Download Chemical Bonding

Chapter
2
In this chapter, you will be
able to
• describe the role of electrons in
ionic and covalent bonding;
• relate the physical and chemical
properties of compounds to the
nature of their chemical bonds;
• predict the nature of bonds by
comparing electronegativity
values;
• use a variety of models to represent the formation and structure
of compounds and molecular elements;
• name a variety of compounds
using common names, classical
names, and IUPAC chemical
nomenclature.
Chemical Bonding
In caves we can sometimes see large structures known as stalagmites (Figure 1).
These formations are made of crystals of calcium carbonate, CaCO3(s), also
known as limestone. Calcium carbonate, as its name and formula suggest, is a
compound made up of three different elements. In addition to its crystalline
structure, calcium carbonate has high melting and boiling points and dissolves
to some extent in water. Many other compounds have similar physical and chemical properties, such as ordinary table salt, sodium chloride. Other compounds,
such as water, H2O, and carbon dioxide, CO2, have significantly different properties. How can we explain these similarities and differences? The answer lies in
an understanding of the special forces of attraction, or bonds, that hold atoms
together in compounds. These forces form the foundation of chemical properties and reactions.
As you know, we often develop models to help us understand abstract concepts. You are already familiar with several models: Bohr’s model of the atom; the
water cycle; the collision model. In this chapter we will develop theories and
models of chemical bonding (including the character of atomic bonds) to
explain the nature and behaviour of matter and to classify compounds.
Reflectyour
on
Learning
1. Why do atoms form compounds? Use examples of compounds you are
familiar with in your explanation.
2. Is there more than one type of force present between the atoms in compounds and, if so, how do these forces compare in strength and other
properties?
3. How could the forces that hold atoms together in a compound determine
the chemical properties of that compound? Again, use examples of
familiar compounds and their properties in your speculations.
Throughout this chapter, note any changes in your ideas as you learn new concepts and develop your skills.
64
Chapter 2
Try This
Activity
Making Models of Compounds
In Chapter 1 you studied various representations, or models, of the
structure of atoms. In this activity, you will build models of a variety of
molecules. This exercise will help you to see that different chemicals
have different molecular structures.
Materials: molecular models kit
• In your molecular model kit, identify the pieces representing each of
the following elements: chlorine, bromine, carbon, nitrogen, oxygen,
iodine, and hydrogen.
• Group the different pieces according to the number of holes present.
(a) To what chemical family or families do these groups of pieces
correspond on the periodic table?
(b) What do the holes in the pieces represent?
(c) What do the sticks represent in your molecular model kit?
Figure 1
Stalagmites are shown extending upward
from this cavern floor. As you can see, these
calcium carbonate formations have the
potential to join floor to ceiling in a
continuous pillar.
• Construct as many different compounds as possible using the
pieces in the kit.
(d) As best you can, name the compounds you have modelled
(later in this chapter you will learn a system for naming compounds).
(e) How do the connections differ among your models?
(f) Classify the models by dividing them into groups. Provide a
rationale for your classification.
(g) Organize the models into two more classification schemes.
Provide a rationale for each classification.
Chemical Bonding 65
2.1
Figure 1
Conductivity is used to distinguish between
aqueous solutions of ionic and molecular
compounds.
ionic compound: a pure substance
formed from a metal and a nonmetal
molecular compound: a pure substance formed from two or more different
nonmetals
electrical conductivity: the ability of a
material to allow electricity to flow through it
electrolyte: a substance that forms a solution that conducts electricity
INQUIRY SKILLS
Questioning
Hypothesizing
Predicting
Planning
Conducting
66
Chapter 2
Recording
Analyzing
Evaluating
Communicating
Classifying Compounds
There are many ways to classify substances. You are already familiar with several
of them. For example, you know that iron is a solid at SATP (not a liquid or a
gas), a metal (not a nonmetal), and an element (not a compound). Similarly,
water is a liquid and a compound. In this section, we will look at ways of further
classifying compounds.
One category of compounds includes table salt (sodium chloride), NaCl(s),
bluestone (copper(II) sulfate), CuSO4(s), and baking soda (sodium bicarbonate),
NaHCO3(s). If you think about their chemical formulas, you might notice that
each one is made up of a metal joined to a nonmetal. These compounds are
called ionic compounds.
A second category includes sulfur dioxide, SO2, carbon dioxide, CO2, and
ammonia, NH3. Look at their formulas. What do you notice? They are all nonmetals combined with nonmetals, and are called molecular compounds.
Many ionic and molecular compounds can be found within your own home.
For example, window cleaners, household bleach, antacid tablets, and milk of
magnesia contain ionic compounds. Vegetable oil, plastics, and sugar contain
molecular compounds.
We can use electrical conductivity—the ability to conduct electricity—to
distinguish between ionic and molecular compounds: Ionic compounds (many
of which dissolve readily in water) form solutions that conduct electricity.
However, molecular compounds (some of which dissolve in water) form solutions that generally do not conduct electricity. Of course, there are exceptions to
this generalization.
Substances that form solutions that conduct electricity are called electrolytes. Their solutions are called electrolytic solutions.
Substances that form solutions that do not conduct electricity are called
nonelectrolytes. Their solutions are called nonelectrolytic solutions (Figure 1).
Investigation 2.1.1
Comparing Ionic and Molecular Compounds
Just as a group of elements can share physical and chemical properties, so too can
a class of compounds. For instance, solid ionic compounds generally have very
high melting points and are brittle—they can often be crushed fairly easily into a
powder. Molecular compounds in solid form tend to have a softer or waxy texture,
and many have melting and boiling points so low that they are gases or liquids at
room temperature. If we can classify a substance as either ionic or molecular, we
should be able to predict some of its properties. Conversely, if we know some of the
properties of a substance, we can classify it as ionic or molecular.
There are several properties that we could take into account. Generally, the
simplest to investigate are physical properties, such as state at SATP, solubility in
water, and ability to conduct electricity in solution.
The purpose of this investigation is to test the generalizations you have
encountered about ionic and molecular compounds. Complete the
Hypothesis/Prediction, Analysis, Evaluation, and Synthesis sections of the lab
report.
2.1
Question
Which of the following substances are ionic, and which are molecular: sodium
nitrate, NaNO3(s); sucrose, C12H22O11(s); sodium chloride, NaCl(s); potassium
sulfate, K2SO4(s); ethanol, C2H5OH(l); sodium bicarbonate, NaHCO3(s); and calcium sulfate, CaSO4(s)?
Hypothesis/Prediction
(a) Answer the Question, considering the state of matter at room temperature
and the chemical formula of each substance. Provide your reasoning for
each decision.
(b) Based on your classification, predict whether each substance will dissolve in
water and whether any solutions formed will conduct electricity.
Experimental Design
The ionic or molecular nature of several compounds will be determined by
applying the tests of solubility and conductivity.
Materials
lab apron
eye protection
8 50-mL beakers
distilled water
wax pencil
scoopula
medicine dropper
stirring rod
sodium nitrate, NaNO3(s)
sucrose, C12H22O11(s)
sodium chloride, NaCl(s)
potassium sulfate, K2SO4(s)
ethanol, C2H5OH(l)
sodium bicarbonate, NaHCO3(s)
calcium sulfate, CaSO4(s)
low-voltage conductivity apparatus
Ethanol is flammable; do not
use near an open flame. Use
only low-voltage conductivity
apparatus.
Wear eye protection and an
apron, and wash hands thoroughly at the end of the
investigation.
Procedure
Part 1: Solubility
1. Obtain a small amount of NaCl(s). Observe and, in a table, record its state
at the ambient temperature.
2. Pour about 10 mL of distilled water into a 50-mL beaker. Add a small
quantity of the chemical to the water (Figure 2).
3. Use a stirring rod to stir the mixture. Note whether the chemical dissolves.
Record your observations. Label the beaker with the name of the chemical
and put it aside for the conductivity test in Part 2.
4. Repeat steps 1 through 3 for each of the other compounds.
Part 2: Conductivity
5. Obtain a small sample of distilled water in a beaker. Test the electrical conductivity of the sample. A reading of zero should be indicated by the apparatus.
6. Test the conductivity of the mixture of sodium chloride and water that was
set aside in Part 1. Record your observations and rinse the probes in distilled water.
7. Repeat step 6 for each of the remaining mixtures.
8. Dispose of the mixtures in the appropriate disposal containers provided by
your teacher. Rinse the beakers.
solid
liquid
Figure 2
For a solid chemical, use a small quantity in a
scoopula as shown. If the chemical is a
liquid, use a dropper full of liquid.
Chemical Bonding 67
Analysis
(c) Use the Evidence you gathered in this experiment to answer the Question
(classify each of the substances as ionic or molecular).
Evaluation
(d) Do you have confidence in your observations? Do you feel that they can be
used to accurately classify the substances? Explain.
(e) Compare your answer from the Analysis with the answer in your
Hypothesis/Prediction (question (a)). How do you account for any
differences?
Synthesis
(f) Were you able to accurately predict the properties of the substances based
on your initial classification? Why or why not?
(g) What assumptions are being made in this investigation?
Electrolytes and Classification
When table salt (sodium chloride), NaCl(s), dissolves in water it forms a solution
that conducts electricity: It is an electrolyte, and dissolving it forms an electrolytic solution. Sodium chloride is a typical ionic compound: Its solution conducts electricity. Sugar will also dissolve in water. However, the sugar solution
will not conduct electricity, so sugar is a nonelectrolyte. It is a typical molecular
compound: Its solution with water does not conduct electricity. In Unit 3, you
will learn more about conductivity and solutions. You will also find out more
about molecular acids: important exceptions to the rule that molecular solutions
do not conduct electricity.
Practice
Understanding Concepts
Table 1: Observations of Five
Unknown Compounds
Compound
State at
SATP
Conductivity
of solution
A
solid
yes
B
liquid
no
C
gas
yes
D
solid
yes
E
solid
no
1. What types of elements combine to form
(a) an ionic compound?
(b) a molecular compound?
2. Briefly describe a diagnostic test for an ionic compound, and give a
theoretical explanation for that test.
Applying Inquiry Skills
3. A student hypothesizes that an unknown substance is composed of
positive and negative ions held together by the attraction of their
opposite charges. Design an experiment that would allow the student
to test this hypothesis.
4. Use the evidence in Table 1 to classify each of the five compounds as
ionic or molecular. Provide your reasoning for each classification.
2.2
Ionic Bonding
When sodium (a metal) is put in a vessel containing chlorine (a nonmetal), the
two elements combine enthusiastically to form the compound sodium chloride,
68
Chapter 2
Analysis
(c) Use the Evidence you gathered in this experiment to answer the Question
(classify each of the substances as ionic or molecular).
Evaluation
(d) Do you have confidence in your observations? Do you feel that they can be
used to accurately classify the substances? Explain.
(e) Compare your answer from the Analysis with the answer in your
Hypothesis/Prediction (question (a)). How do you account for any
differences?
Synthesis
(f) Were you able to accurately predict the properties of the substances based
on your initial classification? Why or why not?
(g) What assumptions are being made in this investigation?
Electrolytes and Classification
When table salt (sodium chloride), NaCl(s), dissolves in water it forms a solution
that conducts electricity: It is an electrolyte, and dissolving it forms an electrolytic solution. Sodium chloride is a typical ionic compound: Its solution conducts electricity. Sugar will also dissolve in water. However, the sugar solution
will not conduct electricity, so sugar is a nonelectrolyte. It is a typical molecular
compound: Its solution with water does not conduct electricity. In Unit 3, you
will learn more about conductivity and solutions. You will also find out more
about molecular acids: important exceptions to the rule that molecular solutions
do not conduct electricity.
Practice
Understanding Concepts
Table 1: Observations of Five
Unknown Compounds
Compound
State at
SATP
Conductivity
of solution
A
solid
yes
B
liquid
no
C
gas
yes
D
solid
yes
E
solid
no
1. What types of elements combine to form
(a) an ionic compound?
(b) a molecular compound?
2. Briefly describe a diagnostic test for an ionic compound, and give a
theoretical explanation for that test.
Applying Inquiry Skills
3. A student hypothesizes that an unknown substance is composed of
positive and negative ions held together by the attraction of their
opposite charges. Design an experiment that would allow the student
to test this hypothesis.
4. Use the evidence in Table 1 to classify each of the five compounds as
ionic or molecular. Provide your reasoning for each classification.
2.2
Ionic Bonding
When sodium (a metal) is put in a vessel containing chlorine (a nonmetal), the
two elements combine enthusiastically to form the compound sodium chloride,
68
Chapter 2
2.2
a substance that you have classified as ionic, as it is an electrolyte. Like some other
ionic compounds that you are familiar with, for example, baking soda (sodium
hydrogen carbonate) and chalk (calcium carbonate), it is also brittle and has a
high melting temperature. How do we explain the formation and properties of
this compound?
You will recall, from Chapter 1, that atomic theory describes electrons
moving about the nucleus of the atom in energy levels, and that the electrons in
the outermost energy level are called the valence electrons. It is the valence electrons of an atom that form chemical bonds.
According to atomic theory, ionic compounds are formed when one or more
valence electrons are transferred from a metal atom to a nonmetal atom. This
leaves the metal atom as a positive ion, or cation, and the nonmetal atom as a
negative ion, or anion. The two oppositely charged ions are attracted to each
other by a force called an ionic bond (Figure 1).
+
Na
Cl
Na
chemical bond: the forces of attraction
holding atoms or ions together
ionic bond: the electrostatic attraction
between positive and negative ions in a compound; a type of chemical bond
–
Cl
Figure 1
An electron is transferred from sodium to
chlorine in the formation of an ionic bond.
Explaining the Properties of Ionic Compounds
Ionic compounds have similar properties: They are solids at SATP with high
melting points, and they are electrolytes. We can hypothesize that these properties might be the result of the bonds formed between the ions, holding them
firmly in a rigid structure.
