CHAPTER
2
Chemical Bonding
Specific Expectations
In this chapter, you will learn how to . . .
• B1.1 analyse on the basis of research,
the properties of a commonly used but
potentially harmful chemical substance
and how that substance affects the
environment, and propose ways to
lessen the harmfulness of the substance
or identify alternative substances that
could be used for the same purpose (2.3)
• B2.1 use appropriate terminology
related to chemical bonding (2.1)
• B2.4 draw Lewis structures to represent
the bonds in ionic and molecular
compounds (2.1)
• B2.5 predict the nature of a bond, using
electronegativity values of atoms (2.1)
• B2.6 build molecular models, and
write structural formulas, for molecular
compounds containing single and
multiple bonds and for ionic crystalline
structures (2.2)
• B2.7 write chemical formulas for binary
and polyatomic compounds, including
those with multiple valences, and
name the compounds using the IUPAC
nomenclature system (2.2)
• B3.4 explain the differences between
the formation of ionic bonds and the
formation of covalent bonds (2.1)
• B3.5 compare and contrast the physical
properties of ionic and molecular
compounds (2.3)
Salt is used in the preparation of many types of food. It is possibly the
most common food additive. The technical name for common table
salt is sodium chloride. As you learned in Chapter 1, sodium is an alkali
metal which is very reactive, and thus dangerous to handle. It reacts
vigorously with water. Chlorine gas is toxic. Exposure to even small
amounts of this gas can cause irritation to the eyes, nose, and throat.
So, why is sodium chloride safe to consume? Sodium and chlorine, in
their elemental form, are chemically very different from the compound
they form when they are bonded together as sodium chloride. In this
chapter, you will learn about the properties of ionic and molecular
compounds, including the nature and formation of chemical bonds.
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Launch Lab
Searching for Clues
In this activity, you will observe four different substances that are normally
found in the kitchen. Then you will examine your observations for clues to
help you determine the ionic or molecular nature of each compound.
Safety Precautions
• Wear safety eyewear throughout this activity.
• Wear a lab coat or apron throughout this activity.
• Do not taste any materials in a laboratory.
Materials
•
•
•
•
table salt
table sugar
baking soda
cornstarch
•
•
•
•
distilled water
watch glass
magnifying lens
5 beakers (100 mL)
•
•
•
•
marker and labels
scoopula
stirring rod
conductivity tester
Procedure
1. Place a small amount of table salt in the watch glass, and observe
it with the magnifying lens. Note whether the particles have a
characteristic shape.
2. Repeat step 1 with the other three solid substances.
3. Label one beaker “control.” Label each of the other four beakers with
the name of one of the substances. Pour 50 mL of distilled water into
each beaker.
4. Add a scoopula of table salt to the appropriate beaker. Stir with the
stirring rod, and observe what happens. Observe whether the salt
does not dissolve, dissolves slowly, or dissolves quickly.
5. Repeat step 3 with the other three substances. Make sure that you
rinse the stirring rod between substances.
6. Test each solution, including the control, for conductivity. Record
your results.
7. Based on each test—shape, solubility, and conductivity—predict
whether you think it indicates that the substance is an ionic or
molecular compound.
Questions
1. Why did you measure the conductivity of the control?
2. Compare all three predictions you made about the nature of each
individual compound. Were your three predictions the same or not?
3. Write down your final conclusion about whether each substance
is an ionic compound or a molecular compound. Compare your
conclusions with the conclusions of the other groups in your class.
4. Which property—shape, solubility, or conductivity—do you think is the
best one for predicting whether the compound is ionic or molecular?
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SECTION
2.1
Key Terms
octet rule
ionic bond
ionic compound
covalent bond
molecular compound
single bond
double bond
The Formation of Ionic and Covalent Bonds
Ninety-two naturally occurring elements combine to form the millions of different
compounds that are found in nature. Very few of these elements, however, are found in
their elemental form in nature. Some of the elements that are found in their elemental
form are the noble gases, as illustrated in Figure 2.1. What property of atoms causes
them to combine with atoms of other elements? Why are some combinations of
elements much more common than others? Answers to these questions are based on
the types of bonds that form between atoms of elements. Over the next few pages,
you will examine some naturally occurring compounds and look for patterns in these
compounds to find clues about the nature of chemical bonds.
triple bond
bonding pair
Clues in Naturally Occurring Compounds
lone pair
Scientists often study patterns in nature to better understand scientific concepts.
Chemists learn a great deal about the nature of chemical bonds by observing trends in
naturally forming compounds. For example, ores are metal compounds that are mined,
as shown in Figure 2.2, to extract the pure metals. Ores are solid and consist of a metal
combined with a non-metal, such as oxygen, sulfur, or a halogen, or with polyatomic
ions such as carbonate ions, CO32-. Very few metals are found in their elemental form
in nature. The few metals that are found in their elemental form, such as gold and
silver, are called precious metals. For a compound, such as an ore, to be in solid form,
some type of strong attractive force must be holding the individual particles together.
Lewis structure
polyatomic ion
polar covalent bond
electronegativity difference
Figure 2.1 The helium that
was used to inflate these
balloons is a noble gas.
Noble gases are some of
the very few elements that
are found in nature in their
elemental form.
Figure 2.2 Ores, consisting of metals combined with non-metals, are sometimes obtained from
open pit mines, such as the copper mine shown here.
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Clues in the Atmosphere
You can gain more insight into the nature of chemical bonds by examining the
atmosphere. It contains the oxygen that you inhale and the carbon dioxide that you
exhale. The atmosphere also contains water vapour that condenses to form clouds,
which can then become snow or rain. The major component of the atmosphere is
nitrogen. As well, there are traces of argon, methane, ozone, and hydrogen. Of these
gases, only argon, a noble gas, is found as individual atoms, not bonded to any other
atoms. The non-metal elements, oxygen, nitrogen, and hydrogen, are found in the
atmosphere as diatomic molecules. This means that they are made up of two identical
atoms bonded together. Carbon dioxide, water, and methane are examples of atoms of
non-metal elements bonded together.
The following patterns can be discerned from these observations:
• Metals usually form bonds with non-metals. The compounds they form are solid.
• Non-metals can bond with one another to form gases, liquids, or solids.
• The only elements that are never found in a combined form in nature are the
noble gases.
