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Chapter 2
Atoms, Molecules, and Ions
Atoms, Molecules, and Ions
• The vocabulary of chemistry is specialized.
• It includes such terms as atom, ion, molecule, isotope, acid,
base, salt, and saturated hydrocarbon.
• Chemistry also involves symbolic notation.
• Atoms of elements are represented by symbols, such as H,
C, N, O, Na, and Cl. These symbols are the alphabet of
chemistry.
• To represent compounds, symbols are combined into
chemical formulas, such as NaCl, and formulas are the words
of chemistry.
Molecule models
S8
C4
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Molecular compounds
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Some molecules
H2 O2
CH3-CH(OH)-CH3
CH3CH2Cl
P4O10
HCOOH
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Molecule models
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New laws are formulated making use
of other laws.
BRICKS ON THE WALL!!!!
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•
Is the mass of the burned match
pictured in the photograph the
same as, less than, or
more than the mass of the
burning match?
The match loses mass in the
reaction. This kind of loss of mass
to carbon dioxide during burning
greatly confused early chemists
who did not measure the mass of
the gases produced.
The rule of conservation of matter
is credited to Lavoisier.
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Gen. Chem. Chapter2
Antoine
Lavoisier (1743 – 1794)
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Another of the 18th century chemists to emphasize the importance of experiments
was Antoine Lavoisier (1743-1794)
Lavoisier studying human respiration. His wife, Marie Anne Paulze, seated at
the table on the right, records the experiment. Lavoisier was the first great
theoretical chemist. Marie was an accomplished artist and made this drawing
herself.
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2.1 Laws of Chemical Combination
• Data obtained by experiment can often be summarized into
laws. Scientific theories are then formulated to explain these
laws.
Lavoisier: The Law of Conservation of Mass
• Modern chemistry dates from the eighteenth century, when
scientists began to make quantitative observations.
•
Antoine Lavoisier made mass measurements precise to about
0.0001 g and established that total mass does not change
during a chemical reaction.
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• For example, Lavoisier heated the red oxide of mercury,
causing it to decompose (break down) to two new products:
mercury metal and oxygen gas.
• By measuring carefully, he found that the total mass of the
products was exactly the same as the mass of the mercury
oxide he started with.
• Lavoisier summarized his findings through the
law of conservation of mass.
The total mass remains constant during a chemical
reaction.
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Proust: The Law of Definite Proportions
• By the end of the eighteenth century, Lavoisier and other scientists
had succeeded in decomposing many compounds into the elements
that form them.
• One of these scientists, Joseph Proust (1754–1826), did careful
quantitative studies by which he established the
Law of Constant Composition
• (also called the Law of Definite Proportions).
All samples of a compound have the same composition; that is, all
samples have the same proportions, by mass, of the elements
present in the compound.
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Joseph Proust (1754 – 1826)
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•
Proust based this law in part on his studies of a substance that we
now call
•
basic copper carbonate ;
•
all samples of this compound have the same composition:
•
57.48% copper,
•
5.43% carbon,
•
0.91% hydrogen,
•
36.18% oxygen
and
by mass.
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A substance known as basic copper carbonate occurs in nature
as the mineral malachite
forms as a patina on copper
roofs and bronze statues
can be synthesized in
the laboratory
The law of definite proportions : Regardless of its source,
basic copper carbonate Cu2(CO3)(OH)2, always has the same
composition.
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The Mineral Malachite
•Chemistry: Cu2(CO3)(OH)2, Copper Carbonate Hydroxide.
•Class: Carbonate
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•
A compound not only has a constant, or fixed, composition; it
also has fixed properties.
•
For example, under normal atmospheric pressure, pure water
always freezes at 0 °C and boils at 100 °C.
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•
The physical and chemical properties of chemical substances—both
elements and compounds—depend on their composition.
•
A twentieth-century modification of this law, based on the work of
Albert Einstein, allows for the possibility of converting mass to energy
(and vice versa).
•
This need not concern us, however, because in chemical rections the
amounts of mass converted to energy are too small to detect.
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2.2 John Dalton and
the Atomic Theory of Matter
In 1803, John Dalton proposed a theory to explain
the laws of conservation of mass and constant
composition.
As he developed what would become known as his atomic theory,
Dalton found evidence of a scientific law describing the composition of
matter.