Although they are composed of ions, pure ionic compounds are electrically
neutral. Therefore, the sum total of the electrical charges on all the ions must be
zero. Ionic compounds are made up of a fixed proportion of positive and negative
ions. Consequently, ionic compounds can only be identified in terms of the
smallest unit of the compound, known as a formula unit, that would still have the
properties of the compound. In the case of sodium chloride, the sodium and chloride ions are present in a 1:1 ratio, as indicated by its chemical formula, NaCl.
The anions and cations in an ionic compound are locked in a regular structure, held by the balance of attractive bonds and electrical repulsion. The most
common model of ions shows them as spheres arranged in a regular threedimensional pattern called a crystal lattice (Figure 2, page 70). We can actually
see the shape of sodium chloride crystals—an observation that supports our
crystal lattice model. We often see similar structures in bridges or scaffolding
where struts are joined one to the other in a repeating pattern.
Not all crystal lattices are square, like that of sodium chloride. Depending on
the sizes and charges of the ions that make up the substance, the crystal lattice
varies in structure. All lattices are arranged so that each ion has the greatest possible number of oppositely charged ions close, while keeping ions with the same
charge as far away as possible. In all cases, each ion is surrounded by ions of
opposite charge. In theory, this arrangement of ions creates strong attractions.
This theory is supported by empirical evidence such as the hard surfaces and
high melting and boiling points of ionic solids.
formula unit: the simplest whole-number
ratio of atoms or ions of the elements in an
ionic compound
crystal lattice: a regular, ordered arrangement of atoms, ions, or molecules
Chemical Bonding 69
It requires a great deal of energy to break the strong electrostatic attractions
within a crystal lattice. The ions resist any movement, as even a slight shift would
cause positive ions to move closer to other positive ions, and negative ions closer
to other negative ions, resulting in strong repulsion. We can use this model to
explain why ionic substances are hard (the ions resist movement), and also why
they are often brittle. Once the lattice is broken, repulsions between ions of the
same charge will cause the substance to split into two crystals.
Finally, our model of the structure of ionic compounds can also explain the
electrical conductivity of their solutions. When ionic compounds are dissolved
in water, the positive and negative ions dissociate:
NaCl(s) → Na+(aq) + Cl–(aq)
The ions are responsible for carrying current when charged electrodes are
placed in the ionic solution.
(a)
The Formation of Ionic Compounds
(b)
Figure 2
(a) The cubic structure of table salt crystals
provides a clue about the internal structure of sodium chloride.
(b) The arrangement of sodium and chlorine
ions in a crystal of sodium chloride. The
sodium and chloride ions occupy positions
in the crystal lattice known as lattice
points.
We classify most simple compounds containing metallic elements as ionic.
Elements within a chemical family (group) tend to participate in similar chemical reactions, producing ionic compounds with the same general formula. For
example, the metals in Groups 13 to 15, except mercury, will form ionic oxides
(compounds composed of a metal and oxygen) when burned in air. In the same
way, elements in Groups 1 and 2 form ionic compounds with oxygen: The oxides
formed by Group 1 elements have the general formula M2O while those of
Group 2 elements (e.g., magnesium oxide, MgO) have the general formula MO,
where M represents a metal ion. Because of its high melting point, magnesium
oxide is used to make objects that are exposed to very high temperatures, such as
crucibles, furnace linings, and thermal insulation.
From our discussion of trends in electronegativity, it makes sense that Group 1
metals readily react with the elements in Group 17 to form ionic compounds
with the general formula MX. These compounds, which are composed of a metal
and a halogen, are collectively referred to as ionic halides. Sodium chloride, an
example of an ionic halide, is found in large underground deposits in various
parts of the world. It is mined from these deposits and used as road salt, table salt,
and as a reactant in many industrial processes.
The elements in Group 2 show a similar trend, as they also react with the
halogens to produce ionic halides with the general formula MX2.
In general, the addition of a metal from Group 1 or Group 2 to water will
produce hydrogen gas and a basic ionic compound. Calcium hydroxide,
Ca(OH)2, is an example of an ionic compound that can be produced in this way.
However, the reaction of calcium and water is quite vigorous. A safer means of
producing this ionic compound is to react calcium oxide with water. Calcium
hydroxide is also referred to as slaked lime and is used to make mortar and
plaster for buildings.
Predicting Common Ions of Atoms
stable octet: a full shell of eight electrons
in the outer energy level of an atom
70
Chapter 2
Noble gases are stable and virtually inert. Similarly, ions with eight valence electrons appear to have a special stability. This arrangement of electrons is known
as a stable octet. To reach this stable state, metal atoms of elements in Groups 1,
2, and 3 will lose electrons to form cations, while elements in Groups 15, 16, and
17 will gain electrons to form anions. By looking at an element’s position in the
periodic table, we can predict the charge on that element’s most stable ion. For
2.2
example, we can predict that a Group 1 element will tend to lose one electron,
becoming a cation with a charge of 1+. A Group 16 element, on the other hand,
will tend to gain two electrons to complete a stable octet, so it will form an anion
with a charge of 2–.
Hydrogen is a special case in that, theoretically, it can either give up an electron to form H+, which is equivalent to a proton, or gain an electron to form H,
which has a filled shell like the noble gas helium.
Practice
Understanding Concepts
1. What properties of ionic compounds suggest that ionic bonds are
strong?
2. What types of elements form ionic bonds with each other?
3. Which of the representative elements tend to form positive ions?
Which tend to form negative ions?
4. What is the minimum number of different ions in the formula of an
ionic compound? Explain.
5. Predict the charge on the most stable ion formed by each of the following elements. Indicate the ion by writing the symbol complete
with charge.
(a) sulfur
(b) barium
(c) bromine
(d) chlorine
(e) calcium
(f) potassium
(g) phosphorus
(h) rubidium
(i) beryllium
Applying Inquiry Skills
6. You have already discovered that solutions of ionic compounds in
water conduct electricity. You might wonder about the conductivity of
pure ionic compounds.
(a) Design an experiment to investigate the conductivity of an ionic
solid. With your teacher’s approval, conduct your investigation.
(b) Research the conductivity of molten (liquid) ionic compounds.
(c) Assemble your findings into a report on the conductivity of ionic
compounds in various states.
(d) Propose a hypothesis for the properties you observe.
Ions and the Human Body
Among the 12 elements that make up more than 99% of the human body (Table 1)
are five metals: calcium, potassium, sodium, magnesium, and iron. These five
metals, which form positively charged ions in solution, are essential for maintaining good health. For example, Mg2+, Na+, and K+ are major constituents of
blood plasma. Ca2+ is vital in the formation of bones and teeth. In addition to
positively charged ions, some elements that form negatively charged ions are also
essential for life. Chloride ions, Cl, are another component of blood, and iodide
ions, I, are required to prevent a condition called goitre, which results in the
enlargement of the thyroid gland.
Table 1: 12 Most Common Elements
in the Human Body
Element
Percentage by
mass
oxygen
65
carbon
19
hydrogen
9.5
nitrogen
3.2
calcium
1.5
phosphorus
1.0
sulfur
0.3
chlorine
0.2
sodium
0.2
magnesium
0.1
iodine
<0.1
iron
<0.1
Chemical Bonding 71
Practice
Making Connections
7. Research and report upon the importance of one of the ions that
make up the human body. Your report should include: a description of
its biological importance; recommended daily minimum requirements; the effects of deficiency or excess on the human body; and
some of the naturally occurring sources. Conclude your report with a
discussion on whether, and under what circumstances, you would
recommend that someone should artificially supplement his/her
intake of this ion.
Follow the links for Nelson Chemistry 11, 2.2.
GO TO
www.science.nelson.com
Representing Ionic Bonds
electron dot diagram or Lewis
symbol: a representation of an atom or
ion, made up of the chemical symbol and
dots indicating the number of electrons in the
valence energy level
We generally find it easier to grasp a new concept if we have a model or mental
image. An American chemist, G. N. Lewis, obviously thought the same thing, and
developed a model of the valence (bonding) electrons of single atoms or
monatomic ions. He represented his models on paper as electron dot diagrams,
or Lewis symbols. These symbols consist of the chemical symbol for the element
plus dots representing the number of valence electrons. For example, a sodium
atom would be represented as
Na
The convention for indicating more electrons is as follows:
Li
Be
B
C
N
O
F
Ne
These diagrams can help us to represent the process of ion formation.
Electron dot diagrams also illustrate the theory that ionic bonds tend to produce full outer orbits of electrons: a configuration exactly the same as that of the
noble gases. Sodium has one valence electron. By transferring this electron to
another entity that has a stronger attraction for the electron, the resulting
sodium ion will have the same electron configuration as neon.
Na
Na + e
A chlorine atom has seven valence electrons. By attracting an electron from
another entity, the resulting chloride ion will have eight electrons in its valence
shell and the same number of electrons as its nearest noble gas, argon.
Representative elements around Groups 13 and 14 of the periodic table can
also achieve this special stability without losing or gaining electrons, as you will
discover later in the chapter.
Cl + e
Cl
(compare with Ar )
Another way to describe the process of forming ionic compounds is to say
that electron transfers result from the large difference in electronegativity between
metals and nonmetals. Nonmetals have a strong tendency to gain electrons and
metals do not. When an atom of low electronegativity, such as a metal atom, and
an atom of high electronegativity, such as a nonmetal atom, are in proximity, one
or more electrons are transferred from the atom with low electronegativity, transforming both atoms into ions. The metal ion will be positive and the nonmetal
ion will be negative. The resulting oppositely charged ions attract each other and
other ions, forming ionic bonds and resulting in a crystal lattice.
72
Chapter 2
2.2
Using electron dot diagrams, we can show the formation of an ionic bond
between sodium and chlorine.
Na + Cl
Na
Cl
When illustrating the formation of an ionic bond, we place square brackets
around the ion to indicate that the charge is not associated with any particular
electron and that all the electrons in the valence shell are equivalent.
We can use the periodic table and electron dot diagrams to predict the formulas of other ionic compounds. By finding out how many electrons they tend
to lose or gain to reach stable octets, we can figure out what ratio of ions will
make an electrically neutral compound.
As an example, suppose we were asked to draw electron dot diagrams to
illustrate the formation of calcium fluoride, state the ratio of ions in the compound, and give the formula for the compound. Calcium is in Group 2, so will
form an ion with a charge of 2+. Fluorine is in Group 17, so will form an ion with
a charge of 1.
Ca
2 F
F
Ca
2
F
To form an electrically neutral ionic compound, the ratio of calcium to fluoride
ions is 1:2. The formula is CaF2.
Sample Problem
Draw electron dot diagrams to illustrate the formation of magnesium oxide.
Write the ion ratio and the chemical formula.
Solution
Mg
O
Mg
2
O
2
The two elements will combine in a ratio of 1:1. The formula is MgO.
Practice
Understanding Concepts
8. (a) How do the electron dot diagrams of metal ions differ from those
of nonmetal ions?
(b) How are the electron dot diagrams of metal ions similar to those
of nonmetal ions?
9. Use electron dot diagrams to illustrate the formation of
(a) lithium iodide
(b) barium chloride
(c) potassium oxide
(d) calcium fluoride
10. Represent each of the following elements using electron dot diagrams:
(a) nitrogen
(b) sulfur
(c) argon
(d) iodine
(e) lithium
(f) cesium
(g) calcium
(h) sodium
Chemical Bonding 73
11. Use electron dot diagrams to determine the ratio in which oxygen
will combine with each of the following elements to form an ionic
compound. Label each diagram with the chemical formula of the
compound.
(a) calcium
(b) rubidium
(c) strontium
(d) aluminum
12. Represent the five halogens, using electron dot diagrams. How are
these diagrams consistent with the concept of a chemical family?
13. Explain, referring to stable octets, how the following ionic compounds are formed from pairs of elements; illustrate the formation of
each compound with electron dot diagrams; and predict the formula
of each compound.
(a) magnesium chloride
(b) sodium sulfide
(c) aluminum oxide (bauxite)
(d) barium chloride
(e) calcium fluoride (fluorite)
(f) sodium iodide
(g) potassium chloride (a substitute for table salt)
14. Give the common names for the following chemicals:
(a) sodium bicarbonate
(b) NaCl
(c) calcium carbonate
(d) Ca(OH)2
Sections 2.1–2.2 Questions
Understanding Concepts
1. Use the concepts of periodic trends and electronegativity to
explain why ionic compounds are abundant in nature.
2. Give a theoretical reason why lithium and oxygen combine in the
ratio 2:1.
3. How are the topics of ion formation and periodic trends related?
4. Give the correct chemical (IUPAC) names for the following chemicals:
(a) chalk
(b) slaked lime
(c) road salt
(d) baking soda
5. A Group 1 metal (atomic number 55) is reacted with the most
reactive of the halogens. A very vigorous reaction results in the
formation of a solid, white compound.
(a) Represent the formation of the compound with electron dot
diagrams.
(b) Write the formula of the compound formed.
(c) What type of compound is formed?
(d) Predict the physical properties of the resulting white compound.
(e) Explain the properties of the compound in terms of the
bonds formed.
(f) Provide a theoretical explanation for the vigorous reaction.
74
Chapter 2
2.3
2.3
Covalent Bonding
Ionic compounds, as you have just learned, contain many ions arranged in a
three-dimensional structure. But not all compounds are brittle ionic solids with
high melting points. Some, such as paraffin wax, are solids at SATP but have relatively low melting points. Others are liquids or gases, such as water and carbon
dioxide. In what way do these substances differ from ionic compounds? Can we
explain these differences with atomic theory? The answer is yes, but first we
should look at their formulas.
We have written the formulas of ionic compounds as simplest formula
ratios. The compound NaCl includes sodium and chloride ions in a 1:1 ratio. We
could build a crystal from 8 ions of each element, or 8 million of each and it
would form the same structure and would have the same properties. There is
only one ionic compound of sodium and chloride ions, and that compound
always contains a 1:1 ratio of the two ions.
