead
Ag silver
Cd
cadmium
Li
lithium
Na sodium
Ca
calcium
Mg
magnesium
Sr
strontium
C
carbon
Mn
manganese
S
sulfur
Cl
chlorine
Hg
mercury
Sn tin
Cr
chromium
Ne
neon
Ti
titanium
Co
cobalt
N
nitrogen
U
uranium
Cu
copper
Ni
nickel
Xe xenon
F
fluorine
O
oxygen
Zn zinc
Bonding in Compounds
 ionic compound
o formed from transfer of electron(s) from one atom (usually a metal) to another (usually
a nonmetal)  Figure 2.10
o composed of ions (+ ion = cation; – ion = anion)
o possesses no net charge (electrically neutral)
o contains ionic bonds: interaction between + and – charges
In formation of ions, A-group elements gain or lose electrons so as to have
the same number of electrons as the nearest noble gas.
 covalent compound
o formed from sharing of electrons between atoms (both usually nonmetals)
o most consist of molecules
o contains covalent bonds: sharing of electrons between nuclei of bonded atoms
 Also, some elements exist naturally as covalently bonded molecules (Fig. 2.14)
H2
N2
O2
P4
S8
Se8
F2
Cl2
Br2
I2
Plus carbon—diamond, graphite, C60, C70, etc.
2-5
Naming Compounds and Writing Formulas (Chemical Nomenclature)
You need to learn the language of chemistry, otherwise you would sound like this….
Backwards words say to used I. Again go I there. [Shucks] oh!
--paraphrased from George Carlin
LEARN the following:
o FIGURE 2.17, Common Monatomic Ions
o TABLE 2.3, Common Polyatomic Ions
o TABLE 2.2, Numerical Prefixes for Hydrates and Binary Covalent Compounds
o SECTION 2.6:
 Rules for naming and writing formulas for Ionic Compounds and Hydrates
 Rules for naming and writing formulas for Inorganic Acids
 Rules for naming and writing formulas for Binary Covalent Compounds
2-6