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The Components of
Matter
Chapter Two
AP Chemistry
Element
Simplest type of substance with unique
physical and chemical properties.
Consists of only one type of atom.
An element cannot be broken down into
any simpler substances by physical or
chemical means.
A few elements exist in nature bonded with
like atoms, e.g. P4, S8, Se8.
The diatomic molecules:
H2, N2, O2, F2, Cl2, Br2, I2
Molecule
Structure that consists of two or more
atoms which are chemically bound
together.
Behaves as an independent unit.
Which is an element and which is a molecule?
i
Molecule
Element
Compound
A substance composed of two or more
elements which are chemically combined.
When a compound is formed, there is
evidence of a chemical reaction.
Evidence of a chemical reaction:
•Energy change
heat, light, sound, motion, electricity, etc.
•Release of a gas
•Color Change
•Precipitate is formed
Evidence of Rxns (Evidence of Rxns.ASF)
Compounds have a fixed mass ratio …
Water is always 2 g H and 16 g O.
Properties of compounds are different
from the individual elements that make
them up.
Na and Cl2 are different from NaCl.
Mixture
A group of two or more elements and/or
compounds that are physically intermingled.
The individual properties of the components of
a mixture are retained.
When a mixture is formed, there is no
evidence of a chemical reaction.
Compound
11
Mixture
Components are present in
one constant ratio only
Components may be
present in a range of ratios
Components may only be
separated chemically
Components may be
separated physically
Components lose their
identities
Components retain their
identities
An energy change between
components and products
takes place
No evidence of an energy
change
Which is a compound and which is a mixture?
Mixture
12
Compound
Distinguishing Elements, Compounds, and
Mixtures at the Atomic Scale
Theses scenes represent an atomic-scale view of three samples of
matter. Describe each sample as an element, compound, or mixture.
(a) mixture
(b) element
(c) compound
The Law of Mass Conservation:
Mass remains constant during a chemical reaction.
reactant 1
+
product
reactant 2
total mass
=
total mass
calcium oxide
+
carbon dioxide
calcium carbonate
CaO
+
CO2
CaCO3
56.08 g
+
44.00 g
100.08 g
Law of Definite (or Constant)
Composition:
No matter what its source, a particular
chemical compound is composed of the
same elements in the same parts
(fractions) by mass.
CaCO3
1 atom Ca
1 atom C
3 atoms O
40.08 amu
12.00 amu
3 x 16.00 amu
100.08 amu
40.08 amu
100.08 amu
= 0.401 parts Ca
12.00 amu
100.08 amu
= 0.120 parts C
48.00 amu
100.08 amu
= 0.480 parts O
• 100 g CaCO3
• …always
contains
–40g of Ca
–12g of C
–48g of O
Analysis by Mass
(grams/20.0 g)
Mass Fraction
(parts/1.00 part)
0.40 calcium
0.12 carbon
0.48 oxygen
8.0 g calcium
2.4 g carbon
9.6 g oxygen
1.00 part by mass
20.0 g
Percent by Mass
(parts/100 parts)
40% calcium
12% carbon
48% oxygen
100% by mass
Calculating the Mass of an
Element in a Compound
Ammonium Nitrate
How much nitrogen (N) is in
455kg of ammonium nitrate?
Formula Mass of Cpd:
Add masses from the P.T.
4 x H = 4 x 1.008 = 4.032 g
2 x N = 2 X 14.01 = 28.02 g
3 x O = 3 x 16.00 = 48.00 g
80.052 g
Therefore g nitrogen/g cpd
28.02g N = 0.3500gN/g cpd
80.052g cpd
Note: .3500 X 100 = 35% N
This is % Composition
455kg x 1000g = 455,000g NH4NO3
kg
455,000g cpd x 0.3500g N
g cpd
= 1.59 x 105g nitrogen
Or:
28.02
kg
N
455 kg NH4NO3 X
80.052 kg NH4NO4
= 159 kg Nitrogen
Calculating the Mass of an
Element in a Compound
Pitchblende (PB) is a cpd that
contains U. Analysis shows that
84.2 g of pitchblende contains
71.4 g of U with O as the only
other element. How many grams
of U can be obtained from 102 kg
of pitchblende?
The mass ratio of U/PB will be
the same no matter the amount
of pitchblende.
102kg PB x 71.4kg U
84.2 kg PB
= 86.5 kg U
86.5 kg U = 86500 g U
Law of Multiple Proportions
003_MULTIPLEPROP.mov
If elements A and B react to
form two compounds, the
different masses of B that
combine with a fixed mass of
A can be expressed as a ratio
of small whole numbers.
The Atomic Basis of the Law
of Multiple Proportions
CO, CO2
CO = 12g C + 16g O
CO2 = 12g C + 32g O
16g O/32g O = 1/2
H2O, H2O2
H2O = 2g H + 16g O
H2O2 = 2g H + 32g O
16g O/32g O = 1/2
NO2, N2O, N2O3, N2O5
Dalton’s Atomic Theory
The Postulates
1. All matter consists of atoms.
2. Atoms of one element cannot be converted into
atoms of another element.
3. Atoms of an element are identical in mass and
other properties and are different from atoms of
any other element.
4. Compounds result from the chemical combination of
a specific ratio of atoms of different elements.
