Atoms, Molecules,and Ions41A Look Ahead We begin with a
historical perspective of the search for the fundamentalunits of
matter. The modern version of atomic theory was laid by JohnDalton
in the nineteenth century, who postulated that elements are
composedof extremely small particles, called atoms. All atoms of a
given element areidentical, but they are different from atoms of
all other elements. (2.1) We note that, through experimentation,
scientists have learned that an atomis composed of three elementary
particles: proton, electron, and neutron. Theproton has a positive
charge, the electron has a negative charge, and theneutron has no
charge. Protons and neutrons are located in a small regionat the
center of the atom, called the nucleus, while electrons are spread
outabout the nucleus at some distance from it. (2.2) We will learn
the following ways to identify atoms. Atomic number is thenumber of
protons in a nucleus; atoms of different elements have
differentatomic numbers. Isotopes are atoms of the same element
having a differentnumber of neutrons. Mass number is the sum of the
number of protons andneutrons in an atom. Because an atom is
electrically neutral, the number ofprotons is equal to the number
of electrons in it. (2.3) Next we will see how elements can be
grouped together according to theirchemical and physical properties
in a chart called the periodic table. Theperiodic table enables us
to classify elements (as metals, metalloids, andnonmetals) and
correlate their properties in a systematic way. (2.4) We will see
that atoms of most elements interact to form compounds, whichare
classifi ed as molecules or ionic compounds made of positive
(cations)and negative (anions) ions. (2.5) We learn to use chemical
formulas (molecular and empirical) to representmolecules and ionic
compounds and models to represent molecules. (2.6) We learn a set
of rules that help us name the inorganic compounds. (2.7) Finally,
we will briefl y explore the organic world to which we will
returnin a later chapter. (2.8)Since ancient times humans have
pondered the nature of matter. Our modernideas of the structure of
matter began to take shape in the early nineteenthcentury with
Daltons atomic theory. We now know that all matter is made ofatoms,
molecules, and ions. All of chemistry is concerned in one way or
anotherwith these species.2.1 The Atomic Theory2.2 The Structure of
the Atom2.3 Atomic Number, MassNumber, and Isotopes2.4 The Periodic
Table2.5 Molecules and Ions2.6 Chemical Formulas2.7 Naming
Compounds2.8 Introduction to OrganicCompoundsChapter OutlineStudent
InteractiveActivityAnimationsCathode Ray Tube (2.2)Millikan Oil
Drop (2.2)Alpha, Beta, and GammaRays (2.2)a-Particle Scattering
(2.2)Media PlayerRutherfords Experiment (2.2)Formation of an
IonicCompound (2.7)Chapter SummaryARISExample Practice ProblemsEnd
of Chapter ProblemsQuantum TutorsEnd of Chapter Problems42 Atoms,
Molecules, and Ions2.1 The Atomic TheoryIn the fi fth century b.c.
the Greek philosopher Democritus expressed the belief thatall
matter consists of very small, indivisible particles, which he
named atomos (meaninguncuttable or indivisible). Although
Democritus idea was not accepted by manyof his contemporaries
(notably Plato and Aristotle), somehow it endured.
Experimentalevidence from early scientifi c investigations provided
support for the notion ofatomism and gradually gave rise to the
modern defi nitions of elements and compounds.In 1808 an English
scientist and school teacher, John Dalton, formulated aprecise defi
nition of the indivisible building blocks of matter that we call
atoms.Daltons work marked the beginning of the modern era of
chemistry. The hypothesesabout the nature of matter on which
Daltons atomic theory is based can besummarized as follows:1.
Elements are composed of extremely small particles called atoms.2.
All atoms of a given element are identical, having the same size,
mass, andchemical properties. The atoms of one element are
different from the atoms ofall other elements.3. Compounds are
composed of atoms of more than one element. In any compound,the
ratio of the numbers of atoms of any two of the elements present is
either aninteger or a simple fraction.4. A chemical reaction
involves only the separation, combination, or rearrangementof
atoms; it does not result in their creation or destruction.Figure
2.1 is a schematic representation of the last three
hypotheses.Daltons concept of an atom was far more detailed and
specifi c than Democritus.The second hypothesis states that atoms
of one element are different from atoms ofall other elements.
Dalton made no attempt to describe the structure or compositionof
atomshe had no idea what an atom is really like. But he did realize
that thedifferent properties shown by elements such as hydrogen and
oxygen can be explainedby assuming that hydrogen atoms are not the
same as oxygen atoms.The third hypothesis suggests that, to form a
certain compound, we need not onlyatoms of the right kinds of
elements, but specifi c numbers of these atoms as well.John Dalton
(17661844). English chemist, mathematician, and philosopher. In
addition to the atomictheory, he also formulated several gas laws
and gave the fi rst detailed description of color blindness,
fromwhich he suffered. Dalton was described as an indifferent
experimenter, and singularly wanting in thelanguage and power of
illustration. His only recreation was lawn bowling on Thursday
afternoons. Perhapsit was the sight of those wooden balls that
provided him with the idea of the atomic theory.(b)Atoms of element
X Atoms of element Y Compounds of elements X and Y(a)Figure 2.1 (a)
According toDaltons atomic theory, atoms ofthe same element are
identical,but atoms of one element aredifferent from atoms of
otherelements. (b) Compound formedfrom atoms of elements X and Y.In
this case, the ratio of theatoms of element X to the atomsof
element Y is 2:1. Note that achemical reaction results only inthe
rearrangement of atoms, notin their destruction or creation.This
idea is an extension of a law published in 1799 by Joseph Proust, a
Frenchchemist. Prousts law of defi nite proportions states that
different samples of the samecompound always contain its
constituent elements in the same proportion by mass.Thus, if we
were to analyze samples of carbon dioxide gas obtained from
differentsources, we would fi nd in each sample the same ratio by
mass of carbon to oxygen.It stands to reason, then, that if the
ratio of the masses of different elements in a givencompound is fi
xed, the ratio of the atoms of these elements in the compound
alsomust be constant.Daltons third hypothesis supports another
important law, the law of multipleproportions . According to the
law, if two elements can combine to form more thanone compound, the
masses of one element that combine with a fi xed mass of the
otherelement are in ratios of small whole numbers. Daltons theory
explains the law ofmultiple proportions quite simply: Different
compounds made up of the same elementsdiffer in the number of atoms
of each kind that combine. For example, carbon formstwo stable
compounds with oxygen, namely, carbon monoxide and carbon
dioxide.Modern measurement techniques indicate that one atom of
carbon combines with oneatom of oxygen in carbon monoxide and with
two atoms of oxygen in carbon dioxide.Thus, the ratio of oxygen in
carbon monoxide to oxygen in carbon dioxide is 1:2.This result is
consistent with the law of multiple proportions ( Figure 2.2
).Daltons fourth hypothesis is another way of stating the law of
conservation ofmass, which is that matter can be neither created
nor destroyed. Because matter ismade of atoms that are unchanged in
a chemical reaction, it follows that mass mustbe conserved as well.
Daltons brilliant insight into the nature of matter was the
mainstimulus for the rapid progress of chemistry during the
nineteenth century.Review of ConceptsThe atoms of elements A (blue)
and B (orange) form two compounds shownhere. Do these compounds
obey the law of multiple proportions?Joseph Louis Proust
(17541826). French chemist. Proust was the fi rst person to isolate
sugar from grapes.According to Albert Einstein, mass and energy are
alternate aspects of a single entity called mass-energy.Chemical
reactions usually involve a gain or loss of heat and other forms of
energy. Thus, when energyis lost in a reaction, for example, mass
is also lost. Except for nuclear reactions (see Chapter 23),
however,changes of mass in chemical reactions are too small to
detect. Therefore, for all practical purposes massis
conserved.Carbon monoxideCarbon dioxideRatio of oxygen incarbon
monoxide tooxygen in carbon dioxide: 1:2OC21OC11_ __ _Figure 2.2 An
illustration of thelaw of multiple proportions.2.2 The Structure of
the AtomOn the basis of Daltons atomic theory, we can defi ne an
atom as the basic unit of anelement that can enter into chemical
combination. Dalton imagined an atom that wasboth extremely small
and indivisible. However, a series of investigations that began
inthe 1850s and extended into the twentieth century clearly
demonstrated that atomsactually possess internal structure; that
is, they are made up of even smaller particles,which are called
subatomic particles. This research led to the discovery of three
suchparticleselectrons, protons, and neutrons.2.2 The Structure of
the Atom 4344 Atoms, Molecules, and IonsThe ElectronIn the 1890s,
many scientists became caught up in the study of radiation , the
emissionand transmission of energy through space in the form of
waves. Information gainedfrom this research contributed greatly to
our understanding of atomic structure. Onedevice used to
investigate this phenomenon was a cathode ray tube, the forerunner
ofthe television tube ( Figure 2.3 ). It is a glass tube from which
most of the air has beenevacuated. When the two metal plates are
connected to a high-voltage source, thenegatively charged plate,
called the cathode, emits an invisible ray. The cathode rayis drawn
to the positively charged plate, called the anode, where it passes
through ahole and continues traveling to the other end of the tube.
When the ray strikes thespecially coated surface, it produces a
strong fl uorescence, or bright light.In some experiments, two
electrically charged plates and a magnet were added tothe outside
of the cathode ray tube (see Figure 2.3 ). When the magnetic fi eld
is on andthe electric fi eld is off, the cathode ray strikes point
A. When only the electric fi eld ison, the ray strikes point C.
When both the magnetic and the electric fi elds are off orwhen they
are both on but balanced so that they cancel each others infl
uence, the raystrikes point B. According to electromagnetic theory,
a moving charged body behaveslike a magnet and can interact with
electric and magnetic fi elds through which it passes.Because the
cathode ray is attracted by the plate bearing positive charges and
repelledby the plate bearing negative charges, it must consist of
negatively charged particles.We know these negatively charged
particles as electrons . Figure 2.4 shows the effectof a bar magnet
on the cathode ray.An English physicist, J. J. Thomson, used a
cathode ray tube and his knowledgeof electromagnetic theory to
determine the ratio of electric charge to the mass of anindividual
electron. The number he came up with was 21.76 3 108 C/g, where
Cstands for coulomb, which is the unit of electric charge.
