Chapter 02
Atoms, Molecules, and Ions
Textbook
Brown, LeMay, Bursten, Murphy, Woodward, Stoltzfus,
“Chemistry, The Central Science”, 15th Ed. (Pearson, 2018)
This electronic presentation is used
with Brown, LeMay, Bursten, Murphy, Woodward, Stoltzfus
“Chemistry, The Central Science”, 14th Ed.
All rights reserved and reproduced by permission.
Text / images may not be modified or reproduced
in any way without prior written permission of the publisher.
https://www.pearson.com/
2
CONTENTS
2.1 The Atomic Theory of Matter
2.2 The Discovery of Atomic Structure
2.3 The Modern View of Atomic Structure
2.4 Atomic Weights
2.5 The Periodic Table
2.6 Molecules and Molecular Compounds
2.7 Ions and Ionic Compounds
2.8 Naming Inorganic Compounds
2.9 Some Simple Organic Compounds
2.1 The Atomic Theory of Matter
• Some Greek philosophers like Democritus described that there was a
smallest particle – “atomos” ( indivisible, uncuttable ) – that made up
all of nature.
Tomas is to cut to divide and a is the pre–syllable saying not.
• Experiments in the
and
centuries led to an organized atomic
theory by John Dalton in the early 1800s, which explained several
laws known at that time:
th
18
th
19
– The Law of Constant Composition
– The Law of Conservation of Mass
– The Law of Multiple Proportions
Democritus
c.460 – c.370
4
The Law of Constant Composition
• Joseph Proust (FRA, 1754–1826)
• “In a given chemical compound, the proportions
by mass of the elements that compose it are fixed,
independent of the origin of the compound or
its mode of preparation.”
• Known as the Law of Definite Proportions.
• In a given compound, the relative numbers and kinds of atoms are
constant.
• Basis of Dalton’s Postulate # 4
1
:
8
:
9
5
The Law of Conservation of Mass ( Matter )
(질량보존의 법칙)
• Antoine Laurent Lavoisier (FRA, 1743–1794)
: Father of modern chemistry
• “In every chemical operation an equal quantity
of matter exists before and after the operation.”
• Can’t create matter in a chemical reaction!
• Basis of Dalton’s Postulate # 3
6
Law of Multiple Proportions (배수비례의 법칙)
• John Dalton (UK, 1766–1844)
• “When two elements form a series of compounds,
the masses of one element that combine with
a fixed mass of the other element are in the ratio
of small integers to each other.”
7
Postulates of Dalton’s Atomic Theory
A New System of Chemical Philosophy (1808)
1) Each element is composed of extremely small particles called atoms.
2) All atoms of a given element are identical to one another in mass
and other properties, but the atoms of one element are different from
the atoms of all other elements.
3) Atoms of an element are not changed into atoms of a different
element by chemical reactions; atoms are neither created nor
destroyed in chemical reactions. → The Law of Conservation of Mass
4) Atoms of more than one element combine to form compounds; a
given compound always has the same relative number and kind of
atoms. → The Law of Constant Composition
8
Postulates of Dalton’s Atomic Theory
9
Limitations of Dalton’s Atomic Theory
1) It does not account for subatomic particles.
( such as protons, electrons, and neutrons )
2) It does not account for isotopes.
1
2
3
( Example: H hydrogen, H deuterium, and H tritium ).
3) It does not account for isobars.
40
40
( Example: Ar and Ca )
4) Elements need not combine in simple, whole-number ratios to form
compounds.
( Example: berthollides, Fe0.85O, sugar/sucrose, C11H22O11 )
5) The theory does not account for allotropes.
( Example: diamond and graphite )
10
2.2 The Discovery of Atomic Structure
Discovery of Subatomic Particles
• In Dalton’s view, the atom was the smallest particle possible.
• Many discoveries → The atom itself was made up of smaller particles.
– Cathode rays and electrons
– Radioactivity
– Nucleus, protons, and neutrons
11
Cathod Rays and Electrons
• Mid 1800’s – scientists studied electrical discharge through partially
evacuated tubes.
• Streams of negatively charged particles were found to emanate from
cathode tubes, causing fluorescence.
• J. J. Thomson is credited
with their discovery (1897).
12
Discovery of Electrons by J. J. Thomson (1897)
• Sir Joseph John Thomson (UK, 1856–1940)
• Proposed that Cathode Rays were actually
particles (negatively charged) that we now know
are electrons.
Thomson measured
the charge/mass ratio of
8
the electron to be 1.76  10
coulombs/gram (C/g).
13
14
Charge of the Electron by Robert Millikan (1909)
• Robert Millikan (US, 1866–1953)
• Millikan Oil‒Drop Experiment
‒ Measured the charges of the oil drops
‒ Calculated the greatest common divisor (최대공약수)
→ Charge of an electron
Robert Millikan
determined the charge
on the electron, equal to
1.602  10–19 C.
