Download Honors Chemistry 2011

Honors Chemistry
2011 – 2012
Summer Assignments
May 25, 2011
Dear ’11 -‘12 Honors Chemistry Students,
I have put together the enclosed summer syllabus and materials to accomplish an introduction to
Honors Chemistry. It is comprised of six reading and problem solving sets covering the first two
chapters of the course text. I estimate that a thoughtful effort would require about 10 hours.
These assignments serve the purpose of introducing preliminary topics that will allow us more
time to cover the College Board’s SAT II Chemistry examination syllabus. However, please
understand that our short school year makes covering the entire SAT II curriculum untenable.
Therefore, students interested in taking the Chemistry SAT II should purchase a commercial testprep book (e.g. Princeton Review, Barron’s, etc.) and develop a plan to bridge the gap in a
gradual and disciplined manner.
These summer assignments are required. The deadline for mailing me your completed work is
August 20. You may send it by regular mail or e-mail scanned copies. Please keep the two text
chapters – do not send them.
These introductory topics represent important foundational material. If you learn them well
(particularly the Chapter 2 material) you will likely fair much better on early assessments.
I thank you in advance for your efforts in this regard.
Best wishes,
Mr. Kemer
P.S. pdf versions of all the materials will be available in our FirstClass conference and the
Library Website. Please let me know if you are not subscribed to this conference by mid-June
Honors Chemistry 2011-2012
Summer Syllabus
Chapter 1. Introduction
Assign.
1.
Reading
Exercises
Chapter 1
1: 1 through 9
Chapter 2. Atoms, Molecules, and Ions
2
Sec. 2.1
2: 1 through 8
3.
Sec. 2.2, 2.3
2: 9 through 15
4.
Sec. 2.4
2: 16, 17, 18, 19, 20, 21, 24, 27, 28, 33, 34
5.
Sec. 2.5- 2.7
2: 36, 39, 40, 41, 43, 44, 47, 49, 50, 51
6.
Sec. 2.8
2: 61 through 70
Chapter 1. Introduction to Chemistry
1.1. The Study of Chemistry
Chemistry's primary goals are to describe the macroscopic properties and transformations of matter and
then explain these in terms of submicroscopic models based on the interactions between atoms.
Chemists have discovered that the incredible variety of matter in the physical world arises from only 92
naturally occurring elements whose smallest unit is an atom (see periodic table provided). Different
atoms link in innumerable ways to form molecules, the smallest unit of compounds. Millions of
different molecules, and therefore unique compounds can be formed from the 92 atoms.
Molecules may be represented using colored spheres to represent each element. Several molecules of are
shown in Figure 1.1. In these examples the red spheres represent oxygen atoms, the black spheres carbon
atoms and the white spheres hydrogen atoms.
First, note that the element oxygen exists as a diatomic molecule and is represented symbolically as O2 .
There are several elements that exist naturally as diatomic atoms. These include hydrogen ( H 2 ),
nitrogen ( N 2 ), fluorine ( F2 ), chlorine ( Cl 2 ), bromine ( Br2 ), and Iodine ( I 2 ). Second, note that ethanol
molecules contain two atoms of carbon (black), six atoms of hydrogen (white), and one atom
€ of oxygen
(red). Ethanol is the type of alcohol found in beverages such as beer and wine. Note that the addition of
€
just one oxygen atom to an ethanol molecule converts it to ethylene glycol, which is the thick toxic liquid
€
€ antifreeze.€The important€lesson to learn here
€ is that minor and subtle changes in
used in automobile
the composition and atomic arrangements of atoms in molecules can cause profound changes in
macroscopic properties.
Molecules are not the only basic
unit of elements and compounds.
Most metal elements exist as
simple atoms aligned in a 3dimensional ordered lattice as
shown in the figure below. Most
compounds formed from the
chemical combinations of metals
and nonmetals, such as sodium
chloride (common salt) exist in a
3-dimensional ordered lattice of
alternating metal and nonmetal
atoms. We will have much more
to say about these different types
of structures later in the course.
Metal atoms packed in a
crystal lattice
1-1
Sodium atoms (purple)
and chlorine atoms (s)
packed in a crystal lattice
1.2 Classifications of Matter
Perhaps the most straightforward way to begin our study of chemistry is to establish a way of classifying
matter according to its properties. In general, classification schemes serve to organize objects and
processes in terms of their similarities and differences. Classification schemes allow us to discern
patterns of behavior that may point to underlying principles or laws. Biologists use this approach in the
study of living creatures which are classified in terms kingdom, phylum, class, order, family, genus, and
species. The simplest way of classifying matter is according to its physical state, or phase.
The Physical States of Matter
Matter generally exists in one of three states (or phases) depending on the temperature and external
pressure. These are the familiar sold, liquid, and gas phases. The physical behaviors associated with each
phase are listed below along with the simple kinetic-molecular model scientists use to explain them. The
figure at the bottom of the page illustrates some features of the model.
Macroscopic Properties:
Submicroscopic Theory:
Solid
Definite shape (rigid, resists deformation), definite volume (little
compressibility or thermal expansion)
Atoms/molecules are closely and tightly bound in an ordered lattice. They
vibrate about fixed position.
freeze ↑↓
Macroscopic Properties:
Submicroscopic Theory:
melt
Liquid
Indefinite shape (flows under its own weight or shear stresses), definite
volume (little compressibility or thermal expansion)
Atoms/molecules are close together but not so tightly bound; they can slide
and jiggle past each other.
condense ↑↓ vaporize
Macroscopic Properties:
Submicroscopic Theory:
Gas
Indefinite shape (fills all available volume), indefinite volume (highly
compressible), expands by significant fractions with temperature
Atoms/molecules are not bound to each other; they fly about randomly and
independently.
The figures above are limited in the sense that they are they are static and two-dimensional. You will
have to imagine the random motions and 3rd dimension” showing (1) molecules in the solid phase
vibrating, (2) molecules in the liquid phase jiggling, rotating, and sliding past each other, and (3)
molecules in the gas phase rotating, vibrating and translating at high speeds through the empty space
between them (and colliding when they run into each other).
1-2
Most compounds can exist in each of the three states if brought to the right conditions of temperature and
pressure. A common example of a single compound that exhibits all three of its phases at temperatures that
humans exist in is water.
melt
boil
→
At 1 atm of pressure:
Water: solid ice 0 C
←
freeze
→
liquid water 100 C
gaseous steam
←
condense
Note that it is possible for a solid to transform directly into gas and vise-versa. These processes are called
sublimation and deposition respectively.
€
€
Classifying Matter by Analysis.
Throughout human history people have interacted with their material world, driven both by the
requirements of survival and by their curiosity. Simple observations revealed that many natural substances
were combinations of simpler materials. These simpler materials seemed to be “pure” in the sense that they
were uniform in all of their properties (such as color, density, hardness, etc.) and seemed to be impervious
to further separation. This experience suggested the concept of elements, the most basic or fundamental
components of matter.
A major goal of early chemists was to identify the elements. This involved dogged efforts to utilize fire and
other materials (such as acids) to separate naturally occurring materials into their fundamental components.
The separation and identification of the components of materials is called chemical analysis.
Over time, the early chemists became quite sophisticated in their analytical methods. Among the greatest
accomplishments was the discovery that many materials could be broken down by the aggressive
application of heat or electricity. Substances that were successfully broken down by these techniques were
called compounds. Substances that could not be broken down further were taken as elements.
These early chemists found that when elements combined to form compounds they did so in definite mass
proportions (or ratios). For example, common salt (NaCl), always contains 39.3% sodium and 60.7%
chlorine. Water contains 12.5% hydrogen and 87.5% oxygen.
By the 19th century, the long road of analysis led chemists to an established classification scheme that went
beyond the simple physical state (solid, liquid, or gas). This scheme was simple and practical. It was based
on characteristic properties and separation methods that took advantage of differences in characterisitc
properties. Before outlining the classification scheme, we must clearly define what is meant by
characteristic properties and then discuss some separation methods based on them.
•
Characteristic properties are those that can be used to distinguish one material from another. There are
two categories of characteristic properties
1.
Physical properties can be observed or measured without altering the material's composition
by chemical reaction.
color
density
solubility
magnetism
electrical conductivity
luster
ductility
boiling point
hardness
thermal conductivity
1-3
specific heat
viscosity
odor
melting point
2.
Chemical properties are those that relate how a material changes in composition upon
interacting with another material. In other words, they relate to chemical reactions and lead
to dramatic changes in physical properties
Reacts with oxygen
Decomposes upon heating
Reacts with acids
Decomposes with electricity
The chief indicators of a chemical change are changes in color, the generation of
significant amounts of thermal energy, and/or the generation of a gas.
In addition to characteristic properties are state properties. State properties cannot be used to distinguish
one type of material from another because any substance may share them. State properties include
temperature
volume
pressure
concentration
mass
A secondary descriptor used for properties is extensive or intensive. Extensive properties are those for
which it makes sense to add or accumulate. Mass and volume are the most common examples. Intensive
properties are those that can vary from point to point within a sample of matter. Examples of intensive
properties include temperature, pressure, and all characteristic properties. Note that some intensive
properties are defined by dividing one extensive property by another. Density is an example of this.
Separation Methods can be classified as either physical or chemical separation.
1.
Physical separation methods exploit differences in the physical properties of materials to achieve
the separation.
Physical separation method
Picking apart
Selective floatation/sedimentation
Distillation (selective vaporization)
Magnetic separation
Filtering
Chromatography
2.