Stability of Atoms and the Octet Rule
Because atoms of the noble gases are always found as monatomic gases, and because
atoms of all other elements are usually found chemically bonded to other atoms, you
can infer that there is something very unique about the chemistry of noble gases, which
prevents the atoms from forming bonds. Recall, from Chapter 1, that the noble gases
are the only elements whose atoms have a filled valence shell, as shown in Figure 2.3.
This leads to the conclusion that atoms that have filled valence shells do not tend to
form chemical bonds with other atoms. Such atoms are referred to as stable.
He
Ne
Ar
Kr
Xe
Rn
Figure 2.3 Atoms of each of the noble gases except helium have eight electrons in their outer
shell, giving them filled valence shells. Because helium is in Period 1, only its first shell, which
holds a maximum of two electrons, is occupied. Thus, for helium, two electrons constitute a filled
valence shell.
The observation that a filled valence shell makes atoms stable led early chemists to
propose that when bonds form between atoms, they do so in a way that gives each
atom a filled valence shell. Because, for most main-group elements, a filled valence shell
contains eight electrons, this configuration is often called an octet. These observations
led to the octet rule for bond formation, which is stated below.
octet rule a “rule of
thumb” that allows you
to predict the way in
which bonds will form
between atoms
The Octet Rule
When bonds form between atoms, the atoms gain, lose, or share electrons in such
a way that they create a filled outer shell containing eight electrons.
As you read in Chapter 1, atoms of the transition elements and inner transition
elements can have complex electron configurations. They can have more than eight
electrons in their valence shells and, therefore, they do not follow the octet rule.
Because main-group elements are much more common on Earth, however, a very large
number of compounds that you study will follow the octet rule. Thus, the octet rule
provides an important basis on which to predict how bonds will form.
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The Formation of Ionic Bonds
ionic bond the
attractive electrostatic
force between a
negative ion and a
positive ion
ionic compound a
chemical compound
composed of ions that
are held together by
ionic bonds
An ionic bond is the attractive electrostatic force between oppositely charged ions.
Thus, before an ionic bond can form, atoms must be ionized. According to the
octet rule, atoms gain or lose electrons to attain a filled valence shell. In Chapter 1,
Section 1.3, you learned that an atom of an element with fewer than four electrons in
its valence shell, especially an alkali metal atom, can lose electrons relatively easily.
You also learned that an atom with more than four electrons in its valence shell can
gain electrons and form a stable ion. Thus, in general, a metal loses all of its valence
electrons and becomes an ion with an octet of electrons in its outer shell. A non-metal
gains enough electrons to fill its valence shell. These oppositely charged ions exert
attractive electrostatic forces on each other, resulting in the formation of an ionic bond.
A compound that is held together by ionic bonds is called an ionic compound.
Because ionic compounds must have an overall charge of zero, the number of
electrons that are lost by the metal atoms must be equal to the number of electrons
gained by the non-metal atoms. Two such examples are shown in Figure 2.4, in the
form of Lewis diagrams. Notice that the electrons of the metals are depicted as open
circles and the electrons of the non-metals are depicted as dots, so you can follow them
throughout the process.
+
Na
Cl
O
Na
-
Cl
2+
Mg
Mg
2-
O
Figure 2.4 When metal atoms, such as sodium and magnesium, lose electrons, they have no
valence electrons remaining. Therefore, there are no dots around the symbols for the metal ions.
In each example in Figure 2.4, the number of electrons gained by the non-metal atom is
exactly the same as the number of electrons lost by the metal atom. It is also possible for
the number of electrons gained by a non-metal atom to be different from the number
of electrons lost by a metal atom. However, the total number of electrons gained by
non-metal atoms must be the same as the total number of electrons lost by metal atoms.
Examples of three such situations are shown in Figure 2.5.
-
Figure 2.5 In each
example, you can see
that the total number of
positive charges (electrons
lost) on the metal ions is
equal to the total number
of negative charges
(electrons gained) on the
non-metal ions.
F
2+
F
Ca
Ca
K
S
-
F
F
K
K
K
+
+
2-
S
2+
Mg
Mg
N
2+
N
Mg
Mg
N
Mg
3-
32+
N
Mg
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Ionic Compounds Containing Transition Metals
All of the examples in Figures 2.4 and 2.5 include only main-group elements. You can
determine the number of valence electrons of an atom of a main-group element by its
group number. Occasionally, however, you will be working with transition metals. In
Chapter 1, you read that the electron configuration of transition metals is quite complex.
Therefore, it is not possible to predict the number of electrons that a transition metal
atom can lose from its group number. In fact, the number of electrons that a transition
metal can lose can vary. For example, an iron atom can lose either two electrons or three
electrons. You can find the number of electrons that atoms of a transition element can
lose by checking the periodic table. Figure 2.6 shows you how to find the possible charges
on the resulting ions after the metal atoms have become ionized. Notice that the possible
charges on the ions are highlighted. The figure shows a few common transition metals
that can form more than one possible ion. As stated above, iron atoms can lose two or
three electrons. Thus iron atoms can form ions with charges of 2+ or 3+.
25
1.6
54.94
2+, 4+
Mn
manganese
26
1.8
27
55.85
3+, 2+
58.93
2+, 3+
1.9
Fe
Co
iron
29
1.9
63.55
2+, 1+
Cu
cobalt
79
2.4
196.97
3+, 1+
Au
1.9
200.59
2+, 1+
Hg
gold
copper
80
The n
the p
you o
mercury
Figure 2.6 These cells are taken directly from the periodic table on page 24. The common ion
charges are highlighted.
When you are working with transition metals, you will be given the charge or enough
information to determine the charge on the ions. For example, you might be told that
two iron atoms have combined with three oxygen atoms and asked to draw a Lewis
diagram of the compound. If you do not know that the oxygen ion has a charge of 2-,
you can find it in the periodic table. Since there are three oxygen ions, the total negative
charge in the compound will be 6-. Thus, the total positive charge on the two iron
ions must be 6+. Therefore, there must be a charge of 3+ on each iron ion. The Lewis
diagram for this compound is shown in Figure 2.7 (A).