Mass of an e– could
be calculated as
9.11  10–28 g.
15
Millikan's oil drop experiment to determine charge of an electron
https://youtu.be/UFiPWv03f6g
16
Radioactivity
Radioactivity is the spontaneous emission of high‒energy radiation by
an atom. Its discovery showed that the atom had more subatomic
particles and energy associated with it.
Discovery of
Po & Ra
Antoine Henri Becquerel
(FRA,1852–1908)
“for his discovery of
spontaneous radioactivity"
Pierre Curie
(FRA,1859–1906)
Marie Sklodowska Curie
(FRA/POL,1867–1934)
“for their joint researches on the radiation phenomena
discovered by Professor Henri Becquerel"
17
Radioactivity
• Three types of radiation were discovered by Ernest Rutherford:
– α particles (positively charged, +2), large mass
– β particles (negatively charged, like electrons, –1), small mass
– γ rays (uncharged), no mass
18
The Nuclear Model of the Atom
The Atom, circa 1900
• The prevailing theory was that of
the “plum pudding” model, put forward
by J. J. Thomson.
• Positive sphere of matter with negative
electrons embedded in it.
19
Discovery of the Nucleus
• Rutherford shot α particles at a thin sheet of gold foil and observed
the pattern of scatter of the particles.
Ernest Rutherford
(NZ/UK, 1871–1937)
Nobel Laureate (Chemistry,1908)
"for his investigations into the
disintegration of the elements,
and the chemistry of
radioactive substances"
20
• Discovery of the Nucleus
21
• The Nuclear Atom
– Since some particles were deflected at large
angles, Thomson’s model could not be correct.
– Rutherford’s nuclear model of the atom: all of the
positive charge and most of the mass is
concentrated at the center – the nucleus.
Electrons occupy the rest of the space (volume)
of the atom.
– Atoms are very small;
1–5 Å or 100–500 pm.
– Other subatomic particles (protons and neutrons
in the nucleus) were discovered.
22
The History of Atomic Chemistry
https://youtu.be/thnDxFdkzZs
23
2.3 The Modern View of Atomic Structure
Subatomic Particles
electronic charge: 1.602 ×
–19
10
C
• Protons (+1) and electrons (–1) have a charge; neutrons are neutral.
• Protons and neutrons have essentially the same mass (relative mass 1).
The mass of an electron is so small we ignore it (relative mass 0).
• Protons and neutrons are found in the nucleus; electrons travel around
the nucleus.
Atomic Mass Unit (amu)
1 amu is defined as 1/12 the mass of an unbound carbon atom carbon–12 at its rest state.
24
Atomic Numbers, Mass Numbers, and Isotopes
Atomic and Mass Numbers
• How do we determine which element an atom is?
• Atomic Number: the number of protons
in the nucleus of an atom.
• Mass Number: the total number of protons and neutrons
in the nucleus of an atom
• Since atoms have no overall charge, the number of protons equals
the number of electrons in an atom.
25
Atoms of an Element
• Elements are represented by a one or two letter symbol, for which
the first letter is always capitalized. C is the symbol for carbon.
• All atoms of the same element have the same number of protons,
which is called the atomic number. It is written as a subscript
BEFORE the symbol. 6 is the atomic number for carbon.
• The mass number is the total number of protons and neutrons
in the nucleus of an atom. It is written as a superscript BEFORE the
symbol.
26
Isotopes
• Isotopes are atoms of the same element with different masses.
• Isotopes have different numbers of neutrons, but the same number of
protons.
• The table below lists four isotopes for carbon.
11
6
C
12
6
C
13
6
C
14
6
C
27
2.4 Atomic Weights
The Atomic Mass Scale
Atomic Mass Unit (amu, u) → relative atomic mass
• The mass of an atom in atomic mass units (amu) is the total number
of protons and neutrons in the atom.
• The atomic mass unit is presently defined by assigning a mass of
12
exactly 12 u to a chemically unbound atom of the C isotope of carbon.
• A carbon (C) atom with 6 protons and 6 neutrons is assigned a mass of
exactly 12 amu. Why carbon? 12C pure isotope in solid, 98.93%
• A mass scale on the atomic level is used, where an atomic mass unit
(amu) is the base unit.
1
235
H
/
1.007825
amu,
U
/
235.04393
amu
1
92
1 amu = 1.66054 ×
–27
10
kg / 1 kg = 6.02214 ×
26
10
amu
28
Atomic Weight
• Isotopic mass is the mass in amu (u), of a particular isotope of an
element.
• Different isotopes of an element all react essentially the same, so a
weighted average of isotopic masses can be used in calculations.
→ Most elements occur in nature as mixtures of isotopes.