This is not true of the bonding of nonmetals with each other. For example,
consider the simplest ratio formula CH. If carbon and hydrogen formed an ionic
compound with this ratio, we would expect that any structure in which these elements were in this ratio would have the same properties. However, this is not the
case. Using a mass spectrometer and combustion analysis we can demonstrate
that there are several compounds that have this simplest ratio formula. The gas
acetylene (ethyne), C2H2(g), and the liquid benzene, C6H6(l), both have the simplest ratio formula CH, but they are otherwise much different in their physical
and chemical properties. Simplest ratio formulas indicate only the relative numbers of atoms in a molecular compound; they give no information about the
actual number of atoms or the arrangement of those atoms in a molecule. To distinguish among molecular compounds, we need to represent them with formulas
that describe the molecules that make them up.
Molecules can be classified by the number of atoms that they contain.
Molecules that contain only two atoms, such as carbon monoxide, CO, are called
diatomic molecules. If they contain more than two atoms, such as ammonia,
NH3, they are called polyatomic molecules.
Some elements also exist as molecules. Hydrogen and oxygen are examples
of elements composed of diatomic molecules. Sulfur, S8, and phosphorus, P4, are
polyatomic molecules.
Figure 1
The pair of shared electrons between the
nuclei of two hydrogen atoms results in a
covalent bond.
diatomic molecule: a molecule consisting
of two atoms of the same or different elements
polyatomic molecule: a molecule consisting of more than two atoms of the same
or different elements
Formation of Covalent Bonds
You already know that hydrogen can form a cation (H+) by losing a valence electron, or it can form an anion (H–) by gaining an electron and filling its electron
shell. However, two hydrogen atoms can each obtain a stable filled energy level
by sharing a pair of electrons. (Remember that the first energy level can only contain two electrons.) The covalent bond that results arises from the simultaneous
attraction of two nuclei for a shared pair of electrons (Figure 1).
We can use the model of the Lewis symbol, or electron dot diagram, to communicate the theory of covalent bonding. When an electron dot diagram is used
to represent covalent bonding, we adapt it slightly and call it a Lewis structure
(because it illustrates the structure of the molecule). A Lewis structure shows the
valence electrons surrounding each of the component atoms as dots, with the
exception of the electrons that are shared: These shared electrons are represented
by a dash. In effect, this dash represents a covalent bond.
Cl
Cl
Cl
covalent bond: the attractive force
between two atoms that results when
electrons are shared by the atoms; a type
of chemical bond
Lewis structure: a representation of
covalent bonding based on Lewis symbols;
shared electron pairs are shown as lines
and lone pairs as dots
Cl
Chemical Bonding 75
lone pair: a pair of valence electrons not
involved in bonding
octet rule: a generalization stating that,
when atoms combine, the covalent bonds are
formed between them in such a way that
each atom achieves eight valence electrons
(two in the case of hydrogen)
According to the Lewis structure for a chlorine molecule, each Cl atom has
three pairs of electrons that are not involved in the formation of a covalent bond.
Each pair is referred to as a lone pair. There is one shared pair of electrons
between the atoms.
Many elements will form bonds that result in a full valence shell for each
atom, so that each atom has the electron structure of an atom of a noble gas
when the shared electrons are included. In other words, it has a stable octet. This
generalization is referred to as the octet rule. There are many exceptions to this
rule, but you will not be learning about these in this course.
Sample Problem 1
Draw Lewis symbols for the reaction of two bromine atoms and a Lewis structure for the resulting bromine molecule.
Solution
Br
structural formula: a representation of
the number, types, and arrangement of atoms
in a molecule, with dashes representing
covalent bonds
Br
Br
Br
This method of representing molecules can be further simplified by not
indicating lone pairs. This representation is referred to as a structural formula.
The structural formula for chlorine would be ClCl, and for bromine BrBr. As
you can see, they are easier to write, and are quite similar to the chemical formulas, Cl2 and Br2.
Each pair of shared electrons results in a single bond. Elements that need
only one more electron to complete their outer shell or energy level tend to form
single bonds. Hydrogen and chlorine are typical examples.
Elements in Group 16 are two electrons short of a full outer shell. How, then,
can they form covalent bonds with each other and still achieve a stable state? The
answer is simple, though not necessarily obvious: They form a double bond, with
each pair of atoms sharing two pairs of electrons between them. Oxygen and
carbon dioxide are examples of molecules that include double bonds. The double
bond is represented, in a Lewis structure and structural formula, as a double dash.
Sample Problem 2
(a) Draw the Lewis structure for a molecule of oxygen.
(b) Give the structural formula for the molecule.
Solution
(a) O + O
(b) O O
O
O
Can you predict how atoms of nitrogen, which require three electrons each
to achieve a stable octet, might form N2? The two atoms form a triple bond by
sharing three electron pairs.
N
N
N
N
The structural formula for a molecule of nitrogen is
N
bonding capacity: the number of electrons lost, gained, or shared by an atom
when it bonds chemically
76
Chapter 2
N
The number of covalent bonds (shared electron pairs) that an atom can form
is known as its bonding capacity (Table 1). Each atom of nitrogen, we have just
2.3
learned, shares three electron pairs, so it has a bonding capacity of three. What is
the bonding capacity of oxygen? It always shares two electron pairs, so has a
bonding capacity of two. It is easy to find the bonding capacity of any element by
looking at the Lewis structure of a molecule containing that element: The number
of dashes associated with the element is the same as the bonding capacity.
Table 1: Bonding Capacities of Some Common Atoms
Atom
Number of
valence electrons
Number of
bonding electrons
Bonding capacity
carbon
4
4
4
nitrogen
5
3
3
oxygen
6
2
2
halogens
7
1
1
hydrogen
1
1
1
So far, we have looked at molecules of elements. Because the atoms are the
same, each has the same bonding capacity and each contributes the same
number of electrons to the covalent bond. The molecules of compounds, however, consist of atoms of two or more different elements, often with different
bonding capacities. How do we decide on their arrangement, when we draw
structural formulas? The central position in the arrangement is often occupied
by the element with the highest bonding capacity. Carbon and nitrogen, for
instance, are commonly at the centre of a structural formula. Electronegativity is
another means of deciding upon the central atom. When there is a choice of
atoms for the central position in the molecule, choose the least electronegative
element. Hydrogen is never the central atom since it can only form a single covalent bond. Halides and oxygen are also usually not the central atom.
There are exceptions to these generalizations, but they meet our needs in
most cases.
Sample Problem 3
(a) Draw a Lewis structure for a molecule of carbon dioxide.
(b) Give the structural formula for the molecule.
Solution
(a) O + C + O
(b) O C O
O
C
O
Practice
Understanding Concepts
1. Draw a Lewis structure and write the molecular formula for each of
the following:
(a) F2(g)
(b) H2O(l)
(c) CH4(g)
(d) PCl3(s)
(e) H2S(g)
(f) SiO2(s)
Chemical Bonding 77
Coordinate Covalent Bonds
polyatomic ion: a covalently bonded
group of atoms with an overall charge
coordinate covalent bond: a covalent
bond in which both of the shared electrons
come from the same atom
Many substances contain a combination of covalent and ionic bonding. Consider
the compound ammonium chloride, NH4Cl. This white, crystalline solid dissolves rapidly in water and is an electrolyte—it dissociates to form a cation and
an anion. It has many of the properties of an ionic compound, but it is composed
only of nonmetals. We explain the properties of ammonium chloride by
describing it as an ionic compound composed of a chloride ion, Cl–, and a polyatomic ion, ammonium, NH4+. The bond holding the chloride and ammonium
ions together is ionic, but the bonds within the polyatomic ammonium ion are
covalent. There are several polyatomic ions, including NH4+, SO42– (sulphate),
and CO32–(carbonate), all of which are covalently bonded groups of atoms carrying an overall charge.
How does this arrangement fit with our description of covalent bonds?
Molecules that are composed of two or more different elements can sometimes
form covalent bonds where both of the electrons making up the bond are provided by the same atom. This type of bond is called a coordinate covalent bond.
Consider the formation of the ammonium ion from the regular covalent molecule ammonia, NH3, and a hydrogen ion, H+. The hydrogen ion does not bring
any electrons with it. To achieve a complete outer shell it can borrow two electrons from the atom with which it bonds.
To explain this bond, we can draw Lewis structures. First, we must establish
the Lewis structure for ammonia. We can arrange the atoms with nitrogen
(which has the highest bonding capacity) at the centre, showing the five valence
electrons of nitrogen and the one valence electron of each of the three hydrogen
atoms.
H
N
H
H
We can show the pairs of shared electrons (covalent bonds) between adjacent
nitrogen and hydrogen atoms as dashes.
H
N
H
H
Notice the lone pair of electrons in this structure. A hydrogen ion, which has
no electrons of its own, can bond to the ammonia molecule by sharing this
unbonded pair of electrons. This is the coordinate covalent bond.
H
H
H
N
H
H
H
N
H
H
Once the hydrogen ion is bonded, there is no way to tell which of the hydrogens was the ion. Each of the four hydrogen atoms in the structure are equivalent: The positive charge is not really associated with any particular hydrogen
atom. To indicate this, square brackets are placed around the entire ammonium
ion and the positive charge is written outside the bracket.
78
Chapter 2
2.3
The Strength of Covalent Bonds
Covalent bonds are strong. A large amount of energy is needed to separate the
atoms that make up molecules. For this reason, molecules tend to be stable at relatively high temperatures: They do not easily decompose upon heating.
The stronger the bonds within the molecule, the greater the energy required
to separate them. The strength of a bond between two atoms increases as the
number of electron pairs in the bond increases. Therefore, triple bonds are
stronger than double bonds, which are stronger than single bonds between the
same two atoms.
SUMMARY
Drawing Lewis Structures and Structural
Formulas for Molecular Compounds
1. Arrange the symbols of the elements of the compound as you would
expect the atoms to be arranged in the compound. The element with the
highest bonding capacity is generally written in the central position
(Figure 2(a)).
2. Add up the number of valence electrons available in each of the atoms
(Figure 2(b)). If the structure is a polyatomic ion, add one electron for
each unit of negative charge, or subtract one for each unit of positive
charge.
3. Place one pair of electrons between each adjacent pair of elements
(forming single covalent bonds) (Figure 2(c)).
4. Place pairs of the remaining valence electrons as lone pairs on the peripheral atoms (not the central atom) (Figure 2(d)).
5. If octets are not complete, move lone pairs into bonding position between
those atoms and the central atom until all octets are complete
(Figure 2(e)).
6. If the peripheral atoms all have complete octets and there are pairs of electrons remaining, place these electrons as lone pairs on the central atom.
7. Count the number of bonds between the central atom and the peripheral
atoms. If this number exceeds the bonding capacity of the central atom,
one or more of the bonds is coordinate covalent. To identify which ones,
try removing the peripheral atoms one at a time. If you can do this and
leave the central structure with complete octets, you have identified coordinate covalent bonds (Figure 2(f)).
8. To give the structural formula, remove the dots representing the lone pairs
and replace bond dots with dashes (Figure 2(g)).
9. If you are representing a polyatomic ion, place brackets around the entire
structure and write the charge outside the brackets.
(a)
O
S
O
O
(b)
3(6) + 6 = 24
O S
(c)
O
S
O
O
(d)
O
S
O
O
(e)
Sulfur atom has an
incomplete octet.
O
S
O
(f)
O
S
O
O
O
SO3 includes two coordinate
covalent bonds.
O
S
O
O
(g)
O
S
O
O
Figure 2
Practice
Understanding Concepts
2. Draw Lewis structures and structural formulas for each of the following molecules:
(a) H2(g)
(d) NF3(g)
(b) O3(g)
(e) N2H2(g)
(c) OF2(g)
(f) P2H4(g)
Chemical Bonding 79
3. Draw Lewis structures and structural formulas for each of the following polyatomic ions:
3–
–
(a) PO4(aq)
(c) BrO3(aq)
–
–
(b) OH(aq)
(d) ClO4(aq)
4. Which of the Lewis structures in questions 2 and 3 include coordinate
covalent bonds?
Explaining the Properties of Molecular Compounds
You are familiar with many molecular compounds: propane, C3H8(g), for the barbecue, water, H2O(l), in your bathtub, and sugar, C12H22O11(s), for your coffee. As
you can see, their physical properties vary greatly. In contrast to ionic compounds, which are all solids at SATP, molecular compounds may be solids, liquids, or gases (Table 2).
Table 2: Comparison of Ionic and Molecular Solids
intramolecular force: the attractive
force between atoms and ions within a compound
intermolecular force: the attractive
force between molecules
80
Chapter 2
Properties
Ionic
Molecular
melting point
high
low
electrical conductivity
in the solid state
in the liquid state
no conductivity
conductivity
no conductivity
no conductivity
consistency of solid
hard, brittle
soft, waxy or flexible
Examples
sodium chloride
copper(II) sulfate
iodine
phosphorus
We have discussed the forces that bond atoms and ions together within a
compound, the intramolecular forces (“intra” means within). These are sufficient to explain the existence of ionic and molecular compounds, and to explain
many of the properties of ionic compounds, but they aren’t sufficient to explain
the physical state of molecular compounds. Why is water a liquid at SATP and a
solid at STP? Why isn’t it a gas at all temperatures? Something, some force, must
hold the molecules together in the solid and liquid states.
Forces between molecules are called intermolecular forces (“inter” means
between). The evidence indicates that these forces are strong enough to cause
molecules to arrange themselves in an orderly fashion to form a lattice structure
(similar to that of ionic solids). Both solid water and solid carbon dioxide (dry
ice) show such structures when examined by crystallography (Figure 3).
However, intermolecular forces must be weak compared to covalent bonds. We
can deduce this from the observation that it is much easier to melt a molecular
solid than it is to cause the same substance to decompose. When water is heated
from –4°C to 104°C it changes state from a solid, to a liquid, and then to a gas,
but it does not decompose to oxygen and hydrogen. The energy added in the
form of heat is sufficient to overcome the intermolecular forces between the molecules, but not the covalent bonds between the atoms. Adding a relatively small
amount of heat will cause a solid molecular compound to change state from a
solid to a liquid, and then to a gas, but it takes much more energy to break the
covalent bonds between the atoms in the compound.
Later in this chapter, you will learn more about intermolecular forces and
how they affect the properties of substances.