Dalton’s Model of the Atom
Explains the Mass Laws
Mass conservation
Atoms cannot be created or destroyed
postulate 1
or converted into other types of atoms.
postulate 2
Since every atom has a fixed mass,
postulate 3
during a chemical reaction atoms are
combined differently, and therefore, there is no
mass change overall.
Definite Composition
Atoms are combined in compounds in
specific ratios
and each atom has a specific mass.
postulate 3
postulate 4
Each element has a fixed fraction of the total mass
in a compound.
Multiple Proportions
Atoms of an element have the same
mass
and atoms are indivisible.
postulate 3
postulate 1
When different numbers of atoms of elements
combine, they must do so in ratios of small, whole
numbers.
Visualizing the Mass Laws
Theses scenes represent an atomic-scale view of a chemical reaction. Which
of the mass laws: mass conservation, definite composition, or multiple
proportions is (are) illustrated?
Mass conservation illustrated if number of each atom before and after
reaction remains constant. Definite composition illustrated by formation of
compounds that always have the same atom ratio. Different compounds
made of same elements have small whole number ratios of those elements
illustrates multiple proportions.
Seven purple and nine green atoms in each circle, mass conserved.
One compound formed has one purple and two green, definite
composition. Law of multiple proportions does not apply.
Properties of Cathode Rays.
Observation
Thompson’s Experiment
Conclusion
Ray bends in magnetic field
Consists of charged particles
Ray bends toward positive plate in
electric field
Consists of negative particles
Ray is identical for any cathode
Particles found in ALL matter
Determined Q/m as 1.756 x 108 C/g
Credited with discovering the electron
+ anode
- cathode
Cathode
Thompson’s Model of the Atom
+ +
+
+
+
+
+
+
+
+
+
+
+
+
Millikan’s Experiment
Millikan
In a replication of Millikan’s oil drop experiment you
observe the following charges, measured in
coulombs, on different oil drops:
-9.6 x 10-19
-1.328 x 10-17
-7.84 x 10-18
-4.16 x 10-18
-4.64 x 10-18
-1.408 x 10-17
-8.16 x 10-18
-1.28 x 10-17
-1.488 x 10-17
-1.456 x 10-17
-1.6 x 10-18
-8.0 x 10-19
-1.12 x 10-18
-1.104 x 10-17
-4.16 x 10-18
-1.6 x 10-19
-3.84 x 10-18
-6.4 x 10-19
-3.2 x 10-19
-1.28 x 10-18
-4.8 x 10-19
-6.24 x 10-18
-9.44 x 10-18
-5.44 x 10-18
-1.104 x 10-17
-1.28 x 10-17
-6.24 x 10-18
-2.72 x 10-18
-1.232 x 10-17
-1.28 x 10-17
What should you conclude is the charge of an
individual electron?
-9.6 x 10-19 (6)
-1.328 x 10-17 (83)
-7.84 x 10-18 (49)
-4.16 x 10-18 (26)
-4.64 x 10-18 (29)
-1.408 x 10-17 (88)
-8.16 x 10-18 (51)
-1.28 x 10-17 (80)
-1.488 x 10-17 (93)
-1.456 x 10-17 (91)
-1.6 x 10-18 (10)
-8.0 x 10-19 (5)
-1.12 x 10-18 (7)
-1.104 x 10-17 (69)
-4.16 x 10-18 (26)
-1.6 x 10-19 (1)
-3.84 x 10-18 (24)
-6.4 x 10-19 (4)
-3.2 x 10-19 (2)
-1.28 x 10-18 (8)
-4.8 x 10-19 (3)
-6.24 x 10-18 (39)
-9.44 x 10-18 (59)
-5.44 x 10-18 (97)
-1.104 x 10-17 (69)
-1.28 x 10-17 (80)
-6.24 x 10-18 (39)
-2.72 x 10-18 (17)
-1.232 x 10-17 (77)
-1.28 x 10-17 (80)
The smallest charge on an oil
-19
drop was 1.60 X 10 C.
Thus he concluded that the
charge of an electron was
1.60 X 10-19 coulombs!
determined by J.J. Thomson and
others
1.6 x
10-19
1
gram
CX
8
1.756 X 10 C
= 9.11 X
-28
10
g
Rutherford’s Experiment
Simple Rutherford
Complex Rutherford
Rutherford
1. The atom is mostly empty space!
2. The charge is concentrated in the center.
Credited with discovering the nucleus
Rutherford’s Model of the Atom
Chadwick’s Experiment
1. The charge on the particle was found to
be zero.
2. The mass of the particle was found to be
1.67 x 10-24 g, or the same as the p+.
Credited with discovering the neutron
Chadwick’s Model of the Atom
General Features of the Atom Today.
•The atom is an electrically neutral, spherical entity composed of a
positively charged central nucleus surrounded by one or more
negatively charged electrons.
•The atomic nucleus consists of protons and neutrons.
The Modern Reassessment
of the Atomic Theory
1. All matter is composed of atoms.
The atom is the smallest body that retains the unique
identity of the element.
2. Atoms of one element cannot be
converted into atoms of another element
in a chemical reaction.
Elements can only be converted into other elements
in nuclear reactions.
3. All atoms of an element have the same
number of p+ and e-, which determines
the chemical behavior of the element.
Isotopes of an element differ in the number of n0, and
thus in mass number. A sample of the element is
treated as though its atoms have an average
mass.
4. Compounds are formed by the chemical
combination of two or more elements in
specific ratios.