Thereafter, in a series ofexperiments carried out between 1908 and
1917, R. A. Millikan succeeded in measuringthe charge of the
electron with great precision. His work proved that the chargeon
each electron was exactly the same. In his experiment, Millikan
examined themotion of single tiny drops of oil that picked up
static charge from ions in the air.He suspended the charged drops
in air by applying an electric fi eld and followed
theirAnimationCathode Ray TubeHigh voltage+Anode
CathodeABCSNFluorescent screenFigure 2.3 A cathode ray tubewith an
electric fi eld perpendicularto the direction of the cathoderays
and an external magneticfi eld. The symbols N and Sdenote the north
and south polesof the magnet. The cathode rayswill strike the end
of the tube atA in the presence of a magneticfi eld, at C in the
presence of anelectric fi eld, and at B when thereare no external
fi elds present orwhen the effects of the electricfi eld and
magnetic fi eld canceleach other.Electrons are normally associated
withatoms. However, they can also be studiedindividually.Joseph
John Thomson (18561940). British physicist who received the Nobel
Prize in Physics in 1906for discovering the electron.Robert Andrews
Millikan (18681953). American physicist who was awarded the Nobel
Prize in Physicsin 1923 for determining the charge of the
electron.AnimationMillikan Oil Dropmotions through a microscope (
Figure 2.5 ). Using his knowledge of electrostatics,Millikan found
the charge of an electron to be 21.6022 3 10219 C. From these
datahe calculated the mass of an electron:mass of an electron
5chargecharge/mass521.6022 3 10219 C21.76 3 108 C/g5 9.10 3 10228
gThis is an exceedingly small mass.RadioactivityIn 1895, the German
physicist Wilhelm Rntgen noticed that cathode rays causedglass and
metals to emit very unusual rays. This highly energetic radiation
penetratedmatter, darkened covered photographic plates, and caused
a variety of substances tofl uoresce. Because these rays could not
be defl ected by a magnet, they could notcontain charged particles
as cathode rays do. Rntgen called them X rays becausetheir nature
was not known.(a) (b) (c)Figure 2.4 (a) A cathode ray produced in a
discharge tube. The ray itself is invisible, but the fl uorescence
of a zinc sulfi de coatingon the glass causes it to appear green.
(b) The cathode ray is bent downward when a bar magnet is brought
toward it. (c) When thepolarity of the magnet is reversed, the ray
bends in the opposite direction.AtomizerViewingmicroscopeCharged
plateX rayto producecharge onoil droplet(_)Smallhole(_)Charged
plate Oil droplets Figure 2.5 Schematic diagramof Millikans oil
drop experiment.2.2 The Structure of the Atom 45Wilhelm Konrad
Rntgen (18451923). German physicist who received the Nobel Prize in
Physics in 1901for the discovery of X rays.46 Atoms, Molecules, and
IonsNot long after Rntgens discovery, Antoine Becquerel, a
professor of physicsin Paris, began to study the fl uorescent
properties of substances. Purely by accident,he found that exposing
thickly wrapped photographic plates to a certain uraniumcompound
caused them to darken, even without the stimulation of cathode
rays. LikeX rays, the rays from the uranium compound were highly
energetic and could not bedefl ected by a magnet, but they differed
from X rays because they arose spontaneously.One of Becquerels
students, Marie Curie, suggested the name radioactivity todescribe
this spontaneous emission of particles and/or radiation. Since
then, any elementthat spontaneously emits radiation is said to be
radioactive.Three types of rays are produced by the decay, or
breakdown, of radioactivesubstances such as uranium. Two of the
three are defl ected by oppositely chargedmetal plates ( Figure 2.6
). Alpha (a) rays consist of positively charged particles, calleda
particles, and therefore are defl ected by the positively charged
plate. Beta (b) rays,or b particles, are electrons and are defl
ected by the negatively charged plate. Thethird type of radioactive
radiation consists of high-energy rays called gamma (g)rays . Like
X rays, g rays have no charge and are not affected by an external
fi eld.The Proton and the NucleusBy the early 1900s, two features
of atoms had become clear: they contain electrons,and they are
electrically neutral. To maintain electric neutrality, an atom must
containan equal number of positive and negative charges. Therefore,
Thomson proposed thatan atom could be thought of as a uniform,
positive sphere of matter in which electronsare embedded like
raisins in a cake ( Figure 2.7 ). This so-called plum-pudding
modelwas the accepted theory for a number of years.Antoine Henri
Becquerel (18521908). French physicist who was awarded the Nobel
Prize in Physics in1903 for discovering radioactivity in
uranium.Marie (Marya Sklodowska) Curie (18671934). Polish-born
chemist and physicist. In 1903 she and herFrench husband, Pierre
Curie, were awarded the Nobel Prize in Physics for their work on
radioactivity. In1911, she again received the Nobel prize, this
time in chemistry, for her work on the radioactive elementsradium
and polonium. She is one of only three people to have received two
Nobel prizes in science. Despiteher great contribution to science,
her nomination to the French Academy of Sciences in 1911 was
rejectedby one vote because she was a woman! Her daughter Irene,
and son-in-law Frederic Joliot-Curie, sharedthe Nobel Prize in
Chemistry in 1935.+Radioactive substanceLead blockFigure 2.6 Three
types of raysemitted by radioactive elements.b rays consist of
negativelycharged particles (electrons) andare therefore attracted
by thepositively charged plate. Theopposite holds true for a
raysthey are positively charged andare drawn to the
negativelycharged plate. Because g rayshave no charges, their path
isunaffected by an externalelectric fi eld.AnimationAlpha, Beta,
and Gamma RaysPositive charge spreadover the entire sphereFigure
2.7 Thomsons model ofthe atom, sometimes describedas the
plum-pudding model,after a traditional English dessertcontaining
raisins. The electronsare embedded in a uniform,positively charged
sphere.In 1910 the New Zealand physicist Ernest Rutherford, who had
studied withThomson at Cambridge University, decided to use a
particles to probe the structure ofatoms. Together with his
associate Hans Geiger and an undergraduate named ErnestMarsden,
Rutherford carried out a series of experiments using very thin
foils of goldand other metals as targets for a particles from a
radioactive source ( Figure 2.8 ). Theyobserved that the majority
of particles penetrated the foil either undefl ected or with onlya
slight defl ection. But every now and then an a particle was
scattered (or defl ected) ata large angle. In some instances, an a
particle actually bounced back in the directionfrom which it had
come! This was a most surprising fi nding, for in Thomsons modelthe
positive charge of the atom was so diffuse that the positive a
particles should havepassed through the foil with very little defl
ection. To quote Rutherfords initial reactionwhen told of this
discovery: It was as incredible as if you had fi red a 15-inch
shell ata piece of tissue paper and it came back and hit
you.Rutherford was later able to explain the results of the
a-scattering experiment interms of a new model for the atom.
According to Rutherford, most of the atom mustbe empty space. This
explains why the majority of a particles passed through the
goldfoil with little or no defl ection. The atoms positive charges,
Rutherford proposed, areall concentrated in the nucleus, which is a
dense central core within the atom. Wheneveran a particle came
close to a nucleus in the scattering experiment, it experienced a
largerepulsive force and therefore a large defl ection. Moreover,
an a particle traveling directlytoward a nucleus would be
completely repelled and its direction would be reversed.The
positively charged particles in the nucleus are called protons . In
separateexperiments, it was found that each proton carries the same
quantity of charge as anelectron and has a mass of 1.67262 3 10224
gabout 1840 times the mass of theoppositely charged electron.At
this stage of investigation, scientists perceived the atom as
follows: The massof a nucleus constitutes most of the mass of the
entire atom, but the nucleus occupiesonly about 1/1013 of the
volume of the atom. We express atomic (and molecular)dimensions in
terms of the SI unit called the picometer ( pm ) , where1 pm 5 1 3
10212 mAnimationa-Particle ScatteringDetecting screen SlitGold
foil(a) (b)ParticleemitterFigure 2.8 (a) Rutherfordsexperimental
design for measuringthe scattering of a particles by apiece of gold
foil. Most of the aparticles passed through the goldfoil with
little or no defl ection. Afew were defl ected at wide
angles.Occasionally an a particle wasturned back. (b) Magnifi ed
view ofa particles passing through andbeing defl ected by
nuclei.2.2 The Structure of the Atom 47Ernest Rutherford
(18711937). New Zealand physicist. Rutherford did most of his work
in England(Manchester and Cambridge Universities). He received the
Nobel Prize in Chemistry in 1908 for hisinvestigations into the
structure of the atomic nucleus. His often-quoted comment to his
students was thatall science is either physics or
stamp-collecting.Johannes Hans Wilhelm Geiger (18821945). German
physicist. Geigers work focused on the structureof the atomic
nucleus and on radioactivity. He invented a device for measuring
radiation that is now commonlycalled the Geiger counter.Ernest
Marsden (18891970). English physicist. It is gratifying to know
that at times an undergraduatecan assist in winning a Nobel Prize.
Marsden went on to contribute signifi cantly to the development
ofscience in New Zealand.A common non-SI unit for atomic length
isthe angstrom (; 1 = 100 pm).Media PlayerRutherfords Experiment48
Atoms, Molecules, and IonsA typical atomic radius is about 100 pm,
whereas the radius of an atomic nucleus isonly about 5 3 1023 pm.
You can appreciate the relative sizes of an atom and itsnucleus by
imagining that if an atom were the size of a sports stadium, the
volumeof its nucleus would be comparable to that of a small marble.
Although the protonsare confi ned to the nucleus of the atom, the
electrons are conceived of as being spreadout about the nucleus at
some distance from it.The concept of atomic radius is useful
experimentally, but we should not inferthat atoms have well-defi
ned boundaries or surfaces. We will learn later that the
outerregions of atoms are relatively fuzzy.The NeutronRutherfords
model of atomic structure left one major problem unsolved. It was
knownthat hydrogen, the simplest atom, contains only one proton and
that the helium atomcontains two protons. Therefore, the ratio of
the mass of a helium atom to that of ahydrogen atom should be 2:1.
(Because electrons are much lighter than protons, theircontribution
to atomic mass can be ignored.) In reality, however, the ratio is
4:1.Rutherford and others postulated that there must be another
type of subatomic particlein the atomic nucleus; the proof was
provided by another English physicist,James Chadwick, in 1932. When
Chadwick bombarded a thin sheet of berylliumwith a particles, a
very high-energy radiation similar to g rays was emitted by
themetal. Later experiments showed that the rays actually consisted
of a third type ofsubatomic particles, which Chadwick named
neutrons, because they proved to beelectrically neutral particles
having a mass slightly greater than that of protons. Themystery of
the mass ratio could now be explained. In the helium nucleus there
aretwo protons and two neutrons, but in the hydrogen nucleus there
is only one protonand no neutrons; therefore, the ratio is
4:1.Figure 2.9 shows the location of the elementary particles
(protons, neutrons, andelectrons) in an atom. There are other
subatomic particles, but the electron, the proton,If the size of an
atom were expanded tothat of this sports stadium, the size of
thenucleus would be that of a marble.James Chadwick (18911972).
British physicist. In 1935 he received the Nobel Prize in Physics
forproving the existence of neutrons.ProtonNeutronFigure 2.9 The
protons andneutrons of an atom are packedin an extremely small
nucleus.Electrons are shown as cloudsaround the nucleus.and the
neutron are the three fundamental components of the atom that are
importantin chemistry. Table 2.1 shows the masses and charges of
these three elementaryparticles.2.3 Atomic Number, Mass Number, and
IsotopesAll atoms can be identifi ed by the number of protons and
neutrons they contain. Theatomic number (Z) is the number of
protons in the nucleus of each atom of an element.In a neutral atom
the number of protons is equal to the number of electrons,so the
atomic number also indicates the number of electrons present in the
atom. Thechemical identity of an atom can be determined solely from
its atomic number. Forexample, the atomic number of fl uorine is 9.
This means that each fl uorine atom has9 protons and 9 electrons.
