Homework Problems

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Homework Problems Chapter 2 Homework Problems
1, 4, 19, 24, 28 (give the number of protons,
neutrons, and electrons), 30, 38, 41, 44, 46, 54,
64, 66, 67, 72, 76, 82, 84, 96, 101, 104, 114,
116, 117, 124
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CHAPTER 2 Atoms, Molecules, and Ions
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Early Theories of Matter The ancient Greeks
discussed two possibilities for the essential
property of matter. Matter is continuous (no
particles of matter) - Plato, Aristotle, and a
majority of Greek philosophers. Matter is
discrete (composed of particles) - Democritus,
Leucippus, and a small number of Greek
philosophers.
Does this process have an end? yes particles
of matter exist no matter is continuous
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General Properties of Chemical Systems As
scientists began studying chemical systems they
discovered several general properties of such
systems. 1) Conservation of Mass. The total
mass in a closed system remains constant, even if
chemical reactions occur. 2) Law of Definite
Proportions. All samples of a particular pure
chemical substance contain the same relative
amounts of each element making up the
substance. Examples methane 74.9 C, 25.1
H water 88.8 O, 11.2 H copper (II)
sulfate 39.8 Cu, 20.1 S, 40.1 O
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3) Law of Multiple Proportions. When two
elements can combine to form several different
chemical compounds, the ratio of the amount of
the second element combining with a fixed amount
of the first element will be the ratio of small
whole numbers. Example There are two common
compounds of carbon and oxygen carbon
monoxide 1.000 g of C reacts with 1.332 g of O
carbon dioxide 1.000 g of C reacts with
2.664 g of O g O in carbon dioxide
2.664 g 2.000 ? 2 g O in carbon monoxide
1.332 g 1 It is easier
to calculate the ratios with the larger number on
top, but that is not required.
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Example The following pure chemical substances
can be formed out of the elements nitrogen and
oxygen nitrogen monoxide 1.000 g of N reacts
with 1.142 g of O nitrogen dioxide 1.000 g of
N reacts with 2.285 g of O nitrous
oxide 1.000g of N reacts with 0.5711 g of O Do
these substances demonstrate the law of multiple
proportions?
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nitrogen monoxide 1.000 g of N reacts with
1.142 g of O nitrogen dioxide 1.000 g of N
reacts with 2.285 g of O nitrous oxide 1.000g
of N reacts with 0.5711 g of O g O in nitrogen
monoxide 1.142 g 2.000 ? 2 g O in
nitrous oxide 0.5711g
1 g O in nitrogen dioxide 2.285 g
4.001 ? 4 g O in nitrous oxide
0.5711 g 1 g O in nitrogen
dioxide 2.285 g 2.001 ? 2 g O in
nitrogen monoxide 1.142 g
1 So yes, these data are consistent with the
law of multiple proportions.
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Daltons Atomic Theory A comprehensive theory
that accounted for the above obser-vations was
proposed by John Dalton, an English chemist, in
1808. There were three parts to the theory. 1)
Elements are composed of particles, called
atoms. a) All atoms of the same element are
identical in size, mass, and chemical
properties. b) Atoms of different elements
differ in their size, mass, and chemical
properties.
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Daltons Atomic Theory (continued) 2) Chemical
compounds are composed of atoms of more than one
element. a) In any particular pure chemical
compound the same kinds of atoms are present in
the same relative numbers. 3) Chemical
reactions can rearrange atoms, but atoms cannot
be created, destroyed, or converted from atoms of
one element to atoms of a different element. We
now know that some of the hypotheses in Daltons
atomic theory are not completely correct
however, the theory represents a good starting
point in understanding the composition of matter.

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Consequences of Daltons Atomic Theory Daltons
theory can be used to explain the observations
cited above. 1) Conservation of mass. Explained
by (1) and (3). 2) Law of definite proportion.
Explained by (1) and (2). Example methane
CH4 Chemical formula - A list of the
elements making up a compound, giving the number
of atoms of each element per molecule or per
formula unit of the compound.
