Download Chapter 2: Atoms, Molecules, Ions 2.1 Atomic Theory • Democritus

Chapter 2: Atoms, Molecules, Ions
2.1
Atomic Theory
•
Democritus (5th century BC) - matter is composed of indivisible
particles call atomos
•
Alchemists (100 - 1600 A. D.) - made many discoveries while
pursuing metaphysical goals
•
Robert Boyle (17th century) - quantitative measurements of the
properties of matter
•
Joseph Priestley and Antoine-Laurent Lavoisier (18th century)
- showed that combustion is a reaction between matter and
oxygen
•
John Dalton - Atomic Theory (ca. 1808)
1.
Elements are composed of extremely small particles call
atoms. All atoms of a given element are identical, having
the same size, mass, and chemical properties. The atoms
of one element are different from the atoms of all other
elements. (Fig.2.1, p. 36)
2.
Compounds are composed of atoms of more than one
element. In any compound, the ratio of the numbers of
atoms of any two of the elements present is either and
integer or a simple fraction. (incorporates laws of definite
proportions and multiple proportions)
3.
•
Law of definite proportions - different samples of the
same compound always contain its constituent elements
in the same proportions by mass.
•
Law of multiple proportions - If two elements can
combine to form more than one compound, the masses
of one element that combine with a fixed mass of the
other element are in ratios of small whole numbers (use
CO and CO2 as an example: 12 to 16 versus 12 to 32.)
•
Law of conservation of mass - matter can be neither
created nor destroyed.
A chemical reaction involves only the separation,
combination, or rearrangement of atoms; it does not result
in their creation or destruction.
2
2.2
The Structure of an Atom
An atom is the basic unit of an element that can enter into chemical
combinations.
But what is an atom made of:
Order of discovery:
Cathode Ray tube (Figure 2.3)
charge to mass ratio
-1.76 × 108 coulomb/g
Millikan Oil drop experiment
• charge of electron = -1.60 × 10-19 coulombs
• mass of the electron is 9.09 × 10-28 g (very small !)
•
Atoms are composed of
•
electrons
negatively charged
mass = 9.1096 × 10-31 kg
•
protons
positively charged
mass = 1.6727 × 10-27 kg
•
neutrons
neutral
mass = 1.6750 × 10-27 kg
• Radioactivity - Radioactive substances emit 3 kinds of particles
and/or radiation
• alpha (α) particles (helium nuclei)
• beta (β) particles (electrons)
• gamma (γ) rays
Radioactive Particles (Fig. 2.6)
Rutherford experiment (Fig. 2.8) relate to α particles
• Proton
Chapter 2
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• Proton
Ernest Rutherford discovered that the positive charge was
not spread out over the entire volume of the atom but was
concentrated in a small, very dense core (the nucleus).
Have the same but opposite charge as the electrons, but
2000 times more mass
• Neutron
discovered by James Chadwick in 1932
2.3. Atomic Number, Mass Number, and Isotopes
• Atomic number (Z) = number of protons in nucleus = number of
electrons in a neutral atom.