Or, viewed another way, every atom in the universe thatcontains 9
protons is correctly named fl uorine.The mass number (A) is the
total number of neutrons and protons present in thenucleus of an
atom of an element. Except for the most common form of
hydrogen,which has one proton and no neutrons, all atomic nuclei
contain both protons andneutrons. In general, the mass number is
given bymass number 5 number of protons 1 number of neutrons5
atomic number 1 number of neutrons(2.1)The number of neutrons in an
atom is equal to the difference between the mass numberand the
atomic number, or (A 2 Z). For example, if the mass number of a
particularboron atom is 12 and the atomic number is 5 (indicating 5
protons in thenucleus), then the number of neutrons is 12 2 5 5 7.
Note that all three quantities(atomic number, number of neutrons,
and mass number) must be positive integers, orwhole numbers.Atoms
of a given element do not all have the same mass. Most elements
havetwo or more isotopes, atoms that have the same atomic number
but different massnumbers. For example, there are three isotopes of
hydrogen. One, simply known ashydrogen, has one proton and no
neutrons. The deuterium isotope contains one protonand one neutron,
and tritium has one proton and two neutrons. The accepted wayto
denote the atomic number and mass number of an atom of an element
(X) is asfollows:mass numberatomic number8n8nZAXChargeParticle Mass
(g) Coulomb Charge UnitElectron* 9.10938 3 10228 21.6022 3 10219
21Proton 1.67262 3 10224 11.6022 3 10219 11Neutron 1.67493 3 10224
0 0*More refi ned measurements have given us a more accurate value
of an electrons mass than Millikans.TABLE 2.1 Mass and Charge of
Subatomic Particles2.3 Atomic Number, Mass Number, and Isotopes
49Protons and neutrons are collectivelycalled nucleons.50 Atoms,
Molecules, and IonsThus, for the isotopes of hydrogen, we write11H
21H 31Hhydrogen deuterium tritiumAs another example, consider two
common isotopes of uranium with mass numbersof 235 and 238,
respectively:23592U 23892UThe fi rst isotope is used in nuclear
reactors and atomic bombs, whereas the secondisotope lacks the
properties necessary for these applications. With the exception
ofhydrogen, which has different names for each of its isotopes,
isotopes of elementsare identifi ed by their mass numbers. Thus,
the preceding two isotopes are calleduranium-235 (pronounced
uranium two thirty-fi ve) and uranium-238 (pronounceduranium two
thirty-eight).The chemical properties of an element are determined
primarily by the protonsand electrons in its atoms; neutrons do not
take part in chemical changes under normalconditions. Therefore,
isotopes of the same element have similar chemistries,forming the
same types of compounds and displaying similar reactivities.Example
2.1 shows how to calculate the number of protons, neutrons, and
electronsusing atomic numbers and mass numbers.1 1H 12 H 13 HReview
of Concepts(a) Name the only element having an isotope that
contains no neutrons.(b) Explain why a helium nucleus containing no
neutrons is likely to be unstable.EXAMPLE 2.1Give the number of
protons, neutrons, and electrons in each of the following
species:(a) 2011Na, (b) 2211Na, (c) 17O, and (d) carbon-14.Strategy
Recall that the superscript denotes the mass number ( A ) and the
subscriptdenotes the atomic number ( Z ). Mass number is always
greater than atomic number.(The only exception is 11H, where the
mass number is equal to the atomic number.) Ina case where no
subscript is shown, as in parts (c) and (d), the atomic number
canbe deduced from the element symbol or name. To determine the
number of electrons,remember that because atoms are electrically
neutral, the number of electrons is equalto the number of
protons.Solution (a) The atomic number is 11, so there are 11
protons. The mass number is20, so the number of neutrons is 20 2 11
5 9. The number of electrons is thesame as the number of protons;
that is, 11.(b) The atomic number is the same as that in (a), or
11. The mass number is 22, so thenumber of neutrons is 22 2 11 5
11. The number of electrons is 11. Note that thespecies in (a) and
(b) are chemically similar isotopes of sodium.(c) The atomic number
of O (oxygen) is 8, so there are 8 protons. The mass number is17,
so there are 17 2 8 5 9 neutrons. There are 8 electrons.(d)
Carbon-14 can also be represented as 14C. The atomic number of
carbon is 6, sothere are 14 2 6 5 8 neutrons. The number of
electrons is 6.Practice Exercise How many protons, neutrons, and
electrons are in the followingisotope of copper: 63Cu?Similar
problems: 2.15, 2.16.2.4 The Periodic TableMore than half of the
elements known today were discovered between 1800 and1900. During
this period, chemists noted that many elements show strong
similaritiesto one another. Recognition of periodic regularities in
physical and chemicalbehavior and the need to organize the large
volume of available informationabout the structure and properties
of elemental substances led to the developmentof the periodic
table, a chart in which elements having similar chemical and
physicalproperties are grouped together. Figure 2.10 shows the
modern periodic tablein which the elements are arranged by atomic
number (shown above the elementsymbol) in horizontal rows called
periods and in vertical columns known as groupsor families,
according to similarities in their chemical properties. Note that
elements112116 and 118 have recently been synthesized, although
they have not yetbeen named.The elements can be divided into three
categoriesmetals, nonmetals, and metalloids.A metal is a good
conductor of heat and electricity while a nonmetal isusually a poor
conductor of heat and electricity. A metalloid has properties that
areintermediate between those of metals and nonmetals. Figure 2.10
shows that the2.4 The Periodic Table
51MetalsMetalloidsNonmetals1H3Li11Na19K37Rb55Cs87Fr20Ca38Sr56Ba88Ra21Sc39Y57La89Ac22Ti40Zr72Hf104Rf23V41Nb73Ta105Db24Cr42Mo74W106Sg25Mn43Tc75Re107Bh26Fe44Ru76Os108Hs27Co45Rh77Ir109Mt28Ni46Pd78Pt29Cu47Ag79Au30Zn48Cd80Hg31Ga49In81Tl32Ge50Sn82Pb33As51Sb83Bi34Se52Te84Po35Br53I85At36Kr54Xe86Rn13Al14Si15P16S17Cl18Ar5B6C7N8O9F10Ne2He4Be12Mg58Ce90Th59Pr91Pa60Nd92U61Pm93Np62Sm94Pu63Eu95Am64Gd96Cm65Tb97Bk66Dy98Cf67Ho99Es68Er100Fm69Tm101Md70Yb102No71Lu103Lr110Ds112
113 114 115 116 (117) 11811A22A33B44B55B66B77B98B111B8 10
122B133A144A155A166A177A188A111RgFigure 2.10 The modern periodic
table. The elements are arranged according to the atomic numbers
above their symbols. With theexception of hydrogen (H), nonmetals
appear at the far right of the table. The two rows of metals
beneath the main body of the tableare conventionally set apart to
keep the table from being too wide. Actually, cerium (Ce) should
follow lanthanum (La), and thorium (Th)should come right after
actinium (Ac). The 118 group designation has been recommended by
the International Union of Pure andApplied Chemistry (IUPAC) but is
not yet in wide use. In this text, we use the standard U.S.
notation for group numbers (1A8A and1B8B). No names have yet been
assigned to elements 112116, and 118. Element 117 has not yet been
synthesized.52C H E M I S T R Yin ActionThe majority of elements
are naturally occurring. How arethese elements distributed on
Earth, and which are essentialto living systems?Earths crust
extends from the surface to a depth of about40 km (about 25 mi).
Because of technical diffi culties, scientistshave not been able to
study the inner portions of Earth as easilyas the crust.
Nevertheless, it is believed that there is a solid coreconsisting
mostly of iron at the center of Earth. Surrounding thecore is a
layer called the mantle, which consists of hot fl uidcontaining
iron, carbon, silicon, and sulfur.Of the 83 elements that are found
in nature, 12 make up99.7 percent of Earths crust by mass. They
are, in decreasingorder of natural abundance, oxygen (O), silicon
(Si), aluminum(Al), iron (Fe), calcium (Ca), magnesium (Mg), sodium
(Na),potassium (K), titanium (Ti), hydrogen (H), phosphorus (P),and
manganese (Mn). In discussing the natural abundance of
theDistribution of Elements on Earth and in Living Systemselements,
we should keep in mind that (1) the elements are notevenly
distributed throughout Earths crust, and (2) most elementsoccur in
combined forms. These facts provide the basisfor most methods of
obtaining pure elements from their compounds,as we will see in
later chapters.The accompanying table lists the essential elements
in thehuman body. Of special interest are the trace elements, such
asiron (Fe), copper (Cu), zinc (Zn), iodine (I), and cobalt
(Co),which together make up about 0.1 percent of the bodys
mass.These elements are necessary for biological functions such
asgrowth, transport of oxygen for metabolism, and defenseagainst
disease. There is a delicate balance in the amounts ofthese
elements in our bodies. Too much or too little over anextended
period of time can lead to serious illness, retardation,or even
death.2900 km 3480 kmCrustCoreMantleStructure of Earths
interior.Essential Elements in the Human BodyElement Percent by
Mass* Element Percent by Mass*Oxygen 65 Sodium 0.1Carbon 18
Magnesium 0.05Hydrogen 10 Iron ,0.05Nitrogen 3 Cobalt ,0.05Calcium
1.6 Copper ,0.05Phosphorus 1.2 Zinc ,0.05Potassium 0.2 Iodine
,0.05Sulfur 0.2 Selenium ,0.01Chlorine 0.2 Fluorine ,0.01*Percent
by mass gives the mass of the element in grams present in a 100-g
sample.(a) Natural abundance of the elementsin percent by mass. For
example, oxygensabundance is 45.5 percent. Thismeans that in a
100-g sample of Earthscrust there are, on the average, 45.5 gof the
element oxygen. (b) Abundanceof elements in the human body in
percentby mass.Magnesium 2.8%Oxygen45.5% Oxygen65%Silicon27.2%
Hydrogen 10%Carbon18%Calcium 4.7%All others 5.3%All others
1.2%Phosphorus 1.2%Calcium 1.6%Nitrogen 3%Iron 6.2%Aluminum 8.3%(a)
(b)majority of known elements are metals; only 17 elements are
nonmetals, and 8 elementsare metalloids. From left to right across
any period, the physical and chemicalproperties of the elements
change gradually from metallic to nonmetallic.Elements are often
referred to collectively by their periodic table group number(Group
1A, Group 2A, and so on). However, for convenience, some element
groupshave been given special names. The Group 1A elements (Li, Na,
K, Rb, Cs, and Fr)are called alkali metals, and the Group 2A
elements (Be, Mg, Ca, Sr, Ba, and Ra) arecalled alkaline earth
metals . Elements in Group 7A (F, Cl, Br, I, and At) are knownas
halogens, and elements in Group 8A (He, Ne, Ar, Kr, Xe, and Rn) are
called noblegases, or rare gases .The periodic table is a handy
tool that correlates the properties of the elementsin a systematic
way and helps us to make predictions about chemical behavior.