Properties of the Three Key Subatomic Particles
Charge
Mass
Location
Name(Symbol) Relative Absolute(C)* Relative(amu)† Absolute(g) in the Atom
Proton (p+)
1+
+1.60218x10-19
1.00727
1.67262x10-24 Nucleus
Neutron (n0)
0
0
1.00866
1.67493x10-24 Nucleus
1-
-1.60218x10-19
Electron (e-)
0.00054858
9.10939x10-28
* The coulomb (C) is the SI unit of charge.
†
The atomic mass unit (amu) equals 1.66054x10-24 g.
Outside
Nucleus
The AMU
• Atomic Mass Unit is a relative unit and
does not represent the mass in real
masses.
• 1 amu = 1/12 C-12 = 1.6606 X 10-24 g.
Modern terminology now uses the unit of
Daltons (D) in place of AMU’s.
Atomic Notation
Mass Number (p+ + n0)
A
X
Z
Atomic Symbol
Z & A are integers
Atomic Number (p+ )
Example:
Manganese
55
Mn
25
A
Z
X
p+ = 25
e- = 25
n0 = 30
Example:
Chlorine
35
Cl
17
A
Z
X
p+ = 17
e- = 17
n0 = 18
Ions
• Ions are charged atoms.
• In many cases, chemical bonding is the
attraction between oppositely charged
particles. (Ionic Bonding)
How do atoms become charged ions?
• By the addition or removal of e-.
(Neutrons are neutral and would have no effect;
changing protons would change the element. The
only particle that is charged and of low enough
energy to be removed is the electron.)
Example:
Lithium Ion
7
Li
3
A
Z
X
+1
p+ = 3
e- = 2
n0 = 4
Example:
Selenide Ion
79
-2
Se
34
A
Z
59
X
p+ = 34
e- = 36
n0 = 45
Isotopes
• Atoms of the same element
which differ slightly in atomic
mass. (A)
–Same element - same
properties.
• Electrons are considered
massless and the number of
protons (Z) can not change.
• Therefore the different
masses (A) are caused by
different numbers of neutrons.
Determining the Number of Subatomic Particles in the Isotopes of an Element
PROBLEM: Silicon (Si) has three naturally occurring isotopes: 28Si, 29Si,
and 30Si. Determine the number of protons, neutrons, and
electrons in each silicon isotope.
PLAN:
Mass number (A), protons + neutrons, is given for the listed
isotopes. Atomic number (Z), number of protons, for each element
is given in the periodic table and equal to the number of electrons.
Number of neutrons is determined using equation 2.2.
SOLUTION: The atomic number of silicon is 14. Therefore
28Si
has 14p+, 14e- and 14n0 (28-14)
29Si
has 14p+, 14e- and 15n0 (29-14)
30Si
has 14p+, 14e- and 16n0 (30-14)
How do we know there are
different isotopes?
• A mass spectrograph is used
to deflect ions of an element
as they move through a
magnetic field.
http://www.colby.edu/chemistry/OChem/DEMOS/MassSpec.html
Mass Spectrometer and Its Data
a
A
a
Heaviest Isotope
Lightest Isotope
Why are the atomic masses
different from the A?
• The atomic weights (masses)
are the average for all the
isotopes.
• B has two isotopes … 10B, 11B.
10B=
10.0129 and 11B= 11.0093.
Are these the masses (A) on the
periodic table?
No!
Why not?
Because 10B is found as 20.00%
and 11B @ 80.00%
A stream of B ions are shot
through a magnetic field. The
following is observed.
Magnet
Ion
Gun
10B
11B
Ion Detector Screen
Calculation
• Average Atomic Mass =
%1 (X1) + %2(X2) + etc
100%
(20.0)10.0129 + (80.0)11.0093
100
= 10.81
Another Type
• An element (Sb) has an
Atomic Mass of 121.76 amu
• There are only two naturally
occuring isotopes.
• Sb –121 @ 120.904 amu
• Sb – 123 @ 122.904 amu
• Since percentages are based
on 100, assume a 100 particle
sample.
121
X =
Sb
(100 - X) = 123Sb
Calculate their percent
abundances.
121.76 = 120.904 (x) + 122.904(100-x)
100
121X
= 57.000% and
123X
= 43.000%
• Silicon has three isotopes.
Masses are Si-28 @ 27.977, Si-29
@ 28.977 and Si-30 @ 29.974 amu.
The % abundance for the middle
isotope is 4.70%. Calculate the %
abundances for the two remaining
isotopes.
•Si-28 = 92.2% and Si-30 = 3.1%
The Modern Periodic Table.
Organization of the Elements
06 Track 6.wma
• 1869: Dmitri Mendeleev –Published an
organizational scheme for
the elements.
–He called it the Periodic
Table.
• Organized his elements by
atomic mass in two rows of 7
elements and two rows of 17
elements.
• Noticed that the properties of
the elements repeated so he
put elements that behaved
the same in the same vertical
column.
• Left blank spots on his
periodic table and predicted
the discovery and properties
of several undiscovered
elements.
__________________________
• Problems were found with the
table.
• Ar -- K Te -- I Co -- Ni
• If Mendeleev’s Table was to
be correct, these elements
should have been reversed.
• Elements could not be
reversed because of their
properties.
• This lead to a new version of
the Periodic Table.
The Modern Periodic Law
• Henry Mosely - The
properties of the elements
are periodic when they are
arranged in increasing order
of their Atomic Number or
Protons.