You might also be asked to draw the Lewis diagram of a compound that contains
one iron ion and two chloride ions. Since a chloride ion has a charge of 1-, the single
iron ion must have a charge of 2+. The Lewis diagram of this compound is shown in
Figure 2.7 (B). It is important to remember that the electron configuration for iron
atoms is complex. Iron atoms do not actually have two valence electrons or three
valence electrons as shown in Figure 2.7. The iron atoms are drawn as though they have
either two or three valence electrons only because these are the numbers of electrons
that they can lose when they become ionized.
A
O
Fe
Fe
O Fe
Fe
O
3+
3+
O
O
O
2-
B
-
2-
2-
Cl
Fe
Cl
Fe
2+
Cl
-
Cl
Figure 2.7 (A) Each of the two iron atoms loses three electrons to oxygen atoms. Thus, the
resulting ions have a charge of 3+. (B) An iron atom loses one electron to one chlorine atom
and a second electron to another chlorine atom. The resulting iron ion has a charge of 2+.
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The Formation of Covalent Bonds
covalent bond the
attraction between
atoms that results from
the sharing of electrons
molecular compound
a chemical compound
that is held together by
covalent bonds
The octet rule states that atoms can also acquire a filled outer shell by sharing electrons.
When the nuclei of two atoms are both attracted to one or more pairs of shared
electrons, the attraction is called a covalent bond. A compound that is held together
by covalent bonds is called a molecular compound. Molecular compounds consist of
non-metal elements only. Examples of molecular compounds are water and carbon
dioxide. Molecular compounds can be solid, liquid, or gas at room temperature.
Only unpaired electrons are likely to participate in chemical bonds. Figure 2.8 shows
how covalent bonds form when (A) hydrogen atoms share their only electrons and when
(B) chlorine atoms share their only unpaired electrons.
A
B
H
H
Cl
Cl
H
H
Cl Cl
H
H
Cl Cl
Figure 2.8 In both (A) and (B), electrons of one atom are shown as open circles and electrons
of the other atom are shown as dots, to help you follow the electrons. In the third row, circles
surrounding each atom in the molecule show the filled shell of electrons for each atom.
Multiple Bonds
In some molecules, there are not enough valence electrons for two atoms to share one
pair of electrons and form filled valence shells. For example, in carbon dioxide, the
carbon atom has four valence electrons and each of the two oxygen atoms has six valence
electrons, as shown in Figure 2.9 (A). If the carbon atom shared one pair of electrons with
each oxygen atom, each of the oxygen atoms would have only seven electrons and the
carbon atom would have only six electrons in the valence shell, as shown in Figure 2.9 (B).
This configuration would not provide all atoms with filled outer shells. Instead, to
complete an octet for each atom, the unpaired electrons on all atoms are rearranged, as
shown in Figure 2.9 (C), to become shared. The atoms now share four electrons as shown
in Figure 2.9 (D). It is important to remember that Lewis diagrams are just that—diagrams.
It would be more correct to show electron clouds overlapping and forming new electron
clouds with different shapes. However, it is more difficult to visualize the number of
electrons in valence shells when using the electron cloud model.
A
B
O
C
O
O
C
O
not found to occur
C
D
O
C
O
O
C
O
Figure 2.9 (A) Each oxygen atom has six valence electrons and the carbon atom has four. (B) If
each oxygen atom shared two electrons with the carbon atom, neither would have a filled outer
shell. (C) Instead, the remaining unpaired electrons in both oxygen atoms and the carbon atom are
rearranged so that they can also be shared by the atoms. (D) When each oxygen atom shares four
electrons (two pair) with the carbon atom, all of the atoms acquire an octet of electrons.
Explain how you could predict the number of bonds that an atom could form with other atoms.
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Types of Covalent Bonds and Electron Pairs
While one pair of shared electrons constitutes a single bond, two pairs of shared
electrons make up a double bond. Compounds can also have triple bonds, which
consist of three pairs of shared electrons. Nitrogen, the gas that makes up most of the
atmosphere, is an example of a molecule that has a triple bond, as shown in Figure 2.10.
N
N
N
N
double bond a
covalent bond that
results from atoms
sharing two pairs of
electrons
Figure 2.10 Two nitrogen atoms must share three pairs of electrons to complete an octet of
electrons around each nitrogen atom.
Although there are rarely any unpaired electrons in molecular compounds, some
electron pairs are shared while others are not. A pair of shared electrons is called a
bonding pair. A pair of electrons that is not involved in a covalent bond is called a
lone pair. These types of electron pairs are labelled in the water molecule in Figure 2.11.
When a Lewis diagram is used to portray a complete molecular compound, as done in
this figure, the diagram is called a Lewis structure.
triple bond a covalent
bond that results from
atoms sharing three
pairs of electrons
bonding pair a pair of
electrons that is shared
by two atoms, thus
forming a covalent bond
lone pair a pair of
electrons that is not part
of a covalent bond
Lewis structure a Lewis
diagram that portrays
a complete molecular
compound
H
bonding
pairs
O
single bond a covalent
bond that results from
atoms sharing one pair
of electrons
H
lone
pairs
Figure 2.11 The oxygen atom in water has two bonding pairs and two lone pairs. The hydrogen
atoms each have one bonding pair.
Drawing Lewis Structures
Although there are no specific steps that you can always follow to draw a Lewis
structure, there are some guidelines that will help you. First, draw a Lewis diagram for
each atom in the structure. Start with the atom that has the most unpaired electrons.
Determine the number of bonds that each atom can form with other atoms. That
number is the same as the number of unpaired electrons. Finally, try to fit the atoms
together in a way that will create a filled outer shell for each atom. The example used
in Figure 2.12 (A) has one carbon atom, two oxygen atoms, and two hydrogen atoms.
Begin with the carbon atom. If both oxygen atoms are bonded to the carbon atom with
double bonds, there will be no way to add the hydrogen atoms. If both hydrogen atoms
are bonded to the carbon atom, there will be bonds for only one oxygen atom. The final
result is shown in Figure 2.12 (B).
A
H
B
C
O
O
H
H
C
O
O
H
Figure 2.12 If you take the atoms in (A) and test different ways of connecting them, you will find
that the Lewis structure in (B) creates filled outer shells for all of the atoms. This compound is
commonly called formic acid, HCOOH.