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3) Law of multiple proportions. Explained by
(1) and (2) .
carbon monoxide (CO)
carbon dioxide (CO2)
1 C atom 1 O atom
1 C atom 2 O atom So for a given amount of
carbon, carbon dioxide will have twice as many
oxygen atoms (and therefore twice the mass of
oxygen) as carbon monoxide.
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Atomic Structure In Daltons atomic theory the
smallest particles (atoms) could not be further
broken down. However, a series of experiments,
beginning in the mid-19th century, demonstrated
that atoms themselves were composed of smaller
particles.
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Radioactivity In 1895, Antoine Becquerel
discovered that some substances (such as radium
and uranium) spontaneously emit radiation, a
process called radioactivity. Three types of
radioactivity were found alpha (?) radiation
positively charged particles, now known to be
He2 nuclei (2 protons 2 neutrons) beta (?)
radiation negatively charged particles, now
known to be electrons
gamma (?) radiation uncharged, now known to be
high energy photons (particles) of light
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Electrons
J. J. Thompson (1897) found that when a high
voltage was applied across two electrodes at low
pressure a beam of particles moved from the
negative to the positive electrode. The
particles, named electrons, were negatively
charged and the same regardless of the gas
between the electrodes or the metal used in the
electrodes. The charge and mass of an electron
were determined experimentally by Millikan
(1909).

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The Plum Pudding Model Since atoms are
electrically neutral, the negative charge of the
electrons in an atom had to be balanced by a
positive charge. Thompson suggested that most of
the space within an atom consisted of a
positively charged substance, with electrons
embedded within, the plum pudding model.
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Rutherford and the Nuclear Atom To test
Thompsons plum pudding model, Ernest Rutherford
(1909) carried out an experiment where a beam of
positively charged particles (alpha particles)
were directed at a thin sheet of gold metal.
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The results of this experiment were inconsistent
with the plum pudding model. Rutherford proposed
a new model, called the nuclear model of the
atom, that did account for the experimental
results.
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Subatomic Particles particle charge
mass Coulombs elementary kg amu proton
(p) 1.60 x 10-19 1 1.673 x
10-27 ? 1 neutron (n) 0
0 1.675 x 10-27 ? 1 electron (e-) -
1.60 x 10-19 - 1 9.11 x 10-31 ?
0 ________ 1 amu 1.6605 x 10-27 kg
mp/me 1836.
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Atomic Structure
nucleus
electron charge cloud
1) The protons and neutrons of the atom are found
in a small region in the center of the atom,
called the nucleus. This region contains most of
the mass of the atom, and all of the positive
charge. 2) Electrons in the atom form a diffuse
cloud of negative charge centered on the nucleus
and occupying most of the volume of the atom. 3)
The size of the charge for the proton and
electron is the same. The charge for the proton
is positive, and the charge for the electron is
negative. Neutrons have no charge.
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4) The type of element for an atom is determined
by the number of protons in the atomic
nucleus. Element (new definition) - An element
is a pure chemical substance composed of atoms,
each of which has the same number of protons in
the nucleus. Hydrogen - one proton per
atom Helium - two protons per atom Lithium -
three protons per atom . .
. . Similarly, we can now define a
compound (new definition) as a pure chemical
substance composed of two or more different kinds
of atoms.
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The Periodic Table
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Atomic Number and Mass Number 1) The atomic
number (Z) is equal to the number of protons in
the atom. 2) Since atoms are electrically
neutral, the number of electrons in an atom is
also equal to Z, the atomic number. 3) The mass
number (A) is equal to the number of protons
neutrons in the atom. a) Because protons and
neutrons have a mass of approximately 1 (in amu)
and electrons have a mass of approximately 0 (in
amu) the mass number is equal to the approximate
mass of the atom in amu. b) Based on the above,
the number of neutrons in an atom is equal to A -
Z. So for an atom protons Z
electrons Z neutrons A - Z
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Notation For Atoms We use the following general
notation to represent isotopes of atoms.
mass number atomic number
symbol for element
Since we can use the symbol for the element and
the periodic table to determine Z, the atomic
number, we often omit Z in giving the symbol for
the atom. Example 3416S 34S We can omit the
subscript because all sulfur atoms contain 16
protons.