“Patterns” on the Periodic
Table.
• Columns:
Group or family - elements
with similar properties.
• Several groups have
descriptive names.
Alkali Metals - Group IA
• React with water to form an
alkaline or basic solution.
• Forms +1 ions (Cations).
“s” sublevel electrons.
Base
Na(S) + H2O(L) --> NaOH(aq) + H2(g)
Alkaline Earth Metals - Group IIA
• Elements also react with
water to form a base.
•Forms +2 ions (Cations).
•“s” sublevel electrons.
Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)
GroupI_And_II_Rxn.rm
Halogens - Group VIIA
• “Salt Formers”
Forms -1 ions (Anions).
“p” sublevel electrons.
Na(S) + Cl2(g) --> NaCl(S)
PhysicalPropsofHalogens.MOV
“Salt”
Noble Gases - Group VIIIA
• Gases that are highly
unreactive because their
electron configuration is
very stable.
• Does not form ions.
• “p” sublevel electrons
completed.
Chalcogens - Group VIA
• any of the elements oxygen,
sulfur, selenium, and
tellurium.
Forms -2 ions (Anions).
“p” sublevel electrons.
2Ca(s) + O2 (g) → 2CaO(s)
The Periodic Table of the Elements
H
Li Be
NaMg
K Ca Sc Ti
R Sr Y Zr
b
Cs Ba La Hf
Fr Ra Ac Rf
B C N
Al Si P
V Cr Mn Fe Co Ni Cu Zn Ga Ge As
NbMo Tc R Rh Pd Ag Cd In Sn Sb
Ta W Re uOs Ir Pt Au Hg Tl Pb Bi
Du Sg Bo HaMe
O
S
Se
Te
Po
He
F Ne
Cl Ar
Br Kr
I Xe
At Rn
Ce Pr Nd PmSm Eu Gd Tb Dy Ho Er TmYb Lu
Th Pa U Np PuAmCmBk Cf Es FmMd No Lr
The Alkali Metals
The Halogens
The Alkaline
Earth Metals
The Noble Gases
The Periodic Table of the Elements
H
Li Be
NaMg
K Ca Sc Ti
Rb Sr Y Zr
Cs Ba La Hf
Fr Ra Ac Rf
B C N
Al Si P
V CrMn Fe Co Ni Cu Zn Ga Ge As
Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb
Ta W Re Os Ir Pt Au Hg Tl Pb Bi
Du Sg Bo Ha Me
O
S
Se
Te
Po
F
Cl
Br
I
At
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np PuAmCm Bk Cf Es FmMd No Lr
Boron family
Nitrogen family
Carbon Family
Chalcogens
He
Ne
Ar
Kr
Xe
Rn
Transition Metals – B
•Generally form +2 cations
•Many are multivalent
•Formed by d-sublevel eScandium
Titanium
Vanadium
Chromium
Manganese
Iron
Cobalt
Nickel
Copper
Zinc
21
22
23
24
25
26
27
28
29
30
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
44.956
4s23d1
47.88
4s2 3d2
50.942
4s2 3d3
51.996
4s1 3d5
54.938
4s2 3d5
55.847
4s23d6
58.933
4s23d7
58.633
4s2 3d8
63.546
4s1 3d10
65.39
4s23d10
Yttrium
Zirconium
Niobium
Molybdenum
Technetium
Ruthenium
Rhodium
Palladium
Silver
Cadmium
39
40
41
42
43
44
45
46
47
48
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
88.906
5s24d1
91.224
5s2 4d2
92.906
5s1 4d4
95.94
5s1 4d5
(98)
5s2 4d5
101.07
5s14d7
102.906
5s14d8
106.42
5s1 4d9
107.868
5s1 4d10
112.411
5s24d10
Lanthanum
Hafnium
Tantalum
Tungsten
Rhenium
Osmium
Iridium
Platinum
Gold
Mercury
57
72
73
74
75
76
77
78
79
80
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
138.906
6s25d1
178.49
6s2 5d2
180.948
6s2 5d3
183.85
6s2 5d4
186.207
6s2 5d5
190.2
6s25d6
192.22
6s25d7
195.08
6s1 5d9
196.967
6s1 5d10
200.59
6s25d10
Actinium
Rutherfordium
Dubnium
Seaborgium
Bohrium
Hassium
Meitnerium
Darmstadtium
Roentgenium
Ununbium
89
104
105
106
107
108
109
110
111
Ac
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
112
Uub
227.028
7s26d1
(261)
7s2 6d2
(262)
7s2 6d3
(263)
7s2 6d4
(262)
7s2 6d5
(265)
7s26d6
(266)
7s26d7
(281)
7s1 6d9
(280)
7s1 6d10
(285)
7s26d10
Inner Transition Metals
•Lanthanide series
•Actinide series
•Formed by f-sublevel eCerium
Praseodymium
Neodymium
Promethium
Samarium
Europium
Gadolinium
Terbium
Dysprosium
Holmium
Erbium
Thulium
Ytterbium
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
140.115
6s24f15d1
140.908
6s24f3
144.24
6s2 4f4
(145)
6s2 4f5
150.36
6s2 4f6
151.965
6s2 4f7
157.25
6s2 4f75d1
158.925
6s2 4f9
162.50
6s2 4f10
164.930
6s2 4f11
167.26
6s2 4f12
168.934
6s2 4f13
173.04
6s24f14
174.967
6s24f145d1
Thorium
Protactinium
Uranium
Neptunium
Plutonium
Americium
Curium
Berkelium
Californium
Einsteinium
Fermium
Mendelevium
Nobelium
Lawrencium
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
232.038
7s26d2
231.036
7s25f26d1
238.029
7s2 5f36d1
237.048
7s2 5f46d1
(244)
7s2 5f6
(243)
7s2 5f7
(247)
7s2 5f76d1
(247)
7s2 5f9
(251)
7s2 5f10
(252)
7s2 5f11
(257)
7s2 5f12
(258)
7s2 5f13
(259)
7s25f14
(260)
7s25f146d1
Lutetium
The Periodic Table of the Elements
H
Li Be
B C N
Na Mg
Al Si P
K Ca Sc Ti V CrMn Fe Co Ni Cu Zn Ga Ge As
Rb Sr Y Zr NbMo Tc Ru Rh PdAg Cd In Sn Sb
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi
Fr Ra Ac Rf Du Sg Bo Ha Me
O
S
Se
Te
Po
He
F Ne
Cl Ar
Br Kr
I Xe
At Rn
Transition Metals
Ce Pr Nd PmSm Eu Gd Tb DyHo Er TmYb Lu
Th Pa U Np PuAmCmBk Cf Es FmMd No Lr
Lanthanides: Rare
Earth Elements
Actinides
Period
• Row on the periodic table.