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Polyatomic Ions and Bond Formation
polyatomic ion a
molecular compound
that has an excess or a
deficit of electrons, and
thus has a charge
When you first look at the structure in Figure 2.13 (A), (ignoring the colour), it appears
to be a typical Lewis structure. However, when you count the electrons, you will find a
new feature in this compound. Count the number of electrons that are the same colour
as the symbol, and you will find that they represent the number of valence electrons
that an atom of that element has. The carbon atom has four black electrons, and each
oxygen atom has six red electrons. The colour-coded electrons account for all of the
valence electrons that are available. If there were no additional electrons, the atoms
would not all have filled valence shells. To fill the shells, two electrons were added,
as shown in green in Figure 2.13 (A). These two electrons give the compound a negative
charge of 2-. Nevertheless, it is a valid Lewis structure. Some molecular compounds,
like non-metal atoms, can gain electrons to complete octets on all of their atoms. Such
compounds are called polyatomic ions because they consist of two or more atoms.
The correct diagram for polyatomic ions includes brackets and a number and sign,
as shown in Figure 2.13 (B).
A
B
2-
O
O
C
C
O
O
O
O
Figure 2.13 (A) If you count the number of electrons, you will get 24. Because this is two more
electrons than the sum of the valence electrons in three oxygen atoms and one carbon atom, the
compound has a charge of 2-. (B) To show that this compound is a polyatomic ion, it is bracketed
and a 2- is placed outside the brackets.
Typically, electrons in Lewis structures are not colour coded so you cannot easily see
whether there are any extra electrons. Nevertheless, you can quickly determine whether
a Lewis structure represents a neutral molecular compound or a polyatomic ion by first
counting the electrons and comparing that number with the total number of valence
electrons that each atom would have. For example, the structure in Figure 2.13 (A) has
24 electrons. Add up the number of valence electrons by reasoning that each oxygen
atom has six valence electrons and the carbon atom has four valence electrons, giving
a total of 22 electrons. You immediately know that you must have two extra electrons,
giving you a negatively charged polyatomic ion.
A negatively charged polyatomic ion can bond to a positively charged ion to form
an ionic compound in the same way that a metal ion and a non-metal ion can bond
to form an ionic compound. Figure 2.14 shows two examples of ionic compounds that
contain the carbonate ion.
2-
2-
O
O
2+
C
Ca
O
O
K
K
+
C
+
O
O
Figure 2.14 Polyatomic ions, like simple ions, must combine with oppositely charged ions that
will give the final compound a neutral charge.
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Positively Charged Polyatomic Ions
Positively charged polyatomic ions also exist, but the only one that is common is the
ammonium ion. The ammonium ion forms when the molecular compound ammonia
combines with a hydrogen ion, as shown in Figure 2.15. Ammonia has three bonded
pairs and one lone pair. The hydrogen ion bonds with the lone pair to form an
ammonium ion.
H
+
+
lone
pair
N
+
H
H
H H
N
H
H
H
ammonia
ammonium ion
Figure 2.15 The ammonium ion is the only common positively charged polyatomic ion. Try not
to confuse it with ammonia, which is a neutral molecular compound.
The ammonium ion forms ionic compounds by bonding with negatively charged ions.
It can bond with a simple negatively charged ion, such as the chloride ion, as shown in
Figure 2.16 (A). It can also bond with a negatively charged polyatomic ion, such as the
carbonate ion shown in Figure 2.16 (B).
A
B
H
+
H
N
H
N
H
2-
O
H
-
H
+
H
Cl
H
H
H
N
C
+
O
O
H
H
Figure 2.16 (A) The ammonium ion behaves like any other positively charged ion and forms
ionic compounds, such as ammonium chloride, by bonding with negatively charged simple ions.
(B) The ammonium ion can also form ionic compounds, such as ammonium carbonate, by
bonding with negatively charged polyatomic ions.
Learning Check
1. State the octet rule, and give one example of how it
can be applied.
2. When a calcium atom becomes ionized, it has
a charge of 2+. When a bromine atom becomes
ionized, it has a charge of 1–. Explain how ionic
bonds can form between calcium and bromine to
produce a compound that has a zero net charge.
3. Given a Lewis structure with four non-metal atoms,
how would you determine whether it is a molecular
compound with no charge or a polyatomic ion?
4. Draw a Lewis structure of two oxygen atoms that
are covalently bonded together to form an oxygen
molecule. Identify the bonding pairs and the
lone pairs.
5. How do double bonds and triple bonds form? Why
do they form?
6. Describe a situation in which two atoms that are
covalently bonded together can be part of an ionic
compound.
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The Importance of Electronegativity in Bond Formation
Based on what you just learned about ionic bonds and covalent bonds, you might
assume that they are two separate and distinct types of connections between atoms.
However, like ionic bonds, covalent bonds also involve electrostatic attractions
between positively charged nuclei and negatively charged electrons. To understand
how electrostatic attraction influences the nature of bonds, recall the concept of
electronegativity, which you learned about in Chapter 1.
Electronegativity is an indicator of the relative ability of an atom of a given element
to attract shared electrons. Shared electrons constitute a covalent bond. Thus, the
relative electronegativities of the elements of the two atoms that are bonded together
should provide information about the nature of the bond. Although Lewis diagrams are
drawn as though no electrons are shared between two nuclei in ionic compounds, the
positively charged nucleus of each ion is attracting the negatively charged electrons of
the other ions. Thus, the concept of electronegativity also applies to ionic compounds.
Electronegativity Difference and Bond Type
What do the relative electronegativities of elements tell you about the nature of bonds?
If the electronegativity of one of the two atoms that are bonded together is greater than
the electronegativity of the other atom, the electrons will be attracted more strongly to
the first atom. In general, electrons spend more time around the atoms with the greater
electronegativity.
Figure 2.17 illustrates a bond between a carbon atom and a chlorine atom. Of course,
the carbon atom is bonded to other atoms as well as the chlorine atom. The electronegativity
of the chlorine atom (3.2) is higher than the electronegativity of the carbon atom
(2.6). The arrow indicates that the shared electrons are more strongly attracted to the
chlorine atom, and thus spend more time there. The Greek letter delta, δ, is often used
to represent “partial.” Therefore, the symbols δ+ and δ− indicate that the carbon atom
is partially positively charged and the chorine atom is partially negatively charged.