• The atomic weight is the weighted average mass, of the naturally
occurring element. It is calculated from the isotopes of an element
weighted by their relative abundances.
• Atomic Weight = Ʃ [(isotope mass) × (fractional natural abundance)]
for ALL isotopes.
• Average atomic mass is known as the atomic weight.
29
Atomic Weight Measurement
• Atomic and molecular masses (actually the mass to charge ratio) can
be measured with great accuracy with a mass spectrometer.
• Mass spectrum: A graph of the intensity of the detector signal versus
ion atomic mass.
→ The masses of the ions and their relative abundances
30
2.5 The Periodic Table
the royal society of chemistry
https://www.rsc.org/periodic-table/
• The periodic table is a systematic
organization of the elements.
• Elements are arranged in order of
IUPAC Periodic Table of the Elements
atomic number.
[ Periods ]
horizontal rows
[ Groups ]
Vertical columns
Containing elements with
similar chemical properties
Elements arranged in
order of increasing
atomic number
31
Periodicity
• A repeating pattern of chemical and physical properties is observed.
• Law of chemical periodicity: the properties of the elements are
periodic functions of atomic number.
32
The Periodic Table
• The arrangement of elements
in order of increasing atomic
number, with elements having
similar properties placed in
vertical columns, is known as the
periodic table.
• The rows on the periodic table
are called periods.
• Columns are called groups.
→ Elements in the same group
have similar chemical and
physical properties.
– A groups: the main group elements
– B groups: the transition elements
33
The Periodic Table: Metals & Nonmetals
• Metals are on the left side of the
periodic table.
– Shiny luster, ductile, malleable
– Conducting heat and electricity
– Solids (except mercury)
• Nonmetals are on the right side
of the periodic table (with the
exception of H).
– Having a wide variety of
properties (solids, liquids and
gases)
– Not conducting electricity well
(except C as graphite).
34
The Periodic Table: Metalloids
• Metalloids or semimetals: some
physical characteristics of metals
and chemical characteristics of
nonmetal.
• Metalloids border the stair–step
line (with the exception of Al and
Po, and At).
35
2.6 Molecules and Molecular Compounds
Molecules and Chemical Formulas
Chemical Formulas
• The subscript to the right of the symbol of an element
tells the number of atoms of that element in one
molecule of the compound.
• Molecular compounds are composed of molecules
and almost always contain only nonmetals.
• The attraction between molecules are often relatively
weak, explaining why gases and liquids are common
among molecular substances.
• Carbon is typically listed first in the formula, unless
C is part of a polyatomic ion.
36
Diatomic Molecules
• These seven elements occur naturally as molecules containing two atoms:
– Hydrogen (H2)
– Nitrogen (N2)
– Oxygen (O2)
– Fluorine (F2)
– Chlorine (Cl2)
– Bromine (Br2)
– Iodine (I2)
cf. O3 (ozone)
37
Molecular and Empirical Formulas / Picturing Molecules
• Empirical formulas (실험식) give the lowest integer ratio of atoms of
each element in a compound. Ethane (C2H6) / EF C1H3 / MF C1H3, C2H6, C3H9
• Molecular formulas (분자식) give the exact number of atoms of each
element in a compound.
• Structural formulas show the order in which atoms are attached.
• Perspective drawings, ball‒and‒stick models, and space‒filling
models show the three‒dimensional order of the atoms.
38
2.7 Ions and Ionic Compounds
Ions
• When atoms lose or gain electrons, they become ions.
• Cations are positive and are formed when at least one electron is lost.
Monatomic cations are formed by metals.
• Anions are negative and are formed when at least one electron is
gained. Monatomic anions are formed by nonmetals, except the noble
gases.
• Sometimes a group of atoms
will gain or lose electrons.
These are polyatomic ions.
+
ex) NH4 , SO4
2–
39
Ioninc Compounds
• Ionic compounds (such as NaCl) are generally formed between
metals and nonmetals.
• Electrons are transferred from the metal to the nonmetal. The
oppositely charged ions attract each other. Only empirical formulas
are written.
40
Writing Formulas cation : magnesium, Mg2+ / anion : nitride, N3‒
To get a neutral salt, 6 electrons are required
as the greatest common factor.
• Because compounds are electrically neutral, one can determine the
formula of a compound this way:
– The charge on the cation becomes the subscript on the anion.
– The charge on the anion becomes the subscript on the cation.
– If these subscripts are not in the lowest whole‒number ratio,
divide them by the greatest common factor.
41
2.8 Naming Inorganic Compounds
Chemical Nomenclature
• The system of naming compounds is called chemical nomenclature.
• We will learn how to name:
1) Ionic compounds
2) Acids
3) Binary Molecular Compounds
4) Simple Organic Compounds
– Alkanes
– Alcohols
42
Names and Formulas of Ionic Compounds
Cations
1. Cations formed from metal atoms have the same name as the metal.
2. If a metal can form cations with different charges, the positive charge
is indicated by a Roman numeral in parentheses following the name
of the metal.