Physical property(s) difference utilized
Color, size
Density relative to a liquid
Boiling point
Attraction to a magnet
Size, solubility
Solubility
Chemical separation methods involve chemical decomposition reactions
Chemical Separation Method
Electrolysis
Thermal Decomposition
Selective precipitation
Chemical Property Utilized
Decomposes with electricity
Thermal stability
Reacts to form an insoluble species
So, we can imagine a program of research undertaken by the early chemists to uncover the simplest forms
of matter. Armed with the concepts of physical and chemical properties, and a wide assortment of
separation methods based on these, they “attacked” samples of naturally occurring material, which in most
occasions was obviously not fundamental (e.g. dirt has many different colored components in it) but
became more challenging the further along they got. Ultimately they arrived at the following classification
scheme
I.
Substances are defined as material samples that are uniform in all their characteristic properties
(below the microscopic level) and cannot be separated into two or more simpler forms by physical
separation methods.
1-4
Substances are further classified according to whether they may still be separated by chemical
means.
1.
Elements are substances that cannot be broken down into simpler substances chemical means.
Elements are the simplest substances from which all matter is composed.
Note: The modern explanation of elements is that each element is made up of one type of atom.
Elements cannot be broken down by chemical means because the atom is the basic chemical
building block.
2.
Compounds are substances that can be separated into two or more elements by chemical means
and have fixed elemental compositions.
Note: The modern explanation for compounds is that a compound is made up of two or more
types of atoms strongly bonded together in fixed number ratios. The interatomic bonds are too
strong to be broken by physical means but can be broken by the more forceful chemical
reactions into more basic elements.
The molecular basis of compounds explains the Law of Definite Mass Proportions, which
states that a particular compound always has a definite mass ratio of elements. The definite
mass ratio arises from the definite number ratio of atoms that make up its molecules. For
example water is a compound made from the elements hydrogen and oxygen. The properties of
water are distinctly different from either hydrogen or oxygen. The chemical composition of
water is 11.11% hydrogen and 88.89 % oxygen by mass. Modern theory explains this fixed
mass ratio by stating that (i) water is made up of molecules comprised of a single atom of
oxygen bonded to two atoms of hydrogen, and (ii) oxygen atoms are 16 times heavier than
hydrogen atoms (oxygen atoms weigh 16 atomic mass units on a scale that defines the mass of
hydrogen to be 1 atomic mass unit) .
Example: Carbon dioxide is a compound of carbon and oxygen. It when decomposed it is
found to be 27.29% carbon and 71.71% oxygen by mass. If a molecule of carbon dioxide is
made up on 1 atom of carbon and 2 atoms of oxygen (CO2) and if carbon atom weighs 12
atomic mass units, what is the mass of an oxygen atom?
Answer: We can set up a ratio to solve this problem with M O and M C representing the
atomic masses of oxygen and carbon and N O and N C representing the number of oxygen and
carbon atoms per molecule
€71.29 2€⋅ M
mass % O N O ⋅ M O
O
= €
=
€⇒
mass % C N C ⋅ M C
27.29
1⋅12
⇒ M O = 16
II. Mixtures are combinations of two or more substances that can be separated by physical means.
They can also exist in a range of mass proportions and, and retain the properties of each component
substance.€
1.
Homogeneous mixtures (or solutions) are mixtures that appear completely uniform in their
properties and composition (below the microscopic level). In this regard, solutions look like
pure substances. The difference is that solutions can be separated by physical means while
substances cannot. Solutions can also exist in a wide range of component proportions while
compounds are made of elements in fixed or definite proportions.
2.
Heterogeneous mixtures are mixtures whose characteristic properties and composition vary
from point to point. These can be identified by the variations in color, texture, or any other
physical property within the material, even if the variations occur on a microscopic scale. For
instance, a mixture of fine white powdered sugar and fine white powdered salt is heterogeneous
even though it may appear uniform in color and texture.
1-5
Most materials encountered in nature are heterogeneous mixtures because it is impossible to completely
remove all impurities from any sample of matter. The purest manufactured material is silicon used in
electronic chips. It is 99.9999% pure. However, still represents billions of billions of impurity atoms per
cubic centimeter. In a more common example, “clear” saltwater will invariably have microscopic dust
particles in it. The ideas of a pure substances and homogeneous solutions is an idealization.
Example: A material sample appears to the naked eye as a gray powder. When it is stirred into water
and filtered through a fine mesh paper into a beaker about half of the original solid sample is trapped
by the filter and appears black. The remainder of the sample is recovered by letting the water evaporate
from the beaker. It appears white. Classify the original sample.
Answer: The original sample was a heterogeneous mixture. It was a mixture because it was
separated by physical means (filtration). It was heterogeneous because black and white powders
were not mixed on a molecular scale. The black particles were larger than the holes in the filter.
The recovered white powder is placed inside an evacuated chamber and heated. When the temperature
reaches 700˚C the powder began to give off a green gas and turned an orange. Classify this
component?
Answer: This component was most likely a compound since there is good evidence of chemical
decomposition. The fact that it was heated in a vacuum is important to this conclusion. If the
reaction occurred in air we could not be sure that the chemical change was a chemical
decomposition. The powder might have been an element reacting with a component of the air.
The definitions for substances and mixtures are examples of operational definitions. This is because they
are based on definite practical operations (procedures). These are outlined schematically represented in the
flowchart on the next page that could serve chemists as a “roadmap” for analysis.
Sample material
Can it be separated
by physical means?
yes
no
Mixture
Substance
In it uniform
throughout (down to
the molecular scale)?
yes
homogeneous
mixture
Can it be separated
by chemical means?
yes
no
compound
heterogeneous
mixture
no
element
Flowchart for the classification of matter incorporating the operational definitions of homogeneous and
heterogeneous mixtures and compound and elemental substances.
The alternative to an operational definition, at least in chemistry, is a theoretical definition. Theoretical
definitions are based on concepts that are based on analogies or models rather than direct or measurable
procedures. For example, the theoretical definition of an element is a sample of matter made up of one type
of atom.
1-6
Our theoretical understanding of the difference between substances (elements and compounds) and
mixtures of substances is based on the notion of chemical bonding. Elements are pure substances that
cannot be separated into simpler components because they are made up of a single type of atom that cannot
be broken apart further1. Compounds are pure substances because they are made up of molecular units,
specific groupings of atoms that are held together by strong forces (chemical bonds) that cannot be
disrupted the relatively gentle forces associated with physical separation methods, but can be separated by
the more aggressive forces associated with chemical separation methods. To separate the chemically
bonded atoms or molecules takes considerable energy.
On the other hand, mixtures are simply separate substances that lie side-by-side and are not
bonded together by particularly strong forces. Each component substance making up a mixture
retains its own identity and can be separated from the other components without much energy
expenditure. This is why they can also have a range of compositions. Physical separation methods
do not involve as much energy.
1
Atoms can be broken apart in nuclear reactions in which their nuclei break apart (radioactive decay or
nuclear fission). But we generally do not consider nuclear reactions in chemistry.
1-7
Honors Chemistry- Summer Assignment 1
Refer to Chapter 1 Reading
8
1.1. Estimate the temperature reading from the dialtype thermometer shown below. Express your
estimate with a reasonable “± absolute error” range.
Be as precise as can be reasonably justified (this can
be better than ±5˚C)
7
6
Figure for Problem 1.7.
1.2. What macroscopic physical characteristics do
solids and liquids share? Explain these similarities in
terms of the molecular model.
1.3. What are the differences between the macroscopic
physical characteristics of liquids and gases. Explain
these differences in terms of the molecular model.
1.4. Consider a solution of sugar and water. (a) Why
is this combination defined as a mixture? (b) Why is
this combination defined as a homogeneous mixture?
(c) Describe how one might separate the sugar from
the water. (d) What physical property does your
method in part (c) exploit?
1.5. A yellow granular material turns black and emits
a colorless gas when heated in air. Eventually it
disappears. Why would it be wrong to conclude from
this process that the original material is a compound
that had decomposed?
1.6. Answer the following True/False questions. (a) A
state property can be used to distinguish one substance
from another. (b) An intensive property of a sample
substance is proportional to the mass of the sample.
(c) The components of a mixture must differ in at least
one physical property in order to separate them.
1.7. Which of the following measurements is the most
appropriate one to report for the temperature indicated
in the thermometer illustrated in the figure below?
Remember that it is possible to interpolate between
scale marks to the nearest one tength of the scale unit.
(a) 7 ˚C
(b) 7.2 ˚C
(c) 7.24 ˚C
1.8. A student partially fills a graduated cylinder of an
unknown liquid to a volume of 2.32±0.02 ml. She
then places the cylinder on an electronic balance and
determines its mass (glass cylinder plus the liquid it
contains) to be 73.41±0.01 grams. Next, she adds
more of the same liquid to a volume of 6.73±0.02 mL
and then determines the new combined mass to be
77.89±0.01 grams. (a) Show that the density of the
liquid is best reported as 1.02 ± 0.02 g/mL (or 1.016 ±
0.014 g/mL) (c) If the student failed to notice that a
bubble of air was trapped under the liquid surface
during the second addition of liquid, how would her
measured density be in error? Would her value be too
great or too small?
1.9. A 32.65-g sample of a solid is placed in a flask.
Liquid toluene, in which the solid is insoluble, is
added to the flask so that the total volume of solid and
liquid together is 50.00 ml. The solid and toluene
together weigh 58.58 g. The density of toluene at the
temperature of the experiment is 0.864 g/ml.
Calculate the density of the solid. Since absolute
errors are not given, report your result with the
appropriate number of significant figures (in this case
there will be 4 significant figures in the final
calculation).
Chapter 2. Atoms, Molecules, and Ions
2.1 The Atomic Theory of Matter: The Early History
As stated in Chapter 1, chemistry's primary goal is to explain the macroscopic behavior of matter in
terms of the atomic scale structures and interactions. This chapter provides brief historical sketch of the
major steps that led our modern atomic theory that accomplishes this goal.