Wewill take a closer look at this keystone of chemistry in Chapter
8.The Chemistry in Action essay on p. 52 describes the distribution
of the elementson Earth and in the human body.Review of ConceptsIn
viewing the periodic table, do chemical properties change more
markedlyacross a period or down a group?2.5 Molecules and IonsOf
all the elements, only the six noble gases in Group 8A of the
periodic table (He,Ne, Ar, Kr, Xe, and Rn) exist in nature as
single atoms. For this reason, they arecalled monatomic (meaning a
single atom) gases. Most matter is composed ofmolecules or ions
formed by atoms.MoleculesA molecule is an aggregate of at least two
atoms in a defi nite arrangement heldtogether by chemical forces
(also called chemical bonds ). A molecule may containatoms of the
same element or atoms of two or more elements joined in a fi xed
ratio,in accordance with the law of defi nite proportions stated in
Section 2.1. Thus, a moleculeis not necessarily a compound, which,
by defi nition, is made up of two or moreelements (see Section
1.4). Hydrogen gas, for example, is a pure element, but it
consistsof molecules made up of two H atoms each. Water, on the
other hand, is amolecular compound that contains hydrogen and
oxygen in a ratio of two H atomsand one O atom. Like atoms,
molecules are electrically neutral.The hydrogen molecule,
symbolized as H2, is called a diatomic molecule becauseit contains
only two atoms. Other elements that normally exist as diatomic
moleculesare nitrogen (N2) and oxygen (O2), as well as the Group 7A
elementsfl uorine (F2),chlorine (Cl2), bromine (Br2), and iodine
(I2). Of course, a diatomic molecule cancontain atoms of different
elements. Examples are hydrogen chloride (HCl) and carbonmonoxide
(CO).The vast majority of molecules contain more than two atoms.
They can be atomsof the same element, as in ozone (O3), which is
made up of three atoms of oxygen,or they can be combinations of two
or more different elements. Molecules containingmore than two atoms
are called polyatomic molecules . Like ozone, water (H2O)
andammonia (NH3) are polyatomic molecules.2.5 Molecules and Ions
53We will discuss the nature of chemicalbonds in Chapters 9 and
10.1A2A 3A 4A 5A 6A 7A8AN O FClBrIHElements that exist as diatomic
molecules.54 Atoms, Molecules, and IonsIonsAn ion is an atom or a
group of atoms that has a net positive or negative charge.The
number of positively charged protons in the nucleus of an atom
remains the sameduring ordinary chemical changes (called chemical
reactions), but negatively chargedelectrons may be lost or gained.
The loss of one or more electrons from a neutralatom results in a
cation, an ion with a net positive charge. For example, a
sodiumatom (Na) can readily lose an electron to become a sodium
cation, which is representedby Na1:Na Atom Na1 Ion11 protons 11
protons11 electrons 10 electronsOn the other hand, an anion is an
ion whose net charge is negative due to an increasein the number of
electrons. A chlorine atom (Cl), for instance, can gain an
electronto become the chloride ion Cl2:Cl Atom Cl2 Ion17 protons 17
protons17 electrons 18 electronsSodium chloride (NaCl), ordinary
table salt, is called an ionic compound because itis formed from
cations and anions.An atom can lose or gain more than one electron.
Examples of ions formed bythe loss or gain of more than one
electron are Mg21, Fe31, S22, and N32. These ions,as well as Na1
and Cl2, are called monatomic ions because they contain only
oneatom. Figure 2.11 shows the charges of a number of monatomic
ions. With very fewexceptions, metals tend to form cations and
nonmetals form anions.In addition, two or more atoms can combine to
form an ion that has a net positiveor net negative charge.
Polyatomic ions such as OH2 (hydroxide ion), CN2(cyanide ion), and
NH14(ammonium ion) are ions containing more than one
atom.11A22A33B44B55B66B77B98B111B8 10
122B133A144A155A166A177A188ALi+Na+K+Rb+Cs+Ca2+Sr2+Ba2+Ag+Ni Zn2+
Se2 Br 2+Ni3+Mn2+Mn3+Cr2+Cr3+Cd2+ Te2 IAl3+ S2 ClC4 N3 O2
FMg2+Fe2+Fe3+Co2+Co3+Cu+Cu2+P3Sn2+Sn4+Pb2+Pb4+Hg2+Hg2+Au+Au3+2Figure
2.11 Common monatomic ions arranged according to their positions in
the periodic table. Note that the Hg221 ion containstwo atoms.In
Chapter 8, we will see why atoms ofdifferent elements gain (or
lose) a specifi cnumber of electrons.2.6 Chemical FormulasChemists
use chemical formulas to express the composition of molecules and
ioniccompounds in terms of chemical symbols. By composition we mean
not only the elementspresent but also the ratios in which the atoms
are combined. Here we areconcerned with two types of formulas:
molecular formulas and empirical formulas.Molecular FormulasA
molecular formula shows the exact number of atoms of each element
in the smallestunit of a substance. In our discussion of molecules,
each example was given withits molecular formula in parentheses.
Thus, H2 is the molecular formula for hydrogen,O2 is oxygen, O3 is
ozone, and H2O is water. The subscript numeral indicates thenumber
of atoms of an element present. There is no subscript for O in H2O
becausethere is only one atom of oxygen in a molecule of water, and
so the number oneis omitted from the formula. Note that oxygen (O2)
and ozone (O3) are allotropes ofoxygen. An allotrope is one of two
or more distinct forms of an element. Two allotropicforms of the
element carbondiamond and graphiteare dramatically differentnot
only in properties but also in their relative cost.Molecular
ModelsMolecules are too small for us to observe directly. An
effective means of visualizingthem is by the use of molecular
models. Two standard types of molecular models arecurrently in use:
ball-and-stick models and space-fi lling models ( Figure 2.12 ). In
balland-stick model kits, the atoms are wooden or plastic balls
with holes in them. Sticksor springs are used to represent chemical
bonds. The angles they form between atomsapproximate the bond
angles in actual molecules. With the exception of the H atom,the
balls are all the same size and each type of atom is represented by
a specifi c color.In space-fi lling models, atoms are represented
by truncated balls held together by snap2.6 Chemical Formulas
55MolecularformulaStructuralformulaBall-and-stickmodelSpace-fillingmodelHydrogenH2HHWaterH2OHOHAmmoniaNH3HNHWHMethaneCH4HWHCHWHFigure
2.12 Molecular and structural formulas and molecular models of four
common molecules.See back endpaper for color codes foratoms.56
Atoms, Molecules, and Ionsfasteners, so that the bonds are not
visible. The balls are proportional in size to atoms.The fi rst
step toward building a molecular model is writing the structural
formula,which shows how atoms are bonded to one another in a
molecule. For example, it isknown that each of the two H atoms is
bonded to an O atom in the water molecule.Therefore, the structural
formula of water is H}O}H. A line connecting the twoatomic symbols
represents a chemical bond.Ball-and-stick models show the
three-dimensional arrangement of atoms clearly,and they are fairly
easy to construct. However, the balls are not proportional to
thesize of atoms. Furthermore, the sticks greatly exaggerate the
space between atoms ina molecule. Space-fi lling models are more
accurate because they show the variationin atomic size. Their
drawbacks are that they are time-consuming to put together andthey
do not show the three-dimensional positions of atoms very well. We
will useboth models extensively in this text.Empirical FormulasThe
molecular formula of hydrogen peroxide, a substance used as an
antiseptic and asa bleaching agent for textiles and hair, is H2O2.
This formula indicates that each hydrogenperoxide molecule consists
of two hydrogen atoms and two oxygen atoms. The ratio ofhydrogen to
oxygen atoms in this molecule is 2:2 or 1:1. The empirical formula
ofhydrogen peroxide is HO. Thus, the empirical formula tells us
which elements are presentand the simplest whole-number ratio of
their atoms, but not necessarily the actualnumber of atoms in a
given molecule. As another example, consider the compoundhydrazine
(N2H4), which is used as a rocket fuel. The empirical formula of
hydrazine isNH2. Although the ratio of nitrogen to hydrogen is 1:2
in both the molecular formula(N2H4) and the empirical formula
(NH2), only the molecular formula tells us the actualnumber of N
atoms (two) and H atoms (four) present in a hydrazine
molecule.Empirical formulas are the simplest chemical formulas;
they are written by reducingthe subscripts in the molecular
formulas to the smallest possible whole numbers.Molecular formulas
are the true formulas of molecules. If we know the
molecularformula, we also know the empirical formula, but the
reverse is not true. Why, then,do chemists bother with empirical
formulas? As we will see in Chapter 3, when chemistsanalyze an
unknown compound, the fi rst step is usually the determination of
thecompounds empirical formula. With additional information, it is
possible to deducethe molecular formula.For many molecules, the
molecular formula and the empirical formula are oneand the same.
Some examples are water (H2O), ammonia (NH3), carbon dioxide(CO2),
and methane (CH4).Examples 2.2 and 2.3 deal with writing molecular
formulas from molecular modelsand writing empirical formulas from
molecular formulas.H2O2The word empirical means derived
fromexperiment. As we will see in Chapter 3,empirical formulas are
determinedexperimentally.Similar problems: 2.47,
2.48.OHCMethanolEXAMPLE 2.2Write the molecular formula of methanol,
an organic solvent and antifreeze, from itsball-and-stick model,
shown in the margin.Solution Refer to the labels (also see back
endpapers). There are four H atoms, oneC atom, and one O atom.
Therefore, the molecular formula is CH4O. However, thestandard way
of writing the molecular formula for methanol is CH3OH because it
showshow the atoms are joined in the molecule.Practice Exercise
Write the molecular formula of chloroform, which is used as
asolvent and a cleansing agent. The ball-and-stick model of
chloroform is shown in themargin on p. 57.Formula of Ionic
CompoundsThe formulas of ionic compounds are usually the same as
their empirical formulasbecause ionic compounds do not consist of
discrete molecular units. For example, asolid sample of sodium
chloride (NaCl) consists of equal numbers of Na1 and Cl2ions
arranged in a three-dimensional network ( Figure 2.13 ). In such a
compound thereis a 1:1 ratio of cations to anions so that the
compound is electrically neutral. As youcan see in Figure 2.13 , no
Na1 ion in NaCl is associated with just one particular Cl2ion. In
fact, each Na1 ion is equally held by six surrounding Cl2 ions and
vice versa.Thus, NaCl is the empirical formula for sodium chloride.
In other ionic compounds,the actual structure may be different, but
the arrangement of cations and anions issuch that the compounds are
all electrically neutral. Note that the charges on thecation and
anion are not shown in the formula for an ionic
compound.ClHCChloroformEXAMPLE 2.3Write the empirical formulas for
the following molecules: (a) acetylene (C2H2), which isused in
welding torches; (b) glucose (C6H12O6), a substance known as blood
sugar; and(c) nitrous oxide (N2O), a gas that is used as an
anesthetic gas (laughing gas) and asan aerosol propellant for
whipped creams.Strategy Recall that to write the empirical formula,
the subscripts in the molecularformula must be converted to the
smallest possible whole numbers.Solution(a) There are two carbon
atoms and two hydrogen atoms in acetylene. Dividing thesubscripts
by 2, we obtain the empirical formula CH.(b) In glucose there are 6
carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms.Dividing the
subscripts by 6, we obtain the empirical formula CH2O. Note that
ifwe had divided the subscripts by 3, we would have obtained the
formula C2H4O2.Although the ratio of carbon to hydrogen to oxygen
atoms in C2H4O2 is the same asthat in C6H12O6 (1:2:1), C2H4O2 is
not the simplest formula because its subscriptsare not in the
smallest whole-number ratio.(c) Because the subscripts in N2O are
already the smallest possible whole numbers, theempirical formula
for nitrous oxide is the same as its molecular formula.Practice
Exercise Write the empirical formula for caffeine (C8H10N4O2), a
stimulantfound in tea and coffee.Similar problems: 2.45,
2.46.Sodium metal reacting with chlorine gas toform sodium
chloride.2.6 Chemical Formulas 57(a) (b) (c)Figure 2.13 (a)
Structure of solid NaCl. (b) In reality, the cations are in contact
with the anions. In both (a) and (b), the smaller spheresrepresent
Na1 ions and the larger spheres, Cl2 ions. (c) Crystals of NaCl.58
Atoms, Molecules, and IonsFor ionic compounds to be electrically
neutral, the sum of the charges on the cationand anion in each
formula unit must be zero. If the charges on the cation and anion
arenumerically different, we apply the following rule to make the
formula electrically neutral:The subscript of the cation is
numerically equal to the charge on the anion, and thesubscript of
the anion is numerically equal to the charge on the cation. If the
chargesare numerically equal, then no subscripts are necessary.