3. Cations formed from molecules composed of nonmetal atoms have
names that end in –ium:
43
Anions
1. The names of monatomic anions are formed by replacing the ending
of the name of the element with –ide.
2. Polyatomic anions containing oxygen have names ending in either
–ate or –ite and are called oxyanions.
+
H
3. Anions derived by adding
to an oxyanion are named by adding as
a prefix the word hydrogen or dihydrogen.
44
Patterns in Oxyanion Nomenclature
• Central atoms on the second row have a bond to, at most, three
oxygens; those on the third row take up to four.
• Charges increase as you go from right to left.
45
Patterns in Oxyanion Nomenclature
• The one with the second fewest oxygens ends in ‒ite.
• The one with the second most oxygens ends in ‒ate.
• The one with the fewest oxygens has the prefix hypo‒ and ends in ‒ite.
• The one with the most oxygens has the prefix per‒ and ends in ‒ate.
46
Common Cations & Anions
• The element’s symbol is followed by a superscript number and a sign that shows the charge on
the ion in electron charge units.
• If the ionic charge is one unit, the number is often omitted, e.g. Na+ is the symbol for a sodium ion.
47
Ionic Compounds
• Write the name of the cation.
• If the anion is an element, change its ending to –ide; if the anion is a
polyatomic ion, simply write the name of the polyatomic ion.
• If the cation can have more than one possible charge (e.g. iron), write
the charge as a Roman numeral in parentheses.
– First, write the name of the cation.
magnesium
Magnesium nitride
– Second, if the anion is an element,
change its ending to –ide.
nitrogen → nitride
48
Name and Formulas of Acids
How to Name Acids
• If the anion in the acid ends in ‒ide, change the ending to ‒ic acid
and add the prefix hydro‒.
– HCl: hydrochloric acid
– HBr: hydrobromic acid
– HI: hydroiodic acid
• If the anion ends in ‒ite,
change the ending to ‒ous acid.
– HClO: hypochlorous acid
– HClO2: chlorous acid
• If the anion ends in ‒ate,
change the ending to ‒ic acid.
– HClO3: chloric acid
– HClO4: perchloric acid
49
Name and Formulas of Binary Molecular Compounds
How to Name Binary Molecular Compounds
• The less electronegative atom (farthest to left) is usually listed first.
• A Greek prefix is used to denote the number of
atoms of each element in the compound.
(mono– is not used on the first element listed.)
• The ending on the more electronegative
element is changed to –ide.
– CCl4: carbon tetrachloride
– CO2: carbon dioxide
50
Name and Formulas of Binary Molecular Compounds
• In nonmetal compounds containing carbon, the C atom is listed first in
the chemical formula, followed by H.
• If the prefix ends with a or o and the name of
the element begins with a vowel, the two
successive vowels are often elided into one:
– CO: carbon monoxide
– N2O4: dinitrogen tetroxide
51
2.9 Some Simple Organic Compounds
Organic Chemistry
• Compounds that contain carbon and hydrogen, often in combination
with oxygen, nitrogen, or other elements.
Alkanes
• The simplest hydrocarbons (compounds containing only carbon and
hydrogen) are alkanes.
• The first part of the names just listed correspond to the number of
carbons (meth‒ = 1, eth‒ = 2, prop‒ = 3, etc.).
• It is followed by ‒ane.
52
Some Derivatives of Alkanes
Alcohols
• When a hydrogen in an alkane is replaced with something else
(a functional group, like –OH in the compounds above),
the name is derived from the name of the alkane.
• The ending denotes the type of compound.
• An alcohol ends in ‒ol.
53
Structural Isomers
• When two or more molecules have the same
chemical formula, but different structures,
they are called isomers.
• 1‒propanol and 2‒propanol have the oxygen
atom connected to different carbon atoms,
but both have the formula C3H8O.
54
What’s Ahead
• 2.1 The Atomic Theory of Matter
Learn how scientists were able to postulate that atoms
are the smallest building block of matter, long before
they could be seen directly.
• 2.2 The Discovery of Atomic Structure
Examine key experiments that led to discovery
of electrons and to the nuclear model of the atom.
• 2.3 The Modern View of Atomic Structure
Learn how atomic number and mass number
can be used to express the number of
each subatomic particle – protons,
neutrons, and electrons – in given atoms.
55
What’s Ahead
• 2.4 Atomic Weights
Learn about the concept of atomic weights and
how they are derived from the masses and
abundances of individual atoms.
• 2.5 The Periodic Table
Examine the organization of the periodic table, in which
elements are put in order of increasing atomic
number and grouped by chemical similarity.
• 2.6 Molecules and Molecular Compounds
Explore the assemblies of atoms called molecules
and learn how their compositions are
represented by empirical and molecular formulas.