Ancient Theories
The ancient Greek philosopher Leucippus and his pupil Democritus (470-380 BC.) were the first to
speculate that matter is made of indivisible particles called atomos (a = "not" + tomos = "to cut"). Their
theory was not widely accepted in their time. Most Greek philosophers, including Aristotle (384-322
BC.), believed that matter was infinitely divisible (i.e., continuous). Aristotle further speculated that all
material derived their properties from the various proportions of just four elements: earth, air, fire, and
water. Their view prevailed for 2000 years.
While the Greeks can be given credit for trying to provide rational explanations for the physical world,
they never adopted the modern scientific approach of using experiments to "test" scientific models or
theories. Their ideas did not amount to scientific theories in the modern sense. Rather, they were mere
speculations aimed at fitting observations to preconceived notions about how the world should be.
From pre-historic times curiosity and practical
needs have led enterprising and curious humans to
manipulate their material environment, beginning
with bone and stone tools, fire, and native metals.
The Medieval alchemists (500 AD-1500 AD)
combined such practical interests with mystical
strivings to search for a "philosopher's stone," a
substance that could transform base metals (copper,
iron, zinc, etc.) into gold. Alchemical "theories" of
transmutation seem strange by today's standards.
Ultimately, none of their ideas could be supported
by experiment.
Although the Alchemists' primary goal was never
Alchemists doing strange experiments.
achieved, their work resulted in many practical
techniques and knowledge that supported the development of modern chemistry later on.
The Dawn of the Modern Atomic Theory
In 1661, Robert Boyle defined an element as a substance that cannot
be broken down into simpler substances by chemical operations (e.g.,
heating). This represented the modern and eminently practical idea of
an operational definition. It was based on practical procedures and
facts and not on preconceived notions or mystical beliefs about the
ultimate nature of matter.
In 1789, Antoine Lavoisier published the first chemistry textbook
that listed 33 elements. While a few of these elements were later
found to be compounds, the notion of fundamental substances with
unique and definite properties was firmly established.
Portrait of Antoine Laurent and MarieAnne Lavoisier, his wife and collaborator
2-1
Lavoisier’s experiments utilized very sensitive balances to measure and compare the mass of reactants
and products in chemical reactions suggested the idea that matter is neither created nor destroyed. That
is, he promoted the Law of Conservation of Mass:
Matter is neither created nor destroyed in any physical or chemical process.
The total mass of the reacting substances equals the mass of the products
Lavoisier’s use of the mass balance placed chemistry on a firm quantitative footing and led to other
important chemical laws. In 1799, Joseph Louis Proust discovered that copper carbonate, whether
obtained from natural sources or synthesized in the labortory, always contained copper, oxygen, and
carbon in the same proportions by mass, (i.e., 5.3 parts Copper to 4.0 parts Oxygen to 1.0 parts Carbon).
Similar results with other substances led to the formulation of the Law of Definite Proportions:
Compounds always contain elements in certain definite proportions by mass.
Dalton's Atomic Theory
In 1803, John Dalton published the first modern atomic theory to
explain the unique identity of elements, what happens in a chemical
reaction, the Law of Conservation of Mass, the Law of Definite
Proportions, and the Law of Multiple Proportions.
Dalton's Atomic Theory consisted of five postulates. These postulates
along with the explanations their provided are:
1.
All elements are made up of tiny, indivisible particles called atoms.
Atoms can neither be created nor destroyed during chemical
reactions.
This postulate explains the conservation of mass.
2.
Dalton imagining atoms.
All atoms of a given element are identical, but the atoms of one
element differ from the atoms of another.
Along with the first postulate, this explains why elements cannot
be broken down into simpler substances.
3.
Atoms of different elements form compounds by joining in
fixed, small, whole number ratios to form molecules.
This postulate explains the law of definite proportions. If
compounds are made of molecules with fixed number ratios of
atoms then they must also have fixed mass ratios of elements.
4.
A chemical reaction involves a change not in the atoms
themselves, but in the way the atoms are combined to form
compounds.
This postulate explains chemical reactions and emphasizes the
fundamental nature of atoms.
Dalton proposed that the atoms of each element are distinguished by
their mass. He assigned the hypothetical hydrogen atom, the lightest,
a relative mass of 1. He and others developed a table of atomic
weights of the other elements from experimentally measured mass
proportions of compounds and guessing the number ratio of atoms in
the compound (i.e., the compound’s formula)
Dalton’s molecular models
2-2
Example: The mass ratio of oxygen to hydrogen in water is 8:1. The mass ratio of carbon to oxygen in
carbon dioxide is 3:8. Determine the atomic masses of oxygen and carbon relative to hydrogen (which is
arbitrarily assigned the value of 1 atomic mass unit) given that the formula of water is H 2O (2 atoms H
to 1 atom O) and the formula of carbon dioxide is CO2 (1 atom C to 2 atoms O).
You can probably figure out the required atomic masses in (please try!). Oxygen would have an atomic
! model can help you
mass of 16 and carbon would then be 12. However, the following mathematical
quickly and confidently calculate them is!very useful for more complicated cases.
Let
mA
be the measured definite mass ratio of element A to element B in a compound of A and B.
mB
NA
be the hypothetical (guessed) number ratio of A atoms to B atoms in a molecule of the
NB
compound.
Let
!
!
!
Let M A and M B be the atomic masses of atoms of elements A and B relative to hydrogen, which is
arbitrarily assigned the value 1 atomic mass unit. Then, Dalton’s Atomic Theory explains the Law
of Definite through the expression
mA N A M A
=
!
mB N B M B
where the left side expresses the empirical measurements that go into the Law of Definition
Proportions and the right side expresses the theoretical justification. Applying this
! in water, assuming the atom number ratio of hydrogen
expression to oxygen and hydrogen
to water is N H N O = 2 /1 leads to
mO N O M O
=
mH N H M H
" MO =
$ 2 ' $8 '
N H mO
#
# M H = & ) # & ) # (1) = 16
%1 ( %1 (
N O mH
!
And applying this formula for carbon in carbon dioxide given that M O = 16 gives
$ 2 ' $ 3'
N O mC
#
# M O = & ) # & ) # (16) = 12
%1 ( %8 (
N C mO
!
Although Dalton's theory successfully explained many chemical laws, it did not prove that atoms existed.
Atoms were considered hypothetical particles. Strong evidence for the actual existence of atoms was not
available
! until the early 1900's.
mC N C M C
=
mO N O M O
!
" MC =
The Law of Multiple Proportions
Soon after proposing his atomic theory John Dalton announced the Law of Multiple
Proportions and was able to employ his theory to provide a convincing explanation.
When two elements form more than one compound, the ratio of the mass ratio of
first element to the second element in one compound to the mass ratio of first
element to the second in the other compound is a simple whole number ratio.
Example: Nitrogen and oxygen form three compounds that show the law of multiple proportions
Compound
number
1
2
3
mass O
mass N
Compound
name
Nitrous oxide
Nitric oxide
Nitrogen dioxide
0.571
1.142
2.284
!
!
2-3
" mass O %
$
'
# mass N &i
" mass O %
$
'
# mass N &1
1
2
4
The Law of Multiple Proportions and its explanation by Dalton’s atomic theory can be mathematically
modeled by simply extending the simple Definite Proportions model outlined above. The expression for
the Law is given on the left side and the following expression and the theory on the right side.
" N (M %
" N %
" mass A %
A
$ A
'
$ A'
$
'
N
(
M
# mass B &cmpd 1
# B
# N B &cmpd 1
B &cmpd 1
=
=
" mass A %
" N (M %
" N %
A
$
'
$ A
'
$ A'
# mass B &cmpd 2
# N B ( M A &cmpd 2
# N B &cmpd 2
All of the symbols have the same meaning as before. That the ratio of mass ratios turns out (empirically)
to be a ratio of simple integers (the Law of Multiple Proportions) follows directly from a model of
compounds!having as their smallest unit molecules that have simple integer number ratios of atoms.
Example: If nitrous oxide molecules (see table above) have the formula N 2O (each molecule contains
two atoms of nitrogen and one atom of oxygen) then what must be the formula of nitric oxide?
Applying the above expression leads to:
" mass
$
# mass
" mass
$
# mass
!
O%
'
N &cmpd
O%
'
N &cmpd
1
2
"
$
#
=
"
$
#
NO %
'
N N &cmpd
NO %
'
N N &cmpd
!
1
2
"1%
$ '
0.571
#2&
(
=
" N %
1.142
$ O'
# N N &cmpd
(
" N %
$ O'
# N N &cmpd
= 1
2
2
Therefore, the simplest possible formula is NO. But it also may be true that the actual molecules are
N 2O2 or N 2O2 , we cannot tel.
Problems with Dalton's Theory
!
! the early 20th century experimental evidence had arisen that showed that Dalton's Theory was
By
incomplete and only approximate. For example
1.
Nonstochiometric compounds having a range of mass ratios of elements were found.
2.
Not all of the atoms of a given element had precisely the same mass. These were called isotopes.
3.
Atoms are not indestructible and have a subatomic structure. The discovery of the “radioactive
decay” of some elements led to the development of nuclear fission and fusion.
These "chinks" in the armor of Dalton's Theory led to the development of the modern theory of the atom
during the late 19th century and early 20th century. Some of the major steps of this development will be
given here, with the rest taken up in Chapter 6.
Still, great progress in chemistry followed Dalton’s program. Dozens of elements were discovered by the
mid 19th century and the relative atomic masses were “weighed” by relating the relative masses of
elements in compounds to the hypothetical number ratio of atoms in the compound.
2.2. The Discovery of Atomic Structure
In 1879, William Crookes discovered cathode rays. These were beams of particles that were
emitted from a highly charged negative metal electrodes embedded inside a partially evacuated
tube. As the cathode rays passed through the gas they made it glow. When the rays struck a
phosphorescent screen, they would cause it to glow also. They could not, however, pass
through the glass walls of the tube.