2.6
δ+
3.2
C
Cl
δ-
Figure 2.17 Because the shared electrons in this bond spend more time near the chlorine
nucleus, the chlorine atom is slightly negatively charged. This leaves the carbon atom slightly
positively charged.
polar covalent bond
a covalent bond around
which there is an
uneven distribution of
electrons, making one
end slightly positively
charged and the other
end slightly negatively
charged
electronegativity
difference the
difference between the
electronegativities of
two atoms
Describe How do you know that the electrons will spend more time near the chlorine atom than
the carbon atom? Describe the data that tell you this.
Covalent bonds, in which the electron distribution is unequal, are called polar covalent
bonds. These bonds are often referred to simply as polar bonds. Because these bonds
have a positive “pole” and a negative “pole,” they are sometimes also called bond dipoles.
Depending on the difference in the electronegativities of the bonded atoms, some
covalent bonds are only slightly polar while others are extremely polar. Chemists have
devised a system for classifying the extent of the polarity of the bonds by calculating
the electronegativity difference (ΔEN) for the two elements involved in the bond.
You can calculate the electronegativity difference for any two elements by finding the
electronegativity of each element in a table, such as the one in Figure 1.22 on page 36,
and then subtracting the smaller electronegativity from the larger electronegativity.
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ΔEN = ENF - ENK
ΔEN = 4.0 - 0.8
3.3
+
-
3.0
2.5
mostly
ionic
2.0
ΔEN
Applying Electronegativity Difference
As shown in Figure 2.18, bonds in which the electronegativity
difference of the atoms is greater than 1.7 are classified as mostly ionic.
The term “mostly” is used because there is always some attraction
between the nucleus of one atom and the electrons of the other atom
involved in the bond. If the electronegativity difference of two atoms
that are bonded together is between 0.4 and 1.7, the bond is classified
as polar covalent. If the electronegativity difference is less than 0.4,
the bond is classified as slightly polar covalent. It is only when the
electronegativity difference is zero that the bond can be classified as
a non-polar covalent bond. The images on the right of Figure 2.18 are
electron cloud models of atoms bonded together. The image at the top
shows a positive ion and a negative ion beside each other, indicating
that the bond is mostly ionic. The next image shows atoms joined
by a polar covalent bond. Chemists often use an arrow, like the one
above this image, to show a polar bond. The tail of the arrow above
this image looks like a plus sign, to signify the slightly positively
charged end of the bond. The arrow points in the direction in which
the electrons spend more time. The bottom image shows two atoms
equally sharing electrons in a non-polar covalent bond.
The following examples show you how to use the information
in Figure 2.18 to calculate the electronegativity difference for two
atoms. The first example involves the bond between a potassium atom
and a fluorine atom. The electronegativity of fluorine is 4.0, and the
electronegativity of potassium is 0.8. The electronegativity difference is
calculated by subtracting the smaller number from the larger number,
as shown below. Because 3.2 is much larger than 1.7, the bond between
potassium and fluorine is mostly ionic.
1.7
1.5
+→
polar
covalent
1.0
0.5
0.4
slightly polar
covalent
non-polar covalent
0.0
Figure 2.18 The shading in the diagram
indicates that the character of bonds
changes gradually from mostly ionic at the
top to non-polar covalent at the bottom.
The electronegativity difference values
on the right are the transition points that
separate the types of bonds. The images on
the far right are models of compounds with
the bond character in the different ranges
of electronegativity difference.
ΔEN = 3.2
Next, consider the bond between two oxygen atoms in an oxygen
molecule. The electronegativity of oxygen is 3.4. The electronegativity
difference, as shown below, is zero. Therefore, the bond between two
oxygen atoms is non-polar covalent.
ΔEN = ENO - ENO
ΔEN = 3.4 - 3.4
ΔEN = 0.0
Finally, consider the bond between a carbon atom and a chlorine atom, discussed on
the previous page. The electronegativity of carbon is 2.6, and the electronegativity of
chlorine is 3.2. The electronegativity difference is 0.6, as shown below. This value is
between 1.7 and 0.4, indicating that the bond is a polar covalent bond.
ΔEN = E NCl - ENC
ΔEN = 3.2 - 2.6
ΔEN = 0.6
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Percent Ionic and Covalent Character
Chemists have devised another approach for describing the bond character, using
percentages of either ionic or covalent character. Table 2.1 relates electronegativity
differences to percent ionic character and percent covalent character. In the following
activity, you will analyze the relationship between electronegativity differences and
percent ionic character.
Table 2.1 Character of Bonds
Electronegativity Difference
0.00
0.65
0.94
1.19
1.43
1.67
1.91
2.19
2.54
3.03
0
10
20
30
40
50
60
70
80
90
100
90
80
70
60
50
40
30
20
10
Percent Ionic Character
Percent Covalent Character
Figure 2.19 The
electronegativity
difference for hydrogen
and chlorine indicates
that a bond between
these atoms results in a
polar covalent molecule
when HCl is in a gaseous
state (A). Its interaction
with water molecules
causes HCl to behave as
an ionic compound (B).
Activity
2.1
Classifying bond type is not always simple. The bond between a hydrogen atom and
a chlorine atom provides a good example of overlap in ionic character and covalent
character. The electronegativity difference for hydrogen and chlorine is 1.0, placing it
in the polar covalent category. As a gas, the compound behaves as a polar molecule.
When the compound is dissolved in water, however, the atoms become separate
ions, both surrounded by water molecules. Thus, the bond type of this compound
varies, depending on whether it is a gas or dissolved in water. Figure 2.19 shows
Lewis diagrams for the two states.
A
-
δ
δ
H
Cl
or
+
HCl
H
-
Cl
Electronegativity Difference versus Percent Ionic Character
Materials
• ruler
HCI dissolved in water
+→
Why did chemists choose the electronegativity differences
of 0.4 and 1.7 for the transition points for slightly polar
covalent, polar covalent, and mostly ionic bonds? Analyzing
the relationship between electronegativity difference
and percent ionic character in this activity will help you
understand the reasons behind the choice of these values.
• graph paper
B
HCI in gaseous state
+
• pencil
Procedure
1. Construct a graph using the data in the first two rows of
Table 2.1. Put electronegativity difference on the x-axis
and percent ionic character on the y-axis. Choose scales
for the axes that will make the graph take up more than
half of a sheet of graph paper.
2. After you plot all the points, draw a smooth curved line
of best fit through the points.
3. Draw a straight, vertical line on the graph through the
point where the electronegativity difference is 1.7. At
the point at which the vertical line crosses the curve,
draw a horizontal line across the graph. Record the value
of the percent ionic character at the point where your
horizontal line touches the axis.