56
What’s Ahead
• 2.7 Ions and Ionic Compounds
Realize that atoms can gain or lose electrons to form ions
and learn how to use the periodic table to predict the
charges on ions as well as the empirical formulas
of ionic compounds.
• 2.8 Naming Inorganic Compounds
Consider the systematic way in which substances are
named, called nomenclature, and how this
nomenclature is applied to inorganic compounds.
• 2.9 Some Single Organic Compounds
Become familiar with simple families of organic
compounds, compounds containing carbon and hydrogen,
and the very basic nomenclature of these compounds.
57
Learning Outcomes
[2.1] List the basic postulates of Dalton’s atomic theory.
[2.2] Describe the key experiments that led to the discovery of
electrons and to the nuclear model of the atom.
[2.3] Describe the structure of the atom in terms of protons,
neutrons, and electrons and express the relative electrical
charges and masses of these subatomic particles.
[2.3] Use chemical symbols together with atomic
number and mass number to express
the subatomic composition of isotopes.
[2.4] Calculate the atomic weight of an element
from the masses of individual atoms and
acknowledge of natural abundance.
58
Learning Outcomes
[2.5] Describe how elements are organized in the periodic table
by atomic number and by similarities in chemical behavior,
giving rise to periods and groups.
[2.5] Identify the locations of metals and nonmetals
in the periodic table.
[2.6, 2.7] Distinguish between molecular substances and
ionic substances in terms of their composition.
[2.6] Distinguish between empirical formulas and
molecular formulas.
[2.6] Describe how molecular formulas and
structural formulas are used to represent
the compositions of molecules.
59
Learning Outcomes
[2.7] Explain how ions are formed by the gain or loss of
electrons and use the periodic table to predict
the charges of common ions.
[2.7] Write the empirical formulas of ionic compounds,
given the charges of their component ions.
[2.8] Write the name of an ionic compound
given its chemical formula or write
the chemical formula given its name.
[2.8] Name or write chemical formulas
for binary inorganic compounds and for acid.
[2.9] Identify organic compounds and
name simple alkanes and alcohols.
60
Sample Exercise 2.1
Atomic Size
The diameter of a U.S. dime is 17.9 mm, and the diameter of a silver atom is 2.88 Å . How many
silver atoms could be arranged side by side across the diameter of a dime?
Solution
The unknown is the number of silver (Ag) atoms. Using the relationship 1 Ag atom = 2.88 Å as a
conversion factor relating number of atoms and distance, we start with the diameter of the dime,
first converting this distance into angstroms and then using the diameter of the Ag atom to
convert distance to number of Ag atoms:
That is, 62.2 million silver atoms could sit side by side across a dime!
61
Sample Exercise 2.1
Atomic Size
Continued
Practice Exercise 1
Which of the following factors determines the size of an atom?
(a) the volume of the nucleus
(b) the volume of space occupied by the electrons of the atom
(c) the volume of a single electron, multiplied by the number of electrons in the atom
(d) the total nuclear charge
(e) the total mass of the electrons surrounding the nucleus.
Practice Exercise 2
The diameter of a carbon atom is 1.54 Å .
(a) Express this diameter in picometers.
(b) How many carbon atoms could be aligned side by side across the width of a pencil line that is
0.20 mm wide?
62
Sample Exercise 2.2
Determining the Number of
Subatomic Particles in Atoms
How many protons, neutrons, and electrons are in an atom of (a) 197Au; (b) strontium‒90?
Solution
(a) The superscript 197 is the mass number (protons + neutrons). According to the list of
elements given on the front inside cover, gold has atomic number 79. Consequently, an atom
of 197Au has 79 protons, 79 electrons, and 197 – 79 = 118 neutrons.
(b) The atomic number of strontium is 38. Thus, all atoms of this element have 38 protons and 38
electrons. The strontium‒90 isotope has 90 – 38 =52 neutrons.
Practice Exercise 1
Which of these atoms has the largest number of neutrons?
(a) 148Eu
(b) 157Dy
(c) 149Nd
(d) 162Ho
(e) 159Gd.
Practice Exercise 2
How many protons, neutrons, and electrons are in an atom of
(a) 138Ba; (b) phosphorus‒31?
63
Writing Symbols for Atoms
Sample Exercise 2.3
Magnesium has three isotopes with mass numbers 24, 25, and 26.
(a) Write the complete chemical symbol (superscript and subscript) for each.
(b) How many neutrons are in an atom of each isotope?
Solution
(a) Magnesium has atomic number 12, so all atoms of magnesium contain 12 protons and 12
electrons. The three isotopes are therefore represented by
,
, and
.
(b) The number of neutrons in each isotope is the mass number minus the number of protons.