2-4
In 1895 Wilhelm Roentgen discovered X-rays. These were later determined to be
high-energy light emitted when cathode rays struck a metal target. X-rays traveled
many meters through air outside the tube. Their ability to pass through flesh, but
not as well bone, led to X-ray photography
In 1896 Antoine Bequeral, Madame Curie, and Pierre Curie observed that
uranium spontaneously emitted X-rays without cathode ray bombardment. They
call it the process radioactivity. Their discovery suggested that "something" was
coming out of the atom. In other words, atoms have "insides" or parts.
In 1897 J. J. Thomson demonstrated that cathode rays are negatively charged
particles by repelling them with a strong negative charge and by deflecting them
sideways with a magnet1. He names them electrons. Thomson also measures the
charge to mass ratio of the electron (q/m). Thomson concluded that electrons are
subatomic particles and a common building block of all matter. The
following experimental results supported this conclusion.
1.
2.
3.
All electrons are identical no matter what metal the cathode material.
Since electrons have negative charge the other parts of the atom must
be positively charged.
The cathode does not lose mass. If cathode rays were whole atoms
the cathode would eventually show a loss of mass.
In 1904 J. J. Thomson developed his "plum pudding" model of atom. It
depicted electrons imbedded like raisins in a continuous positively charged
pudding-like mass. Thomson's model was more sophisticated than this
sounds. He worked out the precise positions the electrons must occupy
within the positive charge to be stable.
In 1909 Robert Millikan determined the charge on an
electron by performing a Newton’s Second law analysis of
the terminal velocities of tiny charged oils drops falling
between electrically charged plates2. This charge was
hypothesized to be the elementary (or smallest) unit of
charge ( 1.6x10"19 Coloumb ). They don’t get smaller!
Millikan’s determination of the electron charge allowed the
mass of the electron to be calculated from Thomson’s q/m
!
ratio. That is:
1
A negatively charged particle moving across a flat-face of the north pole of a magnet will be deflected sideways. This
important effect led to the modern means of measuring atomic masses, as we shall soon see.
2
We will see precisely how this experiment works in Experiment 14. It utilizes several concepts that we learned in
Honors Physics.
2-5
me =
e
"1.6x10"19 Coloumb
=
= 9.11x10"28 g
e / m "1.76x10 8 Coloumb / gram
The mass of the electron was 1/1827 the mass of a hydrogen atom, which supported Thomson’s theory
that the electron is a subatomic particle.
!
In 1907 Ernest Rutherford identified three subatomic particles emitted from radioactive elements. He
named them alpha, beta, and gamma particles after the first three letters of the Greek alphabet. These
particles were later shown to the positive nucleus of a helium atom, electrons, and gamma light rays,
respectively3. This was further evidence that atoms have parts; there exists a subatomic structure.
Particle Name
alpha
beta
gamma
Symbol
"
"
"
Mass (amu)
4
1/837
0
Charge
+2
-1
0
Later Identified as
Helium nucleus
electron
High energy light
!
In 1911 Rutherford and Marsden performed an experiment in which they bombarded a very thin gold
! alpha particles emitted from radioactive uranium (see Figure at top of next
foil with positively charged
page). While Thomson's!“plum pudding” model predicted that the particles would fly straight through
without slowing down, they found that a very small percentage were deflected through very large angles.
The rare deflections suggested that the vast majority of an atom's mass was concentrated in a very tiny
positively charged nucleus. Most positive alpha-particles pass straight through because they miss any
nucleus by wide margins. Only a few pass close enough to a nuclei to suffer a large deflection. A
statistical analysis of the scattering allowed Rutherford to deduce the nuclear diameter to be about
1/10,000 the diameter of the atom! Atoms were mostly empty space. This evidence suggested a
planetary model of the atom in which negatively charged electrons orbited the tiny, but massive,
positively charged nucleus, held in orbit by the attractive electrical forces. To make the atom neutral, the
number of orbiting electrons must equal the nuclear charge.
3
The beta particle is therefore identical to a cathode ray. They are both electrons. The distinction is the source.
Cathode rays are produced by strong electric fields that “pull” them from the outer “orbit” of atoms. Beta particles
originate from inside the atomic nucleus during a radioactive decay process.
2-6
2.3. A Preview of the "Modern" View of Atomic Structure
While Rutherford's planetary model still lingers in the minds of nonscientists, it actually came under
immediate attack because it contradicted the well-established theory of electromagnetic radiation. This
theory predicted that an orbiting electron would continuously convert its kinetic energy into light energy.
As a result it should spiral into the nucleus. Still, there was no doubt that an atom was comprised of a
very small, massive, positive nucleus and electrons that were held close to it by electrical attractions.
Neil's Bohr, a Danish physicist and student of Rutherford, was the first to propose an alternate model.
His model was supported by the discoveries and theories of Albert Einstein and others in the early 20th
century. The major threads of Bohr's contribution to atomic theory, as well its subsequent development
into the most modern form of atomic theory, called quantum mechanics, will be studied up in Chapter 6.
For now it is worth describing just a few of the features of modern atomic theory. Remember, however,
that a more complete picture will be needed to explain the chemical properties and transformations of
matter, which, after all, is the ultimate goal of Chemistry.
Subatomic Particles: Electrons, Protons, and Neutrons
Atoms are not indivisible as Dalton hypothesized. They are made up protons, neutrons, and
electrons which have the properties given in the table below
Particle
Name
electron
proton
neutron
Symbol
ep+
n
Charge
(e)
-1
+1
0
Approx mass
(amu)
1/1837
1
1
1 amu = 1.6606x10"24 g
!
Precise Mass
(g)
9.10953x10"28
1.67265x10"24
1.67495x10"24
!
!
!
The Structure and Size of the Atoms
Size
(m)
“point”
< 1x10"15
< 1x10"15
1 e = 1.6x10"19 C
!
!
!
The protons and neutrons of an atom are located in a very tiny region called the nucleus. The diameter
of atomic nuclei is on the order of 10"15 m . The electrons "swarm" around the nucleus in regions called
orbitals. The electrons determine the overall size of atoms that are on the order of 10"10 m . That is,
atomic size is determined by this region occupied by the electrons. To give an idea of this scale
difference between the nucleus
and the whole atom, if an atom was magnified to the size of a baseball
!
stadium, the nucleus would be the size of a marble.
!
In the late 1980’s the invention of tunneling electron
microscopes allowed images of atoms and dramatically
confirm these scales. Mr. Kemer took the micrograph
below during his sabbatical appointment as a Visiting
Scholar at Amherst College in 2000. It shows neatly
aligned molecules of docosyl-ether and docosyl-sulfide
lying on a flat surface of graphite. The smallest white
dots are hydrogen atoms. The brightest dots are sulfur
atoms. The figure to the right shows models of these
same molecules (the yellow center atom represents
sulfur and the red center represents oxygen). The
blurriness of the images arises in part to the motions of
the molecules during the scan that produced the image.
The image is about 10 nanometers (10 billionths of a
meter) across.
2-7
Atomic Number and Mass Number
All atoms of a particular element have the same number of protons in their nuclei. This number is called
its atomic number (Z). For instance Hydrogen has 1 proton in its nucleus and therefore has the atomic
number 1. It is the atomic number that identifies the element.
The mass number (A) is defined as the total number of protons and neutrons in the nucleus.
Mass number = atomic number + neutron number
(A = Z + A)
A neutral atom has equal numbers of electrons and protons.
Example: Fill in the blanks of the following table. Use a periodic table or other source for the symbols.
Atomic
Number
7
Mass
Number
Number of
Protons
9
39
59
Number of
Neutrons
7
10
Number of
Electrons*
Element
Symbol
Element
name
nitrogen
19
27
2.4. Atomic Weights
Measuring The Masses of Atoms
The mass of atoms and molecules is accomplished using a mass spectrometer. The basis of mass
spectroscopy is the law of magnetic force on a moving charged particle. This law, called the Lorentz Force
law, states that a charged particle will be deflected sideways into a circular path when is moves across the
face of a magnet having its north pole pointing perpendicular to the particles velocity. The magnitude of
the force is proportional to the particle charge and speed and the strength of the magnetic field.
FB = qvB
+
!
ac = v 2 R
F = ma
N
!
!
qvB = m v 2 R
qBR
m=
v
A detailed diagram of a mass spectrometer is shown below. Note that the atoms are stripped of an
! a single positive elementary charge. This loss of mass must be added in to give the
electron to give them
! magnet points down.
correct mass of the atom. Note that the north pole of the
2-8
Isotopes
Two atoms of the same element must have the same number of protons. However, they may have
different numbers of neutrons and therefore different mass numbers and masses. Atoms of a given
element with different mass numbers are called isotopes (“same position”). Although isotopes of the
same element differ in mass, they generally have very similar chemical properties. The symbol system
used to designate an isotope is
A
ZX
where X is the element symbol, Z is the atomic number, and A is the mass number. Often the atomic
number is not written since it is implied by the element symbol.
!
Isotopes are also expressed with the element name followed by a hyphen and the mass number.
35
37
Example: Two isotopes of chlorine: 17
Cl (chlorine-35) and 17
Cl (chlorine-37)
Example: Two isotopes of Carbon: 126 C (carbon-12) and 146 C (carbon-14)
Example: Hydrogen has three
! isotopes. They are the
! only isotopes given specific names.
1
H
1!
2
1 H!(Deuterium)
(Protium)
3
1H
(Tritium)
Example: Complete the following table using the definitions of atomic number, mass number.
! symbol
isotope
12
Carbon-12
6C
!
Atomic
number
! number
Mass
29
61
No. of protons
No. neutrons
92
146
83
38 Sr
!
!
The Atomic Weight Scale and the Mole
The mass spectrometer allows the masses of atoms (minus an electron) to be measured in gram units. For
instance, the 1H isotope has a mass of 1.6735x10"24 g and the 16O isotope has a mass of 2.6560x10"23 g .