4. Repeat step 3 for the point where the electronegativity
difference is 0.4.
Questions
1. What is the percent ionic character when the
electronegativity difference is 1.7? Do you think this is a
reasonable value for the transition point between polar
covalent and mostly ionic bonds? Explain your reasoning.
2. What is the percent ionic character when the
electronegativity difference is 0.4? Do you think this is a
reasonable value for the transition point between polar
covalent and slightly polar covalent bonds? Explain your
reasoning.
3. Why do you think it was important to make your graph
spread out to more than half of the sheet of graph paper?
4. Imagine that you were to draw a graph of percent
covalent character versus electronegativity difference.
Predict the values of percent covalent character that you
would find when the electronegativity differences are
0.4 and 1.7. Explain why you think you would find these
results.
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Section 2.1
RE V I E W
Section Summary
• The octet rule can be used to predict how bonds will form.
• An ionic bond forms when a negatively charged ion and
a positively charged ion are attracted to each other.
• A covalent bond forms when two atoms share one or
more pairs of electrons.
• A polyatomic ion consists of two or more atoms that
are covalently bonded together and carry a charge. A
polyatomic ion can form an ionic compound with a simple
ion or another polyatomic ion of the opposite charge.
• A chemical bond can be non-polar covalent, slightly
polar covalent, polar covalent, or mostly ionic, depending
on the electronegativity difference between the two atoms
that are bonded together.
Review Questions
1.
K/U What property of the noble gases led to the octet
rule? Explain.
2.
K/U Explain why metal atoms tend to lose electrons
to form ions and why non-metal atoms tend to gain
electrons to form ions.
3.
C
Draw Lewis diagrams of calcium and bromine.
Use these diagrams to show how ionic bonds form
between these atoms. Explain how these structures
satisfy the octet rule.
4.
5.
6.
For each of the following, use Lewis diagrams
to predict the number of atoms of each element that
will be present in an ionic compound formed by the
two elements.
a. calcium and fluorine
c. magnesium and nitrogen
b. sodium and oxygen
T/I Predict whether the bond between each pair of
atoms will be non-polar covalent, slightly polar
covalent, polar covalent, or mostly ionic.
a. carbon and fluorine
e. silicon and hydrogen
b. oxygen and nitrogen
f. sodium and fluorine
c. chlorine and chlorine g. iron and oxygen
d. copper and oxygen
h. manganese and oxygen
11.
For each polar and slightly polar covalent bond
in question 10, indicate the locations of the partial
positive and partial negative charges. Explain how you
made each decision.
12.
T/I Arrange the bonds in each group below in order
of increasing polarity.
a. hydrogen bonded to chlorine, oxygen bonded to
nitrogen, carbon bonded to sulfur, sodium bonded
to chlorine
b. carbon bonded to chlorine, magnesium bonded to
chlorine, phosphorus bonded to oxygen, nitrogen
bonded to nitrogen
13.
Make a sketch that shows the relationship
between electronegativity difference and percent ionic
character of a chemical bond. Why do you think that
the transition points between types of chemical bonds
are reported in electronegativity difference rather than
percent ionic character?
14.
K/U Explain the meaning of the
symbol above these chemical symbols.
T/I
T/I Draw Lewis diagrams of two oxygen atoms.
Use your diagrams to show how an oxygen molecule
forms from two oxygen atoms. Explain why there must
be a double bond between the two oxygen atoms.
K/U
How many electrons make up a triple bond?
7.
K/U Draw a Lewis structure of a hydrogen atom
covalently bonded to a fluorine atom. Identify all the
bonding pairs and all the lone pairs.
8.
Assume that you are shown a Lewis structure
with one nitrogen atom and three oxygen atoms. How
would you determine whether the structure represented
a neutral molecule or a polyatomic ion?
9.
10.
K/U
Explain why the following compound can be
considered an ionic compound, even though it does
not contain any metal ions.
T/I
+
H
H
N
H
H
C
+→
NO
15.
A
Toward the beginning of this section, you read
that metals are usually found in combination with
non-metals in nature, and that these compounds are
solid. From what you now know, how would you
classify these compounds? Give an example.
16.
A
The atmosphere consists mostly of nitrogen and
oxygen, along with small amounts of carbon dioxide
and trace amounts of hydrogen. Does the atmosphere
consist almost entirely of polar compounds or
non-polar compounds? Explain your reasoning.
-
I
T/I
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SECTION
2.2
Writing Names and Formulas for Ionic
and Molecular Compounds
Key Terms
“Please pass the sodium chloride.”
alkali
“Do you have enough sucrose for your tea?”
oxoacid
“May I please have some more dihydrogen monoxide?”
structural formula
“These biscuits are so hard and flat! I must have forgotten to put the sodium hydrogen
carbonate in the dough.”
You might have heard statements like these while having a meal with family or friends,
as in Figure 2.20, but the terminology was probably quite different. Four of the terms
in these statements are chemical names for common substances. Do you know what
the substances are? In this section, you will learn how to name and write the formulas
for ionic and molecular compounds. As you read this section, try to figure out the
common names for the chemicals identified in the earlier statements.
Figure 2.20 Would other friends or family members know what you meant if you asked,
“May I please have some more dihydrogen monoxide?” during a meal?
Standardized Naming
Imagine sitting at a table with six people, all of whom speak a different language. If you
said, “Please pass the potatoes,” no one would know what you wanted. Someone else
might say “Kartoffel?” but you would probably not know what he or she meant.
Chemistry is a language that has millions of “words.” Chemists had begun to
recognize the need to “speak the same language” as early as the late 1700s. Several
chemistry organizations began to develop rules for naming compounds. The current
standards are set by the International Union for Pure and Applied Chemistry (IUPAC).
The organization was founded in 1919 and still holds meetings to maintain and
improve on the rules that allow chemists throughout the world to communicate clearly
and concisely.
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Naming Binary Ionic Compounds
Binary compounds are among the simplest compounds to name. A binary ionic
compound is an ionic compound that consists of atoms of only two (bi-) different
elements. Because ionic compounds nearly always consist of metals and non-metals,
one of these two elements must be a metal and the other must be a non-metal. Study
the steps below to review the rules for naming binary ionic compounds.