In some cases, atoms of the same two elements are able to combine to
form two or more different compounds, and the compositions of the
different compounds are related.
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John Dalton
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•
Let's illustrate Dalton's reasoning with a simple example: two compounds
composed of the elements carbon and oxygen.
•
In carbon dioxide (the familiar gas produced in respiration and in the
burning of fuels) the two elements combine in the ratio of 8.0 g of oxygen to
3.0 g of carbon.
•
In carbon monoxide (the poisonous gas formed when a fuel is burned in
limited air) the elements combine in the ratio of 4.0 g of oxygen to 3.0 g of
carbon.
•
Based on evidence such as this, Dalton formulated
The Law of Multiple Proportions
•
which we can state in the following way:
In two or more compounds of the same two elements, the masses of
one element that combine with a fixed mass of the second element
are in the ratio of small whole numbers.
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Law of Multiple Proportions
Ratio of
oxygentocarbon
in CO2 is
exactly
twice
Gen. Chem. Chapter2
the ratio
in CO
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ANIMATION 1:
MULTIPLE PROPORTIONS
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Law of Multiple Proportions
• Four different oxides of nitrogen can be formed,
by combining 28 g of nitrogen with:
• 16 g oxygen, forming Compound I
• 48 g oxygen, forming Compound II
• 64 g oxygen, forming Compound III
• 80 g oxygen, forming Compound IV
1:3:4:5
What is the ratio
16:48:64:80
expressed as small
whole numbers?
• Compounds I–IV are N2O, N2O3, N2O4, N2O5
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• Dalton's Atomic Theory
• The atomic model Dalton developed to explain the laws of chemical
combination is based on four ideas:
- All matter is composed of extremely small, indivisible particles called
atoms.
- All atoms of a given element are alike in mass and other properties, but
the atoms of one element differ from the atoms of every other element.
- Compounds are formed when atoms of different elements unite in fixed
proportions. That is, one atom of A to one of B in AB, two atoms of A to
one of B in A2B and so on.
- A chemical reaction involves a rearrangement of atoms to produce new
compounds. No atoms are created, destroyed, or broken apart in a
chemical reaction.
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2.3 The Divisible Atom
• However, like most other scientific theories, Dalton's model
eventually had to be modified in light of later discoveries.
• Toward the end of the nineteenth century, certain experiments
began to reveal that atoms are made up of smaller parts.
• Of the dozens of subatomic (smaller than atomic) particles now
known, three are of special importance in the study of
chemistry.
• They are the proton, neutron, and electron.
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Subatomic Particles
•
The mass and charge of subatomic particles are so small that they are
conveniently expressed in relative units.
•
The proton has a relative mass of 1. The proton also carries one fundamental
unit of positive electric charge, denoted 1+ .
•
The neutron, as its name implies, is an electrically neutral particle; it has no
charge. Although its mass is slightly greater than that of the proton, for many
purposes we can consider the neutron also to have a relative mass of 1.
•
The third particle, the electron, has a mass that is 1/1836 of the mass of a
proton. An electron has the same quantity of charge as a proton, but it is a
negative charge, denoted as 1- .
•
Protons, neutrons, and electrons are fundamental particles.
•
This means that all protons are alike, all neutrons are alike, and all electrons
are alike in whatever element they are found.
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Subatomic Particles
• Protons and neutrons are densely packed into
a tiny, positively charged core of the atom known
as the nucleus.
• The extremely lightweight electrons are widely
dispersed around the nucleus.
• An atom as a whole is electrically neutral: It has
no net charge because the negative and positive
charges balance each other.
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Atom like
Football Field
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Nuclear Structure
Atomic Diameter 10-8 cm
Nuclear diameter 10-13 cm
=1Å
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Subatomic Particles
•Although the numbers differ from one element
to another, the atoms of every element have
equal numbers of electrons and protons.
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• Dalton believed that the identity of an element is determined by
the mass of one of its atoms.
• We now know that it is not the mass of an atom but rather the
number of protons in the nucleus that determines the kind of
atom and therefore the identity of an element.
• The atomic number (Z) is the number of protons in the nucleus
of an atom of a given element, and it is this number of protons
that defines the element.
• Every atom having two protons in its nucleus has Z = 2 and is
an atom of helium.
• Every atom with 92 protons has Z = 92 and is an atom of
uranium.