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Isotopes The atomic number determines the number
of protons and electrons in an atom. This does
not place any restrictions on the number of
neutrons in the atom. It is possible for atoms
of the same element to have different numbers of
neutrons. These different types of atoms are
called isotopes. Isotopes of Hydrogen
normal hydrogen deuterium
tritium 1H
2H 3H Note
that to a very good approximation isotopes of a
particular element are chemically identical to
one another.
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Example of Notation for Isotopes As an example
of using the above notation, consider the
follow-ing naturally occurring isotopes of
oxygen (Z 8). protons neutrons electrons mass
number symbol 8 8 8 16
16O 8 9 8 17 17O
8 10 8 18 18O Example
How many protons, neutrons, and electrons are
there in one atom of 56Fe? What is the
approximate mass of one atom of 56Fe in amu and
in kg?
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Example How many protons, neutrons, and
electrons are there in one atom of 56Fe? What is
the approximate mass of one atom of 56Fe in amu
and in kg? protons Z 26 neutrons A
Z 56 26 30 electrons Z
26 approximate mass (amu) A 56 approximate
mass (kg) 56 amu 1.6605 x 10-27 kg
1 amu 9.30 x 10-26
kg Note that the actual mass of one atom of
56Fe is 55.934939 amu.
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Atomic Mass Units (amu) The mass of a single
atom of an element, expressed in SI units, is an
extremely small number. For example, the mass of
a single atom of 16O is 2.6560 x 10-26 kg. For
convenience, we often express values for atomic
mass in terms of atomic mass units (amu). Atomic
mass units are defined as follows 12.00 amu
mass of one atom of 12C (exact) From this we
get 1. amu 1.6605 x 10-27 kg (approximate) The
mass of any other atom (or particle) is found
relative to the ratio of its mass to the mass of
a 12C atom, which can be measured
experimentally. Mass of particle (amu) mass
particle (12.00 amu) mass
12C atom
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Example A mass spectrometer is a device for
determining values for mass for atoms and
molecules. In a particular experiment, the ratio
(mass M/mass 12C) is measured and found to be
equal to 7.337. What is the mass of the molecule
M (in amu)?
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Example A mass spectrometer is a device for
determining values for mass for atoms and
molecules. In a particular experiment, the ratio
(mass M/mass 12C) is measured and found to be
equal to 7.337. What is the mass of the molecule
M (in amu)? mass M 7.337
mass 12C atom
Mass M 7.337 (mass 12C atom) 7.337
(12.00 amu) 88.04 amu Note that because of
the way we define atomic mass units, the only
isotope whose mass is exactly equal to its mass
number is 12C. isotope mass (amu) 1H
1.007825 12C 12.000000...
(exact) 238U 238.0508
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Atomic Mass in the Periodic Table Because
different isotopes of an element have different
masses, the question arises as to which mass
should be given in the periodic table. For
short lived radioactive elements the mass number
of the most stable isotope of the element is
listed. Element Z A technetium
(Tc) 43 98 radon (Rn) 86 222
plutonium (Pu) 94 244
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Average Atomic Mass For naturally occurring
elements, the value for mass given in the
periodic table is the average atomic mass, based
on the natural abundance of the isotopes that is
observed. In general, we find the average atomic
mass as follows Mave f1 M1 f2 M2 f3 M3
?i1n fi Mi where f1, f2,...are the
fractions of each isotope observed in nature M1,
M2,are the corresponding masses for each isotope
(in amu) Note the following f1 f2 f3
1 fx X 100
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Non-chemical Example A person has a box of
sandwiches. Half of the sandwiches are 6.0
ounces, and half of the sandwiches are 10.0
ounces. What is the average weight of a
sandwich? Average weight (0.50)(6.0 oz)
(0.50)(10.0 oz) 8.0 ounces We use the same
procedure in finding the average mass of an atom.