Because it would be cumbersome to continually express such small masses in grams, we instead use a unit
called the atomic mass unit, or amu. The amu is defined in as 1/12 the mass of the 12C isotope.
!
!
!
!
1 amu = 1.66054 x10"24 g .
! 5.486 x 10-4 amu.
A proton has a mass of 1.0073 amu, a neutron 1.0087 amu, and an electron
! as the number of 12C atoms in exactly 12.0 grams of 12C . This is also
The mole is the number defined
the number of amu units in 1.0 gram
1 mole = 6.022x10 23
!
!
The clever thing about this definition of a mole is that the mass of 1 mole of an element (or compound) is
the number of grams numerically equal to its mass in amu. This will be made more clear in Chapter 3.
Average Atomic Masses of the!Elements
The atomic weight of an element is a weighted average mass of its isotopes that reflects the relative
abundances of each isotope as they occur naturally.
2-9
35
37
Example: A sample of chlorine contains 75.78% of 17
Cl and 24.22% of 17
Cl by mass. Calculate the
35
37
atomic weight of chlorine given that the atomic mass of 17
Cl is 34.969 amu and 17
Cl is 36.966 amu.
(0.7578)(34.969 amu)
! + (0.2423)(36.966!amu) = 35.45 amu
35
37
!
Example: A naturally occurring sample of Chlorine,
which contains of!17
Cl and 17
Cl , has an average
atomic weight of 35.45 amu. What is the relative abundance of each isotope?
35
Let x be the fractional relative abundance of 17
Cl . Approximate the atomic mass of these two isotopes
!
!
to be 35 and 37.
(x)(35 amu) + (1 - x)(37 amu) = 35.45 amu
!
Solving for x leads to x = 0.76. Therefore,
35
% 17
Cl = (100%) x = 76%,
37
% 17
Cl = (100%)(1-x) = 24%
The relative abundances of isotopes are determined by mass spectroscopy. On a mass spectrograph the
relative heights of the peaks is directly related to the mass percentages of each isotope.
!
!
Example: The mass spectrograph for chlorine on the top right figure of page 2-9 shows the two peaks
for the isotopes of chlorine. The relative heights of the peaks give the relative mass % abundances. In
this case the relative heights of the 35Cl and 37Cl peaks are 5.0 and 1.6, respectively. Therefore, naturally
occurring chlorine contains 5/(5+1.6) = 76% 35Cl and 1.6/(5+1.6) = 24% 37Cl by mass.
2.5. The Periodic Table
In 1814, J. J. Berzelius invented modern element symbols. Today we know that each element is made
up of a unique type of atom. 89 elements (atoms) occur naturally on the earth. 20 elements have been
synthesized since 1937 by smashing smaller atoms (nuclei) together at nearly the speed of light in
particle accelerators. Most synthesized atoms radioactively decay into stable elements in very short
times
Abundant and Rare Elements
The matter of the known universe is 93% H, 7% He, and < 0.1% all else. The solar system is 85% H,
15% He, and < 0.1% all else. Strangely, more than 60% of the mass of the Universe is of unknown
composition; it is simply called “dark matter.”
Periodic Table of the Elements
In 1869 Dimitri Mendeleev, a Russian Chemist, organized the known elements into a table in order of
increasing atomic weights. He chose to start new horizontal rows (or periods) of the table in such a way
that elements with similar chemical and physical properties lined up in vertical columns (or groups).
2-10
Mendeleev's ordering scheme resulted in some "empty" places that he took for undiscovered elements.
Indeed, on this basis he predicted the existence of scandium (Sc), gallium (Ga) and germanium (Ge), all
of which were discovered shortly afterward.
Modern atomic theory has determined that the true basis for the ordering the elements in the periodic
table is increasing atomic number and that the vertical groups are were comprised of elements whose
atoms have similar arrangements on electrons in the atoms. In other words, modern atomic theory
relates chemical properties to the electronic structure of atoms.
Elements are commonly classified as follows:
Type
Physical properties
Metals
.
Solid at room temperature (except Hg), good conductors of heat and electricity, shiny,
ductile, have a wide range of hardness.
Nonmetals May be solid, liquid, and gas at room temperature, are poor conductors of heat and
electricity, nonmetal solids have dull surfaces and are brittle.
Metalloids have properties are intermediate to metals and nonmetals. Some have unique electrical
properties (semiconductors)
Diatomic
Elements
These exist naturally as diatomic molecules. They include : H 2 , N 2 , and O2 plus
the halogens: F2 , Cl 2 , Br2 , and I 2
Noble:
Gases
These do not ordinarily form compounds with other elements. They are all
monatomic gases at room temperature. They include
! He,
! Ne, Ar,
! Kr, Xe, Rn
! ! !
!
The group designations include:
•
The Group IA metals are called the Alkali metals
•
The Group IIA metals are called the Alkaline Earth metals
•
The Group 1B-8B metals (center portion of the periodic table) are called the Transition Metals
•
The Group VIIA nonmetals are called the Halogens
•
The section of elements 57-70 and 89-102 are called the Inner Transition Metals
•
The other Groups are simply referred to by their Group number or by the element at the top.
2-11
Modern periodic tables include a wealth of physical and chemical information for each element
including the average atomic weight. It is important to review three things about the atomic weight
listed in periodic tables
1. It represents the weighted average of all the naturally occurring isotopes of that element.
2. If interpreted as amu units, it gives the average mass of a single atom.
3. If interpreted as gram units, it gives the mass of 1 mole of atoms of that element.
2.6. Molecules and Molecular Compounds
The atom is the smallest representative sample of an element. However, only the noble gas elements (He,
Ne, Ar, Rn) are normally found in nature as isolated atoms. Most matter is composed of molecules or
ions, both of which are formed from atoms. We examine molecules here and ions in the next section.
A molecule is an assembly of two or more atoms tightly bound together. The resultant "package" of
atoms behaves in many ways as a single, distinct object. We will discuss the forces that hold the atoms
together (the chemical bonds) in Chapters 8 and 9.
Molecules and Chemical Formulas
Many elements are found in nature in molecular form; that is, two or more of the same type of atom are
bound together. For example, the oxygen normally found in air consists of molecules that contain two
oxygen atoms. We represent this molecular form of oxygen by the chemical formula O2 (read "oh two").
The subscript in the formula tells us that two oxygen atoms are present in each molecule. Any molecule
that is made up of two atoms is said to be a diatomic molecule. The other diatomic elements include
hydrogen (H2), nitrogen (N2), fluorine (F2), chlorine (Cl2), bromine (Br2), and iodine (I2)
Oxygen also exists in another molecular form known as ozone. Molecules of ozone consist of three
oxygen atoms, so its chemical formula is O3. Even though "normal" oxygen (O2) and ozone are both
composed only of oxygen atoms, they exhibit very different chemical and physical properties. For
example, O2 is essential for life, but O3 is toxic; O2 is odorless, whereas O3 has a sharp, pungent smell.
Elements that exist in more than one form are called allotropes.
Substances that are composed of molecules made up or two or more types of atoms are called molecular
compounds. For example, water is a molecular compound comprised of molecules made of two
hydrogen atoms and one oxygen atom. It is represented by the chemical formula H2O. Lack of a subscript
on the O indicates one atom of O per water molecule. Another molecular compound composed of these
same elements is hydrogen peroxide, H2O2. The properties of these two compounds are very different.
Several molecules of other molecular compounds are shown in Figure 2.16. Notice how the composition
of each compound is given by its chemical formula. Notice also that these substances are composed only
of nonmetallic elements. Molecular substances contain only nonmetals.
Figure 2.16. Molecular models of some simple molecules.
2-12
Molecular and Empirical Formulas
Chemical formulas that indicate the actual numbers and types of atoms in a molecule are called
molecular formulas (the formulas in Figure 2.16 are molecular formulas). Chemical formulas that give
only the relative number of atoms of each type in a molecule are called empirical formulas. The
subscripts in an empirical formula always give the smallest whole-number ratios of the atoms.
Example: The molecular formula for hydrogen peroxide is H2O2; its empirical formula is HO.
Example: The molecular formula for ethylene is C2H4; its empirical formula is CH2.
For many substances, the molecular formula and the empirical formula are identical, as in the case of
water, H2O.
Whenever we know the molecular formula of a compound, we can determine its empirical formula.
However, the converse is not true; if we know the empirical formula of a substance, we can't determine
its molecular formula unless we know more information. So why do chemists concern themselves with
empirical formulas? As we will see in Chapter 3, common methods of analyzing substances lead to the
empirical formula only. Once the empirical formula is known, additional experiments can give the
information needed to convert the empirical formula to the molecular one. In addition, there are
substances such ionic compounds that do not exist as molecules. For these substances only empirical
formulas are meaningful.
Example: Write the empirical formulas for the following molecules:
(a) glucose (a.k.a. blood sugar or dextrose) whose molecular formula is C6H12O6 : CH2O
(b) nitrous oxide (a.k.a laughing gas) whose molecular formula is N2O: Same as empirical formula
Picturing Molecules
The molecular formula of a substance summarizes its composition but does not show how its atoms join
together. For many molecules it is possible to draw structural formulas that show the connectivity of
the atoms. For example the structural formulas of water, hydrogen peroxide, and methane (CH4) are:
Usually a structural formula does not depict the actual geometry of the molecule. That is, it cannot
show the actual angles at which atoms are joined together in three-dimensions. A structural formula
can be written, however, as a perspective drawing to give some sense of three-dimensional shape, as
shown in Figure 2.17.
2-13
Scientists also rely on various other types of models to help them visualize molecules. Ball-and-stick
models show atoms as spheres and the bonds as sticks, and they give a good sense of the angles at which
the atoms are attached to one another within the molecule. Often, all atoms are represented by balls of
the same size. In other cases, however, the relative sizes of the balls reflect the relative sizes of the atoms.