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Isotopes
• Atoms that have the same number of protons but different
numbers of neutrons are called isotopes.
• The atomic number (Z) is the number of protons in the nucleus
of a given atom of a given element.
• The mass number (A) is an integral number that is the sum of
the numbers of protons and neutrons in an atom.
• The number of neutrons = A – Z.
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Isotopes
• For example, there are three isotopes of hydrogen.
• The most abundant isotope, occasionally called protium, has a
single proton and no neutrons in its nucleus.
• About one in every 6700 hydrogen atoms, however, has a
neutron as well as a proton. Because the mass of a neutron is
essentially the same as the mass of a proton, the mass of this
hydrogen isotope, called deuterium, is about twice that of
protium.
• A third, very rare isotope of hydrogen, called tritium, has two
neutrons and one proton in the nucleus. A tritium atom has
about three times the mass of a protium atom.
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Isotopes of Hydrogen
Atomic number (Z) = number of protons
Protium has 1 proton, 0 neutrons :
Z=1
Deuterium has 1 proton, 1 neutron :
Z=1
Tritium has 1 proton, 2 neutrons :
Z=1
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• Some elements have only one naturally occurring isotope; these
include fluorine-19, sodium-23, and phosphorus-31.
• Most elements, however, have two or more isotopic forms.
• Tin has the greatest number of naturally occurring isotopes: 10.
• The naturally occurring isotopes are always found in certain
precise proportions. In natural sources of chlorine, for example,
75.77% of the atoms are chlorine-35 and 24.23% are chlorine-37.
• Chemical symbols for isotopes are commonly written in the form
with A being the mass number and Z the atomic number of the
element E.
A
Z
E
• For the two naturally occurring isotopes of chlorine, we can write
35
Cl
17
and
37
Cl
17
indicating mass numbers of 35 and 37
for chlorine-35 and chlorine-37, respectively.
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Chemical Symbol
To represent a particular atom we use the symbolism:
A= mass number
Z = atomic number
To represent an ion
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Isotopes
Atoms can be represented using the element’s
symbol and the mass number (A) and atomic
number (Z):
A
E
Z
35
Cl
17
37
Cl
17
• How many protons are in chlorine-35?
• How many protons are in chlorine-37?
• How many neutrons are in chlorine-37?
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Other Examples of Isotopes
The number of neutrons = A – Z
Carbon-12
Carbon-13
Carbon-14
Z=6
Z=6
Z=6
so 6 neutrons
so 7 neutrons
so 8 neutrons
Chlorine-35
Chlorine-37
Z = 17
Z = 17
so 18 neutrons
so 20 neutrons
Uranium-234
Uranium-235
Uranium-238
Z = 92
Z = 92
Z = 92
so 142 neutrons
so 143 neutrons
so 146 neutrons
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C, H, O Analyser
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2.4 Atomic Masses
• By international agreement, the current atomic mass standard is
the pure isotope carbon-12, which is assigned a mass of exactly
12 atomic mass units (12 u).
•
Based on this standard, we can define an atomic mass unit,
(abbreviated amu and having the unit u)
as exactly one-twelfth the mass of a carbon-12 atom.
•
In more familiar units of mass, 1 u = 1.66054 x 10-24 g.
•
Because we are interested in the masses of atoms and not their weights,
we will now shift to the term atomic mass in place of the older atomic
weight, except in a few cases of historical interest.
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•
The atomic mass of an element is defined as the weighted average of the
masses of the naturally occurring isotopes of that element.
•
Consider carbon, which consists of a mixture of two naturally occurring
isotopes.
•
The much more abundant isotope is carbon-12. The other is carbon-13,
with a mass of 13.00335 u. Both isotopes are present in substances
containing carbon atoms, and in proportions that generally do not vary
from one carbon-containing substance to another.
•
To describe the atomic mass of carbon, then, we need to use an average
value, but not the simple average (12 + 13) / 2 = 12.5. Because carbon-12 is
much more abundant than carbon-13, the average we seek lies much
closer to the mass of carbon-12 than to that of carbon-13. We say that it is
"weighted" toward the mass of carbon-12.
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•
To calculate the atomic mass of an element, we need two quantities—
•
(1) the atomic masses of the isotopes of the element and
•
(2) the naturally occurring fractional abundances of the isotopes.
•
Let us explain how these quantities are obtained experimentally
•
To illustrate, let's return to the atomic mass of carbon.