• There are 7 periods on the
periodic table.
–One period for each
“layer” or energy level of
electrons.
The Periodic Table of the Elements
H
Li Be
Na Mg
K Ca Sc
Rb Sr Y
Cs Ba La
Fr Ra Ac
Ti
Zr
Hf
Rf
B C N
Al Si P
V Cr Mn Fe Co Ni Cu Zn Ga Ge As
NbMo Tc Ru Rh Pd Ag Cd In Sn Sb
Ta W Re Os Ir Pt Au Hg Tl Pb Bi
Du Sg Bo Ha Me
O
S
Se
Te
Po
F
Cl
Br
I
At
He
Ne
Ar
Kr
Xe
Rn
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er TmYb Lu
Th Pa U Np PuAmCm Bk Cf Es FmMdNo Lr
Metals
Non Metals
Semi - metals
Metalloids
A Biological Periodic Table.
Bonding
• Most elements occur
combined … very few are
found in elemental form.
• Noble gases are found
uncombined.
• Au, Ag,Cu, and Pt do not react
easily and are often found
uncombined.
•Some elements combine with
themselves …
Diatomics
S8 and P4
Types of Chemical Formulas
A chemical formula is comprised of element symbols and numerical
subscripts that show the type and number of each atom present in
the smallest unit of the substance.
An empirical formula indicates the relative number of atoms of
each element in the compound. It is the simplest type of formula.
The empirical formula for hydrogen peroxide is HO.
A molecular formula shows the actual number of atoms of
each element in a molecule of the compound.
The molecular formula for hydrogen peroxide is H2O2.
A structural formula shows the number of atoms and the
bonds between them, that is, the relative placement and
connections of atoms in the molecule.
The structural formula for hydrogen peroxide is H-O-O-H.
H2O
Ammonia - NH3
Methane - CH4
Salt Crystal
Cl
Na
Single Bond
Double Bond
Which molecule
would be hardest
to pull apart?
Triple Bond
Types of Bonds
Ionic: attraction of oppositely charged
ions resulting from the transfer of e-’s
Reaction is typically between a metal and a
nonmetal
Covalent: attraction based on localized epair sharing between two atoms
Reaction is typically between two nonmetals
Metallic: attraction based on delocalized
e-’s and metal ions
Reaction is typically between two metals
Ionic Bond
• Electrons transferred—usually between
metals and nonmetals.
• Metals lose e-’s and form + ions (cations)
• Nonmetals gain e-’s and form – ions
(anions)
• Metals and nonmetals will lose or gain
e-’s to form a noble gas configuration
• Na → Na+1 + e- (isoelectronic with Ne)
• Cl + e- → Cl-1
(isoelectronic with Ar)
•Sodium is isoelectronic with Ne
•Iso = same
•Chlorine is isoelectronic with Ar
•Almost exclusively in ionic
bonding, the energy of the newly
formed compound is MUCH
lower than the original elements.
The Relationship Between Ions
Formed and the Nearest Noble Gas.