Sometimes the chemical symbols of the elements are superimposed on the balls, but often the atoms are
identified simply by color.
Space-filling models give a semi-accurate depiction of what the molecule would look like if it were
scaled up in size. These models show the relative sizes of the atoms, but the angles between atoms, which
help define their molecular geometry, are often more difficult to see than in ball-and-stick models. As in
ball-and-stick models, the identities of the atoms are indicated by their colors, but they may also be
labeled with the element's symbol.
You might wonder how chemists determine all of the details of these structures. This is a fascinating
question that we will address later in the course. If you do not care to wait, try doing a little research
yourself. Google is a convenient place to start.
2.7. Ions and Ionic Compounds
The nucleus of an atom is unchanged by ordinary chemical processes, but atoms can readily gain or lose
electrons. If electrons are removed or added to a neutral atom, a charged particle called an ion is formed.
An ion with a positive charge is called a cation (pronounced CAT-ion); a negatively charged ion is called
an anion (AN-ion).
For example, the neutral sodium atom, which has 11 protons and 11 electrons, easily loses one electron.
The resulting cation has 11 protons and 10 electrons, and hence has a net charge of 1+. The net charge on
an ion is represented by a superscript; +, 2+, and 3+ mean a net charge resulting from the loss of one,
two, or three electrons, respectively. The superscripts -, 2-, and 3- represent net charges resulting from
the gain of one, two, or three electrons, respectively. The formation of the Na+ ion from a Na atom is
shown schematically below:
Chlorine, with 17 protons and 17 electrons, gains an electron in chemical reactions, producing the Cl- ion:
<
(In general, metal atoms tend to lose electrons; nonmetal atoms tend to gain electrons)
In addition to simple ions such as Na+ and Cl-, there are polyatomic ions such as NO3" (nitrate ion) and
SO4 2" (sulfate ion). These ions consist of atoms joined as in a molecule, but they have a net positive or
negative charge. We will consider further examples of polyatomic ions in Section 2.8. The chemical
properties of ions are greatly different from those of the atoms from which
! they are derived.
!
2-14
Determining Ionic Charges
Many atoms gain or lose electrons so as to end up with the same number of electrons as the noble gas
closest to them in the periodic table. The members of the noble gas family are chemically very nonreactive and form very few compounds. We might deduce that this is because their electron arrangements
are somehow complete and, therefore, energetically stable. The other elements can obtain these same
stable electron arrangements by losing or gaining electrons. For example, the loss of one electron from
an atom of sodium leaves it with the same number of electrons as the neutral neon atom (atomic number
10). Similarly, when a chlorine atom gains an electron, it ends up with 18, the same as argon (atomic
number 18). We will use this simple rule to predict the formation of ions.
Example: Predict the charges expected for the most stable ions of barium (Ba) and oxygen (O).
From the periodic table, barium has atomic number 56. The nearest noble gas is xenon, atomic number
54. Barium can most readily obtain the stable arrangement of 54 electrons by losing two of its electrons,
forming the Ba 2+ cation. Oxygen has atomic number 8. The nearest noble gas is neon, atomic number
10. Oxygen can most readily obtain this stable electron arrangement by gaining two electrons, thereby
forming an anion of 2- charge, O 2" . In general, metals tend to lose electrons to form positive cations
and
! nonmetals tend to gain electrons to form negative anions.
The periodic table can be used to determine the charges of ions. The table below shows how the charges
!
of these ions relate in a simple way to their positions in the table. On the left side of the table, we see that
the group 1A elements (the alkali metals) form 1+ ions, and the group 2A elements (the alkaline earths)
form 2+ ions. On the other side of the table, the group 7A elements (the halogens) form 1- ions, and the
group 6A elements form 2- ions. The transition metals do not lend themselves to such simple rules and
many can form more than one charge.
Ionic Compounds
A great deal of chemical activity involves the transfer of electrons between substances. Ions form when
one or more electrons transfer from one neutral atom to another. In Figure 2.19 below we see that when
elemental sodium is allowed to react with elemental chlorine, an electron transfers from a neutral sodium
+
-
atom to a neutral chlorine atom. We are left with a Na ion and a Cl ion. But that is not the end of the
+
-
matter. Because objects of opposite charge attract, the Na and Cl ions bind together to form NaCl, an
ionic compound. An ionic compound is a compound that contains positively charged ions and
negatively charged ions.
We can often tell whether a compound is ionic or molecular (consisting of molecules) from its
composition. Ionic compounds are generally combinations of metals and nonmetals, as in NaCl. In
contrast, molecular compounds are generally composed of nonmetals only, as in H 2O , C 6 H 12O6 ,
and NO2 .
!
!
2-15
!
Ionic compounds are also identified by the fact that they do not conduct electricity in solid form, but do
conduct electricity when molten (liquid) or dissolved in water. This behavior is good evidence that ionic
compounds are made of charged ions. When in the molten state or dissolved in water, the charged ions
are mobile and free to respond to an external electric field; in other words, conduct electricity. For this
reason ionic compounds are called electrolytes.
Example: Which of the following compounds would you expect to be ionic: Na 2O , CaCl 2 , SF4 ?
Answer: Na 2O and CaCl 2 are ionic compounds since they are made of a metal and non-metal. SF4 is
a molecular compound since S and F are both non-metals.
!
!
!
The ions in an ionic compound are arranged in an extended three-dimensional structure called a lattice.
! arrangement
! of Na+ and Cl- ions in NaCl is shown in Figure 2.19. Note how each positive
!
The
ion is
surrounded by negative ion nearest neighbors and vise versa. This extended, well-ordered, threedimensional structure distinguishes ionic compounds from molecular compounds. The basic unit of a
molecular compound is a molecule - a self-contained neutral group of atoms linked together in a
particular order (see Figure 2-16). Ionic crystals do not have smallest units per se. They are built up
from these extended three-dimensional arrangements of ions held together by electrostatic forces. How
many positive and negative ions must get together before you can say there is a piece of an ionic solid?
That’s hard to say. It is certainly many times more than a few of each ion. The tiniest crystal just visible
under a microscope would have trillions. On the hand, a single molecule constitutes a piece of a
molecular compound.
Because we cannot identify ionic compounds with molecules, the formulas we use to represent them
communicate only the relative numbers of positive and negative ions. For example, when we represent
magnesium chloride by the formula MgCl 2 we are NOT referring to a molecule containing one Mg and
2 Cl atoms. This formula means that the ration of Mg 2+ ions to Cl " ions 1:2. In other words, formulas
for ionic compounds are empirical formulas.
!
It is a simple matter to write the empirical formula for an ionic compound if we know the charges of the
!
!
ions of which it is composed. Chemical compounds
are always electrically neutral. Consequently, the
ions in an ionic compound always occur in such a ratio that the total positive charge is equal to the total
negative charge. Thus, there is one Na+ to one Cl- giving NaCl, one Ba2+ to two Cl- giving BaCl2, and so
forth. As you consider these and other examples, you will see that if the charges on the cation and anion
are equal, the subscript on each ion will be 1 (and left unwritten). If the charges are not equal, the charge
on one ion (without its sign) will become the subscript on the other ion:
2-16
Example: What are the empirical formulas of the compounds formed by (a) Al 3+ and Cl " ions;
(b) Al 3+ and O 2" ions; (c) Mg 2+ and NO3" ions?
SOLUTION (a) AlCl 3 (b) Al 2O3 (c) Mg( NO3) . Note that in this
! last case,
! the formula for the
2
!
"
!
in parentheses
so that it is clear that the subscript 2 applies to all the atoms of
NO!
!
3 must be enclosed
that ion. Otherwise it might look like there were 32 oxygen atoms! ( MgNO32 )
!
!
!
!
!
2.8. Naming Inorganic Compounds (Chemical Nomenclature)
The astronomical number of compounds that exist makes it essential that a system of chemical
nomenclature be used. The International Union of Pure and Applied Chemistry (IUPAC) has
developed such a system for assigning names and formulas to compounds. In this section you will learn
this system as it applies to inorganic compounds.
A. Naming Methods for Monatomic Ions
1. Cations of representative metal elements having just one charge (or oxidation) state:
element name ion
Example: Mg 2+ is named the "magnesium ion."
2. Anions of representative non metal elements having just one charge state:
element root- ide ion
!
Cl " is named the "chloride ion."
Example:
P 3" is named the "phosphide" ion
!
3. Cations with more than one possible charge states:
!
!
Several transition metals can form cations of different charge. Examples include copper ( Cu + and
Cu 2+ ) and iron ( Fe 2+ and Fe 3+ ). There are two methods for naming ions that exist in more than
one charge state:
!
• The
A Roman numeral is put in parentheses after the element name to indicate
! Stock System:
!
the particular value of the cation charge.
Examples: Cu + = copper (I) ion (read "copper one ion")
Cu 2+ = copper (II) ion (read "copper two ion")
!
•
Fe 2+ = iron (II) ion (read "iron two ion")
Fe 3+ = iron (III) ion (read "iron three ion")
!
!
Classical system: The suffixes -ous and -ic are added to the element root name. The -ous is
!
used for the lower possible charge state while the -ic is used for the higher charge.
Examples: Cu + = cuprous ion; Cu 2+ = cupric ion
Fe 2+ = ferrous ion; Fe 3+ = ferric ion
In!this course we will!
stick to the modern Stock system, that is, we will use the Roman
!
!
numerals.
2-17
B. Naming Polyatomic Ions
Polyatomic ions are tightly bound groups of atoms that behave as a unit and carry a charge. They are
"charged molecules." Common examples include the nitrate ion ( NO3" ), the sulfate ion ( SO4 2" ), and
the phosphate ion ion ( PO4 3" ). The only common polyatomic cation is the ammonium ion ( NH 4 + ).