Rules for Naming Binary Ionic Compounds
1. The name of the metal ion is first, followed by the name of the non-metal ion.
2. The name of the metal ion is the same as the name of the metal atom.
3. If the metal is a transition metal, it might have more than one possible charge.
In these cases, a roman numeral is written in brackets after the name of the metal
to indicate the magnitude of the charge.
4. The name of the non-metal ion has the same root as the name of the atom,
but the suffix is changed to -ide.
Table 2.2 Names of Some
Common Non-metal Ions
Formula
for Ion
The names of several common non-metal ions are listed in Table 2.2.
As you learned in Section 2.1, when forming an ionic compound, the positive and
negative ions must combine in numbers that result in a zero net charge. There is no
need to indicate these numbers in the name, however, because they are determined
by the charges on the ions. The examples in Table 2.3 will help you review the rules for
naming binary ionic compounds, starting with Lewis diagrams.
Name
of Ion
F−
fluoride
Cl−
chloride
Br−
bromide
I−
iodide
O2−
oxide
S2−
sulfide
N3−
nitride
Table 2.3 Examples of Naming Binary Ionic Compounds
2-
Compound
+
Na
-
Cl
2+
F
Ca
3+
O
Fe
-
F
23+
O
Fe
2-
O
Steps
1. Name the metal
ion first.
The metal is sodium.
The metal is calcium.
The metal is iron.
2. The name of the ion
is the same as the
name of the metal.
The name of the sodium ion is
sodium.
The name of the calcium ion is
calcium.
The name of the iron ion is iron.
3. If the metal ion can
have more than one
charge, indicate the
charge with a roman
numeral in brackets.
Sodium ions always have a charge
of 1+ so no roman numeral
is needed.
Calcium ions always have a charge Iron ions can have a charge of 2+
of 2+ so no roman numeral is
or 3+. The Lewis diagram shows
needed.
that the iron ions have a charge
of 3+, so the roman numeral III
must be added. The name of the
metal ion becomes iron(III).
4. Name the non-metal
ion second. Use the
root name of the
atom with the suffix
-ide.
The non-metal is chlorine.
Change chlorine to chloride.
Add the name chloride to sodium.
The name of the compound is
sodium chloride.
The non-metal is fluorine.
Change fluorine to fluoride.
Add the name fluoride to calcium.
The name of the compound is
calcium fluoride.
The non-metal is oxygen.
Change oxygen to oxide.
Add the name oxide to iron(III).
The name of the compound is
iron(III) oxide.
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Naming Ionic Compounds with Polyatomic Ions
The rules for naming ionic compounds that have polyatomic ions are fundamentally
the same as the rules for naming binary ionic compounds. Each polyatomic ion that
you encounter has its own name and is treated as a single unit in a compound. The
names and structures of all of the polyatomic ions that you are likely to encounter are
listed in Table 2.4. Recall that the only common positively charged polyatomic ion is
the ammonium ion, NH4+, which you saw in Figure 2.15.
Table 2.4 Some Common Polyatomic Ions
Name
Formula
Name
Formula
ammonium
NH4+
nitrate
NO3−
acetate or ethanoate
CH3COO−
nitrite
NO2−
benzoate
C6H5COO-
oxalate
OOCCOO2−
borate
BO3
hydrogen oxalate
HOOCCOO−
carbonate
CO32−
permanganate
MnO4−
hydrogen carbonate
HCO3−
phosphate
PO43−
3−
−
perchlorate
ClO4
hydrogen phosphate
HPO42−
chlorate
ClO3−
dihydrogen phosphate
H2PO4−
chlorite
ClO2−
sulfate
SO42−
hypochlorite
ClO−
hydrogen sulfate
HSO4−
chromate
CrO42−
sulfite
SO32−
dichromate
Cr2O72−
hydrogen sulfite
HSO3−
cyanide
CN−
cyanate
CNO−
hydroxide
OH−
thiocyanate
SCN−
iodate
IO3−
thiosulfate
S2O32−
There are no comprehensive rules for naming polyatomic ions, so it is best just to learn
the names. There are some generalizations, however, that will help you remember some
of the names. If you read through the names and structures in Table 2.4, you will notice
that several groups, or families, of polyatomic ions have names with similar roots and
have compositions that vary only in the number of oxygen atoms. Table 2.5 lists the
prefixes and suffixes and shows how they are assigned to each family of ions.
Table 2.5 Prefixes and Suffixes for Families of Polyatomic Ions
Relative Number of Oxygen Atoms
Prefix
Suffix
Example
Family of Four
most
per-
-ate
ClO4−
perchlorate
second most
(none)
-ate
ClO3−
chlorate
-ite
ClO2
−
chlorite
-ite
ClO−
hypochlorite
-ate
NO3−
nitrate
-ite
−
nitrite
second fewest
fewest
(none)
hypo-
Family of Two
most
fewest
(none)
(none)
NO2
It is important to notice that the suffix, -ate or -ite, does not specify a certain number
of oxygen atoms. Instead it indicates the relative number of oxygen atoms. For example,
nitrate has three oxygen atoms and nitrite has two oxygen atoms, whereas sulfate has
four oxygen atoms and sulfite has three oxygen atoms.
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In Table 2.4, you will also notice that some polyatomic ions with a charge of 2− or
1− have hydrogen or dihydrogen at the beginning of their name. This term describes
the number of hydrogen ions added to the original polyatomic ion. For example,
the phosphate ion, PO43−, has no hydrogen ions. Hydrogen phosphate, HPO42−, has
one hydrogen ion and one less negative charge than the phosphate ion. Dihydrogen
phosphate, H2PO4−, has two hydrogen ions and two fewer negative charges than the
phosphate ion.
As well, you will notice the prefix thio- in front of two of the polyatomic ions. This
prefix indicates that a sulfur atom has taken the place of an oxygen atom. For example,
the sulfate ion, SO42−, has one sulfur atom and four oxygen atoms. The thiosulfate ion,
S2O32−, has two sulfur atoms and three oxygen atoms.
Writing Chemical Formulas for Ionic Compounds
A chemical name provides all the information that you need to write the chemical
formula for a compound. The following steps summarize the rules for writing the
chemical formulas for ionic compounds.
Rules for Writing Chemical Formulas for Ionic Compounds
1. Identify the positive ion and the negative ion.
2. Find the chemical symbols for the ions, either in the periodic table or in the table of
polyatomic ions. Write the symbol for the positive ion first and the symbol for the
negative ion second.