This rule follows from the factthat because the formulas of ionic
compounds are usually empirical formulas, the subscriptsmust always
be reduced to the smallest ratios. Let us consider some examples.
Potassium Bromide. The potassium cation K1 and the bromine anion
Br2 combineto form the ionic compound potassium bromide. The sum of
the charges is11 1 (21) 5 0, so no subscripts are necessary. The
formula is KBr. Zinc Iodide. The zinc cation Zn21 and the iodine
anion I2 combine to form zinciodide. The sum of the charges of one
Zn21 ion and one I2 ion is 12 1 (21) 511. To make the charges add
up to zero we multiply the 21 charge of the anionby 2 and add the
subscript 2 to the symbol for iodine. Therefore the formulafor zinc
iodide is ZnI2. Aluminum Oxide. The cation is Al31 and the oxygen
anion is O22. The followingdiagram helps us determine the
subscripts for the compound formed by thecation and the anion:Al2
O3Al 3 _ O 2 _The sum of the charges is 2(13) 1 3(22) 5 0. Thus,
the formula for aluminumoxide is Al2O3.Note that in each of the
above three examples,the subscripts are in the smallest ratios.When
magnesium burns in air, it forms bothmagnesium oxide and magnesium
nitride.EXAMPLE 2.4Write the formula of magnesium nitride,
containing the Mg21 and N32 ions.Strategy Our guide for writing
formulas for ionic compounds is electrical neutrality;that is, the
total charge on the cation(s) must be equal to the total charge on
theanion(s). Because the charges on the Mg21 and N32 ions are not
equal, we know theformula cannot be MgN. Instead, we write the
formula as MgxNy, where x and y aresubscripts to be
determined.Solution To satisfy electrical neutrality, the following
relationship must hold:(12)x 1 (23)y 5 0Solving, we obtain x/y 5
3/2. Setting x 5 3 and y 5 2, we writeMg3 N2Mg N 2 _ 3 _Check The
subscripts are reduced to the smallest whole number ratio of the
atomsbecause the chemical formula of an ionic compound is usually
its empirical formula.Practice Exercise Write the formulas of the
following ionic compounds: (a) chromiumsulfate (containing the Cr31
and SO422 ions) and (b) titanium oxide (containing the Ti41and O22
ions).Refer to Figure 2.11 for charges of cationsand anions.Similar
problems: 2.43, 2.44.2.7 Naming CompoundsWhen chemistry was a young
science and the number of known compounds was small,it was possible
to memorize their names. Many of the names were derived from
theirphysical appearance, properties, origin, or applicationfor
example, milk of magnesia,laughing gas, limestone, caustic soda,
lye, washing soda, and baking soda.Today the number of known
compounds is well over 20 million. Fortunately, itis not necessary
to memorize their names. Over the years chemists have devised
aclear system for naming chemical substances. The rules are
accepted worldwide,facilitating communication among chemists and
providing a useful way of labelingan overwhelming variety of
substances. Mastering these rules now will prove benefi -cial
almost immediately as we proceed with our study of chemistry.To
begin our discussion of chemical nomenclature, the naming of
chemical compounds,we must fi rst distinguish between inorganic and
organic compounds. Organiccompounds contain carbon, usually in
combination with elements such as hydrogen,oxygen, nitrogen, and
sulfur. All other compounds are classifi ed as inorganic
compounds.For convenience, some carbon-containing compounds, such
as carbon monoxide(CO), carbon dioxide (CO2), carbon disulfi de
(CS2), compounds containing thecyanide group (CN2), and carbonate
(CO322) and bicarbonate (HCO32) groups areconsidered to be
inorganic compounds. Section 2.8 gives a brief introduction
toorganic compounds.To organize and simplify our venture into
naming compounds, we can divideinorganic compounds into four
categories: ionic compounds, molecular compounds,acids and bases,
and hydrates.Ionic CompoundsIn Section 2.5 we learned that ionic
compounds are made up of cations (positive ions)and anions
(negative ions). With the important exception of the ammonium ion,
NH41,all cations of interest to us are derived from metal atoms.
Metal cations take theirnames from the elements. For example,For
names and symbols of the elements,see front end papers.1A2A 3A 4A
5A 6A 7A8AN OAl SFClBrILiNaKRbCsMgCaSrBaThe most reactive metals
(green) and themost reactive nonmetals (blue) combine toform ionic
compounds.Review of ConceptsMatch each of the diagrams shown here
with the following ionic compounds:Al2O3, LiH, Na2S, Mg(NO3)2.
(Green spheres represent cations and red spheresrepresent
anions.)(a) (b) (c) (d)2.7 Naming Compounds 59Element Name of
CationNa sodium Na1 sodium ion (or sodium cation)K potassium K1
potassium ion (or potassium cation)Mg magnesium Mg21 magnesium ion
(or magnesium cation)Al aluminum Al31 aluminum ion (or aluminum
cation)Many ionic compounds are binary compounds, or compounds
formed from justtwo elements. For binary compounds, the fi rst
element named is the metal cation,followed by the nonmetallic
anion. Thus, NaCl is sodium chloride. The anion is namedMedia
PlayerFormation of an Ionic Compound60 Atoms, Molecules, and Ionsby
taking the fi rst part of the element name (chlorine) and adding
-ide. Potassiumbromide (KBr), zinc iodide (ZnI2), and aluminum
oxide (Al2O3) are also binary compounds.Table 2.2 shows the -ide
nomenclature of some common monatomic anionsaccording to their
positions in the periodic table.The -ide ending is also used for
certain anion groups containing different elements,such as
hydroxide (OH2) and cyanide (CN2). Thus, the compounds LiOH andKCN
are named lithium hydroxide and potassium cyanide, respectively.
These and anumber of other such ionic substances are called ternary
compounds, meaning compoundsconsisting of three elements. Table 2.3
lists alphabetically the names of anumber of common cations and
anions.Certain metals, especially the transition metals, can form
more than one type of cation.Take iron as an example. Iron can form
two cations: Fe21 and Fe31. An older nomenclaturesystem that is
still in limited use assigns the ending -ous to the cation with
fewerpositive charges and the ending -ic to the cation with more
positive charges:Fe21 ferrous ionFe31 ferric ionThe names of the
compounds that these iron ions form with chlorine would thus
beFeCl2 ferrous chlorideFeCl3 ferric chlorideThis method of naming
ions has some distinct limitations. First, the -ous and -icsuffi
xes do not provide information regarding the actual charges of the
two cationsinvolved. Thus, the ferric ion is Fe31, but the cation
of copper named cupric hasthe formula Cu21. In addition, the -ous
and -ic designations provide names foronly two different elemental
cations. Some metallic elements can assume three ormore different
positive charges in compounds. Therefore, it has become
increasinglycommon to designate different cations with Roman
numerals. This is called theStock system. In this system, the Roman
numeral I indicates one positive charge,II means two positive
charges, and so on. For example, manganese (Mn) atoms canassume
several different positive charges:Mn21: MnO manganese(II)
oxideMn31: Mn2O3 manganese(III) oxideMn41: MnO2 manganese(IV)
oxideThese names are pronounced manganese-two oxide,
manganese-three oxide, andmanganese-four oxide. Using the Stock
system, we denote the ferrous ion and theGroup 4A Group 5A Group 6A
Group 7AC carbide (C42)* N nitride (N32) O oxide (O22) F fl uoride
(F2)Si silicide (Si42) P phosphide (P32) S sulfi de (S22) Cl
chloride (Cl2)Se selenide (Se22) Br bromide (Br2)Te telluride
(Te22) I iodide (I2)*The word carbide is also used for the anion C2
22.TABLE 2.2 The -ide Nomenclature of Some Common Monatomic
AnionsAccording to Their Positions in the Periodic TableAlfred E.
Stock (18761946). German chemist. Stock did most of his research in
the synthesis and characterizationof boron, beryllium, and silicon
compounds. He was the fi rst scientist to explore the dangersof
mercury poisoning.3B 4B 5B 6B 7B 8B 1B 2BThe transition metals are
the elements inGroups 1B and 3B8B (see Figure 2.10).FeCl2 (left)
and FeCl3 (right).Keep in mind that the Roman numeralsrefer to the
charges on the metal cations.Cation Anionaluminum (Al31) bromide
(Br2)ammonium (NH14) carbonate (CO3 22)barium (Ba21) chlorate
(ClO32)cadmium (Cd21) chloride (Cl2)calcium (Ca21) chromate (CrO4
22)cesium (Cs1) cyanide (CN2)chromium(III) or chromic (Cr31)
dichromate (Cr2O7 22)cobalt(II) or cobaltous (Co21) dihydrogen
phosphate (H2PO42)copper(I) or cuprous (Cu1) fl uoride
(F2)copper(II) or cupric (Cu21) hydride (H2)hydrogen (H1) hydrogen
carbonate or bicarbonate (HCO32)iron(II) or ferrous (Fe21) hydrogen
phosphate (HPO4 22)iron(III) or ferric (Fe31) hydrogen sulfate or
bisulfate (HSO42)lead(II) or plumbous (Pb21) hydroxide (OH2)lithium
(Li1) iodide (I2)magnesium (Mg21) nitrate (NO32)manganese(II) or
manganous (Mn21) nitride (N32)mercury(I) or mercurous (Hg2 21)*
nitrite (NO22)mercury(II) or mercuric (Hg21) oxide (O22)potassium
(K1) permanganate (MnO42)rubidium (Rb1) peroxide (O2 22)silver
(Ag1) phosphate (PO4 32)sodium (Na1) sulfate (SO4 22)strontium
(Sr21) sulfi de (S22)tin(II) or stannous (Sn21) sulfi te (SO3
22)zinc (Zn21) thiocyanate (SCN2)*Mercury(I) exists as a pair as
shown.TABLE 2.3 Names and Formulas of Some Common Inorganic
Cationsand Anions2.7 Naming Compounds 61ferric ion as iron(II) and
iron(III), respectively; ferrous chloride becomes iron(II)chloride;
and ferric chloride is called iron(III) chloride. In keeping with
modern practice,we will favor the Stock system of naming compounds
in this textbook.Examples 2.5 and 2.6 illustrate how to name ionic
compounds and write formulasfor ionic compounds based on the
information given in Figure 2.11 and Tables 2.2and 2.3 .EXAMPLE
2.5Name the following compounds: (a) Cu(NO3)2, (b) KH2PO4, and (c)
NH4ClO3.Strategy Note that the compounds in (a) and (b) contain
both metal and nonmetalatoms, so we expect them to be ionic
compounds. There are no metal atoms in (c) butthere is an ammonium
group, which bears a positive charge. So NH4ClO3 is also
an(Continued)62 Atoms, Molecules, and IonsSimilar problems:
2.57(b), (e), (f).EXAMPLE 2.6Write chemical formulas for the
following compounds: (a) mercury(I) nitrite, (b) cesiumsulfi de,
and (c) calcium phosphate.Strategy We refer to Table 2.3 for the
formulas of cations and anions. Recall that theRoman numerals in
the Stock system provide useful information about the charges ofthe
cation.Solution(a) The Roman numeral shows that the mercury ion
bears a 11 charge. According toTable 2.3 , however, the mercury(I)
ion is diatomic (that is, Hg2 21) and the nitrite ionis NO22.