The names of common polyatomic anions are given Table 2.5 at the end of the chapter.
!
!
C. Names and!Formulas of Binary Ionic Compounds
!
A binary ionic compound is made up of a monatomic cation and a monatomic anion.
1. Writing Formulas of Binary Ionic Compounds
Consider the ionic compound formed between Mg and Cl.
•
Mg atoms lose two electrons to become cations with a +2 charge. This is written as Mg 2+
Cl atoms gain 1 electron each to become anions with a -1 charge. This is written as Cl "
Because ionic compounds are electrically neutral, there must be two Cl " for each Mg 2+ in the compound.
Therefore the formula for this compound is MgCl 2 .
!
x+
! A "y is given by
In general, the ionic compound formed from the metal ion M and a nonmetal ion
!
!
•
•
•
!
M y Ax
!
Example: Sodium oxide: The ions are Na + , O 2" " Na 2O
!
Example: Iron (III) oxide: The ions are Fe 3+ and O 2" " Fe2O3
!
!
!
!
Example: Magnesium nitride: The ions are Mg 2+ and N 3" " Mg 3N 2
!
!
!
2. Naming Binary Ionic Compounds
!
!
!
Binary ionic compounds are named by writing the cation name followed by the anion name (-ide ending).
Example:
NaCl , sodium chloride
AlN , aluminum nitride
Na 3P , sodium phosphide
When a cation has more than one possible charge state the particular cation charge that it exists in must be
!
indicated by a Roman numeral in parentheses or by the proper cation suffix (-ous or -ic)
!
Example:
CuO, copper (II) oxide or cupric oxide
Cu 2O , copper (I) oxide or cuprous oxide
The anion charge along with the condition of electrical neutrality determines the cation charge.
!
Example:
SnO2 , must be tin (IV) oxide since the oxygen anion always has a -2 charge.
Roman numerals are used on an "as needed" basis only. That is, tehy are only used when the cation is one
that can have more than one charge.
!
2-18
D. Names and Formulas of Ionic Compounds Containing Polyatomic Ions
1. Formulas of Ionic Compounds with Polyatomic Ions
The procedure for writing the formulas of ionic compounds having polyatomic ions is similar to
that used for binary ionic compounds.
Example: Calcium nitrate: Ca 2+ , NO3"1 " Ca( NO3) 2
Example: Ammonium phosphate: NH 4 + , PO4 3"
#
(NH ) PO
4
+
3
4
!
! be placed
! around the polyatomic ion when there is more than one of
Note that parentheses must
them in the formula. This is necessary to make the subscripts clear. Parentheses are never
placed around monatomic!ions.
2. Naming Ionic Compounds Having Polyatomic Ions
Simply name the ions, cation first and anion second.
Example: NaC 2 H 3O2 , sodium acetate
Example: Cu(C 2 H 3O2 ) 2 , copper (II) acetate
! the necessity of using Roman numeral after any cation that is one of those that can have
Note
more than one charge.
!
E. Binary Molecular Compounds
Binary molecular compounds are composed of two nonmetallic elements. Prefixes are used to
describe the number of each atom in molecules.
1. Naming Binary Molecular Compounds
Since molecular compounds have nothing to do with ionic charges, the formulas cannot be predicted
and prefixes are needed to show how many atoms of a given element are in a given molecule.
Prefix
monoditritetra-
Number
1
2
3
4
Prefix
pentahexaheptaocta-
Number
5
6
7
8
Prefix
nonadeca-
Number
9
10
To name a binary molecular compound
1. Recognize the compound as being binary molecular. It must contain no metal atoms
2. Add prefixes to each element to indicate the number of atoms of each element.
3. The second element is written with an -ide ending
4. The "o" at the end of the prefix mono- is dropped when the element name begins with a vowel.
5. The mono- prefix is omitted entirely when there is a single atom of the first element.
2-19
Examples:
N 2O :
dinitrogen monoxide (note that the "o" is dropped from mono- before "oxide."
PCl 3 : phosphorous trichloride
!
!
!
!
AlCl 3
aluminum chloride (WATCH OUT, this is an ionic compound so prefixes are not used!)
CO2 :
carbon dioxide
CO:
carbon monoxide (Note: drop the “o” on mono when the element name starts with a
vowel. You don’t want to write monoxide)
!
Two important facts should be kept in mind when dealing with molecular compounds
•
They are not made of ions so ionic charges cannot be used in determine or imply the formula.
•
Nonmetallic elements can often combine in more than one way. For example, there are many
molecules that have C and H each having different numbers of each element. That is why
prefixes are needed
2. Writing Formulas of Binary Molecular Compounds.
The prefixes give the numbers of atoms and therefore the subscripts to use.
Write the element symbols with the numbers of atoms of each as subscripts (1 is understood)
Examples:
Carbon disulfide: CS2
Dichlorine pentoxide: Cl 2O5
F. Nomenclature of Acids
!
!
Acids are molecular compounds that contain a hydrogen atom that breaks away from the rest of its
molecule when it dissolves in water. When not dissolved in water they are neutral molecules. When
dissolved, they break up into an H + ion and a negative anion. For this reason, when dissolved in
water, acids can be thought of as ionic compounds with a hydrogen ion H+ playing the role of the
metal cation.
Examples:
!
In its pure gaseous state, hydrogen chloride is a molecular compound and is symbolically
represented as HCl(g). When it dissolved in water hydrogen chloride dissociates into H + (aq) and
Cl " (aq) ions. In this dissolved state HCl(aq) is called hydrochloric acid.
!
!
In its pure gaseous state H 2 S (g) is called hydrogen sulfide. When dissolved in water H 2 S (aq) is
called hydrosulfuric acid
! that a compound is dissolved in water the symbol (aq) is placed
!
Note that to indicate
after the
compound symbol. This stands for aqueous solution (dissolved in water).
2-20
1. Naming Acids
The general formula for an acid is:
H a Ab (aq)
where H is hydrogen, A is an anion, and the symbol (aq) indicates that it is dissolved in water.
! acids.
The following are the rules for naming
1. When the anion is derived from a single atom, write:
Examples:
HBr (aq): Br " = bromide ion
HI (aq): I " = iodide ion
!
H 2 S (aq) :
!
"
!
S 2" = sulfide ion
!
"
!
!
hydrobromic acid
hydroiodic acid
hydrosulfic acid
"
2. When the anion
-ite , write:
! name ends in !
!
Examples:
H 2 SO3 (aq) : SO32" = sulfite ion "
hydro-anion stem-ic acid
anion stem-ous acid
sulfurous acid
HClO(aq) : ClO" = hypochlorite ion "
!
!
hypochlorous acid
3. When the anion ends in -ate, write anion stem-ic acid
!
!
Examples:
HNO3 (aq) : NO3" = nitrate ion "
!
nitric acid
H 3PO4 (aq) : PO43" = phosphate ion " phosphoric acid
!
!
!2. Writing Acid
! Formulas: Use the!above rules in reverse
2-21
G. Nomenclature of Hydrates
Many ionic compounds incorporate a specific number of water molecules per formula unit into their
crystal structures. These water molecules form part of the crystal and do not make the salt appear
“wet.” Ionic compounds that contain water in this manner are called hydrates and the water that is
incorporated into the crystal is called water of hydration.
The water of hydration can usually be removed by heating at temperatures near 100 ˚C . An ionic
compound that has had its water of hydration removed is called an anhydrous salt.
The formula of a hydrate consists of the formula of the ionic compound followed
by ". x H 2O " where
!
x is the number of water molecules attached to each formula unit of ionic compound.
The name of the hydrate is given by the ionic compound name followed by prefix-hydrate where the
!
prefix is that corresponding to the x.
Examples: Copper (II) sulfate pentahydrate: CuSO4 " 5H 2O
!
The images above show the hydrate copper (II) sulfate on the left. In the center is the anhydrous
copper (II) sulfate attained after gentle heating. The right photograph shows a drop of water added
to the anhydrous copper (II) sulfate which results in the return to the hydrated form.
The “dehyradtion” of copper (II) sulfate pentahydrate may be written as a chemical equation:
CuSO4 " 5H 2O(s)
heat
#
(hydrate)
+
5H 2O(l)
(anhydrous salt)
CuSO4 (s)
(water)
!
Another example of a hydrated salt is barium chloride dihydrate, BaCl 2. " 2H 2O .
!
2-22
Table of Common Ions
Cations
Name
Symbol
Aluminum
Al 3+
Ammonium
Barium
Anions
Hydrated
color
-
Name
Symbol
Acetate
CH 3COO"
NH 4 +
-
Arsenide
As 3"
-
Ba 2+
Cd 2+
-
Bromide
yellow
Calcium !
Chromium!(II)
Ca 2+
Cr 2+
-
Carbide
!
Br "
BrO3"
-
Cadmium!
Carbonate
!
Chromium!(III)
C 4"
CO32"
-
varies
Cr 3+
blue
Hydrogen carbonate
!
HCO3"
-
Bromate!
-
Cobalt (II)!
!
Cobalt (III)
Co 2+
Co 3+
light red
Chlorate† !
Copper (I)!
Cl "
ClO3"
-
-
Cu +
-
Chlorite† !
-
!
Copper (II)
Cu 2+
H+
pale blue
Hypochlorite† !
ClO2"
ClO"
-
-
Perchlorate† !
ClO4 "
-
Fe 2+
pale green
Chromate !
CrO4 2"
yellow
Fe 3+
brown
Dichromate !
Cr2O7 2"
orange
Pb 2+
Pb 4+
Li +
Mg 2+
-
CN "
SCN "
-
F"
H"
-
Mn 2+
Mn 3+
Hg 2 2+
light pink
-
-
Nitrate
OH "
N 3"
NO3"
Hg 2+
-
Nitrite
NO2"
-
Ni 2+
green
Oxalate
!
!
-
Oxide
!