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•
The percentage abundances and fractional abundances of the
carbon isotopes are as follows:
To obtain a weighted average atomic mass, we calculate the contribution
of each isotope to the weighted average from the relationship
The term atomic weight is still widely used, however, as by the Commission
on Atomic Weights of the International Union of Pure and Applied Chemistry
(IUPAC).
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Use the data cited below to determine the weighted average
atomic mass of carbon.
The contribution of each isotope to the weighted average atomic mass is
given by Equation.
Solution
The contributions are
Contribution of carbon-12 = 0.98892 x 12.00000 u = 11.867 u
Contribution of carbon-13 = 0.01108 x 13.00335 u =
0.1441 u
The weighted average atomic mass is the sum of the two contributions.
Atomic mass of carbon = 11.867 u + 0.1441 u = 12.011 u
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Carbon – 14 ( C-14 Dating )
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2.5 The Periodic Table: Elements Organized
• In the nineteenth century, chemists discovered dozens of new
elements.
• By 1830, 55 elements were recognized, but there was no
apparent pattern in their properties. Chemists badly needed a
way to organize the growing collection of chemical data.
• One way to do this was to arrange the elements in a manner
that would establish categories of elements having similar
physical and chemical properties.
• Dimitri Mendeleev published the first successful arrangement,
called a periodic table, in 1869.
• In its modern form, the periodic table organizes a vast array of
chemical knowledge.
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Dimitri Mendeleev
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Mendeleev's Periodic Table
Mendeleev arranged the elements in order of
increasing atomic weight, from left to right in rows
and from top to bottom in columns (or groups).
In this arrangement, elements that most closely
resemble one another in physical and chemical
properties tend to fall in the same vertical group.
This group similarity repeats periodically, hence the name periodic table.
•
There would be no exceptions to the principle that all the elements in a
group display similar properties, therefore Mendeleev placed some
elements out of order, that is, not in the strict order of increasing atomic
weight.
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For example, he correctly
placed tellurium (atomic weight
127.6) before iodine (atomic
weight 126.9) so that tellurium
would be in the same column
as the similar elements sulfur
and selenium.
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•
When Mendeleev placed elements with similar properties in
the same vertical group, a few gaps were created in his table.
Instead of seeing these gaps as defects, he boldly predicted
the existence of undiscovered elements to fill the gaps.
•
Furthermore, because the table was based on patterns of
properties, he was able to predict some properties of the
missing elements.
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Mendeleev’s Table
Mendeleev’s Table
Dimitri Mendeleev created this, the original periodic table.
Stowe's table
Spiral form
Triangular form
Mendeleevs’s early table was published in 1872.
• Using his table he was able to correct several values of atomic
masses.
Folded Table
• Because of its obvious usefullness his periodic table was almost
universally adopted, and it remains one of the most valuable tools at the
chemist’s use.
• The only fundemantal difference between todays table and that of his
Periodic Table
is that in the current table the elements are ordered by atomic number
rather than by atomic mass.
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•
For example, he left a blank
space for an undiscovered
element that he called "ekasilicon" and used its location
between silicon and tin to
predict an atomic weight of 72
and other properties.
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•
Table 2.2 shows just how accurate his predictions were, when compared
to the properties of the actual element germanium discovered 15 years
later.
•
The predictive nature of Mendeleev's periodic table led to its wide
acceptance as a tremendous scientific accomplishment.
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Alkali Metals
The Periodic table
Alkaline Earths
Halogens
Noble Gases
Main Group
Transition Metals
Main Group
Lanthanides and Actinides
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(2) PERIODIC
PROPERTIES
ANIMATION
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Empirical and Molecular
Formulas
Empirical formula: the simplest whole number
ratio of elements in a compound
Example:
Molecular formula of glucose – C6H12O6
The elemental ratio C:H:O is 1:2:1, so the
empirical formula is CH2O
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Molecular compounds
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Introduction to Compounds
A molecule is a group of two or more atoms held
together by covalent bonds.
A chemical formula
is a symbolic
representation of the
composition of a
compound in terms
of its constituent
elements.
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Structural Formulas
Shows how atoms are attached to one another.