• Mass number (A) = sum of number of protons plus number of
neutrons in nucleus
A
Z
X
Examples of symbols: H, C, O, U-235 (92)
• Element - a form of matter in which all of the atoms have the
same atomic number
• Two atoms of the same element can have different mass
numbers, however. These are isotopes--- the same number of
protons but a different number of neutrons
• Examples: 1-H, 2-H, 3-H, C-12,13,14; U-235, 238
• Example: How many protons, neutrons, and electrons are in
the following isotope of copper? Cu-63
Answer: 29 protons, 34 neutrons
Chapter 2
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2.4. The Periodic Table
• Figure 2.10
• organized by atomic number
• General types of elements
• metals, metalloids, nonmetals (locate)
• main group elements (all 1A, 2A, 3 - 8 A) (locate)
• Special names
• 1A - alkali metal
• 2A - alkaline earth metals
• 7A - halogens
• 8A - noble gases
2.5. Molecules and Ions
1. Molecule -- an aggregated of at least two atoms in a definite
arrangement held together by chemical forces (bonds)
• 7 diatomic elements: H, N, O, F, Cl, Br, I
H2, N2, O2, F2, Cl2, Br2, I2 diatomic
ball and stick model (Fig 2.12)
space filling model (Fig 2.12)
• Water molecule ("polyatomic") – draw
2. Ion - charged species formed from a neutral atom or molecule
when electrons are gained or lost as a result of a chemical change
• cations - positive
• anions - negative
• Examples Na and Cl
# of protons and electrons in each
• Common cations - Na+1, Ca+2, Al+3, Fe+2 or +3
• common anions: nonmetals: Cl-, O2Chapter 2
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2.6. Chemical Formulas - express the composition of molecules and
ionic compounds in terms of chemical symbols
•
Levels of Structure
• elemental composition
• empirical formula
• molecular formula
• constitution
• configuration
• conformation
• Molecular formula - shows the exact number of atoms of each
element in the smallest unit of a substance
Methane CH4
structural formula, ball and stick, space filling (Fig 2.12)
Allotrope: one of two or more forms of an element, e.g.,
oxygen, O2 and ozone, O3
Empirical formula- tells which atoms are present and the
simplest whole number ratio of the atoms (determined
by experiment)
glucose contains C, H, O in 1:2:1 ratio
Empirical formula is CH2O
other experiments show that molecular formula is six time
the empirical form. i.e., C6H12O6
Chapter 2
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• Ionic compounds - discrete molecules are not present; ionic
compounds are represented by their empirical formulas
• Fig. 2.13 --- Structure of solid NaCl Ions are stacked in
cube and touching each other
• sum of positive charges equals the sum of negative
charges
• subscript of cation equals charge of anion and vice versa
potassium bromide
KBr one K+ and one Brzinc iodide ZnI2
aluminum oxide Al2O3
one Zn+2 and two Itwo Al+3 and three O-2
2.7. Naming compounds (Nomenclature)
Four types of compounds:
A.
ionic
B.
molecular
C.
acids
D.
bases, hydrates
Types of names
• Common names - water is a common name: there is
nothing in this name that tell us that it contains hydrogen and
oxygen
Formula
Common name Systematic name
AgCl
Lunar caustic
silver chloride
H2SO4
MgSO4
Oil of vitriol
sulfuric acid
Epsom salts
magnesium sulfate
⇒ Systematic names - sodium chloride is a systematic name;
the name indicates what species are present
Chapter 2
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Binary Compounds
(two elements)
IONIC Cmpds
COVALENT Cmpds
(molecular)
(bonding- electrostatic)
bonding - via shared electrons
metals and nonmetals
elements from left and right side
of the periodic table
nonmetals
elements on same side (right side)
of the periodic table
ion charges show element ratios
prefixes must be used to
show element ratios
+ ion - name same as element
- ion - drop ending & add ide
Chapter 2
element on left named as element
element on right - drop ending and add ide
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General rules for naming binary compounds (2 elements)
e.g., NaCl, FeBr3, Al2O3, N2O5, P4O10
A.
Binary IONIC compounds : (One metal and one nonmetal)
2 elements- 2 words
• 1st word
metals named first - use name of element (no extra suffix)
• 2nd word -nonmetal - use name of element plus ide suffix
e.g., (remember that + charges = - charges to get a
neutral complex)
sodium bromide NaBr
calcium oxide
CaO
barium chloride BaCl2
aluminum oxide Al2O3
Problem
Write the formulas for potassium sulfide .