2.3
Practice
Understanding Concepts
5. Distinguish between bonding electrons and lone pairs.
6. Are the following pairs of atoms more likely to form ionic or covalent
bonds?
(a) sulfur and oxygen
(b) iodine and iodine
(c) calcium and chlorine
(d) potassium and bromine
7. (a) List six examples of molecular elements and compounds and six
examples of ionic compounds.
(b) Compare the two lists, referring to bond types to explain the contrasting physical properties.
8. (a) How does the bonding capacity of nitrogen differ from that of
chlorine?
(b) Give a theoretical explanation for your answer to (a).
Figure 3
The molecules in solid carbon dioxide, CO2,
form a crystal lattice.
9. How are coordinate covalent bonds similar to covalent bonds? How
are they different?
10. (a) Use an electron dot diagram to explain the formula for nitrogen, N2.
(b) Draw the Lewis structure for nitrogen.
(c) Nitrogen is a fairly inert (unreactive) gas. Explain this, referring to
the bonds involved.
11. Illustrate the formation of each of the following molecular compounds, using electron dot diagrams and Lewis structures:
(a) HCl
(c) H2S
(b) NH3
(d) CO2
12. Use the octet rule to develop a table that lists the bonding capacities
for carbon, nitrogen, oxygen, hydrogen, fluorine, chlorine, bromine,
and iodine.
13. Use bonding capacities (Table 1, page 77) to draw the structural formula of each of the following molecules.
(a) O2
(b) Br2
(c) H2O2
(d) C2H4
(e) HCN
(f) C2H5OH
(g) CH3OCH3
(h) CH3NH2
14. (a) Illustrate the structure of a hydronium ion (H3O+) by drawing its
Lewis structure.
(b) Name the bonds within the hydronium ion.
(c) What kinds of bonds is this ion likely to form with other entities?
15. Is it correct for the structural formula of H2S to be written as H—H—S?
Explain, using a diagram.
Chemical Bonding 81
Section 2.3 Questions
Understanding Concepts
1. Compound A is formed when the element with atomic number 3
combines with the element of atomic number 9. Compound B is
formed when the element with atomic number 7 combines with
the element of atomic number 9.
(a) Compare the properties of compounds A and B.
(b) What types of compounds are A and B? Give reasons for
your answer.
(c) Clearly show the structure of each compound formed, using
electron dot diagrams and Lewis structures.
2. Oxygen forms ionic bonds with aluminum to form bauxite (or
aluminum oxide). However, oxygen forms covalent bonds with
carbon to form carbon dioxide. Use the concepts of electronegativity and periodic trends to explain these differences in bonding.
3. An alkali metal M reacts with a halogen X to form a compound
with the formula MX.
(a) Would this compound have ionic or covalent bonds? Explain.
(b) Predict the physical properties of the compound MX.
4. The compound NaCl and the element Cl2 are both held together
by chemical bonds.
(a) Classify each substance as either an ionic or a covalent compound.
(b) Explain, in terms of chemical bonds and attractive forces,
why NaCl is a solid at SATP while Cl2 is a gas at SATP.
5. Using only oxygen and sulfur atoms, create as many compounds
as you can and draw Lewis diagrams for each. Which of the compounds contain only multiple bonds?
Applying Inquiry Skills
6. Design an experiment to test your classification of compounds A
and B in question 1.
7. Design an experiment to test your predictions of the physical
properties of the compound MX in question 3.
(a)
2.4
Electronegativity, Polar Bonds,
and Polar Molecules
Why are some molecular substances solid, some liquid, and some gaseous at
SATP? Why do different liquids have different boiling points? Why can water
striders walk across the surface of a pond without falling in? Believe it or not, all
these phenomena depend on the bonds and forces between molecules.
(b)
Polar Covalent Bonds
Figure 1
Electron densities of the bonding electrons in
two covalent molecules
(a) Cl2
(b) HCl
82
Chapter 2
So far, we have discussed models for two types of chemical bonding: ionic and
covalent. However, when a chemical bond is formed, it is not always exclusively
one or the other.
When electrons are shared between two atoms, a covalent bond results.
When the atoms are identical, such as in a chlorine molecule, the electrons are
shared equally (Figure 1(a)). However, this is not the case for a compound like
hydrogen chloride, where electrons are shared between two different elements. In
Section 2.3 Questions
Understanding Concepts
1. Compound A is formed when the element with atomic number 3
combines with the element of atomic number 9. Compound B is
formed when the element with atomic number 7 combines with
the element of atomic number 9.
(a) Compare the properties of compounds A and B.
(b) What types of compounds are A and B? Give reasons for
your answer.
(c) Clearly show the structure of each compound formed, using
electron dot diagrams and Lewis structures.
2. Oxygen forms ionic bonds with aluminum to form bauxite (or
aluminum oxide). However, oxygen forms covalent bonds with
carbon to form carbon dioxide. Use the concepts of electronegativity and periodic trends to explain these differences in bonding.
3. An alkali metal M reacts with a halogen X to form a compound
with the formula MX.
(a) Would this compound have ionic or covalent bonds? Explain.
(b) Predict the physical properties of the compound MX.
4. The compound NaCl and the element Cl2 are both held together
by chemical bonds.
(a) Classify each substance as either an ionic or a covalent compound.
(b) Explain, in terms of chemical bonds and attractive forces,
why NaCl is a solid at SATP while Cl2 is a gas at SATP.
5. Using only oxygen and sulfur atoms, create as many compounds
as you can and draw Lewis diagrams for each. Which of the compounds contain only multiple bonds?
Applying Inquiry Skills
6. Design an experiment to test your classification of compounds A
and B in question 1.
7. Design an experiment to test your predictions of the physical
properties of the compound MX in question 3.
(a)
2.4
Electronegativity, Polar Bonds,
and Polar Molecules
Why are some molecular substances solid, some liquid, and some gaseous at
SATP? Why do different liquids have different boiling points? Why can water
striders walk across the surface of a pond without falling in? Believe it or not, all
these phenomena depend on the bonds and forces between molecules.
(b)
Polar Covalent Bonds
Figure 1
Electron densities of the bonding electrons in
two covalent molecules
(a) Cl2
(b) HCl
82
Chapter 2
So far, we have discussed models for two types of chemical bonding: ionic and
covalent. However, when a chemical bond is formed, it is not always exclusively
one or the other.
When electrons are shared between two atoms, a covalent bond results.
When the atoms are identical, such as in a chlorine molecule, the electrons are
shared equally (Figure 1(a)). However, this is not the case for a compound like
hydrogen chloride, where electrons are shared between two different elements. In
2.4
this situation, the sharing is unequal, as the bonding electrons spend more time
near one atom than near the other (Figure 1(b)). The electrons in the H–Cl
bond in a hydrogen chloride molecule spend more time near the chlorine atom
than near the hydrogen atom. This is because of chlorine’s greater attraction for
electrons. Due to this unequal sharing of electrons, the hydrogen atom is, on
average, slightly positively charged while the chlorine atom is slightly negatively
charged (Figure 2). We indicate these slight charges by d and d. (The Greek
letter delta, d, indicates “a small difference.”) While chlorine shows the greater
attraction for the bonding electrons, the attraction is not strong enough to actually bring about an electron transfer, as in an ionic compound. The bond is
somewhere between an ionic bond and a covalent bond and is called a polar
covalent bond.
We can predict which parts of the molecule will have d and d charges by
comparing the electronegativities of the atoms (given in the periodic table at the
end of this text). The most electronegative atoms will attract the electron pair
strongly, and so will tend to have a d charge in the covalent compounds they
form.
–
+
Figure 2
Slight positive and negative charges are indicated by the Greek symbol d, delta.
polar covalent bond: a covalent bond
formed between atoms with significantly different electronegativities; a bond with some
ionic characteristics
Sample Problem
(a) Draw the structural formula for methane (CH4, the main constituent
of natural gas).
(b) Indicate which atoms are slightly positive and which are slightly negative.
Solution
(a)
H
H
C
H
H
H+
(b)
H+
C-
H+
H+
Each hydrogen atom carries a partial positive charge while the carbon atom carries a partial negative charge.
To predict whether a chemical bond between two atoms will be ionic, polar
covalent, or covalent, we must consider the electronegativities of the elements
involved. The absolute value of the difference in electronegativities of two
bonded atoms provides a measure of the polarity in the bond: the greater the difference, the more polar the bond (Table 1). According to the periodic trends in
electronegativity described in Chapter 1, fluorine atoms have the greatest tendency to gain or attract electrons and cesium atoms have the least. The difference
between their electronegativities is 3.3 (cesium 0.7 and fluorine 4.0). Since the
difference in electronegativity for cesium and fluorine is large, they should (and
do) form an ionic bond. In other words, the fluorine atom is strong enough, relative to the cesium atom, to “steal” a cesium electron.
Table 1: Electronegativity Differences
of a Selection of Bonds
Bond
Elecronegativity difference
Cl—Cl
0.0
C—Cl
0.3
C—H
0.4
Be—H
0.6
N—H
0.9
C—O
1.0
O—H
1.4
Mg—Cl
1.8
Na—Cl
2.1
Mg—O
2.3
Ba—O
2.6
Ca—F
3.0
Chemical Bonding 83
The difference in electronegativities of hydrogen (2.1) and chlorine (3.0) is
much less. They instead form a polar covalent bond. The greater the difference
in electronegativity, the more polar, and then more ionic, the bond becomes.
This relationship describes a continuum, a range from covalent to ionic, rather
than three different kinds of bonds (Figure 3). How do we know when to classify
a chemical bond as fully ionic? By convention, a difference in electronegativity
greater than 1.7 indicates an ionic bond. However, there are many binary compounds in which the two atoms have an electronegativity difference of less than
1.7, but the compound still has ionic properties. For example, MgI2 (electronegativity difference 1.3) is an electrolyte, i.e., is an ionic compound.
ionic
Figure 3
In this model of a bonding continuum, the
colour change from right to left indicates the
increasing difference in electronegativity and
bond polarity.
3.3
polar covalent
1.7
Electronegativity Difference
covalent
0
Practice
Understanding Concepts
1. Describe how we can use electronegativity values to predict the types
of bonds that will form within a compound.
2. Which type of bond—ionic or covalent—will form between each of
the following pairs of atoms? Of the covalent bonds, indicate which
would be the most polar.
(a) H and Cl
(b) Si and O
(c) Mg and Cl
(d) Li and O
(e) N and O
(f) O and O
(g) I and Cl
(h) Cr and O
(i) C and Cl
3. Identify the more polar bond in each of the following pairs.
(a) H—F; H—Cl
(b) N—O; C—O
(c) S—H; O—H
(d) P—Cl; S—Cl
(e) C—H; N—H
(f) S—O; P—O
(g) C—N; N—N
4. Draw Lewis structures for the following substances. In each case, use
appropriate notation to indicate the atoms that are slightly positive
and those that are slightly negative.
(a) H2O
(b) Br2
(c) HBr
(d) PCl3
(e) OF2
84
Chapter 2
2.4
SUMMARY
Bonding Characteristics
Table 2: Summary of Bonding
Intramolecular force
Bonding model
ionic bond
• involves an electron transfer, resulting in the formation of cations
and anions
• cations and anions attract each other
polar covalent bond
• involves unequal sharing of pairs of electrons by atoms of two
different elements
• bonds can involve 1, 2, or 3 pairs of electrons, i.e., single (weakest),
double, or triple (strongest) bonds
covalent bond
• involves equal sharing of pairs of electrons
• bonds can involve 1, 2, or 3 pairs of electrons, i.e., single (weakest),
double, or triple (strongest) bonds
Polar Molecules
Molecules of water, ammonia (a reactant in the production of nitrogen fertilizers), and sulfur dioxide (an industrial pollutant contributing to acid rain formation) all have polar covalent bonds holding their atoms together. If a molecule
contains polar covalent bonds, the entire molecule may have a positive end and
a negative end, in which case it would be classified as a polar molecule. Not all
molecules containing polar covalent bonds are polar molecules, however.
Carbon tetrachloride, CCl4(l), contains four polar covalent bonds and HCl(g) has
one. However, HCl(g) is a polar molecule (Figure 4(a)), while CCl4(l) is not
(Figure 4(b)).
With practice you will be able to predict which molecules are likely to be
polar by looking at their chemical or structural formulas. Table 3 shows which
types of compounds tend to be polar molecules.
polar molecule: a molecule that is
slightly positively charged at one end and
slightly negatively charged at the other
because of electronegativity differences
(a)
–
+
Table 3: Guidelines for Predicting Polar and Nonpolar Molecules
Polar
Nonpolar
Type
Description
Examples
AB
HAx
AxOH
OxAy
NxAy
diatomic compounds
any molecule with a single H
any molecule with an OH at one end
any molecule with an O at one end
any molecule with an N at one end
CO(g)
HCl(g)
C2H5OH(l)
H20(l), OCl2(g)
NH3(g), NF3(g)
Ax
CxAy
all elements
most carbon compounds (including organic
solvents, fats, and oils)
Cl2(g), N2(g)
CO2(g), CH4(g)
(b)
–
+
–
–
–
Practice
Understanding Concepts
5. Some molecules contain polar covalent bonds but are not themselves
polar. Explain, with diagrams, how this is possible.
6. Use Lewis structures and electronegativity values to explain why
methane, CH4(g), is not a polar substance.
Figure 4
(a) A molecule of HCl acquires oppositely
charged ends because of the polar covalent nature of the H-Cl bond.
(b) Since a molecule of CCl4 is quite symmetrical, it lacks oppositely charged ends and
is not polar.
Chemical Bonding 85
Intermolecular Forces
van der Waals forces: weak intermolecular attractions, including London dispersion
forces and dipole–dipole forces
dipole–dipole force: an attractive force
acting between polar molecules
London dispersion force: an attractive
force acting between all molecules, including
nonpolar molecules
H
H
(a)
O
(b)
+
+
–
Figure 5
(a) The oxygen atom in a water molecule has
two single polar covalent bonds and two
lone pairs of electrons. This results in a
V-shaped molecule.