We multiply the fraction of each isotope by the
mass of that isotope, and then add the results to
find the average mass.
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Chemical Example There are three naturally
occurring isotopes of the element magnesium.
Based on the information below, find the average
atomic mass of a magnesium atom. Isotope percent
f M(amu) 24Mg 78.70 23.98504 25Mg 1
0.03 24.98584 26Mg 11.17 25.98259
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Chemical Example There are three naturally
occurring isotopes of the element magnesium.
Based on the information below, find the atomic
mass of a magnesium atom. Isotope percent
f M(amu) 24Mg 78.70 0.7870 23.98504 25Mg 1
0.03 0.1003 24.98584 26Mg 11.17 0.1117 25.98
259 So Mave (0.7870)(23.98504 amu)
(0.1003)(24.98584 amu) (0.1117)(25.98259
amu) 24.30 amu, the value given in the
periodic table.
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Periodic Table The periodic table is an
arrangement of the chemical elements based on
similarities in their physical and chemical
properties The periodic table contains a large
amount of useful information about the chemical
elements. Organization There are several ways
in which the elements in the periodic table may
be classified. Rows Periods Columns Groups
This is the more important classification.
Elements in the same group usually have similar
physical and chemical properties.
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Simplified Periodic Table
1A 2A

3A 4A 5A 6A 7A 8A
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1A 2A

3A 4A 5A 6A 7A 8A
You are responsible for knowing the names/symbols
for elements 1-57, 72-86, and 92.
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Major Groups in the Periodic Table
1A 2A

3A 4A 5A 6A 7A 8A
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Metals, Nonmetals, and Metalloids Metals Usually
solid at room temperature (exceptions Cs, Fr,
Hg) Shiny metallic luster Good
conductors of electricity and heat
Malleable (can be hammered into thin sheets)
Ductile (can be drawn into thin
wires) Nonmetals Can be solid, liquid, or gas at
room temperature Dull colored (as solids)
Poor conductors of electricity and heat
Not malleable, not ductile Metalloids
(semimetals) Intermediate between metals and
nonmetals
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Metals, Nonmetals, Metalloids in the Periodic
Table
1A 2A

3A 4A 5A 6A 7A 8A
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Examples of Elements (as found in nature)
nickel
germanium sulfur
(metal) (metalloid)
(nonmetal)
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Formation of Ions Ions are charged particles.
Ions can be formed from an atom by either adding
electrons (to form an anion) or removing
electrons (to form a cation). Ions cannot be
formed by changing the number of protons in the
atom. cations anions particle Z
electrons particle Z electrons Na 11
11 Cl 17 17 Na 11
10 Cl- 17 18 Ca 20 20
S 16 16 Ca2 20 18
S2- 16 18 Note that the charge of an
ion is indicated by a superscript to the right of
the symbol for the ion. Metals usually form
cations, while nonmetals usually form anions.
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Just as we can predict the number of protons,
neutrons, and electrons from the symbol for an
atom, we can do the same thing for cations and
anions formed from atoms. We do this using the
atomic number (Z) the mass number (A) and the
charge of the ion. Example How many protons,
neutrons and electrons are there for a 31P3- ion
and a 25Mg2 ion?
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Example How many protons, neutrons and
electrons are there for a 31P3- ion and a 25Mg2
ion? 31P3- protons Z 15 neutrons A -
Z 31 - 15 16 Charge is 3-, so there are 3
more electrons than protons, and so the number of
electrons 18. 25Mg2 protons Z 12
neutrons A - Z 25 - 12 13 Charge is 2, so
there are two more protons than electrons, and so
the number of electrons is 10.