C 2O4 2"
purple
Hydrogen!
Iron (II) !
Iron (III) !
!
Lead (II)
!
Lead (IV)
Lithium !
Magnesium
!
Manganese
! (II)
Manganese
! (III)
Mercury (I)
!
Mercury (II)
!
Nickel (II)!
!
Nickel (III)
!
Hydrated
Color
-
Chloride†
Cyanide
-
!
Thiocyanate!
-
Fluoride
!
!
Hydroxide
!
-
Hydride
-
Nitride
!
!
-
-
-
Ni 3+
K+
-
Permanganate!
Sodium !
Tin (II) !
Tin (IV) !
O 2"
MnO4 "
Na +
-
Peroxide !
O2 2"
-
Sn 2+
Sn 4+
-
Phosphide
Phosphate !
Strontium!
P 3"
PO4 3"
-
-
Sr 2+
-
Phosphite !
-
!
!
PO33"
Ag +
-
Selenide
Se 2"
-
-
Sulfide
2"
S
SO4 2"
-
HSO4 "
-
SO32"
-
HSO3"
-
S2O32"
-
Potassium!
Silver
Zinc
!
Zn
2+
!
!
!
Sulfate
!
Hydrogen sulfate!
!
!
!
Sulfite
† Ions with bromine and iodine in place of chlorine
are named similarly
!
Hydrogen sulfite
Thiosulfate !
!
!
!
2-23
-
-
Chapter 2. Homework
Section 2-1.
2.1. What is the main goal of chemistry?
2.2. Why are the chemical theories of both the ancient
Greeks and the medieval alchemists not considered
scientific today?
2.3. (a) What is the general meaning of the term
operational definition? (b) In what sense was Boyle’s
definition of an element an “operational definition.”
2.4. What basic laws of matter did Dalton’s Theory
provide explanations for?
2.5. If we assume that nitrous oxide molecules are
comprised of 2 atoms of nitrogen (N) and 1 atom of
oxygen (O), use the Law of Multiple Proportions to
determine the number ratio of nitrogen to oxygen
atoms in nitric oxide and nitrogen dioxide.
2.6. The mass ratio of nitrogen to hydrogen in
ammonia is 4.67. Taking the atomic mass of hydrogen
to be 1 atomic mass unit (amu), determine the atomic
mass of nitrogen if the formula of ammonia is
assumed to be (a) NH (1 atom N: 1 atom H), (b) NH2
(1 atoms N: 2 atoms H), (c) NH3 (1 atom N to 3 atoms
H). Use the general expression found in the reading
relating bulk mass ratio to atomic mass ratio to derive
your results.
2.12. What is similar between a cathode ray and a
beta particle? What is the difference?
2.13. Imagine that you begin dropping pennies into a
deep circular well having a diameter of 1-meter, You
release each penny from random locations about the
opening. For each of the first 78 drops you note that
the pennies make no sound until they “plunk” in the
water 3.0 seconds after release. Then, on the 79th
penny, you hear a distinct metallic “clink” 2.0 seconds
after release. After 10,000 additional random drops
you find that the randomly dropped pennies produce
the metallic “clink” once every 100 drops, on average.
It would be easy to conclude from this experiment that
there is an obstruction part way down the well. (a)
Estimate the obstruction’s cross sectional area from
the data described. (b) Calculate in meters the depth
of the well and the depth of the obstruction.
Section 2-3.
2.14. The ratio of the diameter of a typical atomic
nucleus to the diameter of a typical atom is 1/10,000.
(a) What is the approximate ratio of the volume of the
nucleus to the volume of the atom? (b) If the diameter
of the nucleus were 1 inch, what would be the
diameter of the atom in feet?
2.15. Complete the table at the end of Section 2-3
(Reproduce it on your own worksheet to hand in)
Section 2-4.
2.7. The mass ratio of carbon to oxygen in carbon
dioxide ( CO2 ) is 3:8. Predict the mass ratios would
be for the compounds (a) CO, (b) CO3 and (c) C 2O .
2.8. Use Google to find an example of a
!nonstoichiometric compound.
!
!
Section 2-2.
2.9. State two differences in the behavior of cathode
rays (electrons) and X-rays (high energy light), both of
which were generated in a Crooke’s tube.
2.10. What experimental evidence supported
Thomson’s hypothesis that electrons are subatomic
particles (i.e., universal parts or building blocks of
atoms)
2.11. What piece of data did J.J. Thomson need in
order to calculate the mass of the electron from the
charge to mass ratio that he calculated? Who provided
this information?
2-16. Is there an error in the direction that the positive
Cl ion’s deflection as shown in the detailed diagram of
the mass-spectrometer in the text? (Hint: Compare the
mass spectrometer diagram with the preceding
diagram showing the path of + and – charged particles
across the face of a north pole of a magnet (pointing
out of the page).
2-17. Complete the isotope table in Section 12-4
(Reproduce it on your answer sheet to hand in).
in the form 1H "1O"1H and make sure you don’t
duplicate.
!
Section 2-5 and 2-6
(Reproduce this table on your answer sheets)
Section 2-7
2.47. Fill in the gaps of the following table
(Hint: This requires taking a weighted average)
Symbol
protons
neutrons
electrons
!
Ion charge
59
Co 3+
34
46
36
76
116
80
120
78
2+
(Reproduce this table on your answer sheets)
2.34. How many peaks will be found in a mass
spectrograph of water molecules assuming that
there are detectable amounts of all three isotopes
of hydrogen and detectable amounts of two
isotopes of oxygen. Hint: There are more than
you may at first think! Write out each possibility
(Note: In each case the positive charge on the
metal is determined by the group (or column)
number (e.g. Na = +1, Al = +3, Cl = 7 - 8 = -1)
Section 2-8 Questions
Note: Take the time needed to get them right the first time! Refer explicitly to the nomenclature rules in the text. Be
particularly alert to those cations that require a roman numeral in the naming of the compounds they form, such as iron (III)
oxide and nickel (II) chloride. You will find these in the table of ions at the end of Chapter 2.
ON YOUR PAPER WRITE OUT YOUR ANSWERS IN THE SAME ARRANGEMENT (COLUMN AND ROWS) AS
THE QUESTIONS. THINK ABOUT AND INCLUDE THE ONES FOR WHICH THE ANSWERS ARE GIVEN.
2-61. State differences in the types of elements that combine to form ionic vs. molecular compounds.
2-62. Would you expect the following compounds to be ionic or molecular? (Think about metals and non metals)
(a) CO
(c) Mn 3 ( PO4 ) 2
(b) KBr
(d) C 3H 8
2-63. Using only the periodic table, write the formula of the ions formed from these representative elements. Include
the charge as a superscript.
!
!
!
!
(a) lithium ( Li + )
(b) fluorine ( F " )
(c) oxygen
(d) potassium
(e) barium
(f) nitrogen ( N 3" )
(g) beryllium
(h) magnesium
!
(i) aluminum ( Al 3+ )
!
(j) chlorine
(k) calcium ( Ca 2+ )
!
2-64. Write the formulas for compounds formed from these pairs of ions.
!
(a) Sr 2+ , Se 2" ( SrSe) (b) K + , O 2"
!
(c) Ca 2+ , N 3"
(l) sulfur
(d) Co 3+ , I "
2-65.
Write
for these
compounds (Note the
!
! formulas
! !
! names
! that need Roman!numerals.
! Why is that?)
!
(a) tin (IV) bromide
(b) ammonium dichromate
(d) sodium hydrogen sulfate (e) sodium bicarbonate
(c) lithium hydrogen sulfate
(f) copper (II) carbonate
2-66. Name the following binary molecular compounds
(a) OF2
(b) Cl 2O8
(c) SO3
(d P4 O10
(e) CCl 4
2-67. Write formulas for the following binary molecular compounds
!
(a) dinitrogen!tetroxide
(c) nitrogen trifluoride
!
!
!
(b) phosphorous
pentachloride
(d) disulfur dichloride
2-68. Name each of the following acids (Be sure you understand the restricted use of “hydro-“ in naming an
acid.)
(a) H 2C 2O4 (aq)
(b) HF (aq)
(c) H 2 SO4 (aq)
(d) HClO2 (aq)
(e) H 2CO3 (aq)
(f) HCl (aq)
(g) HNO3 (aq)
(h) HC 2 H 3O2 (aq)
!
!
!
!
!
!
2-69. Name each of the following substances. Be careful to distinguish molecular compounds, ionic compounds, and acids.
Acids are indicated by (aq).
(a) CaO
(b) CuC 2 H 3O2
(c) Ba 3( PO4 ) 2
(d) HClO4 (aq)
(e) I 2
(f) Cl 2O
(g) BaSO4
(h) HgF2
(i) Mg(OH ) 2
!
!
(m) Li 2 HPO4
(q) SnCl 4
!
!
!
(j) ( NH 4 ) 2 C 2O4
(n) FeCO3
(r) HgS
!
!
!
(k) NO2
(o) H 2CrO4 (aq)
(s) N 2O5
!
!
!
(l) Ni 3 ( PO4 ) 2
(p) H 2O2
(t) CS2
!
!
!
!
2-70. Write the chemical formula for each of the following substances. Keep in mind that these include both
! ionic, molecular compounds, and acids. Be alert to
! those metal ions that can!form ions of two different
positive charges and require roman numerals to name (example: copper (I) and copper (II)). Finally,
remember to that (aq) must be places after the formula to designate an acid.
(a) calcium carbonate
(d) barium hydroxide
(b) nitrogen gas
(e) iron (III) sulfate
(c) sodium bromide
(f) hydrobromic acid
(g) magnesium sulfide
(j) sulfite ion
(h) sulfuric acid
(k) sulfur trioxide
(i) potassium permanganate
(l) ammonium perchlorate
(m) phosphorous pentabromide
(n) copper (II) iodide
(o) phosphoric acid