• Common cations of metals use periodic table
• +1
Li+1, Na+1, K+1, Ag+1
• +3
Mg+2, Ca+2,...Cu+2, Fe+2, Pb+2, Mn+2, Hg+2,
Sn+2, Zn+2
Al+3, Cr+3, Fe+3
• others
H+, NH4+, Hg2+
• +2
• Common anions of nonmetals
• -1
F-1, Cl-1, Br-1, I-1
• -2
• -3
Chapter 2
O-2, S-2
N-3, P-3
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• Polyatomic Anions –
two or more atoms combined in a single charged unit
NH4+
H3O+
NO3PO43HCO3-
ammonium ion
hydronium ion
nitrate ion
phosphate ion
hydrogen carbonate (or bicarbonate ion)
LEARN THE 29 POLYATOMIC ANIONS
(see Table at end of these notes, page 13)
• But, some metals form more than one kind of ion. Then what?
(usually a transition metal)
• To name these types of metals, indicate the type of cation
(charge) in parentheses using Roman numerals.
e.g., MnO
manganese (II) oxide
Mn2O3
manganese (III) oxide
MnO2
manganese (IV) oxide
avoids problem with ous and ic
• common metals with more than one type of cation
Chapter 2
• +1, +2
Cu, Hg
• +2, +3
Fe, Co, Ni
• +2, +4
Sn, Pb
• +2, +3, +6
Cr
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Problem: Write the formulas for the following compounds:
tin(II) fluoride
SnF2
mercury (II) oxide
HgO
mercury (I) iodide
Hg+ exists as Hg2+2 so Hg2I2 is the answer
B. Binary MOLECULAR compounds : (Two nonmetals)
• If 2 nonmetals are in different groups, the one to the left in
the periodic table is named first
• If 2 nonmetal are in the same group, the one farthest down
that group in the periodic table is named first
• The suffix -ide is added to the element named second
• Counting prefixes (Table 2.4) are used with each
name (Never use prefixes with IONIC compounds!!!!)
• (mono is not used with the first name- it's "understood")
Prefixes
mono
1
hexa
6
di
2
hepta
7
tri
3
octa
8
tetra
4
nona
9
penta
5
deca
10
Examples
Chapter 2
CO
carbon monoxide
CO2
SO2
carbon dioxide
SO3
PCl3
sulfur trioxide
PCl5
phosphorus pentachloride
NO2
N2O4
nitrogen dioxide
sulfur dioxide
phosphorus trichloride
dinitrogen tetroxide
Chem 1303
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Practice exercises
NF3
Cl2O7
nitrogen trifluoride
sulfur tetrafluoride
SF4
dinitrogen pentoxide
N 2 O5
dichlorine heptoxide
C. Acids and Bases
• An acid is a substance that yields hydrogen ions (H+) when
dissolved in water
1. Some SIMPLE acids
Table 2.5
Anion
F- (fluoride)
Corresponding Acid
HF (hydrofluoric acid)
Cl- (chloride)
HCl (hydrochloric acid)
Br- (bromide)
HBr (hydrobromic acid)
I- (iodide)
HI (hydroiodic acid)
CN- (cyanide)
HCN (hydrocyanic acid)
S2- (sulfide)
H2S (hydrosulfuric acid)
• Acids that contain hydrogen, oxygen and another element
are called oxoacids.
Chapter 2
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2. OXOACIDS
Naming rules:
• remove the ate and add ic
• remove the ite and add ous
• NO3- (nitrate) and NO2- (nitrite);
• SO4-2(sulfate) and SO3-2 (sulfite),
• other halogens follow chlorine example below
Acetate ion
acetic acid
OCH3
OH
CH3
C
O
O
sulfite SO3-2
sulfurous acid
H2SO3 or
HOSOOH
sulfate SO4-2
sulfuric acid
H2SO4
HOSO2OH
hypochlorous acid
HOCl
hypochlorite ClOchlorite
ClO2-
chlorous acid
HClO2
HOClO
chlorate
ClO3-
chloric acid
HClO3
HOClO2
perchloric acid
HClO4
HOClO3
perchlorate ClO4-
Chapter 2
C
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• A base is a substance that gives hydroxide anions (OH-) when
dissolved in water.