Therefore, the formula is Hg2(NO2)2.(b) Each sulfi de ion bears two
negative charges, and each cesium ion bears one positivecharge
(cesium is in Group 1A, as is sodium). Therefore, the formula is
Cs2S.(c) Each calcium ion (Ca21) bears two positive charges, and
each phosphate ion (PO4 32)bears three negative charges. To make
the sum of the charges equal zero, we mustadjust the numbers of
cations and anions:3(12) 1 2(23) 5 0Thus, the formula is
Ca3(PO4)2.Practice Exercise Write formulas for the following ionic
compounds: (a) rubidiumsulfate and (b) barium hydride.Note that the
subscripts of this ioniccompound are not reduced to the
smallestratio because the Hg(I) ion exists as a pairor
dimer.Similar problems: 2.59(a), (b), (d), (h), (i).ionic compound.
Our reference for the names of cations and anions is Table 2.3 .
Keepin mind that if a metal atom can form cations of different
charges (see Figure 2.11 ), weneed to use the Stock
system.Solution(a) The nitrate ion (NO32) bears one negative
charge, so the copper ion must have twopositive charges. Because
copper forms both Cu1 and Cu21 ions, we need to usethe Stock system
and call the compound copper(II) nitrate.(b) The cation is K1 and
the anion is H2PO42 (dihydrogen phosphate). Becausepotassium only
forms one type of ion (K1), there is no need to use potassium(I)
inthe name. The compound is potassium dihydrogen phosphate.(c) The
cation is NH41 (ammonium ion) and the anion is ClO32. The compound
isammonium chlorate.Practice Exercise Name the following compounds:
(a) PbO and (b) Li2SO3.Molecular CompoundsUnlike ionic compounds,
molecular compounds contain discrete molecular units. Theyare
usually composed of nonmetallic elements (see Figure 2.10 ). Many
molecularcompounds are binary compounds. Naming binary molecular
compounds is similar tonaming binary ionic compounds. We place the
name of the fi rst element in the formulafi rst, and the second
element is named by adding -ide to the root of the element
name.Some examples areHCl hydrogen chlorideHBr hydrogen bromideSiC
silicon carbideIt is quite common for one pair of elements to form
several different compounds.In these cases, confusion in naming the
compounds is avoided by the use of Greekprefi xes to denote the
number of atoms of each element present ( Table 2.4 ). Considerthe
following examples:CO carbon monoxideCO2 carbon dioxideSO2 sulfur
dioxideSO3 sulfur trioxideNO2 nitrogen dioxideN2O4 dinitrogen
tetroxideThe following guidelines are helpful in naming compounds
with prefi xes: The prefi x mono- may be omitted for the fi rst
element. For example, PCl3 isnamed phosphorus trichloride, not
monophosphorus trichloride. Thus, the absenceof a prefi x for the
fi rst element usually means there is only one atom of that
elementpresent in the molecule. For oxides, the ending a in the
prefi x is sometimes omitted. For example, N2O4may be called
dinitrogen tetroxide rather than dinitrogen tetraoxide.Exceptions
to the use of Greek prefi xes are molecular compounds
containinghydrogen. Traditionally, many of these compounds are
called either by their common,nonsystematic names or by names that
do not specifi cally indicate the number of Hatoms present:B2H6
diboraneCH4 methaneSiH4 silaneNH3 ammoniaPH3 phosphineH2O waterH2S
hydrogen sulfi deNote that even the order of writing the elements
in the formulas for hydrogen compoundsis irregular. In water and
hydrogen sulfi de, H is written fi rst, whereas it appearslast in
the other compounds.Writing formulas for molecular compounds is
usually straightforward. Thus, thename arsenic trifl uoride means
that there are three F atoms and one As atom in eachmolecule, and
the molecular formula is AsF3. Note that the order of elements in
theformula is the same as in its name.Binary compounds containing
carbon andhydrogen are organic compounds; they donot follow the
same naming conventions.We will discuss the naming of
organiccompounds in Chapter 24.2.7 Naming Compounds 63EXAMPLE
2.7Name the following molecular compounds: (a) SiCl4 and (b)
P4O10.Strategy We refer to Table 2.4 for prefi xes. In (a) there is
only one Si atom so we donot use the prefi x mono.Solution (a)
Because there are four chlorine atoms present, the compound is
silicontetrachloride.(b) There are four phosphorus atoms and ten
oxygen atoms present, so the compound istetraphosphorus decoxide.
Note that the a is omitted in deca.Practice Exercise Name the
following molecular compounds: (a) NF3 and (b) Cl2O7.Similar
problems: 2.57(c), (i), (j).Greek Prefi xes Used inNaming
MolecularCompoundsTABLE 2.4Prefi x Meaningmono- 1di- 2tri- 3tetra-
4penta- 5hexa- 6hepta- 7octa- 8nona- 9deca- 1064 Atoms, Molecules,
and IonsFigure 2.14 summarizes the steps for naming ionic and
binary molecularcompounds.EXAMPLE 2.8Write chemical formulas for
the following molecular compounds: (a) carbon disulfi deand (b)
disilicon hexabromide.Strategy Here we need to convert prefi xes to
numbers of atoms (see Table 2.4 ). Becausethere is no prefi x for
carbon in (a), it means that there is only one carbon atom
present.Solution (a) Because there are two sulfur atoms and one
carbon atom present, theformula is CS2.(b) There are two silicon
atoms and six bromine atoms present, so the formula is
Si2Br6.Practice Exercise Write chemical formulas for the following
molecular compounds:(a) sulfur tetrafl uoride and (b) dinitrogen
pentoxide.Cation hasonly one chargeIonicCation: metal or NH4+Anion:
monatomic orpolyatomic Binary compoundsof
nonmetalsCompoundMolecularCation has morethan one
chargeNamingNamingNaming Alkali metal cations Alkaline earth metal
cations Ag+, Al3+, Cd2+, Zn2+ Other metal cations Name metal first
If monatomic anion,add ide to theroot of the elementname If
polyatomic anion,use name of anion(see Table 2.3) Name metal first
Specify charge ofmetal cation withRoman numeralin parentheses If
monatomic anion,add ide to theroot of the elementname If polyatomic
anion,use name of anion(see Table 2.3) Use prefixes forboth
elements present(Prefix monousually omitted forthe first element)
Add ide to theroot of the secondelementFigure 2.14 Steps for naming
ionic and binary molecular compounds.Similar problems: 2.59(g),
(j).2.7 Naming Compounds 65Anion Corresponding AcidF2 (fl uoride)
HF (hydrofl uoric acid)Cl2 (chloride) HCl (hydrochloric acid)Br2
(bromide) HBr (hydrobromic acid)I2 (iodide) HI (hydroiodic acid)CN2
(cyanide) HCN (hydrocyanic acid)S22 (sulfi de) H2S (hydrosulfuric
acid)TABLE 2.5 Some Simple AcidsNote that these acids all exist as
molecularcompounds in the gas phase.Acids and BasesNaming AcidsAn
acid can be described as a substance that yields hydrogen ions (H1)
when dissolvedin water. (H1 is equivalent to one proton, and is
often referred to that way.)Formulas for acids contain one or more
hydrogen atoms as well as an anionic group.Anions whose names end
in -ide form acids with a hydro- prefi x and an -icending, as shown
in Table 2.5 . In some cases two different names seem to be
assignedto the same chemical formula.HCl hydrogen chlorideHCl
hydrochloric acidThe name assigned to the compound depends on its
physical state. In the gaseous orpure liquid state, HCl is a
molecular compound called hydrogen chloride. When it isdissolved in
water, the molecules break up into H1 and Cl2 ions; in this state,
thesubstance is called hydrochloric acid.Oxoacids are acids that
contain hydrogen, oxygen, and another element (the centralelement).
The formulas of oxoacids are usually written with the H fi rst,
followedby the central element and then O. We use the following fi
ve common acids as ourreferences in naming oxoacids:H2CO3 carbonic
acidHClO3 chloric acidHNO3 nitric acidH3PO4 phosphoric acidH2SO4
sulfuric acidOften two or more oxoacids have the same central atom
but a different number of Oatoms. Starting with our reference
oxoacids whose names all end with -ic, we usethe following rules to
name these compounds.1. Addition of one O atom to the -ic acid: The
acid is called per . . . -ic acid.Thus, adding an O atom to HClO3
changes chloric acid to perchloric acid,HClO4.2. Removal of one O
atom from the -ic acid: The acid is called -ous acid. Thus,nitric
acid, HNO3, becomes nitrous acid, HNO2.3. Removal of two O atoms
from the -ic acid: The acid is called hypo . . . -ousacid. Thus,
when HBrO3 is converted to HBrO, the acid is called
hypobromousacid.ClHClH3O+When dissolved in water, the HCl
moleculeis converted to the H1 and Cl2 ions.The H1 ion is
associated with one ormore water molecules, and is
usuallyrepresented as H3O1.OHNHNO3OHCH2CO366 Atoms, Molecules, and
IonsThe rules for naming oxoanions, anions of oxoacids, are as
follows:1. When all the H ions are removed from the -ic acid, the
anions name ends with-ate. For example, the anion CO322 derived
from H2CO3 is called carbonate.2. When all the H ions are removed
from the -ous acid, the anions name endswith -ite. Thus, the anion
ClO22 derived from HClO2 is called chlorite.3. The names of anions
in which one or more but not all the hydrogen ions havebeen removed
must indicate the number of H ions present. For example,
considerthe anions derived from phosphoric acid:H3PO4 phosphoric
acidH2PO42 dihydrogen phosphateHPO422 hydrogen phosphatePO432
phosphateNote that we usually omit the prefi x mono- when there is
only one H in theanion. Figure 2.15 summarizes the nomenclature for
the oxoacids and oxoanions, andTable 2.6 gives the names of the
oxoacids and oxoanions that contain chlorine.OPHH3PO4hypo ous
acidous acidReference ic acidper ic aciditeateper ateRemoval ofall
H+ ionshypo iteOxoacid Oxoanion+[O][O][O]Figure 2.15 Naming
oxoacidsand oxoanions.Acid AnionHClO4 (perchloric acid) ClO42
(perchlorate)HClO3 (chloric acid) ClO32 (chlorate)HClO2 (chlorous
acid) ClO22 (chlorite)HClO (hypochlorous acid) ClO2
(hypochlorite)TABLE 2.6 Names of Oxoacids and Oxoanions That
Contain ChlorineExample 2.9 deals with the nomenclature for an
oxoacid and an oxoanion.2.7 Naming Compounds 67EXAMPLE 2.9Name the
following oxoacid and oxoanion: (a) H3PO3 and (b) IO42.Strategy To
name the acid in (a), we fi rst identify the reference acid, whose
nameends with ic, as shown in Figure 2.15 . In (b), we need to
convert the anion to itsparent acid shown in Table 2.6 .Solution
(a) We start with our reference acid, phosphoric acid (H3PO4).