The numbers of neutrons in an atom of each isotope are therefore 12, 13, and 14,
respectively.
Practice Exercise 1
Which of the following is an incorrect representation for a neutral atom?
(a)
(b)
(c)
(d)
(e)
.
Practice Exercise 2
Give the complete chemical symbol for the atom that contains 82 protons, 82 electrons, and 126
neutrons.
64
Sample Exercise 2.4
Calculating the Atomic Weight of
an Element from Isotopic Abundances
Naturally occurring chlorine is 75.78% 35Cl (atomic mass 34.969 amu) and 24.22% 37Cl (atomic
mass 36.966 amu). Calculate the atomic weight of chlorine.
Solution
We can calculate the atomic weight by multiplying the abundance of each isotope by its mass and
summing these products. Because 75.78% = 0.7578 and 24.22% = 0.2422, we have
This answer makes sense: The atomic weight, which is actually the average atomic mass, is
between the masses of the two isotopes and is closer to the value of 35Cl, the more abundant
isotope.
65
Sample Exercise 2.4
Calculating the Atomic Weight of
an Element from Isotopic Abundances
Continued
Practice Exercise 1
There are two stable isotopes of copper found in nature, 63Cu and 65Cu. If the atomic weight of
copper Cu is 63.546 amu, which of the following statements are true?
(a) 65Cu contains two more protons than 63Cu.
(b) 63Cu must be more abundant than 65Cu.
(c) All copper atoms have a mass of 63.546 amu.
Practice Exercise 2
Three isotopes of silicon occur in nature: 28Si (92.23%), atomic mass 27.97693 amu; 29Si
(4.68%), atomic mass 28.97649 amu; and 30Si (3.09%), atomic mass 29.97377 amu. Calculate
the atomic weight of silicon.
66
Sample Exercise 2.5
Using the Periodic Table
Which two of these elements would you expect to show the greatest similarity in chemical and
physical properties: B, Ca, F, He, Mg, P?
Solution
Elements in the same group of the periodic table are most likely to exhibit similar properties. We
therefore expect Ca and Mg to be most alike because they are in the same group (2A, the
alkaline earth metals).
Practice Exercise 1
A biochemist who is studying the properties of certain sulfur (S)–containing compounds in the
body wonders whether trace amounts of another nonmetallic element might have similar
behavior. To which element should she turn her attention?
(a) F
(b) As
(c) Se
(d) Cr
(e) P.
Practice Exercise 2
Locate Na (sodium) and Br (bromine) in the periodic table. Give the atomic number of each and
classify each as metal, metalloid, or nonmetal.
67
Sample Exercise 2.6
Relating Empirical and Molecular
Formulas
Write the empirical formulas for (a) glucose, a substance also known as either blood sugar or
dextrose‒molecular formula C6H12O6; (b) nitrous oxide, a substance used as an anesthetic and
commonly called laughing gas‒molecular formula N2O.
Solution
(a) The subscripts of an empirical formula are the smallest whole‒number ratios. The smallest
ratios are obtained by dividing each subscript by the largest common factor, in this case 6.
The resultant empirical formula for glucose is CH2O.
(b) Because the subscripts in N2O are already the lowest integral numbers, the empirical formula
for nitrous oxide is the same as its molecular formula, N2O.
Practice Exercise 1
Tetracarbon dioxide is an unstable oxide of carbon with the following molecular structure:
What are the molecular and empirical formulas of this substance?
(a) C2O2, CO2
(b) C4O, CO
(c) CO2, CO2
(d) C4O2, C2O
(e) C2O, CO2.
Practice Exercise 2
Give the empirical formula for decaborane, whose molecular formula is B10H14.
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Sample Exercise 2.7
Writing Chemical Symbols for Ions
Give the chemical symbol, including superscript indicating mass number, for
(a) the ion with 22 protons, 26 neutrons, and 19 electrons; and
(b) the ion of sulfur that has 16 neutrons and 18 electrons.
Solution
(a) The number of protons is the atomic number of the element. A periodic table or list of
elements tells us that the element with atomic number 22 is titanium (Ti). The mass number
(protons plus neutrons) of this isotope of titanium is 22 + 26 = 48. Because the ion has three
more protons than electrons, it has a net charge of 3+ and is designated 48Ti3+.
(b) The periodic table tells us that sulfur (S) has an atomic number of 16. Thus, each atom or ion
of sulfur contains 16 protons. We are told that the ion also has 16 neutrons, meaning the mass
number is 16 + 16 = 32. Because the ion has 16 protons and 18 electrons, its net charge is 2–
and the ion symbol is 32S2–.
In general, we will focus on the net charges of ions and ignore their mass numbers unless the
circumstances dictate that we specify a certain isotope.
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Sample Exercise 2.7
Writing Chemical Symbols for Ions
Continued
Practice Exercise 1
In which of the following species is the difference between the number of protons and the number
of electrons largest?