(b) A water molecule has a slight negative
charge on the oxygen atom and slight
positive charges on the hydrogens.
hydrogen bond: a relatively strong
dipole–dipole force between a positive
hydrogen atom of one molecule and a highly
electronegative atom (F, O, or N) in another
molecule
86
Chapter 2
As mentioned earlier, the properties of molecular compounds cannot be
explained simply by covalent bonds. If covalent bonds were the only forces at
work, most molecular compounds would be gases, as there would be no attraction between the molecules strong enough to order the molecules into solids or
liquids. The concepts of the polar molecule and small charges on atoms help
explain why these molecular compounds are not all gases at SATP.
The existence of intermolecular forces was first suggested by Johannes van
der Waals, toward the end of the 19th century, to explain why gases liquefy when
they are cooled. Many different observations, such as surface tension and heat of
vaporization, provide evidence that there are three kinds of intermolecular
forces, each with different strengths. Two of these are classified as van der Waals
forces (in honour of Johannes van der Waals): dipole–dipole forces and London
dispersion forces. The third intermolecular force—hydrogen bonding—is generally not grouped with the other two.
Dipole–dipole forces are the forces of attraction between oppositely charged
ends of polar molecules (e.g., HCl). The positive end of each molecule attracts
the negative ends of neighbouring molecules—rather like a weak version of ionic
bonds. London dispersion forces, by contrast, exist between all molecules—both
polar and nonpolar—so are the only intermolecular forces acting between nonpolar molecules. Chemists believe that the weak attractive forces are the result of
temporary displacements of the electron “cloud” around the atoms in a molecule, resulting in extremely short-lived dipoles. Because the dipoles last for only
tiny fractions of a second, the attraction is continually being lost, so the forces are
very weak. Dipole–dipole forces, when they exist, tend to be much stronger than
London dispersion forces. You will learn more about these forces later in this
course.
Hydrogen Bonds
Water, a polar molecule, consists of one atom of oxygen bound by single covalent bonds to two hydrogen atoms. Its structure is simple, but water exhibits
some rather unusual properties: higher than expected melting and boiling
points, high vapour pressure, high surface tension, and the ability to dissolve a
large number of substances. To explain these properties, we must consider the
intermolecular forces that exist between water molecules. As a result of the large
difference in electronegativity between hydrogen and oxygen, the O–H bonds in
a molecule of water are highly polar covalent. The oxygen atom in a molecule of
water carries a slight negative charge while the hydrogen atoms carry a small positive charge (Figure 5). As a result, the hydrogen atoms of one water molecule
exert a strong force of attraction on the oxygen atom of neighbouring water molecules. This kind of intermolecular force is referred to as a hydrogen bond.
Hydrogen bonds occur among highly polar molecules containing F–H, O–H,
and N–H bonds. Although a hydrogen bond is similar to a dipole–dipole force,
it is stronger than any of the van der Waals forces.
Water molecules have a tendency to “stick together” because of hydrogen
bonding. This is one of the reasons why water appears to “climb” the sides of a
narrow tube (Figure 6). We can also explain water’s high melting and boiling
points using the concept of the hydrogen bond: Large amounts of energy are
required to break the hydrogen bonds in the solid and liquid states.
2.4
Try This
Activity
Molecular Models
Molecular models provide a way of representing molecules in three dimensions. In this activity, you will practise various ways of modelling molecules.
Materials: molecular model kits
(a) Draw Lewis structures to represent the following molecules:
water, hydrogen sulfide, hydrogen fluoride, methane, carbon
dioxide, ammonia, and nitrogen.
(b) Write the structural formula for each compound.
• Assemble models to represent each of the substances.
(c) Which of the molecules required double or triple covalent
bonds? Explain why.
(d) Which of the molecules would you expect to be polar?
(e) Give examples of other compounds that you would expect to
have similar shapes to water, ammonia, and hydrogen fluoride.
Explain your answers using diagrams and models.
Most substances are more dense in the solid state than in the liquid state.
However, you have probably seen solid ice cubes floating in a glass of water. If
they float, they must be less dense than the liquid. Hydrogen bonds enable us to
explain this unusual behaviour: ice is less dense than liquid water because it
forms an open lattice structure when it freezes, with a great deal of empty space
between the molecules (Figure 7). This lower density of water in the solid state
enables aquatic life to survive the winter. Instead of freezing from the bottom up,
lakes freeze from the top down, creating an insulating layer of ice at the top. This
is one of the many reasons why scientists consider water to be indispensable for
the existence of life.
Figure 6
The curved surface of water in a container is
called the meniscus.
H
O
H
H
H
H
O
O
H
H
O
H
O
H
H
O
H
H
Figure 7
In ice, hydrogen bonds between the molecules result in a regular hexagonal crystal
structure. The • • • H– represents a hydrogen
(proton) being shared between two pairs of
electrons. In liquid water these bonds constantly break and reform as molecules move
past each other.
Chemical Bonding 87
Practice
Understanding Concepts
7. Water is known for its many anomalous properties. Use your knowledge of intermolecular forces and intramolecular bonding to explain
theoretically why lakes freeze from top to bottom.
8. Using Table 3 (page 85), predict whether each of the following molecules would be polar or nonpolar.
(a) CH3OH(l)
(d) PCl3(s)
(b) I2(s)
(e) HC2H3O2(aq)
(c) HBr(g)
(f) CCl4(l)
9. Which compound(s) in question 8 were classified as nonpolar but
contain polar covalent bonds? Explain how this is possible.
10. How are hydrogen bonds different from other dipole–dipole forces?
11. Predict the type of intermolecular force that would exist between
molecules of the following:
(a) I2(s)
(b) H2O(l)
(c) NH3(g)
12. Explain each of your predictions in question 11 theoretically.
Applying Inquiry Skills
13. A student was provided with seven sample liquids and asked to
investigate whether each liquid is affected by a positively or negatively charged object. Complete the Prediction, Evidence, Analysis,
and Synthesis sections of the investigation report.
Question
Does a charged object have an effect on a thin stream of the following liquids: NCl3, H2O, Br2, CCl4, CH3OH, H2O2, and vegetable oil?
Prediction
(a) Using the generalizations of polar and nonpolar molecules, predict an answer to the Question.
Experimental Design
A thin stream of each liquid was tested by holding a positively
charged object near the liquid stream. The procedure was repeated
with a negatively charged object.
Evidence
Table 4: Effects of Charged Objects on Seven Liquids
Samples
Positive charge
Negative charge
1-3
no effect
no effect
4-7
stream moved toward charged object
stream moved toward charged object
(b) Which of the seven substances could be samples 1 to 3? Which
could be 4 to 7?
Analysis
(c) Use the Evidence to answer the Question.
Synthesis
(d) Provide a theoretical explanation to justify your answer in (b).
(e) Speculate as to why the liquids were affected by both positive
and negative charges.
88
Chapter 2
2.5
Section 2.4 Questions
Understanding Concepts
1. Explain why the bonding in compounds cannot be described only
in terms of covalent bonds and ionic bonds.
2. How does the sharing of the bonding electrons in a molecule of
Cl2 compare to that in a molecule of HCl? Explain your answer,
including diagrams.
3. Both boron and nitrogen form compounds with chlorine. In each
case, the formula has the general form AX3.
(a) Classify each of these compounds as ionic or molecular.
Justify your answer.
(b) How are the B—Cl bonds and N—Cl bonds similar? How are
they different?
(c) What other properties can you predict for the two compounds? Use the concepts of electronegativity and Lewis
structures to justify your answers.
4. Carbon tetrachloride, CCl4(l), is a nonpolar substance, although
carbon and chlorine have different electronegativities. Explain
this lack of polarity.
2.5
The Names and Formulas of
Compounds
Prior to the late 1700s, there was no systematic method of naming compounds.
Substances were named in a variety of ways. In some cases, the name referred to
the use of the compound; in other cases, it incorporated an obvious property, or
perhaps referred to the sources of the substance. These common names gave
little, if any, information about the composition of the compound.
The classical system of nomenclature based on the Latin names of elements
was devised in 1787 by French chemist Guyton de Morveau (1737–1816). As
more and more compounds were discovered, however, it became apparent that a
more comprehensive and consistent system of nomenclature was needed.
Today, we are aware of a very large number of compounds, and there is every
possibility that more will be discovered in the future. We need a system that is
easy to use and provides information on the composition of every compound. As
a result, scientists now use the system of chemical nomenclature chosen by the
International Union of Pure and Applied Chemistry (IUPAC) (sometimes just
called the IUPAC system) when naming compounds. Just as many immigrants to
Canada have learned a new language, you will soon become familiar with the language of chemical nomenclature. Chemical nomenclature provides us with a systematic means of both naming and identifying compounds.
In describing a chemical compound, chemists may use its name or its formula. The formula of an ionic compound tells us the ratio of elements present.
The formula of a molecular compound indicates the number of atoms of each
element present in a molecule of the compound. Every chemical compound has
a unique name. Knowing the names of common chemicals and being able to
write their formulas correctly are useful skills in chemistry. The common and
IUPAC names for some familiar compounds are shown in Table 1.
DID YOU KNOW ?
Salt of the Gods
Ammonium chloride was at one time referred
to as sal ammoniac. The Greeks built a
temple dedicated to the God Ammon after
parts of Egypt had been conquered by
Alexander the Great (356–323 B.C.). Dried
camel manure was used as fuel for fires in
this temple. Over many years of manure
burning, a white, crystalline, saltlike material
was gradually deposited on the walls of the
temple. This deposit became known as “salt
of Ammon” or sal ammoniac.
chemical nomenclature: a system,
such as the one approved by IUPAC, of
names used in chemistry
Table 1: Common and IUPAC Names
for Some Compounds
Common name
IUPAC name
quicklime
calcium oxide
laughing gas
dinitrogen monoxide
saltpetre
sodium nitrate
potash
potassium carbonate
muriatic acid
hydrochloric acid
baking soda
sodium hydrogen carbonate
(sodium bicarbonate)
cream of tartar
potassium hydrogen tartrate
grain alcohol
ethanol or ethyl alcohol
sal ammoniac
ammonium chloride
Chemical Bonding 89
2.5
Section 2.4 Questions
Understanding Concepts
1. Explain why the bonding in compounds cannot be described only
in terms of covalent bonds and ionic bonds.
2. How does the sharing of the bonding electrons in a molecule of
Cl2 compare to that in a molecule of HCl? Explain your answer,
including diagrams.
3. Both boron and nitrogen form compounds with chlorine. In each
case, the formula has the general form AX3.
(a) Classify each of these compounds as ionic or molecular.
Justify your answer.
(b) How are the B—Cl bonds and N—Cl bonds similar? How are
they different?
(c) What other properties can you predict for the two compounds? Use the concepts of electronegativity and Lewis
structures to justify your answers.
4. Carbon tetrachloride, CCl4(l), is a nonpolar substance, although
carbon and chlorine have different electronegativities. Explain
this lack of polarity.
2.5
The Names and Formulas of
Compounds
Prior to the late 1700s, there was no systematic method of naming compounds.
Substances were named in a variety of ways. In some cases, the name referred to
the use of the compound; in other cases, it incorporated an obvious property, or
perhaps referred to the sources of the substance. These common names gave
little, if any, information about the composition of the compound.
The classical system of nomenclature based on the Latin names of elements
was devised in 1787 by French chemist Guyton de Morveau (1737–1816). As
more and more compounds were discovered, however, it became apparent that a
more comprehensive and consistent system of nomenclature was needed.
Today, we are aware of a very large number of compounds, and there is every
possibility that more will be discovered in the future. We need a system that is
easy to use and provides information on the composition of every compound. As
a result, scientists now use the system of chemical nomenclature chosen by the
International Union of Pure and Applied Chemistry (IUPAC) (sometimes just
called the IUPAC system) when naming compounds. Just as many immigrants to
Canada have learned a new language, you will soon become familiar with the language of chemical nomenclature. Chemical nomenclature provides us with a systematic means of both naming and identifying compounds.
In describing a chemical compound, chemists may use its name or its formula. The formula of an ionic compound tells us the ratio of elements present.
The formula of a molecular compound indicates the number of atoms of each
element present in a molecule of the compound. Every chemical compound has
a unique name. Knowing the names of common chemicals and being able to
write their formulas correctly are useful skills in chemistry. The common and
IUPAC names for some familiar compounds are shown in Table 1.
DID YOU KNOW ?
Salt of the Gods
Ammonium chloride was at one time referred
to as sal ammoniac. The Greeks built a
temple dedicated to the God Ammon after
parts of Egypt had been conquered by
Alexander the Great (356–323 B.C.). Dried
camel manure was used as fuel for fires in
this temple. Over many years of manure
burning, a white, crystalline, saltlike material
was gradually deposited on the walls of the
temple. This deposit became known as “salt
of Ammon” or sal ammoniac.
chemical nomenclature: a system,
such as the one approved by IUPAC, of
names used in chemistry
Table 1: Common and IUPAC Names
for Some Compounds
Common name
IUPAC name
quicklime
calcium oxide
laughing gas
dinitrogen monoxide
saltpetre
sodium nitrate
potash
potassium carbonate
muriatic acid
hydrochloric acid
baking soda
sodium hydrogen carbonate
(sodium bicarbonate)
cream of tartar
potassium hydrogen tartrate
grain alcohol
ethanol or ethyl alcohol
sal ammoniac
ammonium chloride
Chemical Bonding 89
Binary Ionic Compounds
binary compound: a compound composed of two kinds of atoms or two kinds of
monatomic ions
valence: the charge on an ion
The simplest compounds are called binary compounds. Binary ionic compounds
consist of two types of monatomic ions (ions consisting of one charged atom).
In the formula of a binary ionic compound, the metal cation is always written
first, followed by the nonmetal anion. (This reflects the periodic table: metals to
the left, nonmetals to the right.) The name of the metal is stated in full and the
name of the nonmetal ion has an -ide suffix; for example, NaCl(s) is sodium chloride and LiBr(s) is lithium bromide. Binary ionic compounds can be made up of
more than two ions, providing they are of only two kinds: aluminum oxide is
Al2O3(s).