The Periodic Table of the Elements
Most Probable Oxidation State
+1
+4
0
+3 -4 -3 -2 -1 He
H +2
Li Be
Na Mg +3
K Ca Sc
Rb Sr Y
Cs Ba La
B C N O F Ne
+1 +2 Al Si P S Cl Ar
+4 +5
Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge A Se Br Kr
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn sSb Te I Xe
Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac Rf Du Sg Bo Ha Me
+3
+3
Ce Pr Nd PmSm Eu Gd Tb Dy Ho Er TmYb Lu
Th Pa U Np Pu AmCmBk Cf Es FmMd No Lr
Common Monoatomic Ions
Cations
Charge
Formula
1+
2+
3+
H+
Li+
Na+
K+
Cs+
Ag+
Mg2+
Ca2+
Sr2+
Ba2+
Zn2+
Cd2+
Al3+
Name
hydrogen
lithium
sodium
potassium
cesium
silver
magnesium
calcium
strontium
barium
zinc
cadmium
aluminum
Anions
Charge
Formula
1-
HFCl Br I-
Name
hydride
fluoride
chloride
bromide
iodide
2-
O2S2 -
oxide
sulfide
3-
N 3-
nitride
Listed by charge; those in boldface are most common
Metals With Several Oxidation States
Element
Copper
Cobalt
Iron
Manganese
Tin
Lead
Ion Formula
Systematic Name
Common Name
Cu+1
copper(I)
cuprous
Cu+2
copper(II)
cupric
Co+2
cobalt(II)
Co+3
cobalt (III)
Fe+2
iron(II)
ferrous
Fe+3
iron(III)
ferric
Mn+2
manganese(II)
Mn+3
manganese(III)
Sn+2
tin(II)
stannous
Sn+4
tin(IV)
stannic
Pb+2
lead(II)
Pb+4
lead(IV)
Coulomb’s Law
The magnitude of the electrostatic force
between two point electric charges is
directly proportional to the product of the
magnitudes of each charge and inversely
proportional to the square of the distance
between the charges.
k ⋅ Q1 ⋅ Q 2
F=
2
d
charge1× charge2
Forceα
2
Distance
charge1× charge2
Energyα
Distance
Covalent Bond
• Electrons shared—usually between two
nonmetals.
• e-’s are attracted to both nuclei
• More common than ionic bonds
• e-’s are shared in pairs
• Localized—does not extend to other
molecules
Covalent Bonding
Formation of a Covalent Bond between
Two Hydrogen Atoms
Covalently bonded
substances exist as individual
entities.
Ionic substances are
continuous arrays of oppositely
charged ions.
No individual
“molecules”
exist in ionic
compounds.
Metallic Bond
• Electrons shared among all of the atoms
in the substance
• e-’s form an electron sea
The reason metals deform.
metal is deformed
Structure of Metals
Polyatomic Ions
• A charged ion whose atoms are
covalently bonded groups
• Behave as a monatomic –ion (anion)
when joined with a metal.
• Most end in the suffixes –ate or –ite
• Oxygen is the prevalent atom
Polyatomic Ions
•Many important ions:
Carbonate CO3-2
Nitrate NO3Nitrite NO2+
Ammonium NH4 …
A Polyatomic Ion
Some Common Polyatomic Ions
Formula
Name
Formula
Cations
NH4+
H3O+
Anions Cont.
Ammonium
Hydronium
Anions
CH3COOCNOHClOClO2ClO3ClO4NO2NO3-
Name
Acetate
Cyanide
Hydroxide
Hypochlorite
Chlorite
Chlorate
Perchlorate
Nitrite
Nitrate
MnO4CO3-2
HCO3CrO4-2
Cr2O7-2
O2-2
PO4-3
HPO4-2
H2PO4SO3-2
SO4-2
HSO4-
Permanganate
Carbonate
Hydrogen Carbonate
Chromate
Dichromate
Peroxide
Phosphate
Hydrogen Phosphate
Dihydrogen Phosphate
Sulfite
Sulfate
Hydrogen Sulfate
Writing Ionic Formulas
Write the symbol of the
positive metal ion or
polyatomic group first and
include the oxidation
number.
• Write the symbol of the
negative non-metal ion or
polyatomic group second
with the oxidation number.
Using the LCD, write subscripts
to balance the oxidation
numbers and remember:
The algebraic sum of all
charges must be ZERO.
The formulas must be in the
simplest whole number ratio.
The Criss -Cross
Method
• Caution: This method does
not always produce the
SIMPLEST WHOLE NUMBER
ratio.
Aluminum and sulfur:
Al+3 S-2
Al2 S3
Hence Al2+3S3-2
EXAMPLES
• Aluminum and chlorine
Al+3 Cl-1
Al+3Cl3-1
• Magnesium and phosphorus
+2
-3
Mg P
+2
-3
Mg3 P2
Metal Ions With Multiple
Oxidation Numbers.
Numbers
Lead (IV) with oxygen:
+4
-2
Pb O
+4
-2
Pb2 O4
However the answer is:
Pb+4O2-2 … The simplest
whole number ratio.
Lead (II) with oxygen:
+2
-2
+2
-2
Pb O gives Pb2 O2
...or
Pb+2O-2 after it is
reduced.
Polyatomic Ions
• Groups of atoms that act
like a single ion.
• If more than one polyatomic
group is needed, the group
must be wrapped in
parathensis.
Try these:
• calcium acetate
Ca+2(C2H3O2-1)2
• aluminum hydroxide
+3
-1
Al (OH )3
• aluminum sulfate
Al2+3(SO4-2)3
• ammonium phosphate
+1
-3
(NH4 )3PO4
Naming Ionic
Compounds
• IONIC COMPOUNDS = METAL + NON METAL
1. Write the name of the positive metal first
2. Write the ROOT of the negative non metal
then drop the last syllable, sometimes two,
and add IDE .
• Try…
Try… Al2O3
–Aluminum
Aluminum oxide
oxide
• Try… NaCl
–Sodium
Sodium Chloride
Chloride
• Try … Ba3P2
–Barium
Barium Phosphide
Phosphide
Naming ionic compounds
with metals that have more
than one oxidation number.
• Use the “ROMAN “ method.
• Iron has two oxidation
numbers, Fe+2 and Fe+3.