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Molecular Compounds
Ball-and-stick model vs. Spacefilling model
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Binary Molecular Compounds
Molecular formulas are usually written with the more
“metallic” first – “metallic” means farther left in the
period and lower in the group
e.g., NaCl, KBr
Compounds that are typically comprised of two
nonmetallic elements:
e.g., CO, NO, HF
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Formulas and Subscripts
Subscripts are used when a given atom is
used more than once
e.g., H2O, CO2, N2O, HF, B2O3
The presence of
subscripts is
reflected in the
names of
compounds
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Names of Binary Compounds
Consider the compounds CO and CO2
The compound name consists of two words,
one for each element in the compound
Name the element that appears first in the formula:
CARBON
The second element has an altered name: retain the stem
of the element name and replace the ending by -ide
OXYGEN  OXIDE
However, both compounds cannot be carbon oxide
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Names of Binary Compounds
The names are further modified by adding prefixes to
denote the numbers of atoms:
CO (Carbon Mon-oxide), CO2 (Carbon Di-oxide)
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2.7 Ions and Ionic Compounds
In an isolated atom, the number of protons equals the number of
electrons, and the atom is therefore electrically neutral.
In some chemical reactions, however, an individual atom or a
group of bonded atoms may lose or gain one or more electrons,
thereby acquiring a net electric charge and becoming an ion.
Ions are formed only through the loss or gain of electrons; there is
no change in the number of protons in the nucleus of the atom(s).
If electrons are lost, there are more protons than electrons in the
resulting ion and so it has a positive charge.
If electrons are gained, there are more electrons than protons in
the resulting ion and so it has a negative charge.
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Ions and Ionic Compounds
Atoms that gain or lose electrons are called ions
Positive ions: CATIONS
Atoms that lose
electrons form cations
Na  Na+ + e–
Negative ions: ANIONS
Atoms that gain
electrons form anions
Cl + e–  Cl–
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Monatomic Ions
Group A metals usually lose the number of electrons equal to their
Group number.
Nonmetal atoms usually gain electrons and have a charge equal to
their Group number minus eight.
The periodic table cannot be used to determine the charge on
Group B metals.
For naming, Group B metals capable of multiple charges have
the corresponding Roman numeral in parentheses added after
the element name.
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Common Monatomic Ions
FIGURE 2.10 Symbols and periodic table locations of some monatomic ions
Three general observations can be made of these data:
(1) Aluminum and the metals of groups 1A and 2A form just one cation, which carries a
positive charge equal in magnitude to the A-group number.
(2) Most of the metals of the B groups form two or more cations of different charges,
though in some cases only one of these cations is commonly encountered.
(3) The nonmetals of groups 7A and 6A, along with nitrogen and phosphorus of group
5A, form anions that have a charge equal to the group number minus 8.
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Formulas and Names for
Ionic Compounds
Ionic compounds form when oppositely charged
ions are attracted to each other
NaCl
Resulting compound is
electrically neutral
Na+
Cl–
(+1) +
(–1) = 0
Ionic compound names use the cation name
followed by the anion name
Sodium chloride
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Now consider the formula for aluminum oxide.
We cannot simply combine one Al3+ and one O2- , because
this would produce a formula unit with a net charge of 1+.
The combination of two Al3+ ions and three O2- ions, though,
is an electrically neutral formula unit:
2 (3+) + 3 (2-) = + 6 - 6 = 0
Therefore, the formula for aluminum oxide is Al2O3 .