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Ion Charges For Main Group Elements Main group
elements tend to form ions by adding or removing
electrons so that the number of electrons
remaining in the ion is equal to the number of
electrons in one atom of the nearest noble
gas. cations (metals) group 1A (Li, Na, K,
Rb, Cs) form 1 ions group 2A (Mg, Ca, Sr, Ba)
form 2 ions group 3A (Al) form 3 ions anions
(nonmetals) group 5A (N, P) form 3-
ions group 6A (O, S, Se, Te) form 2-
ions group 7A (F, Cl, Br, I) form 1- ions
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Transition Metal Ions Transition metals form
cations. Most transition metals, as well as a
few main group metals like tin (Sn) and lead
(Pb), can form ions with several different
charges, but a few transition metals, such as
silver, usually form only one type of cation (Ag
for silver, Cd2 for cadmium, Zn2 for zinc).
Example Iron (Fe) Fe2, Fe3
Copper (Cu) Cu, Cu2 Chromium (Cr)
Cr3, Cr6 It is generally not easy to
predict which cations a transition metal will
form.
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Chemical Formula The chemical formula for a
substance provides information concerning the
composition of the substance. We can divide
substances into two general types. 1) Substances
that exist as collections of molecules. In this
case the chemical formula indicates the number of
atoms of each elements present per molecule.
water phosphorus
pentachloride nitrous acid
(H2O) (PCl5)
(HNO2)
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For organic molecules, the chemical formula is
often given in a way that indicates how the
molecule is put together.
ethyl alcohol
dimethyl ether acetone
CH3CH2OH C2H6O CH3OCH3 C2H6O
CH3COCH3 C3H6O Notice that this longer
notation makes it possible to distinguish among
different forms (isomers) of organic molecules.
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2) Substances that exist as collections of atoms
or ions in the form of a crystal structure, a
regular arrangement of the particles making up
the substance. For these substances, the formula
that is given is usually the empirical formula.
An empirical formula gives the relative number of
atoms of each element making up the compound,
reduced to the smallest set of whole number
coefficients.
Ca2
F-
sodium chloride polonium
calcium fluoride
NaCl Po
CaF2
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Empirical Formula for Molecular Compounds The
empirical formula for a substance gives the
relative number of atoms of each element making
up a pure chemical substance, reduced to the
smallest set of integer values. For substances
that exist as molecules, the molecular formula
must be an integer multiple of the empirical
formula. Substance Chemical formula Empirical
formula water H2O H2O hydrogen
peroxide H2O2 HO benzene C6H6 CH dichloroet
hane C2H4Cl2 CH2Cl acetic acid CH3COOH CH2O
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Molecular Compound A molecular compound is a
compound composed of individual particles called
molecules. The bonding between atoms is in such
compounds is due to the sharing of one or more
pairs of electrons. Molecular compounds are
usually made up of one or more nonmetallic
elements. Note that these compounds are
sometimes called covalent compounds since the
molecules are held together by covalent bonding.
Examples HBr hydrogen bromide SF6 sulfur
hexafluoride CS2 carbon disulfide N2O4 dinitrogen
tetroxide CH3Cl chloromethane SO2 sulfur
dioxide Some elemental substances, such as O2
and N2, exist as indi-vidual molecules, but are
not compounds, since they are composed of atoms
of a single element.
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Ionic Compound An ionic compound is a compound
formed from positive ions (cations) and negative
ions (anions) held together by electrostatic
attraction. For these compounds we usually give
the empirical formula, (sometimes called the
formula unit) or smallest electrically neutral
collection of ions making up the compound.
Binary ionic compounds (compounds formed from
ions of two different elements) are usually a
combination of a metal cation and a nonmetal
anion. Examples NaCl sodium chloride (Na
and Cl-) Fe2O3 iron (III) oxide (Fe3 and
O2-) Na2O sodium oxide (Na and
O2-) CuS copper (II) sulfide (Cu2 and S2-)
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Chemical Formulas For Main Group Ionic
Compounds Because main group elements form ions
with a particular charge (depending on which
group the element is from) we can predict the
chemical formula for a main group ionic compound.