If we know what ions make up a compound, we can often predict the compound’s formula. First, we determine the charges on each type of ion making up
the compound. The charge on an ion is sometimes called the valence. We then
balance the charges to determine the simplest ratio in which they combine. We
can predict the charge on the most common ion formed by each representative
element by counting the number of electrons that would have to be gained or lost
to obtain a stable octet. (Ionic charges for the most common ions of elements are
shown in the periodic table at the back of this text.) For example, the compound
magnesium chloride, a food additive used to control the colour of canned peas,
is composed of the ions Mg2+ and Cl–. For a net charge of zero, the ratio of magnesium to chlorine ions must be 1:2. The formula for magnesium chloride is
therefore MgCl2(s).
Sample Problem 1
Predict the formula for magnesium oxide—a source of dietary magnesium when
used as a food additive.
Solution
Charge on magnesium ion: 2+
Charge on oxygen ion: 2–
The ratio of magnesium ions to oxide ions that produces a net charge of zero is 1:1.
The formula is therefore MgO.
You may be familiar with a method known as the crisscross rule for predicting the formula of an ionic compound. The crisscross rule works as follows:
1. Write the symbol of each of the elements in the order in which they appear
in the name of the compound.
2. Write the valence number (electrons lost or gained in forming that element’s most stable ion) above the symbol of each of the elements.
3. Crisscross the numbers written above the symbols such that the valence
number of one element becomes a subscript on the other.
4. Divide each subscript by the highest common factor. The resulting subscripts indicate the ratio of ions present in the compound.
5. Omit any subscript equal to 1 from the formula.
For example, if we were to use the crisscross rule to predict the formula of
magnesium chloride, we would go through these steps.
1. Write symbols in order:
Mg Cl
90
Chapter 2
2.5
2. Write valences above symbols:
2
1
Mg
Cl
3. Crisscross the valences, making them subscripts:
Cl2
Mg1
4. Divide subscripts by highest common factor (= 1):
Cl2
Mg1
5. Remove 1 subscripts:
MgCl2
Sample Problem 2
Use the crisscross rule to predict the formula of barium sulfide.
Solution
2
2
Ba
S
S2
Ba2
The formula of barium sulfide is BaS.
Most transition metals and some representative metals can form more than
one kind of ion. Metals that can have more than one valence, or charge, are classified as multivalent. For example, iron can form an Fe2+ ion or an Fe3+ ion,
although Fe3+ is more common. The periodic table at the end of this text shows
the most common ion of each element first, with one alternative ion charge
below. It does not list all of the possible ions of the element.
Copper is a multivalent element. It is capable of bonding with chlorine in
two different ratios to form two different chloride compounds: CuCl and CuCl2.
How are the names of these compounds different? The IUPAC system of naming
compounds containing multivalent ions is very simple. The name of the metal
ion includes the charge on the ion, indicated by Roman numerals in brackets.
Consequently, CuCl(s) (in which copper has a charge of 1+) is copper(I) chloride,
and CuCl2(s) (in which copper has a charge of 2+) is copper(II) chloride. This
system of naming is sometimes referred to as the Stock system.
To determine the chemical name of a compound containing a multivalent metal
ion, we have to figure out the necessary charge on that ion to yield a net charge of zero.
If we are given the formula, we simply have to calculate the equivalent negative charge.
The metal’s charge is then written, in Roman numerals, after the name of the metal.
For example, if you were asked to name the compound MnO2, using IUPAC nomenclature, you would first look at the charge on the nonmetal ions. In this case, the charge
on each O is 2–, so the total negative charge is 4 –. The charge on the Mn ion must be
4+. Consequently, the IUPAC name for MnO2 is manganese(IV) oxide.
If the ion of a multivalent metal is not specified in a name, it is assumed that
the charge on the ion is the most common one.
multivalent: the property of having more
than one possible valence
Chemical Bonding 91
Sample Problem 3
The formula of a compound is found to be SnCl2. What is its IUPAC name?
Solution
Charge on each Cl ion: 1–
Total negative charge is 2–
Charge on Sn ion: 2+
The IUPAC name for SnCl2 is tin(II) chloride.
Sample Problem 4
The formula of a compound is found to be Fe2O3. What is its IUPAC name?
Solution
Total negative charge: 6–
Total positive charge: 6+
Charge on each Fe ion: 3+
The IUPAC name of Fe2O3 is iron(III) oxide.
The classical nomenclature system has, in the past, been used for naming
compounds containing multivalent metals with no more than two possible
charges. In this system, the Latin name for the element along with the suffix -ic
was applied to the larger charge, and the suffix -ous was applied to the smaller
charge. The compounds formed by copper and chlorine were therefore known as
cuprous chloride (CuCl(s)) and cupric chloride (CuCl2(s)).
In many industries, the classical system is still used extensively. Table 2 shows
a comparison of the classical and IUPAC names of multivalent metal ions.
Table 2: Classical and IUPAC Names of Common Multivalent Metal Ions
92
Chapter 2
Metal
Ion
Classical name
IUPAC name
iron
Fe2+
Fe3+
ferrous
ferric
iron(II)
iron(III)
copper
Cu+
Cu2+
cuprous
cupric
copper(I)
copper(II)
tin
Sn2+
Sn4+
stannous
stannic
tin(II)
tin(IV)
lead
Pb2+
Pb4+
plumbous
plumbic
lead(II)
lead(IV)
antimony
Sb3+
Sb5+
stibnous
stibnic
antimony(III)
antimony(V)
cobalt
Co2+
Co3+
cobaltous
cobaltic
cobalt(II)
cobalt(III)
gold
Au+
Au2+
aurous
auric
gold(I)
gold(II)
mercury
Hg+
Hg2+
mercurous
mercuric
mercury(I)
mercury(II)
2.5
Practice
Understanding Concepts
1. How were compounds named before the advent of systematic means
of chemical nomenclature?
2. State the common names of each of the following chemicals:
(a) hydrochloric acid
(b) sodium hydrogen carbonate
(c) dinitrogen monoxide
(d) ethanol
3. How many elements are there in each binary compound?
4. Describe the IUPAC system of naming a binary ionic compound.
5. If an element is described as“multivalent,” what characteristic does it
have?
6. Write the formula for each of the following compounds:
(a) calcium fluoride
(l) mercury(I) oxide
(b) sodium sulfide
(m) nickel(II) bromide
(c) aluminum nitride
(n) zinc oxide
(d) aluminum chloride
(o) cobalt(III) chloride
(e) potassium oxide
(p) strontium bromide
(f) calcium chloride
(q) gold(I) fluoride
(g) copper(II) sulfide
(r) lithium chloride
(h) lead(II) bromide
(s) strontium nitride
(i) silver iodide
(t) barium bromide
(j) barium nitride
(u) tin(IV) iodide
(k) iron(II) fluoride
7. Write the names of each of the following binary ionic compounds,
using the IUPAC system of chemical nomenclature:
(a) table salt, NaCl(s)
(h) chalcocite, Cu2S(s)
(b) lime, CaO(s)
(i) galena, PbS2(s)
(c) road salt, CaCl2(s)
(j) hematite, Fe2O3(s)
(d) magnesia, MgO(s)
(k) molybdite, MoO3(s)
(e) bauxite, Al2O3(s)
(l) argentite, Ag2S(s)
(f) zinc ore, ZnS(s)
(m) zincite, ZnO(s)
(g) cassiterite, SnO2(s)
8. Write the IUPAC names for the following binary ionic compounds:
(a) Na2O(s)
(g) Ni2O3(s)
(b) SnCl4(s)
(h) Ag2S(s)
(c) ZnI2(s)
(i) FeCl2(s)
(d) SrCl2(s)
(j) KBr(s)
(e) AlBr3(s)
(k) CuI2(s)
(f) PbCl4(s)
(l) NiS(s)
9. Write the chemical formulas and names (using IUPAC chemical
nomenclature) for the binary ionic compounds formed by each of the
following pairs of elements:
(a) strontium and oxygen
(b) sodium and sulfur
(c) silver and iodine
(d) barium and fluorine
(e) calcium and bromine
(f) lithium and chlorine
10. Write the chemical formulas for the following ionic compounds:
(a) mercury(II) sulfide, cinnabar ore
(b) molybdenum(IV) sulfide, molybdenite ore
(c) manganese(IV) oxide, pyrolusite ore
(d) nickel(II) bromide
Chemical Bonding 93
(e) copper(II) chloride
(f) iron(III) iodide
11. Rename each of the following compounds, using the IUPAC system
of nomenclature:
(a) ferrous sulfide
(b) plumbic bromide
(c) stannous chloride
Reflecting
12. Why is it sometimes helpful to understand a system of nomenclature
that has been replaced by another, more comprehensive, system?
Compounds with Polyatomic Ions
tertiary compound: a compound composed of three different elements
oxyanion: a polyatomic ion containing
oxygen
determines the
stem of the
ion name
ion
charge
ClO3 –
determines the suffix
and prefix to the stem
of the ion name
Figure 1
The chlorate anion
Many familiar compounds (such as sulfuric acid, H2SO4(aq), used in car batteries,
and sodium phosphate, Na3PO4(aq), a food additive typically found in processed
cheese) are composed of three different elements. Compounds of this type are
classified as tertiary compounds. (Many compounds containing polyatomic ions
consist of more than three different elements, but we will not be dealing with
these at this stage.) Tertiary ionic compounds are composed of a metal ion and
a polyatomic ion (a covalently bonded group of atoms, possessing a net charge).
We treat polyatomic ions much like regular monatomic ions when we write them
in formulas or chemical equations.
Polyatomic ions that include oxygen are called oxyanions. One example is
the nitrate ion, NO3–. (Compounds involving the nitrate ion are often used in the
processing of foods, particularly cured meats, where they are often used to control colour. Potassium nitrate and sodium nitrate are added to foods to control
the growth of microorganisms.) In determining the name of a compound containing an oxyanion, the first part of the name is easy: It is the name of the metal
cation. The second part requires more thought: We have to consider the three
parts of the ion indicated in Figure 1.
There are four polyatomic ions formed from combinations of chlorine and
oxygen. Note that all of these oxyanions have the same charge, despite the fact
that their formulas are different.
• ClO is the hypochlorite ion;
• ClO2– is the chlorite ion;
• ClO3– is the chlorate ion; and
• ClO4– is the perchlorate ion.
Note that in each name the stem is –chlor–. The suffixes and prefixes vary
according to the number of oxygen atoms in the ion, as described below.
• The per–ate oxyanion has one more oxygen atom than does the -ate
oxyanion.
• The -ite oxyanion has one fewer oxygen atom than does the -ate oxyanion.
• The hypo–ite oxyanion has one fewer oxygen atom than the -ite oxyanion.
Table 3 indicates some polyatomic ions commonly found in compounds.
You do not need to memorize each entry in this table. If you become familiar
with the -ate oxyanions, the most stable and common combination of nonmetal
and oxygen, then you need only remember the simple relationship between the
names of the ions.
The crisscross method can be applied to predict the formula of ionic compounds involving polyatomic ions. From the ion charges (whether for a single
94
Chapter 2
2.5
ion or for a polyatomic ion), determine the number of each ion necessary to yield
a net charge of zero.
Sample Problem 5
Write the formula of copper(II) nitrate. (Use Table 3 to find the charge of a
nitrate ion.)
Solution
Charge on each Cu ion: 2+
Charge on each NO3 ion: 1
2
1
Cu
NO3
Cu1
(NO3)2
The formula of copper(II) nitrate is Cu(NO3)2.
Notice how brackets are used in the formula if there is more than one of the
polyatomic ions. Brackets are not required with one polyatomic ion or with
simple compounds.
Sample Problem 6
Use IUPAC nomenclature to name CaCO3, which is commonly added to breakfast cereals as a source of calcium.
Solution
CaCO3 is calcium carbonate.
Hydrates
Consumers are sometimes surprised to find a tiny white pouch when they open
the box containing a newly purchased pair of boots. What is it? What purpose
does it serve? The pouch contains a white, crystalline powdered desiccant (a
substance that absorbs water) called silica gel (SiO2(s)). The pouch keeps the air
inside the box dry so mildew and other moulds will not grow. Similarly
absorbent compounds are included in such diverse products as powdered foods,
talcum powder, and cat litter. Many tertiary ionic compounds form crystals that
contain molecules of water within the crystal structure. Such compounds are
referred to as hydrates. When heat is applied to a hydrate, it will decompose to
produce water vapour and an associated ionic compound, indicating that the
water is loosely held to the ionic compound. The water molecules are assumed
to be electrically neutral in the compound. When this water, called water of
hydration, is removed, the product is referred to as anhydrous. Bluestone,
hydrated copper(II) sulfate, is an example of a hydrate. Its formula is written
CuSO4·5 H2O(s). Notice how the chemical formula includes both the formula of
the compound and the formula for water. This is true of the chemical formulas
for all hydrated compounds. The formula for bluestone indicates the association
of five water molecules with each unit of copper(II) sulfate (Figure 2, page 96).
The IUPAC names for ionic hydrates indicate the number of water molecules by
a Greek prefix (Table 4, page 96), so bluestone, or CuSO4·5H2O(s), is called
copper(II) sulfate pentahydrate.
Table 3: IUPAC Names and Formulas
of Some Common
Polyatomic Ions
Name
Formula
acetate
C2H3O2
bromate
BrO3
carbonate
CO32
hydrogen carbonate
(bicarbonate)
HCO3
hypochlorite
ClO–
chlorite
ClO2
chlorate
ClO3
perchlorate
ClO4
chromate
CrO42
dichromate
Cr2O7 2
cyanide
CN–
hydroxide
OH–
iodate
IO3
permanganate
MnO4
nitrite
NO2
nitrate
NO3
phosphate
PO43
hydrogen phosphite
HPO3 2
hydrogen phosphate
HPO42
dihydrogen phosphite
H2PO3
dihydrogen phosphate
H2PO4
sulfite
SO32
sulfate
SO4 2
hydrogen sulfide
(bisulfide)
HS–
hydrogen sulfite
(bisulfite)
HSO3
hydrogen sulfate
(bisulfate)
HSO4
thiosulfate
S2O3
ammonium
NH4+
hydrate: a compound that contains water
as part of its ionic crystal structure (theoretical definition); a compound that decomposes
to an ionic compound and water vapour when
heated (empirical definition)
DID YOU KNOW ?