BecauseH3PO3 has one fewer O atom, it is called phosphorous
acid.(b) The parent acid is HIO4. Because the acid has one more O
atom than our referenceiodic acid (HIO3), it is called periodic
acid. Therefore, the anion derived from HIO4is called
periodate.Practice Exercise Name the following oxoacid and
oxoanion: (a) HBrO and(b) HSO42.Similar problem: 2.58(f).Naming
BasesA base can be described as a substance that yields hydroxide
ions (OH2) when dissolvedin water. Some examples areNaOH sodium
hydroxideKOH potassium hydroxideBa(OH)2 barium hydroxideAmmonia
(NH3), a molecular compound in the gaseous or pure liquid state,
isalso classifi ed as a common base. At fi rst glance this may seem
to be an exceptionto the defi nition of a base. But note that as
long as a substance yields hydroxideions when dissolved in water,
it need not contain hydroxide ions in its structureto be considered
a base. In fact, when ammonia dissolves in water, NH3
reactspartially with water to yield NH41 and OH2 ions. Thus, it is
properly classifi ed asa base.HydratesHydrates are compounds that
have a specifi c number of water molecules attached tothem. For
example, in its normal state, each unit of copper(II) sulfate has
fi ve watermolecules associated with it. The systematic name for
this compound is copper(II)sulfate pentahydrate, and its formula is
written as CuSO4 ? 5H2O. The water moleculescan be driven off by
heating. When this occurs, the resulting compound is CuSO4,which is
sometimes called anhydrous copper(II) sulfate; anhydrous means that
thecompound no longer has water molecules associated with it (
Figure 2.16 ). Some otherhydrates areBaCl2 ? 2H2O barium chloride
dihydrateLiCl ? H2O lithium chloride monohydrateMgSO4 ? 7H2O
magnesium sulfate heptahydrateSr(NO3)2 ? 4H2O strontium nitrate
tetrahydrate68 Atoms, Molecules, and IonsFamiliar Inorganic
CompoundsSome compounds are better known by their common names than
by their systematicchemical names. Familiar examples are listed in
Table 2.7 .2.8 Introduction to Organic CompoundsThe simplest type
of organic compounds is the hydrocarbons, which contain onlycarbon
and hydrogen atoms. The hydrocarbons are used as fuels for domestic
andindustrial heating, for generating electricity and powering
internal combustion engines,and as starting materials for the
chemical industry. One class of hydrocarbons is calledthe alkanes.
Table 2.8 shows the names, formulas, and molecular models of the fi
rstten straight-chain alkanes, in which the carbon chains have no
branches. Note that allthe names end with - ane . Starting with
C5H12, we use the Greek prefi xes in Table 2.4to indicate the
number of carbon atoms present.The chemistry of organic compounds
is largely determined by the functionalgroups, which consist of one
or a few atoms bonded in a specifi c way. For example,when an H
atom in methane is replaced by a hydroxyl group (}OH), an
aminoFigure 2.16 CuSO4 ? 5H2O (left)is blue; CuSO4 (right) is
white.Formula Common Name Systematic NameH2O Water Dihydrogen
monoxideNH3 Ammonia Trihydrogen nitrideCO2 Dry ice Solid carbon
dioxideNaCl Table salt Sodium chlorideN2O Laughing gas Dinitrogen
monoxideCaCO3 Marble, chalk, limestone Calcium carbonateCaO
Quicklime Calcium oxideCa(OH)2 Slaked lime Calcium hydroxideNaHCO3
Baking soda Sodium hydrogen carbonateNa2CO3 ? 10H2O Washing soda
Sodium carbonate decahydrateMgSO4 ? 7H2O Epsom salt Magnesium
sulfate heptahydrateMg(OH)2 Milk of magnesia Magnesium
hydroxideCaSO4 ? 2H2O Gypsum Calcium sulfate dihydrateTABLE 2.7
Common and Systematic Names of Some CompoundsCH3OHCH3NH2CH3COOHName
Formula Molecular ModelMethane CH4Ethane C2H6Propane C3H8Butane
C4H10Pentane C5H12Hexane C6H14Heptane C7H16Octane C8H18Nonane
C9H20Decane C10H22TABLE 2.8 The First Ten Straight-Chain
Alkanesgroup (}NH2), and a carboxyl group (}COOH), the following
molecules aregenerated:HH C OHHMethanolHH NH2 CHMethylamineHH C C
OHH OAcetic acid2.8 Introduction to Organic Compounds 6970 Atoms,
Molecules, and IonsThe chemical properties of these molecules can
be predicted based on the reactivityof the functional groups.
Although the nomenclature of the major classes of organiccompounds
and their properties in terms of the functional groups will not be
discusseduntil Chapter 24, we will frequently use organic compounds
as examples to illustratechemical bonding, acid-base reactions, and
other properties throughout the book.Acid, p. 65Alkali metals, p.
53Alkaline earth metals, p. 53Allotrope, p. 55Alpha (a) particles,
p. 46Alpha (a) rays, p. 46Anion, p. 54Atom, p. 43Atomic number ( Z
), p. 49Base, p. 67Beta (b) particles, p. 46Beta (b) rays, p.
46Binary compound, p. 59Cation, p. 54Chemical formula, p.
55Diatomic molecule, p. 53Electron, p. 44Empirical formula, p.
56Families, p. 51Gamma (g) rays, p. 46Groups, p. 51Halogens, p.
53Hydrate, p. 67Inorganiccompounds, p. 59Ion, p. 54Ionic compound,
p. 54Isotope, p. 49Key WordsKey Equationmass number 5 number of
protons 1 number of neutrons5 atomic number 1 number of neutrons
(2. 1)1. Modern chemistry began with Daltons atomic theory,which
states that all matter is composed of tiny, indivisibleparticles
called atoms; that all atoms of the sameelement are identical; that
compounds contain atoms ofdifferent elements combined in
whole-number ratios;and that atoms are neither created nor
destroyed inchemical reactions (the law of conservation of mass).2.
Atoms of constituent elements in a particular compoundare always
combined in the same proportions by mass(law of defi nite
proportions). When two elements cancombine to form more than one
type of compound, themasses of one element that combine with a fi
xed massof the other element are in a ratio of small whole
numbers(law of multiple proportions).3. An atom consists of a very
dense central nucleus containingprotons and neutrons, with
electrons movingabout the nucleus at a relatively large distance
from it.4. Protons are positively charged, neutrons have no
charge,and electrons are negatively charged. Protons and
neutronshave roughly the same mass, which is about 1840times
greater than the mass of an electron.5. The atomic number of an
element is the number of protonsin the nucleus of an atom of the
element; it determinesthe identity of an element. The mass number
isSummary of Facts and Conceptsthe sum of the number of protons and
the number ofneutrons in the nucleus.6. Isotopes are atoms of the
same element with the samenumber of protons but different numbers
of neutrons.7. Chemical formulas combine the symbols for the
constituentelements with whole-number subscripts to showthe type
and number of atoms contained in the smallestunit of a compound.8.
The molecular formula conveys the specifi c numberand type of atoms
combined in each molecule of a compound.The empirical formula shows
the simplest ratiosof the atoms combined in a molecule.9. Chemical
compounds are either molecular compounds(in which the smallest
units are discrete, individual molecules)or ionic compounds, which
are made of cationsand anions.10. The names of many inorganic
compounds can be deducedfrom a set of simple rules. The formulas
can bewritten from the names of the compounds.11. Organic compounds
contain carbon and elements likehydrogen, oxygen, and nitrogen.
Hydrocarbon is thesimplest type of organic compound.Media
PlayerChapter SummaryLaw of conservation ofmass, p. 43Law of defi
niteproportions, p. 43Law of multipleproportions, p. 43Mass number
( A ), p. 49Metal, p. 51Metalloid, p. 51Molecular formula, p.
55Molecule, p. 53Monatomic ion, p. 54Neutron, p. 48Noble gases, p.
53Nonmetal, p. 51Nucleus, p. 47Organic compound, p. 59Oxoacid, p.
65Oxoanion, p. 66Periods, p. 51Periodic table, p. 51Polyatomic ion,
p. 54Polyatomic molecule, p. 53Proton, p. 47Radiation, p.
44Radioactivity, p. 46Structural formula, p. 56Ternary compound, p.