(a) Ti2+
(b) P3–
(c) Mn
(d) Se2–
(e) Ce4+.
Practice Exercise 2
How many protons, neutrons, and electrons does the 79Se2– ion possess?
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Sample Exercise 2.8
Predicting Ionic Charge
Predict the charge expected for the most stable ion of barium and the most stable ion of oxygen.
Solution
We will assume that barium and oxygen form ions that have the same number of electrons as the
nearest noble‒gas atom. From the periodic table, we see that barium has atomic number 56. The
nearest noble gas is xenon, atomic number 54. Barium can attain a stable arrangement of 54
electrons by losing two electrons, forming the Ba2+ cation.
Oxygen has atomic number 8.
The nearest noble gas is neon, atomic number 10.
Oxygen can attain this stable electron arrangement
by gaining two electrons, forming the O2– anion.
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Sample Exercise 2.8
Predicting Ionic Charge
Continued
Practice Exercise 1
Although it is helpful to know that many ions have the electron arrangement of a noble gas, many
elements, especially among the metals, form ions that do not have a noble‒gas electron
arrangement. Use the periodic table, Figure 2.14, to determine which of the following ions has a
noble‒gas electron arrangement, and which do not. For those that do, indicate the noble‒gas
arrangement they match:
(a) Ti4+
(b) Mn2+
(c) Pb2+
(d) Te2–
(e) Zn2+.
Practice Exercise 2
Predict the charge expected for the most stable ion of
(a) aluminum and (b) fluorine.
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Sample Exercise 2.9
Identifying Ionic and Molecular
Compounds
Which of these compounds would you expect to be ionic: N2O, Na2O, CaCl2, SF4?
Solution
We predict that Na2O and CaCl2 are ionic compounds because they are composed of a metal
combined with a nonmetal. We predict (correctly) that N2O and SF4 are molecular compounds
because they are composed entirely of nonmetals.
Practice Exercise 1
Which of these compounds are molecular: CBr4, FeS, P4O6, PbF2?
Practice Exercise 2
Give a reason why each of the following statements is a safe prediction:
(a) Every compound of Rb with a nonmetal is ionic in character.
(b) Every compound of nitrogen with a halogen element is a molecular compound.
(c) The compound MgKr2 does not exist.
(d) Na and K are very similar in the compounds they form with nonmetals.
(e) If contained in an ionic compound, calcium (Ca) will be in the form of the doubly charged ion,
Ca2+.
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Using Ionic Charge to Write Empirical
Sample Exercise 2.10
Formulas for Ionic Compounds
Write the empirical formula of the compound formed by
(a) Al3+ and Cl– ions,
(b) Al3+ and O2– ions,
(c) Mg2+ and NO3– ions.
Solution
(a) Three Cl– ions are required to balance the charge of one Al3+ ion, making the empirical
formula AlCl3.
(b) Two Al3+ ions are required to balance the charge of three O2– ions. A 2:3 ratio is needed to
balance the total positive charge of 6+ and the total negative charge of 6–. The empirical
formula is Al2O3.
(c) Two NO3– ions are needed to balance the charge of one Mg2+, yielding Mg(NO3)2. Note that
the formula for the polyatomic ion, NO3–, must be enclosed in parentheses so that it is clear
that the subscript 2 applies to all the atoms of that ion.
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Using Ionic Charge to Write Empirical
Sample Exercise 2.10
Formulas for Ionic Compounds
Continued
Practice Exercise 1
Which of the following nonmetals will form an ionic compound with Sc3+ that has a 1:1 ratio of
cations to anions?
(a) Ne
(b) F
(c) O
(d) N.
Practice Exercise 2
Write the empirical formula for the compound formed by
(a) Na+ and PO43–; (b) Zn2+ and SO42–; (c) Fe3+ and CO32–.
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Determining the Formula of
Sample Exercise 2.11
an Oxyanion from Its Name
Based on the formula for the sulfate ion, predict the formula for
(a) the selenate ion and (b) the selenite ion.
(Sulfur and selenium are both in group 6A and form analogous oxyanions.)
Solution
(a) The sulfate ion is SO42–. The analogous selenate ion is therefore SeO42–.
(b) The ending ‒ite indicates an oxyanion with the same charge but one O atom fewer than the
corresponding oxyanion that ends in ‒ate. Thus, the formula for the selenite ion is SeO32–.
Practice Exercise 1
Which of the following oxyanions is incorrectly named?
(a) ClO2–, chlorate
(b) IO4–, periodate
(c) SO32–, sulfite
(d) IO3–, iodate
(e) NO2–, nitrite
Practice Exercise 2
The formula for the bromate ion is analogous to that for the chlorate ion. Write the formula for the
hypobromite and bromite ions.