Naming Hydrates
When researching scientific references, you
may encounter an alternative naming system
for hydrates, in which CuSO4·5H2O(s) is called
copper(II) sulfate-5-water.
Chemical Bonding 95
2–
SO 4
Cu
2+
Figure 2
In a model of the compound copper(II) sulfate
pentahydrate, the copper(II) ions are surrounded by four water molecules. The fifth
water molecule is hydrogen-bonded to the
sulfate ion.
Table 4: Prefixes Used When Naming
Hydrated Compounds
Number of water
molecules
in chemical formula
Prefix in chemical
nomenclature
1
mono
2
di
3
tri
4
tetra
5
penta
6
hexa
7
hepta
8
octa
9
nona
10
deca
Practice
Understanding Concepts
13. How many elements are there in a tertiary compound?
14. Use each of the following terms correctly in a sentence about the formation of compounds:
(a) polyatomic ion
(b) oxyanion
(c) hydrate
15. Write the IUPAC name for each of the following ionic compounds:
(a) NaNO3(s) (found in tobacco)
(b) NaNO2(s) (a meat preservative)
(c) Cu(NO3)2(s) (forms a blue solution in water)
(d) CuNO3(s) (forms a green solution in water)
(e) Al2(SO3)3(s) (a food additive in pickles)
(f) Ca(OH)2(s) (firming agent in fruit products)
(g) PbCO3(s) (cerussite, a mineral popular with collectors)
(h) Sn3(PO4)2(s) (use to fix paints to silk)
(i) Fe2(SO4)3(s) (a mineral found on Mars)
16. Write the chemical formula for each of the following ionic compounds:
(a) calcium carbonate (active ingedient in antacids)
(b) sodium bicarbonate (a foaming agent added to foods)
(c) sodium hypochlorite (a component of bleach)
(d) calcium sulfate (plaster of Paris)
(e) ammonium nitrate (used in fertilizers)
(f) ammonium phosphate (a leavening agent added to foods)
(g) copper(II) sulfate (used as a fungicide)
(h) sodium hydroxide (a strong base used as a washing agent)
(i) potassium permanganate (a traditional antiseptic)
17. Use IUPAC chemical nomenclature to name each of the following
ionic compounds containing polyatomic ions:
(a) LiClO3(s)
(n) Ag2SO4(s)
(b) BaSO4(s)
(o) Hg(BrO3)2(s)
(c) Hg2CO3(s)
(p) Fe2(CO3)3(s)
(d) Mg(NO3)2(s)
(q) NH4ClO(s)
(e) Fe(BrO3)3(s)
(r) Au(NO3)3(s)
(f) Na3PO4(s)
(s) Mg(BrO3)2(s)
(g) NH4IO3(s)
(t) NaIO(s)
(h) AuC2H3O2(s)
(u) Zn(ClO2)2(s)
(i) Zn3(PO4)2(s)
(v) SnCO3(s)
96
Chapter 2
2.5
(j)
(k)
(l)
(m)
Sb(ClO3)5(s)
MnSO3(s)
KBrO(s)
AlPO5(s)
(w)
(x)
(y)
(z)
SrSO3(s)
NiPO4(s)
Cu(C2H3O2)2(s)
Ba3(PO5)2(s)
18. Write the IUPAC name and chemical formula for each of the following
ionic compounds containing polyatomic ions:
(a) cuprous hypophosphite
(b) stannic chlorite
(c) ferrous bromate
(d) ferric chlorite
(e) plumbic sulfate
19. Name each of the following hydrated ionic compounds:
(a) bluestone, CuSO4·5 H2O(s)
(b) Na2SO4·10 H2O(s)
(c) MgSO4·7 H2O(s)
20. Write the chemical formulas for the following ionic hydrates:
(a) iron(III) oxide trihydrate (rust)
(b) aluminum chloride hexahydrate (component of antiperspirant)
(c) sodium thiosulfate pentahydrate (photographic “hypo”)
(d) cadmium(II) nitrate tetrahydrate (photographic emulsion)
(e) lithium chloride tetrahydrate (in fireworks)
(f) calcium chloride dihydrate (deicer)
21. How would you convert a hydrate to an anhydrous compound?
Naming Molecular Compounds
So far, in this section, we have been looking at the names and formulas of ionic
compounds. We now turn to molecular compounds, for which a different system
is used.
If a binary compound is formed from two nonmetals, it is classified as a molecular compound. Even though there are fewer nonmetals than metals, there is a wide
variety of compounds formed from the combination of two nonmetals, because
two nonmetals may combine to form more than one compound. For example, N2O,
NO, and NO2 are three of the several binary compounds that can be formed from
nitrogen and oxygen. Each compound has different properties, and so different uses.
In naming compounds formed from two nonmetals, a Greek prefix is
attached to the name of each element in the binary compound indicating the
number of atoms of that element in the molecule. The common prefixes and
their numerical equivalences are shown in Table 5. If there is only one of the first
type of atom, we leave out the prefix “mono.”
Suppose you are asked to write the IUPAC name for the chemical compound
represented by the formula N2O. Looking at the first element, you can see that
the subscript after the nitrogen is two, so the prefix for nitrogen is “di.” Looking
at the second element, you can see that there is only one oxygen atom, so the
prefix for oxygen will be “mono.” Therefore, the formula’s IUPAC name is dinitrogen monoxide.
Table 5: Prefixes Used When Naming
Binary Covalent Compounds
Subscript in
chemical formula
Prefix in chemical
nomenclature
1
mono
2
di
3
tri
4
tetra
5
penta
6
hexa
7
hepta
8
octa
9
nona
10
deca
Sample Problem 7
What is the IUPAC name for the chemical compound CF4?
Solution
C: carbon (not monocarbon)
F: tetrafluoride
The IUPAC name for CF4 is carbon tetrafluoride.
Chemical Bonding 97
Once again, hydrogen is an exception to this rule: The common practice is
not to use the prefix system for hydrogen. For example, we do not call H2S dihydrogen sulfide, but simply hydrogen sulfide.
Practice
Understanding Concepts
22. Write the chemical formula for each of the following molecules:
(a) nitrogen
(n) sulfur tetrafluoride
(b) carbon dioxide
(o) phosphorus pentachloride
(c) carbon monoxide
(p) disulfur dichloride
(d) nitrogen dioxide
(q) carbon tetrachloride
(e) nitrogen monoxide
(r) sulfur trioxide
(f) dinitrogen oxide
(s) sulfur hexafluoride
(g) dinitrogen tetroxide
(t) chlorine dioxide
(h) sulfur dioxide
(u) dinitrogen pentoxide
(i) diiodine pentoxide
(v) phosphorus trichloride
(j) silicon tetrafluoride
(w) silicon tetrachloride
(k) boron trifluoride
(x) carbon disulfide
(l) phosphorus triiodide
(y) phosphorus pentabromide
(m) diphosphorus pentoxide
(z) carbon tetrafluoride
23. Name the compound indicated by each of the following formulas:
(a) SF6(g)
(f) IF7(g)
(b) N2O3(g)
(g) BF3(g)
(c) NO2(g)
(h) P2S5(s)
(d) PCl3(l)
(i) P2O5(s)
(e) PCl5(s)
Naming Acids
You are already familiar with many acids and bases. You may have taken acetylsalicylic acid (also known as ASA, or Aspirin) to treat a headache, poured a little
acetic acid solution (vinegar) on your fish and chips, or used an ammonia solution to clean the smears from a mirror. You also know one of the tests for acids
and bases: acids turn blue litmus red, and bases turn red litmus blue. But you
might have been somewhat mystified about the naming of these substances.
Acids are well-known, long-established chemicals. They were originally named
decades or even centuries ago, and the use of traditional names persists. The
International Union of Pure and Applied Chemistry suggests that names of acids
should be derived from the IUPAC name for the compound. According to this rule,
sulfuric acid would be called aqueous hydrogen sulfate. However, it is difficult to get
people to change to new names when the old, familiar names are so widely used.
Let us take, as a simple example, HCl(g). It is a binary compound formed
from a combination of hydrogen and a halogen. When a gas, it is named
hydrogen chloride. When it is dissolved in water, the resulting aqueous solution
displays a set of specific properties called acidic, and the name of the substance
changes. Binary acids are classically named by using the prefix hydro- with the
stem of the name of the most electronegative element and the ending -ic. The
name “hydrogen” does not appear. Instead, the word “acid” is added after the
hydro-stem-ic combination, as indicated in Table 6. Consequently, the classical
name for HCl(aq) is hydrochloric acid. Its IUPAC name is aqueous hydrogen
chloride. Similarly, HBr(aq) is hydrobromic acid or aqueous hydrogen bromide.
Note that the difference between the solution and the pure binary compound
is indicated by the presence or absence of the subscript (aq) in the formula.
98
Chapter 2
2.5
Table 6: Naming Systems for Binary Acids
Formula
Classical name
IUPAC name
HF(aq)
hydrofluoric acid
aqueous hydrogen fluoride
HCl(aq)
hydrochloric acid
aqueous hydrogen chloride
HBr(aq)
hydrobromic acid
aqueous hydrogen bromide
HI(aq)
hydroiodic acid
aqueous hydrogen iodide
H2S(aq)
hydrosulfuric acid
aqueous hydrogen sulfide
A second group of acids is named (by the classical system) in the same way as
binary acids. In this group, the IUPAC names for the polyatomic ions end in -ide
(e.g., the cyanide ion, CN–). Looking at Table 6, you can see that the classical
name for the acidic solution HCN(aq) will be hydrocyanic acid.
A third group of acids is formed from various combinations of oxyanions (negative polyatomic ions consisting of a nonmetal plus oxygen) with hydrogen. Perhaps
the best-known example is H2SO4(aq), or sulfuric acid, which is one of the most
widely produced industrial chemicals in the world. It is used to make pharmaceuticals, detergents, and dyes, and is a component of car batteries. Phosphoric acid,
H3PO4(aq), is another example of an acid formed from an oxyanion and hydrogen.
Phosphoric acid is an ingredient in soft drinks, is a reagent in the manufacture of
fertilizers, and can be used as a rust remover. These acids are classified as oxyacids
because they incorporate oxyanions. Table 7 compares the classical name to the
IUPAC name for a series of acids derived from a chlorine-based oxyanion.
oxyacid: an acid containing oxygen,
hydrogen, and a third element
Table 7: Classical and IUPAC Nomenclature System for ChlorineBased Oxyacids
Classical name
IUPAC name
Formula
perchloric acid
aqueous hydrogen perchlorate
HClO4(aq)
chloric acid
aqueous hydrogen chlorate
HClO3(aq)
chlorous acid
aqueous hydrogen chlorite
HClO2(aq)
hypochlorous acid
aqueous hydrogen hypochlorite
HClO(aq)
hydrochloric acid
aqueous hydrogen chloride
HCl(aq)
The classical names for oxyacids can be derived according to the simple rules
in Table 8.
Table 8: Rules for Naming Acids and Oxyanions
Name of
oxyanion
Example
Formula
Classical
name of acid
Example
per–ate
persulfate
SO52
per–ic acid
persulfuric acid
-ate
sulfate
-ic acid
sulfuric acid
-ite
sulfite
-ous acid
sulfurous acid
hypo–ite
hyposulfite
SO42
SO32
SO22
hypo–ous acid
hyposulfurous acid
When naming oxyacids, we omit the word hydrogen and add the word
“acid.” For example, to name the acidic solution with the formula HNO2(aq), we
would first consider the IUPAC name: hydrogen nitrite (from Table 3, or from
the Table of Polyatomic Ions in Appendix C). “Nitrite” changes to “nitrous,” we
drop the “hydrogen” from the front of the name, and add “acid” to the end. Thus,
HNO2(aq) is called nitrous acid.
Chemical Bonding 99
If you were asked to write the formula for an acid, you would first have to
figure out the names of the ions involved, then their symbols or formulas, then
their ratio. For example, what is the formula for phosphoric acid? The -ic ending
indicates the presence of the -ate oxyanion of phosphorus: phosphate. The phosphate oxyanion is PO43– with a charge of 3–. The cation in oxyacids is always
hydrogen, which has a charge of 1+. To find the ratio of the ions, use the crisscross method:
1
3
H
(PO4)
This gives the subscripts for each ion.
(PO4)1
H3
Divide each subscript by the highest common factor—in this case, 1:
(PO4)1
H3
The hydrogen ions and the phosphate oxyanions will combine in a ratio of 3:1.
Therefore, the correct formula is H3PO4(aq).
Sample Problem 8
What is the formula for the oxyacid sulfurous acid?
Solution
sulfurous indicates sulfite ion: SO32
1
2
(SO3)1
H2
The formula for sulfurous acid is H2SO3(aq).
Sample Problem 9
What is the formula for the oxyacid hypochlorous acid?
Solution
hypochlorous indicates one fewer oxygen atom than a chlorite ion: ClO–
1
1
(ClO)1
H1
The formula for hypochlorous acid is HClO(aq).
Naming Bases
Chemists have discovered that all aqueous solutions of ionic hydroxides are
bases. (You will learn more about bases in Chapter 8.) Other solutions have also
been classified as bases, but for the time being we will restrict our exploration of
bases to aqueous ionic hydroxides such as NaOH(aq) and Ba(OH)2(aq). Notice
that these bases are formed of a combination of a metal cation with one or more
hydroxide anions. The name of the base is the name of the ionic hydroxide: in
this case, aqueous sodium hydroxide and aqueous barium hydroxide.