We do this by first finding the charges of the
ions formed, and then combining them to get an
overall neutral compound using the smallest set
of whole number coefficients. Examples What is
the chemical formula for the ionic compound
formed from magnesium and chlorine, from sodium
and sulfur, and from calcium and oxygen?
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Examples What is the chemical formula for the
ionic compound formed from magnesium and
chlorine, from sodium and sulfur, and from
calcium and oxygen? Mg and Cl Ions are
Mg2 and Cl-, so formula is MgCl2, magnesium
chloride. Na and S Ions are Na and S2-, so
formula is Na2S, aluminum sulfide. Ca and O
Ions are Ca2 and O2-, so formula is CaO,
calcium oxide.
magnesium chloride sodium sulfide
calcium oxide
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Transition Metal Cations Transition metals can
usually form ions with several different charges
(this is also true for a few main group metallic
elements like tin (Sn) and lead (Pb)). While we
cannot easily predict which compounds will form
between a transition metal and a main group
nonmetal, we can usually figure out the charge of
the transition metal cation if we know the
chemical formula for the compound. This is
indicated in the name of the compound. Examples
CuCl CuCl2 TiO2 MnS2
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Examples CuCl Cl- ion, so Cu ion copper (I)
chloride CuCl2 Cl- ion, so Cu2 ion copper
(II) chloride TiO2 O2- ion, so Ti4
ion titanium (IV) oxide MnS2 S2- ion, so Mn4
ion manganese (IV) sulfide
CuCl
CuCl2 TiO2
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Polyatomic Ions A polyatomic ion is a group of
atoms which collectively has a charge and acts as
an ion in an ionic compound. ion name example
of ionic compound NO3- nitrate ion NaNO3,
Ca(NO3)2, Ni(NO3)2 SO42- sulfate ion CuSO4,
K2SO4, Al2(SO4)3 Note that when more than one
polyatomic ion is present in a chemical compound
the ion is placed in parentheses and the number
of ions per formula unit of compound is given as
a subscript outside the parentheses.
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Hydrates Some ionic compounds can exist in forms
where there is a specific number of water
molecules associated with every formula unit of
the ionic compound. Such substances are called
hydrates.
Cobalt (II) chloride hexahydrate
Cobalt (II) chloride
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Naming Rules For Simple Compounds Ionic
compounds. a) Main group metal (1A, 2A, metals
and aluminum (3A), Ag, Cd2, Zn2 main group
nonmetal name of metal name of nonmetal
ide Examples K2S potassium sulfide NaCl sodium
chloride CaI2 calcium iodide ZnF2 zinc
fluoride Al2O3 aluminum oxide b) Transition
group metal (or group 4A metal) main group
nonmetal name of metal (charge of metal)
name of nonmetal ide Examples FeCl2 iron (II)
chloride NiO nickel (II) oxide FeCl3 iron (III)
chloride PbS2 lead (IV) sulfide Cu2S copper (I)
sulfide
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c) Cation group NH4 ammonium ion Hg22
mercury (I) ion Use name of cation group
name of nommetal ide. Example
NH4Br ammonium bromide Hg2Cl2 mercury (I)
chloride (NH4)2S ammonium sulfide Hg2O mer
cury (I) oxide HgO mercury (II) oxide
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d) Anion group C2H3O2- is acetate ion CO32-
is carbonate ion CN- is cyanide ion SCN- is
thiocyanate ion OH- is hydroxide ion C2O42-
is oxalate ion N3- is azide ion O22- is
peroxide ion CrO42- is chromate ion Cr2O72- is
dichromate ion MnO4- is permanganate
ion ClO3- is chlorate ion NO3- is nitrate
ion BrO3- is bromate ion SO42- is sulfate
ion IO3- is iodate ion PO43- is phosphate
ion 1 oxygen changes the name to per ________
ate -1 oxygen changes the name to
________ite -2 oxygen changes the name to hypo
________ ite
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Example ClO3- is chlorate ion, so 1 O ClO4-