• … then Fe+2Cl2-1 is named
Iron(II
Iron(II)
II) chloride
1
+3
• … and Fe Cl3 is named
Iron(III
Iron(III)chloride
III)chloride
• … and PbCl4 is named
Lead(IV)chloride
Naming Compounds With
Polyatomic Ions
• 1. Write the name of the
positive metal ion or
polyatomic ion.
• 2. Write the name of the
negative polyatomic ion.
Examples
• (NH4+1)3PO4-3
• Ammonium Phosphate
• Fe2(SO4)3
• Fe2+3(SO4-2)3
• Iron(III) Sulfate
Hydrated Compounds
• Some ionic salts often trap
water molecules within their
structures.
• Mechanically held water in a
specific ratio by weight is a
“hydrated” compound.
• CuSO4 . 5 H2O; 36.0% water.
•BaCl2 . 2 H2O; 14.7% water.
How are these hydrates
named?
Numerical Prefixes for
Hydrates and Binary Covalent
Compounds
1 = mono 5 = penta 9 = nona
2 = di
6 = hexa 10 = deca
3 = tri
7 = hepta
4 = tetra
8 = octa
Hence:
CuSO4 . 5 H2O is:
Copper (II) sulfate pentahydrate
BaCl2 . 2 H2O is:
Barium chloride dihydrate
Naming Covalent Compounds.
“Greek “ Method
For the 1st element … name that
element using a prefix indicating the
number of atoms present.
For the 2nd element … name the
second element using a prefix
indicating the number of atoms
present and add root and -ide
ending.
P2O5 is named:
diphosphorus pentaoxide
N2O3 is named:
dinitrogen trioxide
P2S is named:
diphosphorus monosulfide
SiO2 is named:
silicon dioxide …
when there is one atom of the
first element the prefix mono
is not used.
Then CO & CO2 are ….
carbon monoxide and
carbon dioxide
The Following May Have
Mistakes … Correct Them
• SO2 is sodium dioxide
–sulfur dioxide
• Cl2O is chlorine dioxide
–dichlorine monoxide
• SF10 is disulfur decafluoride
–sulfur decafluoride
•MgCl2 is magnesium dichloride
–magnesium chloride (ionic)
Something is wrong with the
second part of each statement.
Provide the correct name or
formula.
Ba(C2H3O2)2 = barium diacetate
Barium is always a +2 ion and acetate is -1. The
“di-” is unnecessary.
Sodium sulfide = (Na)2SO3
An ion of a single element does not need
parentheses. Sulfide is S2-, not SO32-. The
correct formula is Na2S.
Iron(II) sulfate = Fe2(SO4)3
Since sulfate has a 2- charge, only 1 Fe2+ is
needed. The formula should be FeSO4.
Cesium carbonate = Cs2(CO3)
The parentheses are unnecessary. The
correct formula is Cs2CO3.
Basic Oxoanions
• Combos With Two
Possibilities:
The most common are SO4-2
and SO3-2 and are named
sulfate and sulfite.
The ion with the greater # of
oxygen atoms has the ate
ending.
The lesser # of oxygen
atoms has the -ite ending!
Two other common
oxoanions:
NO3- and NO2-
Naming2Seriesof2Oxyanions.swf
Nitrate and nitrite
More Basic Oxoanions Ions
• Combos With Four Possibilities:
– This would be any halogen (group 7A) with
4,3,2,1 oxygen atoms respectively.
BrO4-1 = perbromate
BrO3-1 = bromate
BrO2-1 = Bromite
BrO-1 = hypobromite
NamingASeriesof4Oxyanions.swf
NaClO3 = Sodium chlorate
KNO2 = Potassium Nitrite
BaSO3 =
Barium Sulfite
LiClO4 = Lithium Perchlorate
Ca(BrO)2 = Calcium Hypobromite
NaClO3
Sodium Chlorate
Al(IO2)3
Aluminum Iodite
RbClO2
Rubidium Chlorite
KBrO3
Potassium Bromate
CsClO
Cesium Hypochlorite
LiIO4
Lithium Periodate
Calcium Nitrate = Ca(NO3)2
Ammonium Sulfite = (NH4)2SO3
Lithium Nitrite = LiNO2
Potassium Hypochlorite = KClO
Lithium Perbromate = LiBrO4
Magnesium Hypoiodite = Mg(IO)2
Naming Acids
• Two types:
• 1. Binary acids -two
elements
•Forms when certain
gaseous compounds
dissolve in water.
Naming Binary Acids
Hydro + root + ic
HCl = hydro + chlor + ic
hydrochloric acid
HBr = hydrobromic acid
HF = hydrofluoric acid
H2S = hydrosulfuric acid
2. Ternary (Oxoacid )acids- those
with three or more elements.
• Naming ternary:
–use the polyatomic name
but change the:
-ate to ic
-ite to ous
Or …
Anion “-ate” suffix becomes an
“-ic” suffix in the acid.
Anion “ -ite” suffix becomes an
“ -ous” suffix in the acid.
The oxoanion prefixes
“hypo-” and “per-” are retained.
Examples:
HClO3 is named after the
chlorate ion, and thus
becomes chloric acid.
HBrO4 is perbromic acid and
is named after the
perbromate ion.