60Questions and ProblemsStructure of the AtomReview Questions2.1
Defi ne the following terms: (a) a particle, (b) b particle,(c) g
ray, (d) X ray.2.2 Name the types of radiation known to be emitted
byradioactive elements.2.3 Compare the properties of the following:
a particles,cathode rays, protons, neutrons, electrons.2.4 What is
meant by the term fundamental particle?2.5 Describe the
contributions of the following scientiststo our knowledge of atomic
structure: J. J. Thomson,R. A. Millikan, Ernest Rutherford, James
Chadwick.2.6 Describe the experimental basis for believing that
thenucleus occupies a very small fraction of the volumeof the
atom.Problems2.7 The diameter of a helium atom is about 1 3 102
pm.Suppose that we could line up helium atoms side byside in
contact with one another. Approximately howmany atoms would it take
to make the distance fromend to end 1 cm?2.8 Roughly speaking, the
radius of an atom is about10,000 times greater than that of its
nucleus. If anatom were magnifi ed so that the radius of its
nucleusbecame 2.0 cm, about the size of a marble, what wouldbe the
radius of the atom in miles? (1 mi 5 1609 m.)Atomic Number, Mass
Number, and IsotopesReview Questions2.9 Use the helium-4 isotope to
defi ne atomic number andmass number. Why does a knowledge of
atomic numberenable us to deduce the number of electrons presentin
an atom?2.10 Why do all atoms of an element have the same atomicnum
ber, although they may have different massnumbers?2.11 What do we
call atoms of the same elements withdifferent mass numbers?2.12
Explain the meaning of each term in the symbolZAX.Problems2.13 What
is the mass number of an iron atom that has28 neutrons?2.14
Calculate the number of neutrons of 239Pu.2.15 For each of the
following species, determine the numberof protons and the number of
neutrons in the nucleus:2 3He, 24 He, 1224Mg, 2512Mg, 4822Ti,
7935Br, 19578Pt2.16 Indicate the number of protons, neutrons, and
electronsin each of the following species:157N, 3316S, 6329Cu,
8438Sr, 13056Ba, 18674W, 20280Hg2.17 Write the appropriate symbol
for each of the followingisotopes: (a) Z 5 11, A 5 23; (b) Z 5 28,
A 5 64.Questions and Problems 71Electronic Homework ProblemsThe
following problems are available at www.aris.mhhe.com if assigned
by your instructor as electronic homework.Quantum Tutor problems
are also available at the same site.ARIS Problems: 2.13, 2.15,
2.22, 2.32, 2.35, 2.36,2.43, 2.44, 2.46, 2.48, 2.49, 2.50, 2.58,
2.59, 2.60, 2.63,2.65, 2.77, 2.90, 2.91, 2.96, 2.97, 2.100, 2.101,
2.102.Quantum Tutor Problems: 2.43, 2.44, 2.45, 2.46,2.57, 2.58,
2.59, 2.60.72 Atoms, Molecules, and Ions2.18 Write the appropriate
symbol for each of the followingisotopes: (a) Z 5 74, A 5 186; (b)
Z 5 80; A 5 201.The Periodic TableReview Questions2.19 What is the
periodic table, and what is its signifi cancein the study of
chemistry?2.20 State two differences between a metal and a
nonmetal.2.21 Write the names and symbols for four elements in
eachof the following categories: (a) nonmetal, (b) metal,(c)
metalloid.2.22 Defi ne, with two examples, the following terms:(a)
alkali metals, (b) alkaline earth metals, (c) halogens,(d) noble
gases.Problems2.23 Elements whose names end with ium are usually
metals;sodium is one example. Identify a nonmetal whosename also
ends with ium.2.24 Describe the changes in properties (from metals
tononmetals or from nonmetals to metals) as we move(a) down a
periodic group and (b) across the periodictable from left to
right.2.25 Consult a handbook of chemical and physical data(ask
your instructor where you can locate a copy ofthe handbook) to fi
nd (a) two metals less dense thanwater, (b) two metals more dense
than mercury,(c) the densest known solid metallic element, (d)
thedensest known solid nonmetallic element.2.26 Group the following
elements in pairs that you wouldexpect to show similar chemical
properties: K, F, P,Na, Cl, and N.Molecules and IonsReview
Questions2.27 What is the difference between an atom and
amolecule?2.28 What are allotropes? Give an example. How
areallotropes different from isotopes?2.29 Describe the two
commonly used molecular models.2.30 Give an example of each of the
following: (a) a monatomiccation, (b) a monatomic anion, (c) a
polyatomiccation, (d) a polyatomic anion.Problems2.31 Which of the
following diagrams represent diatomicmolecules, polyatomic
molecules, molecules that arenot compounds, molecules that are
compounds, or anelemental form of the substance?2.32 Which of the
following diagrams represent diatomicmolecules, polyatomic
molecules, molecules that arenot compounds, molecules that are
compounds, or anelemental form of the substance?(a) (b) (c)(a) (b)
(c)2.33 Identify the following as elements or compounds:NH3, N2,
S8, NO, CO, CO2, H2, SO2.2.34 Give two examples of each of the
following: (a) adiatomic molecule containing atoms of the
sameelement, (b) a diatomic molecule containing atoms ofdifferent
elements, (c) a polyatomic molecule containingatoms of the same
element, (d) a polyatomic moleculecontaining atoms of different
elements.2.35 Give the number of protons and electrons in each
ofthe following common ions: Na1 , Ca21, Al31, Fe21,I2, F2, S22,
O22, and N32.2.36 Give the number of protons and electrons in each
ofthe following common ions: K1, Mg21, Fe31, Br2,Mn21, C42,
Cu21.Chemical FormulasReview Questions2.37 What does a chemical
formula represent? What is theratio of the atoms in the following
molecular formulas?(a) NO, (b) NCl3, (c) N2O4, (d) P4O62.38 Defi ne
molecular formula and empirical formula.What are the similarities
and differences between2.49 Which of the following compounds are
likely to beionic? Which are likely to be molecular? SiCl4,
LiF,BaCl2, B2H6, KCl, C2H42.50 Which of the following compounds are
likely to beionic? Which are likely to be molecular? CH4,
NaBr,BaF2, CCl4, ICl, CsCl, NF3Naming Inorganic CompoundsReview
Questions2.51 What is the difference between inorganic compoundsand
organic compounds?2.52 What are the four major categories of
inorganiccompounds?2.53 Give an example each for a binary compound
and aternary compound.2.54 What is the Stock system? What are its
advantagesover the older system of naming cations?2.55 Explain why
the formula HCl can represent two differentchemical systems.2.56
Defi ne the following terms: acids, bases, oxoacids,oxoanions, and
hydrates.Problems2.57 Name these compounds: (a) Na2CrO4, (b)
K2HPO4,(c) HBr (gas), (d) HBr (in water), (e) Li2CO3,(f) K2Cr2O7,
(g) NH4NO2, (h) PF3, (i) PF5, (j) P4O6,(k) CdI2, (l) SrSO4, (m)
Al(OH)3, (n) Na2CO3 ? 10H2O.2.58 Name these compounds: (a) KClO,
(b) Ag2CO3,(c) FeCl2, (d) KMnO4, (e) CsClO3, (f) HIO, (g) FeO,(h)
Fe2O3, (i) TiCl4, (j) NaH, (k) Li3N, (l) Na2O,(m) Na2O2, (n) FeCl3
? 6H2O.2.59 Write the formulas for the following compounds:(a)
rubidium nitrite, (b) potassium sulfi de, (c) sodiumhydrogen sulfi
de, (d) magnesium phosphate, (e) calciumhydrogen phosphate, (f)
potassium dihydrogen phosphate,(g) iodine heptafl uoride, (h)
ammonium sulfate,(i) silver perchlorate, (j) boron trichloride.2.60
Write the formulas for the following compounds:(a) copper(I)
cyanide, (b) strontium chlorite, (c) perbromicacid, (d) hydroiodic
acid, (e) disodium ammoniumphosphate, (f) lead(II) carbonate, (g)
tin(II) fl uoride,(h) tetraphosphorus decasulfi de, (i) mercury(II)
oxide,(j) mercury(I) iodide, (k) selenium hexafl uoride.Questions
and Problems 73HCNOthe empirical formula and molecular formula of
acompound?2.39 Give an example of a case in which two moleculeshave
different molecular formulas but the same empiricalformula.2.40
What does P4 signify? How does it differ from 4P?2.41 What is an
ionic compound? How is electrical neutralitymaintained in an ionic
compound?2.42 Explain why the chemical formulas of ionic
compoundsare usually the same as their empirical
formulas.Problems2.43 Write the formulas for the following ionic
compounds:(a) sodium oxide, (b) iron sulfi de (containingthe Fe21
ion), (c) cobalt sulfate (containing theCo31 and SO422 ions), and
(d) barium fl uoride. ( Hint:See Figure 2.11 .)2.44 Write the
formulas for the following ionic compounds:(a) copper bromide
(containing the Cu1 ion), (b) manganeseoxide (containing the Mn31
ion), (c) mercuryiodide (containing the Hg221 ion), and (d)
magnesiumphosphate (containing the PO432 ion). ( Hint: See
Figure2.11 .)2.45 What are the empirical formulas of the
followingcompounds? (a) C2N2, (b) C6H6, (c) C9H20, (d) P4O10,(e)
B2H62.46 What are the empirical formulas of the followingcompounds?
(a) Al2Br6, (b) Na2S2O4, (c) N2O5,(d) K2Cr2O72.47 Write the
molecular formula of glycine, an aminoacid present in proteins. The
color codes are: black(carbon), blue (nitrogen), red (oxygen), and
gray(hydrogen).HCO2.48 Write the molecular formula of ethanol. The
colorcodes are: black (carbon), red (oxygen), and gray(hydrogen).74
Atoms, Molecules, and IonsAdditional Problems2.61 A sample of a
uranium compound is found to be losingmass gradually. Explain what
is happening to thesample.2.62 In which one of the following pairs
do the two speciesresemble each other most closely in chemical
properties?Explain. (a) 11H and 11H1, (b) 147N and 147N32,(c) 126C
and 136C.2.63 One isotope of a metallic element has mass number
65and 35 neutrons in the nucleus. The cation derivedfrom the
isotope has 28 electrons. Write the symbolfor this cation.2.64 One
isotope of a nonmetallic element has mass number127 and 74 neutrons
in the nucleus. The anionderived from the isotope has 54 electrons.
Write thesymbol for this anion.2.65 Determine the molecular and
empirical formulas ofthe compounds shown here. (Black spheres are
carbonand gray spheres are hydrogen.)2.71 Explain why anions are
always larger than the atomsfrom which they are derived, whereas
cations arealways smaller than the atoms from which they
arederived. ( Hint: Consider the electrostatic attractionbetween
protons and electrons.)2.72 (a) Describe Rutherfords experiment and
how it led tothe structure of the atom. How was he able to
estimatethe number of protons in a nucleus from the scatteringof
the a particles? (b) Consider the 23Na atom. Giventhat the radius
and mass of the nucleus are3.04 3 10215 m and 3.82 3 10223 g,
respectively,calculate the density of the nucleus in g/cm3. The
radiusof a 23Na atom is 186 pm. Calculate the density of thespace
occupied by the electrons in the sodium atom.Do your results
support Rutherfords model of an atom?[The volume of a sphere of
radius r is (4/3)pr3.]2.73 What is wrong with the name (in
parentheses) for eachof the following compounds: (a) BaCl2 (barium
dichloride),(b) Fe2O3 [iron(II) oxide], (c) CsNO2 (cesiumnitrate),
(d) Mg(HCO3)2 [magnesium(II) bicarbonate]?2.74 What is wrong with
the chemical formula for each ofthe following compounds: (a)
(NH3)2CO3 (ammoniumcarbonate), (b) CaOH (calcium hydroxide), (c)
CdSO3(cadmium sulfi de), (d) ZnCrO4 (zinc dichromate)?2.75 Fill in
the blanks in the following table:Symbol 5426Fe21Protons 5 79
86Neutrons 6 16 117 136Electrons 5 18 79Net charge 23 02.76 (a)
Which elements are most likely to form ionic compounds?(b) Which
metallic elements are most likelyto form cations with different
charges?2.77 Write the formula of the common ion derived fromeach
of the following: (a) Li, (b) S, (c) I, (d) N, (e) Al,(f) Cs, (g)
Mg2.78 Which of the following symbols provides more
informationabout the atom: 23Na or 11Na? Explain.2.79 Write the
chemical formulas and names of binary acidsand oxoacids that
contain Group 7A elements. Do thesame for elements in Groups 3A,
4A, 5A, and 6A.2.80 Of the 117 elements known, only two are liquids
atroom temperature (25C). What are they? ( Hint: Oneelement is a
familiar metal and the other element is inGroup 7A.)2.81 For the
noble gases (the Group 8A elements), 42He,2010Ne, 4018Ar, 8436Kr,
and 13254Xe, (a) determine the numberof protons and neutrons in the
nucleus of each atom,and (b) determine the ratio of neutrons to
protons inthe nucleus of each atom. Describe any general trend(a)
(b) (c) (d)2.66 What is wrong with or ambiguous about the
phrasefour molecules of NaCl?2.67 The following phosphorus sulfi
des are known: P4S3,P4S7, and P4S10. Do these compounds obey the
law ofmultiple proportions?2.68 Which of the following are
elements, which are moleculesbut not compounds, which are compounds
butnot molecules, and which are both compounds andmolecules? (a)
SO2, (b) S8, (c) Cs, (d) N2O5, (e) O,(f) O2, (g) O3, (h) CH4, (i)
KBr, (j) S, (k) P4, (l) LiF2.69 The following table gives numbers
of electrons, protons,and neutrons in atoms or ions of a numbe