is perchlorate ion -1 O ClO2- is chlorite
ion -2 O ClO- is hypochlorite ion So we get
(use name of cation name of nonmetal
group) NaClO4 sodium perchlorate NaClO3 sodium
chlorate NaClO2 sodium chlorite NaClO sodium
hypochlorite If more than one of a group is
present, we place the group in parentheses with a
number outside indicating how many of the group
are present. Mg(NO3)2 magnesium nitrate
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Hydrogen containing anions HPO42- is hydrogen
phosphate ion H2PO4- is dihydrogen phosphate
ion HCO3- is hydrogen carbonate (bicarbonate)
ion HSO4- is hydrogen sulfate (bisulfate)
ion Examples NaNO3 sodium nitrate CuSO4 copp
er (II) sulfate Zn(ClO3)2 zinc
chlorate Zn(ClO2)2 zinc chlorite KH2PO4 potas
sium dihydrogen phosphate
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Acids. A substance that produces H ions when
added to water (Arrhenius definition). a)
Binary acids (hydrogen nonmetal) hydro
nonmetal ic acid (when in aqueous
phase) Examples HBr hydrobromic acid
H2S hydrosulfuric acid b) Ternary acids
(hydrogen oxygen nonmetal) hypo ________ite
ion becomes hypo ________ ous acid ________ ite
ion becomes ________ ous acid ________ ate ion
becomes ________ ic acid per ________ ate ion
becomes per ________ ic acid So -ate is changed
to ic acid, -ite is changed to -ous acid.
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Example ClO- hypochlorite ion HClO hypochlorous
acid ClO2- chlorite ion HClO2 chlorous
acid ClO3- chlorate ion HClO3 chloric
acid ClO4- perchlorate ion HClO4 perchloric
acid Binary molecular compounds (compounds are
usually two non-metals) a) Left or lower element
is named first, second element is given an ide
ending, prefix is used to indicate the number of
atoms per molecule (but the prefix mono is never
used for the first element) mono 1 tri
3 penta 5 hepta 7 di 2 tetra 4 hexa
6 octa 8 Example NO nitrogen
monoxide NO2 nitrogen dioxide N2O dinitrogen
monoxide (nitrous oxide) b) A few molecules have
common names H2O is water, NH3 is ammonia, CH4
is methane.
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c) If a binary molecular compound forms an acid
when added to water the naming of the compound
depends on whether it is in the gas phase or
aqueous phase. HCl(g) hydrogen
chloride HCl(aq) hydrochloric acid
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Organic Molecules There are systematic rules
for naming organic molecules, but you will not be
responsible for naming such compounds. CH3CH2CH2
CH2CH3 n-pentane CH3CH2OH ethyl
alcohol CH3CH2OCH2CH3 diethyl ether
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Hydrocarbons Hydrocarbons are molecules that
contain only carbon and hydrogen. One kind of
hydrocarbon in an alkane. Alkanes have the
general formula CnH2n2, where n 1, 2, 3,
n Formula Name 1 CH4 methane 2 C2H6
ethane 3 C3H8 propane 4 C4H10
butane 5 C5H12 pentane



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Functional Groups Organic molecules can also be
classified based on whether or not they contain a
particular group of atoms, called a functional
group. Molecules containing the same functional
group often have similar physical and chemical
properties.
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End of Chapter 2
the ultimate particles of all homogeneous
bodies are perfectly alike in weight, figure, and
so forth. - John Dalton, A New System of
Chemical Philosophy (1808) Elements arranged
according to the size of their atomic weights
show clear periodic properties. - D. I.
Mendeleev (1869) I dont believe that atoms
exist! - Ernst Mach (1897) One of the wonders
of this world is that objects so small can have
such consequences Any visible lump of matter -
even the merest speck - contains more atoms than
there are stars in our galaxy. - P. W. Atkins