100 Chapter 2
2.5
Practice
Understanding Concepts
24. Write the chemical formulas for
(a) aqueous hydrogen chloride
(b) hydrochloric acid
(c) aqueous hydrogen sulfate
(d) sulfuric acid
(e) aqueous hydrogen acetate
(f) acetic acid
the following acids:
(g) aqueous hydrogen nitrite
(h) nitric acid
(i) hydrobromic acid
(j) hyposulfurous acid
(k) hydroiodic acid
(l) aqueous hydrogen perchlorate
25. Name each of the following compounds,
the IUPAC nomenclature systems:
(a) H2SO3(aq)
(d) H2CO3(aq)
(b) H3PO4(aq)
(e) H2S(aq)
(c) HCN(aq)
(f) HCl(aq)
using both the classical and
(g) HCN(aq)
(h) H2SO4(aq)
(i) H3PO4(aq)
26. Write the names of the following bases:
(a) KOH(aq)
(b) CaOH2(aq)
27. Write the formulas of the following bases:
(a) aqueous magnesium hydroxide
(b) aqueous sodium hydroxide
(c) aqueous aluminum hydroxide
Section 2.5 Questions
Understanding Concepts
1. The atmosphere of Saturn contains traces of ammonia, while
hydrogen cyanide is a component of the atmosphere of one of
Saturn’s many moons.
(a) Write formulas for each of these compounds.
(b) Classify each substance as ionic or covalent.
(c) Which compound contains both ionic and covalent bonds?
How could you verify this prediction?
(d) If hydrogen cyanide is added to water, what substance is
formed? If the substance were tested with litmus paper, what
would you expect the results to be?
2. Caustic potash (KOH) is a component of oven cleaners. Muriatic
acid (HCl) is used to clean mortar from bricks.
(a) Classify each of these compounds. Provide your reasoning.
(b) Create electron dot diagrams and structural formulas for
each compound.
(c) Use the IUPAC system to name each compound.
(d) What can you predict about the reactivity of these substances? How would this affect the way you handle the substances? Explain your answer.
3. (a) Write the IUPAC names and formulas for as many compounds as you can, using only the following elements: K, C,
H, F, Mg, O, Cl, and Na.
(b) State the common name, if there is one, for each compound.
(c) Classify each compound as polar or nonpolar. In each case,
use Lewis structures and periodic trends to justify your
classification.
(d) Classify the compounds as binary, tertiary, acidic, or basic.
(e) Indicate which compounds, if any, contain both ionic and
covalent bonding.
Chemical Bonding 101
Summary
Chapter 2
Key Terms
Key Expectations
Throughout this chapter, you have had the
opportunity to do the following:
• Identify chemical substances and reactions in everyday
use or of environmental significance. (all sections)
• Explain how different elements combine to form
covalent and ionic bonds, using the octet rule. (2.2)
• Demonstrate an understanding of the formation of
ionic and covalent bonds and explain the properties of
the products. (2.2, 2.3)
• Predict the ionic character or polarity of a given bond
using electronegativity values, and represent the formation of ionic and covalent bonds using diagrams. (2.2,
2.3, 2.4)
• Draw Lewis structures, construct molecular models,
and give the structural formulas for compounds containing single and multiple bonds. (2.3)
• Relate common names of substances to their systematic
(IUPAC) names. (all sections)
• Write, using IUPAC or traditional (classical) systems,
the formulas of binary and tertiary compounds,
including those containing elements with multiple
valences, and recognize the formulas in various contexts. (2.5)
• Use appropriate scientific vocabulary to communicate
ideas related to chemical reactions. (all sections)
binary compound
bonding capacity
chemical bond
chemical nomenclature
coordinate covalent
bond
covalent bond
crystal lattice
diatomic molecule
dipole–dipole force
electrical conductivity
electrolyte
electron dot diagram
formula unit
hydrate
hydrogen bond
intermolecular force
intramolecular force
ionic bond
ionic compound
Make a
Summary
Summarize the concepts presented in this chapter by constructing and
completing a table like Table 1. Try to include as many examples of compounds as possible for each type of intramolecular bond.
Table 1: Summarizing Bonds and Forces
Compound
Properties
Lewis
structure
Intramolecular
bond type
Polarity
Reflectyour
on
Learning
Revisit your answers to the Reflect on Your Learning questions
at the beginning of this chapter.
• How has your thinking changed?
• What new questions do you have?
102 Chapter 2
Intermolecular
forces
Lewis structure
Lewis symbol
London dispersion force
lone pair
molecular compound
multivalent
octet rule
oxyacid
oxyanion
polar covalent bond
polar molecule
polyatomic ion
polyatomic molecule
stable octet
structural formula
tertiary compound
valence
van der Waals forces
Chapter 2
Review
Understanding Concepts
1. Describe the relationship between the positions of two
reacting elements in the periodic table and the type of
compound and bond they form.
2. How does an ionic bond differ from a covalent bond?
3. (a) Briefly summarize and explain the properties of
ionic and molecular compounds.
(b) Explain why electrical conductivity is generally a
suitable test for ionic compounds.
4. (a) What is intramolecular bonding? Describe the main
types with examples.
(b) What is the relationship between the type of bond
and the properties of a substance?
5. For each chemical bond or force listed below, indicate
which types of entities are involved.
(a) covalent bond
(b) dipole–dipole force
(c) hydrogen bonds
(d) ionic bond
6. How does the information conveyed by the chemical
formulas for ionic compounds differ from the information conveyed by the chemical formulas for molecular
compounds? Include examples with your explanation.
7. Explain why halogens tend to form diatomic molecules.
8. Sketch the bonding continuum. Indicate where Na—Cl
and Cl—Cl bonds would fit along the continuum.
Provide a theoretical explanation of your answer.
9. Draw Lewis symbols (electron dot diagrams) to represent a single atom of each of the following elements:
(a) Ca
(e) S
(b) Al
(f) Br
(c) K
(g) Ne
(d) N
10. For each of the following substances, predict whether
the bonds present are ionic, covalent, or polar covalent:
(d) Fe2O3(s)
(a) I2 (s)
(b) SO2(g)
(e) KBr(s)
(c) OCl2(g)
(f) SrO(s)
11. The most common oxides of Period 3 elements are as
follows: Na2O(s), MgO(s), Al2O3(s), SiO2(s), P2O5(s),
SO2(g), and Cl2O(g).
(a) Classify the oxides as either ionic or molecular.
(b) Use electron dot diagrams or Lewis structures to
show the formation of each compound.
(c) What differences would you expect to observe in
the properties of each of the compounds?
(d) How is the difference in electronegativity of the
constituent elements related to the properties of the
compound?
12. Create a table like Table 1 and complete it for the following compounds: HF(g), BCl3(g), SiH4(g), CCl4(l),
NCl3(g), H2O2(l), CO2(g), HCN(g)
Table 1: Structures of Covalent Compounds
Compound
Lewis
structure
Structural
formula
Types
of bonds
13. Use electronegativity values to predict the polarity of
the bond formed by each of the following pairs of elements. In each case, represent the formation of the
bonds using diagrams.
(a) carbon and chlorine
(b) calcium and fluorine
(c) aluminum and chlorine
(d) silicon and oxygen
(e) carbon and oxygen
14. Use Lewis structures and structural formulas to represent the following compounds. Indicate which compounds, if any, involve coordinate covalent bonds.
(d) H2O(l)
(a) NH4Cl(s)
(b) BF3(g)
(e) NH3BF3(g)
(c) NH3(g)
15. Write the chemical formula (plus state) for each of the
following substances:
(a) sodium hydrogen sulfate (toilet bowl cleaner)
(b) sodium hydroxide (lye, drain cleaner)
(c) carbon dioxide (dry ice, soda pop)
(d) acetic acid (vinegar)
(e) sodium thiosulfate pentahydrate (photographic
“hypo”)
(f) sodium hypochlorite (laundry bleach)
(g) octasulfur (vulcanizing rubber)
(h) potassium nitrate (meat preservative)
(i) phosphoric acid (rust remover)
(j) iodine (disinfectant)
(k) aluminum oxide (alumina, aluminum ore)
(l) potassium hydroxide (caustic potash)
(m) aqueous hydrogen carbonate (carbonated beverages)
16. Write the chemical formula for each of the following
substances:
(a) magnesium bromide
(b) carbon disulfide
(c) mercury(II) nitrite
Chemical Bonding 103
(d) hydrochloric acid
(e) lithium hydroxide
(f) silver carbonate
(g) aluminum perchlorate
(h) copper(II) sulfate
(i) sulfur trioxide
(j) nickel(III) phosphate
(k) magnesium oxide
(l) dinitrogen monoxide
(m) iron(II) persulfate
(n) carbonic acid
(o) calcium hydroxide
(p) zinc hypochlorite
(q) lead(IV) perchlorate
(r) phosphorous pentabromide
(s) arsenic(V) chloride
(t) bismuth(III) nitrate
(u) sodium hypochlorite
(v) oxygen dichloride
(w) tin(II) bromide
(x) sulfuric acid
(y) potassium hydroxide
(z) barium carbonate
17. Write the chemical formula for each of the following
substances:
(a) ammonium dihydrogen phosphite
(b) lithium hydrogen sulfite
(c) potassium hydrogen sulfate
(d) barium chloride trihydrate
(e) sodium dihydrogen phosphate
(f) sodium hydrogen carbonate
18. Give the names of the following substances, using
IUPAC chemical nomenclature:
(a) CaCO3(s) (marble, limestone, chalk)
(b) P2O5(s) (fertilizer)
(c) MgSO4 • 7 H2O(s) (Epsom salts)
(d) N2O(g) (laughing gas, an anesthetic)
(e) Na2SiO3(s) (water glass)
(f) Ca(HCO3)2(s) (hard-water chemical)
(g) HCl(aq) (muriatic acid, gastric fluid)
(h) CuSO4 • 5 H2O(s) (copperplating, bluestone)
(i) H2SO4(aq) (acid in car battery)
(j) Ca(OH)2(s) (slaked lime)
(k) SO3(g) (a cause of acid rain)
(l) NaF(s) (toothpaste additive)
19. Give the IUPAC names of the following substances:
(d) Pb(C2H3O2)2(s)
(a) NaCl(s)
(b) P2O3(s)
(e) NH4OCl(s)
(c) HNO3(aq)
(f) Sn(BrO3)4(s)
104 Chapter 2
(g) Sb2O3(s)
(q) Ba(C2H3O2)2(s)
(h) Zn(IO3)2(s)
(r) ICl(s)
(i) Fe(NO4)2(s)
(s) AuCl3(s)
(j) Ca(OH)2(s)
(t) MgS(s)
(k) KI(s)
(u) N2F2(g)
(l) SF2(g)
(v) NiSO4(s)
(m) HBr(aq)
(w) H2S(aq)
(n) CuCO3(s)
(x) AgBrO3(s)
(o) Al2(SO3)3(s)
(y) LiClO4(s)
(p) NH4OH(l)
20. Give the names of the following substances, using
IUPAC chemical nomenclature:
(d) LiHCO3(s)
(a) CaHPO4(s)
(b) CuSO4• 7 H2O(s) (e) KHSO4(s)
(c) Na2HPO4(s)
21. Write the name and formula (with state at SATP) for
the compound formed by each of the following pairs of
elements. Where a molecular compound is formed, give
the structural formula. For ionic compounds, assume
the most common ion charges for the ions.
(a) potassium and bromine
(b) silver and iodine
(c) lead and oxygen
(d) zinc and sulfur
(e) copper and oxygen
(f) lithium and nitrogen
Applying Inquiry Skills
22. A forensic chemist was given samples of four unknown
solutions, the identity of which could affect the outcome of a court case involving an electrocution. The
chemist had reason to believe that the four substances
were KCl(aq), C2H5OH(aq), HCl(aq), and Ba(OH)2(aq).
The investigation was designed to identify the chemicals. Complete the Analysis and Synthesis sections of
the report.
Question
What is the identity of each of the four substances,
labelled 1 to 4?
Experimental Design
Each of the samples was dissolved in water and tested
with a conductivity apparatus and litmus paper. Taste
tests were not carried out.
Evidence
Table 2: Evidence for Identifying Unknown Solutions
Solution
Conductivity
Litmus
water
none
no change
1
high
no change
2
high
blue to red
3
none
no change
4
high
red to blue
24. Boron nitride (BN) is a compound that involves a large
number of covalent bonds. Use the Internet to research
the structure of boron nitride. Use your findings to
make predictions about the physical properties of
boron nitride. How do the properties of boron nitride
and diamond compare? Could boron nitride serve as a
substitute for diamond in some practical way? Provide a
theoretical argument that supports your answer.
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Analysis
(a) Answer the Question.
Synthesis
(b) Why was the water used to prepare the solutions
also tested?
(c) Which of the solutions, 1, 2, 3, or 4, could have
been involved in somebody getting electrocuted?
Use electron dot diagrams and the concept of electronegativity to explain your conclusion.
Making Connections
23. In historical dramas you may have seen someone
brought back to consciousness by having a bottle of
“smelling salts” waved under their noses. Smelling salts
were made by mixing perfume and the active ingredient, which was once called sal volatile.
(a) What is the IUPAC name and chemical formula for
sal volatile?
(b) Based on the nature of the bonds in this substance,
what properties would you predict that it would
demonstrate?
(c) Does your prediction fit with the evidence that
smelling salts have a strong, sharp, ammonia smell?
Explain.
(d) Using the Internet, find out how smelling salts used
to be administered, how they worked, and comment on the safety concerns that this use would
raise today.
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25. Glycerol is a sweet-tasting, colourless, viscous liquid
that is considered (by many national food safety organizations) to be quite safe to eat. The food industry adds
glycerol (sometimes called glycerine) to a wide range of
processed cheese and meat, baked goods, and candies.
(a) Look for glycerol on ingredients labels at home or
on the shelves of stores. Make a list of foods in
which glycerol is used.
(b) Research the physical and chemical properties of
glycerol.
(c) Research the chemical formula and molecular
shape of glycerol. Use this information, as well as
your knowledge of intra- and intermolecular
bonding, to explain at least three of the properties
of glycerol.
(d) Write a short report, outlining why glycerol is so
useful to the food industry.
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Review.
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Exploring
26. London dispersion forces are named for Fritz London,
who explained the origin of intermolecular forces
between nonpolar molecules in 1928. Research
London’s work and write a brief report on his major
findings.
Follow the links for Nelson Chemistry 11, Chapter 2
Review.
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Chemical Bonding 105