HFO =
Hypoflourous acid
H2SO4 =
Sulfuric acid
Nitric acid = HNO3
Chlorous acid = HClO2
Determining Names and Formulas
of Anions and Acids
Problem: Name the following anions and give the names and
a) I -
formulas of the acid solutions derived from them:
b) BrO2c) SO3-2 d) NO3e) CN -
a) The anion is Iodide; and the acid is Hydroiodic acid, HI
b) The anion is Bromite; and the acid is Bromous acid, HBrO2
c) The anion is Sulfite; and the acid is Sulfurous acid, H2SO3
d) The anion is Nitrate; and the acid is Nitric acid, HNO3
e) The anion is Cyanide; and the acid is Hydrocyanic acid, HCN
Naming Alkanes
Alkanes are hydrocarbons that are called “saturated”
hydrocarbons, they contain only single bonds, no
multiple bonds!
Alkanes have the general formula:
C n H 2n+2
Alkanes are found in three distinct groups:
a) Straight chain hydrocarbons
b) Branched chain hydrocarbons
c) Cyclic hydrocarbons
Each carbon atom has four bonds to
others atoms !
The names for alkanes all end in –
ane.
The First 10 Straight-Chain Alkanes
Name
Formula
Structural Formulas
Methane
Ethane
CH4
C2H6
Propane
Butane
C3H8
C4H10
CH3-CH2-CH3
Pentane
Hexane
C5H12
C6H14
CH3-(CH2)3-CH3
H
H
C
H
H
H
H
H
C C
H
H
H
CH3-(CH2)2-CH3
CH3-(CH2)4-CH3
Heptane
Octane
C7H16
C8H18
Nonane
Decane
C9H20
C10H22
CH3-(CH2)5-CH3
CH3-(CH2)6-CH3
CH3-(CH2)7-CH3
CH3-(CH2)8-CH3
Branched Chain Hydrocarbons
Cyclic Hydrocarbons
Calculating the Molecular Mass of a Compound
PROBLEM:
Using the data in the periodic table, calculate the molecular (or
formula) mass of:
(a) tetraphosphorous trisulfide
PLAN:
(b) ammonium nitrate
Write the formula and then multiply the number of atoms by the
respective atomic masses. Add the masses for the compound.
SOLUTION:
(a) P4S3
molecular mass = (4 x atomic mass of P) + (3 x atomic mass of S)
= (4 x 30.97 amu) + (3 x 32.07 amu) = 220.09 amu
(b) NH4NO3
molecular mass = (2 x atomic mass of N) + (4 x atomic mass of H) +
(3 x atomic mass of O)
= (2 x 14.01 amu) + (4 x 1.008 amu) + (3 x 16.00 amu)
= 80.05 amu
Naming Compounds from Their Depictions
PROBLEM:
Each circle contains a representation of a binary compound.
Determine its formula, name, and molecular (formula) mass.
PLAN:
Each compound contains only two elements. Find simplest whole number
ratio of one atom to the other to determine formula, name, and mass.
SOLUTION:
(a) There is 1 sodium (brown) for every fluorine (green), so the formula is NaF.
formula mass = (1x atomic mass of Na) +
=
(1x atomic mass of F)
22.99 amu + 19.00 amu = 41.99 amu
(b) There are 3 fluorines (green) for every nitrogen (blue), so the formula is NF3.
molecular mass =
=
(3x atomic mass of F)
+
(1x atomic mass of N)
(3x 19.00 amu) + 14.01 amu = 71.01 amu
The Distinction Between Mixtures and Compounds.
S2-
Fe2+
Physically mixed therefore can be
separated by physical means; in this
case by a magnet.
Allowed to react chemically
therefore cannot be separated by
physical means.
Mixtures
Heterogeneous mixture: has one or more visible
boundaries between the components.
Homogeneous mixture: has no visible boundaries
because the components are mixed as individual
atoms, ions, and molecules.
Solutions: A homogeneous mixture is also called a
solution. Solutions in water are called aqueous
solutions, and are very important in chemistry.
Although we normally think of solutions as liquids,
they can exist in all three physical states.
Solutions
• A solution that mixes completely is
miscible. (alcohol and water, gases)
• Solutions that mix only within certain
limits are expressed as a solubility.
(sugar in water, carbon in iron)
• Substance that is dissolved is the
solute.
• Substance that does the dissolving is
the solvent.
170
Nine Types of Solutions
171
Solute
Gas
Gas
Gas
Liquid
Liquid
Liquid
Solvent
Gas
Liquid
Solid
Gas
Liquid
Solid
Example
Air
Carbonated water
Catalytic converter
Water in air
Alcohol in water
Mercury in gold
Solid
Gas
Smoke
Solid
Solid
Liquid
Solid
Salt in water
Copper in zinc
Tools of the Laboratory
Basic Separation Techniques
Filtration: Separates components of a mixture based upon
differences in particle size. Normally separating a
precipitate from a solution, or particles from an air stream.
Filtration
Crystallization: Separation is
based upon differences in solubility
of components in a mixture.
Crystallization
Distillation: separation is based upon differences in
volatility.
Extraction: Separation is based upon differences in solubility
in different solvents (major material).
Centrifugation
Separates
components of a
mixture based
upon differences
in density.
176
Precipitation
Separates
components of a
mixture based
upon the formation
of a solid.
177
Chromatography: Separation is based upon differences in
solubility in a solvent versus a stationary phase.
Procedure for Column Chromatography
Separation by Gas - Liquid Chromatography
PaperChromatographyofInk.MOV
The End