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Determining the Names of
Sample Exercise 2.12
Ionic Compounds from Their Formulas
Name the ionic compounds
(a) K2SO4; (b) Ba(OH)2; (c) FeCl3.
Solution
In naming ionic compounds, it is important to recognize polyatomic ions and to determine the
charge of cations with variable charge.
(a) The cation is K+, the potassium ion, and the anion is SO42–, the sulfate ion, making the name
potassium sulfate. (If you thought the compound contained S2– and O2– ions, you failed to
recognize the polyatomic sulfate ion.)
(b) The cation is Ba2+, the barium ion, and the anion is OH–, the hydroxide ion: barium hydroxide.
(c) You must determine the charge of Fe in this compound because an iron atom can form more
than one cation. Because the compound contains three chloride ions, Cl–, the cation must be
Fe3+, the iron(III), or ferric, ion. Thus, the compound is iron(III) chloride or ferric chloride.
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Determining the Names of
Sample Exercise 2.12
Ionic Compounds from Their Formulas
Continued
Practice Exercise 1
Which of the following ionic compounds is incorrectly named?
(a) Zn(NO3)2, zinc nitrate
(b) TeCl4, tellurium(IV) chloride
(c) Fe2O3, diiron oxide
(d) BaO, barium oxide
(e) Mn3(PO4)2, manganese(II) phosphate.
Practice Exercise 2
Name the ionic compounds
(a) NH4Br; (b) Cr2O3; (c) Co(NO3)2.
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Relating the Names and Formulas of
Sample Exercise 2.13
Acids
Name the acids
(a) HCN; (b) HNO3; (c) H2SO4; (d) H2SO3.
Solution
(a) The anion from which this acid is derived is CN–, the cyanide ion. Because this ion has an
‒ide ending, the acid is given a hydro‒ prefix and an ‒ic ending: hydrocyanic acid. Only water
solutions of HCN are referred to as hydrocyanic acid. The pure compound, which is a gas
under normal conditions, is called hydrogen cyanide. Both hydrocyanic acid and hydrogen
cyanide are extremely toxic.
(b) Because NO3– is the nitrate ion, HNO3 is called nitric acid (the ‒ate ending of the anion is
replaced with an ‒ic ending in naming the acid).
(c) Because SO42– is the sulfate ion, H2SO4 is called sulfuric acid.
(d) Because SO32– is the sulfite ion, H2SO3 is sulfurous acid (the ‒ite ending of the anion is
replaced with an ‒ous ending).
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Relating the Names and Formulas of
Sample Exercise 2.13
Acids
Continued
Practice Exercise 1
Which of the following acids are incorrectly named? For those that are, provide a correct name or
formula.
(a) hydrofluoric acid, HF
(b) nitrous acid, HNO3
(c) perbromic acid, HBrO4
(d) iodic acid, HI
(e) selenic acid, H2SeO4.
Practice Exercise 2
Give the chemical formulas for
(a) hydrobromic acid; (b) carbonic acid.
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Relating the Names and Formulas of
Sample Exercise 2.14
Binary Molecular Compounds
Name the compounds
(a) SO2; (b) PCl5; (c) Cl2O3.
Solution
The compounds consist entirely of nonmetals, so they are molecular
rather than ionic. Using the prefixes in Table 2.6, we have
(a) sulfur dioxide,
(b) phosphorus pentachloride,
(c) dichlorine trioxide.
Practice Exercise 1
Give the name for each of the following binary compounds of carbon:
(a) CS2; (b) CO; (c) C3O2; (d) CBr4; (e) CF.
Practice Exercise 2
Give the chemical formulas for
(a) silicon tetrabromide; (b) disulfur dichloride; (c) diphosphorus hexaoxide.
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Writing Structural and Molecular
Sample Exercise 2.15
Formulas for Hydrocarbons
Assuming the carbon atoms in pentane are in a linear chain, write
(a) the structural formula and (b) the molecular formula
for this alkane.
Solution
(a) Alkanes contain only carbon and hydrogen, and each carbon is
attached to four other atoms. The name pentane contains the
prefix penta‒ for five (Table 2.6), and we are told that the carbons
are in a linear chain. If we then add enough hydrogen atoms to
make four bonds to each carbon, we obtain the structural formula.
This form of pentane is often called n‒pentane, where the
n‒ stands for “normal” because all five carbon atoms are in one
line in the structural formula.
(b) Once the structural formula is written, we determine the molecular
formula by counting the atoms present. Thus, n‒pentane has the
molecular formula C5H12.
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Writing Structural and Molecular
Sample Exercise 2.15
Formulas for Hydrocarbons
Continued
Practice Exercise 1
(a) What is the molecular formula of hexane, the alkane with six carbons?
(b) What are the name and molecular formula of an alcohol derived from hexane?
Practice Exercise 2
These two compounds have “butane” in their name. Are they isomers?
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