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Polyatomic Ions
Polyatomic
ions are
charged
groups of
covalently
bonded
atoms
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Hydrates
A hydrate is an ionic compound in which the formula unit
includes a fixed number of water molecules associated with
cations and anions
Examples:
BaCl2 . 2 H2O
CuSO4 . 5 H2O
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Acids
• Taste sour
• Turn blue litmus paper red
• React with metals to form
hydrogen gas
• Neutralize a base
IntroToAcids
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Bases
• Taste bitter
• Turn red litmus paper blue
• Feel slippery on skin
• Neutralize an acid
IntroToBases
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Arrhenius Concepts
Acids are compounds that ionize in water
to form a solution of H+ ions and anions
Bases are compounds that ionize in water to
–
form solutions of OH and cations
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Arrhenius Concepts
Acids and bases react to form a salt and water
= neutralization
HCl + NaOH  “Salt” + Water
Acid
Base
NaCl
HOH
cation/anion
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Animation (3) & (4) :
Acids, Bases
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Formulas and Names for
Acids
Binary acids start with hydro
and end with “ic” plus the
word acid
Ternary acids simply use the
polyatomic anion name with
“ate” changing to “ic” plus
the word acid
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Formulas and Names for Bases
Arrhenius bases always have hydroxide ions
The name follows ionic compound convention
e.g., NaOH – sodium hydroxide
Molecular bases form OH– after reacting with
water
NH3 + HOH  NH4OH
Ammonia  ammonium hydroxide
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Formulas and Names for Salts
Binary salts use the “ide”
ending on the anion name
e.g., sodium chloride
Polyatomic salts use “ate”
ending on the anion name
e.g., sodium sulfate
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Organic Compounds
• Organic Chemistry is the study of carbon and its
compounds
• Carbon compounds containing one or more of the
elements H, O, N, or S are especially common
• Most organic compounds are molecular compounds
• Can exist as acids, bases, and salts
• Compounds have systematic names AND common
names
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Representations of Molecules
Condensed Structural Formula
CH3CH2CH3
Structural Formula
Ball and Stick
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Saturated Hydrocarbons
Hydrocarbons have only hydrogen and carbon
atoms
Saturated hydrocarbon: has
the maximum number of hydrogen
atoms possible for each carbon atom
Alkanes are saturated hydrocarbons
Methane (CH4) is the first molecule in the alkane
series
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Prefixes for Number of Carbon
Used for simple
organic molecules
Combined with
alkane ending
“ane”
e.g., propane is a
3-carbon alkane
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Isomers
Compounds with the same molecular formula but
different structural formulas
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FIGURE 2.14 Butane and isobutane
Ball-and-stick and structural models illustrate the
difference between
the two
isomers of butane.
Gen. Chem.
Chapter2
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Cyclic Alkanes
Alkane compounds that have carbons arranged in a
ring structure are called cycloalkanes.
use the prefix cyclo-
methylcyclopropane
cyclohexane
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FIGURE 2.15 Cyclohexane
The molecular and line-angle formulas for cyclohexane indicate the bonding of atoms
within the molecule. The ball-and-stick model further indicates that all of the carbon atoms
of the cyclohexane molecule do not lie in the same plane. Rather, it can assume several
different arrangements or conformations. The most stable arrangement, shown here, is
called the chair conformation because the six carbon atoms outline a structure that
somewhat resembles a reclining chair.
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Functional Groups
Specific groupings of atoms attached to a carbon
chain that give the compound unique properties
Most-common functional groups include:
•
•
•
•
•
Alcohols
Ethers
Carboxylic Acids
Esters
Amines
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Alcohols
Alcohols are molecules that contain a hydroxyl
group (OH)
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Alcohols
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Ethers
Ethers are molecules in which two alkane groups
(R-) are attached to a central oxygen atom
The general formula is R-O-R´
R and R´ may be the same
or different groups
CH3CH2OCH2CH3
CH3CH2OCH2CH2CH3
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Carboxylic Acids
Carboxylic acids are alkanes that also contain
a carboxyl group and are weak acids
Acts like an Arrhenius
acid, loses a hydrogen ion
HCOO– + H+
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Esters
Esters are molecules in which two alkanes are
attached to each side of a carboxyl group
(R’-COO-R)
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Amines
Amines are molecules in which alkanes and
hydrogen(s) are attached to a central nitrogen
Amines are weak bases
NH2(CH2)4NH2
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Summary of Concepts
• The basic laws of chemical combination are the laws of
conservation of mass, constant composition, and multiple
proportions.
• The three main subatomic particles are the protons, neutrons,
and electrons.
• Atoms with the same number of protons but different numbers
of neutrons are called isotopes.
• A chemical formula indicates the relative numbers of atoms of
each type in a compound.
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Summary (cont.)
• The periodic table is an arrangement of the elements by
atomic number that places elements with similar properties
into the same vertical group.
• Ions are formed by the gain or loss of electrons. Positive ions
are cations and negative ions are anions.
• Many compounds are classified as either acids (H+), bases
(OH–), or salts (neutralization of acid and base).
• Organic compounds are based on the element carbon.
• Functional groups confer distinctive properties on an organic
molecule.
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Arrhenius Concepts
Acids and bases react to form a salt and water
= neutralization
HCl + NaOH  “Salt” + Water
Acid
Base
NaCl
HOH
cation/anion
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