NaOH
sodium hydroxide
KOH
potassium hydroxide
Ba(OH)2
barium hydroxide
NH3
ammonia
NH3 + H2O
D.
NH4+ + OH-
Hydrates
• Hydrates are compounds that have a specific number of
water molecules attached.
BaCl2 . 2 H2O
Sr(NO3)2 . 4 H2O
barium chloride dihydrate
strontium nitrate tetrahydrate
2.7. Introduction to Organic Compounds
Hydrocarbons – compounds of carbon and hydrogen
alkanes –Table 2.8
A.
Alkanes
CH4
C2H6
C3H8
C4H10
C5H12
Chapter 2
CnH2n+2
methane
ethane
propane
butane
pentane
C6H14
C7H16
C8H18
C8H20
C10H22
hexane
heptane
octane
nonane
decane, etc......
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Selected Polyatomic Ions (Memorize these!)
+1
-1
-2
NH4+ (ammonium)
CN- (cyanide)
O2-2 (peroxide)
H3O+ (hydronium)
OCN- (cyanate)
C2O4-2 (oxalate)
-3
SCN- (thiocyanate)
OH- (hydroxide)
CrO4-2 (chromate)
C2H3O2- (acetate)
Cr2O7-2 (dichromate)
MnO4- (permanganate)
NO3- (nitrate)
NO2- (nitrite)
H2PO4- (dihydrogen
phosphate)
HPO4-2 (hydrogen
phosphate)
PO4-3
(phosphate)
HCO3- (hydrogen
carbonate or
bicarbonate)
CO3-2 (carbonate)
AsO4-3
(arsenate)
HSO4- (hydrogen
sulfate
SO4-2 (sulfate)
HSO3- (hydrogen
sulfite)
SO3-2 (sulfite)
ClO4- (perchlorate)
S2O3-2 (thiosulfate)
ClO3- (chlorate)
ClO2- (chlorite)
ClO- (hypochlorite)
+ bromo, iodo relatives
Chapter 2
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Periodic Table of the Elements
1
2
3
4
5
6
7
IA
VIIIA
(1)
(18)
1
2
H
IIA
IIIA
IVA
VA
VIA
VIIA
He
1.0080
(2)
(13)
(14)
(15)
(16)
(17)
4.0026
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
6.9410
9.0122
10.811
12.011
14.007
15.999
18.998
20.179
11
12
13
14
15
16
17
18
Na
Mg
IIIB
IVB
VB
VIB
VIIB
. . VIIIB . . .
IB
IIB
Al
Si
P
S
Cl
Ar
22.990
24.305
(3)
(4)
(5)
(6)
(7)
(8)
(9)
(10)
(11)
(12)
26.982
28.086
30.974
32.066
35.453
39.948
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
39.098
40.078
44.956
47.880
50.942
51.996
54.938
55.847
58.933
58.690
63.546
65.380
69.723
72.610
74.922
78.960
79.904
83.800
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
85.468
87.620
88.906
91.224
92.906
95.940
98.907
101.07
102.91
106.42
107.87
112.41
114.82
118.71
121.75
127.60
126.90
131.29
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
132.91
137.33
138.91
178.49
180.95
183.85
186.21
190.20
192.22
195.09
196.97
200.59
204.38
207.20
208.98
208.98
209.99
222.02
87
88
89
104
105
106
107
62
63
64
65
66
67
68
69
70
71
Fr
Ra
Ac
Unq
Unp
Unh
Uns
223.02
226.03
227.03
261.11
262.11
263.12
262.12
58
59
60
61
Chapter 2
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
140.12
140.91
144.24
145.91
150.36
151.97
157.25
158.93
162.50
164.93
167.26
168.93
173.04
174.97
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
232.04
231.04
238.03
237.05
244.06
Am Cm
243.06
247.07
Bk
Cf
Es
Fm
Md
No
Lr
247.07
242.06
252.08
257.10
258.10
259.10
260.11
Chem 1303
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Chapter 2
Chem 1303