3. Determine the charges of the ions. If you do not know the charges, you can find them
in the periodic table on page 24.
4. Check to see if the charges differ. If the magnitudes of the charges are the same, the
formula is complete. If they differ, determine the number of each ion that is needed
to create a zero net charge. Write the numbers of ions needed as subscripts beside
the chemical symbols, with one exception. When only one ion is needed, leave the
subscript blank. A blank means one. If a polyatomic ion needs a subscript, the formula
for the ion must be in brackets and the subscript must be outside the brackets.
When the charges of the ions are not the same, you have to determine the number
of each ion that is needed to create a zero net charge. To do this, you could simply “guess
and check.” However, the cross-over method shown in Figure 2.21 is a more direct way
to determine the number of each ion that is needed. As shown in Figure 2.21, use the
magnitude of the charge of each ion as the subscript for the opposite ion. Below each
diagram is a calculation that demonstrates why the subscripts always give you the
numbers of ions that result in a zero net charge for the compound. Table 2.6, on the next
page, shows examples of applying the rules for writing formulas for ionic compounds.
3+
Al
2-
SO4
2+
Ca
3-
N
3+
Co
Cl
1-
Al2(SO4)3
Ca3N2
Co3Cl3
2(+3) + 3(-2) = 6 - 6 = 0
3(+2) + 2(-3) = 6 - 6 = 0
1(+3) + 3(-1) = 3 - 3 = 0
Figure 2.21 When you make
the number of ions of each
element (the subscript) equal
in magnitude to the charge
of the opposite ion, you will
create a compound with a
zero net charge.
Notice, in the first example in Figure 2.21, that you use only the net charge and ignore
the subscript on the polyatomic ion. Also note that the new subscript that indicates the
number of polyatomic ions in the compound goes outside the brackets.
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Examples of Using Rules for Formulas for Ionic Compounds
The following examples will show you how the rules are applied to writing formulas
for ionic compounds.
Table 2.6 Examples of Writing Formulas for Ionic Compounds
Name
Steps
aluminum chloride
calcium iodide
potassium permanganate
1. Identify the positive ion and
the negative ion.
Aluminum is first, so it is
the positive ion. Chloride is
second and has the suffix -ide,
so it is the negative ion.
Calcium is first, so it is the
positive ion. Iodide is second
and has the suffix -ide, so it is
the negative ion.
Potassium is first, so it is the
positive ion. Permanganate is
second, so it is the negative ion.
It does not end with -ide, so it
is a polyatomic negative ion.
2. Find the chemical symbols
for the ions. Write the
symbol for the positive ion
first and the symbol for the
negative ion second.
The symbol for aluminum is
Al, and the symbol for chloride
is Cl. The formula without
subscripts is Al_Cl_.
The symbol for calcium is Ca,
and the symbol for iodide is I.
The formula without subscripts
is Ca_I_.
The symbol for potassium
is K, and the symbol for
permanganate is MnO4. The
formula without subscripts is
K_MnO4_.
3. Determine the charges
of the ions.
The aluminum ion has a
charge of 3+, and the chloride
ion has a charge of 1−.
The calcium ion has a charge
of 2+, and the iodide ion has a
charge of 1−.
The potassium ion has a charge
of 1+, and the permanganate
ion has a charge of 1−.
4. Check to see if the charges
differ. If the charges are
the same, the formula is
complete. If they differ,
determine the number of
each ion that is needed to
create a zero net charge.
Write the numbers of ions
needed as subscripts beside
the chemical symbols.
The charges differ, so use the
method in Figure 2.21 to find
the number of ions needed.
The charges differ, so use the
method in Figure 2.21 to find
the number of ions needed.
The charges are the same,
so the formula is KMnO4.
3+
1-
Al
2+
Cl
Al3Cl3
You need three chloride ions
for one aluminum ion. The
formula is AlCl3.
1-
Ca
I
Ca3I2
You need two iodide ions for
one calcium ion. The formula
is CaI2.
Writing Names and Formulas for Acids and Bases
Acids are compounds that ionize, or come apart, in water and release a hydrogen ion,
H+. Thus, the positive ion in acids is the hydrogen ion. For example, on page 62, you
read that when HCl dissolves in water, it separates into a hydrogen ion, H+, and a
chloride ion, Cl-.
Bases are compounds that produce a hydroxide ion, OH-, when they dissolve in
water. Thus, the negative ion in bases is usually the hydroxide ion. Notice that the
hydroxide ion was listed in Table 2.4 among the polyatomic ions. Sodium hydroxide,
NaOH(s), is a common example of a base. When NaOH(s) dissolves in water, it
separates into a sodium ion, Na+, and a hydroxide ion, OH-.
Naming and Writing Formulas for Bases
The rules for naming bases and for writing their formulas are the same as the rules for
naming and writing formulas for all other ionic compounds. For example, NaOH is
sodium hydroxide because Na+ is the metal ion and the name of the ion is the same as
the name of the metal. Hydroxide is the name of the polyatomic ion, OH-. Thus, any
compound with a metal ion or a positively charged polyatomic ion combined with the
hydroxide ion is a base. Pure bases are often solids. You can distinguish between the
pure base and the basic solution simply by noting its state, which is often indicated by a
symbol in brackets after the formula. For example, NaOH(s), where (s) represents solid,
is the pure compound. NaOH(aq), where (aq) means “aqueous solution,” is the solution
of sodium hydroxide in water.
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Terminology Involving Bases
The term alkali is often used to refer to a base that is soluble in water. This term comes
from soap making in the Middle East thousands of years ago. Warm water was poured
on the ashes from burnt wood or dried plants, dissolving the bases that were in the
ashes. The resulting solution was then boiled with animal fats to make soap. The Arabic
word al-qali means “the ashes.” The earliest records of soap making date from around
2800 bce in ancient Babylonia, which is now part of Iraq.
You probably recognize the term alkali from the Group 1 metals in the periodic
table, which are called alkali metals. As you know, the alkali metals react violently with
water. One of the products of that reaction is the hydroxide of the alkali metal, which is
a base that is dissolved in water.
alkali a base that is
soluble in water
Naming and Writing Formulas for Acids
Acids, in their pure form, are molecular compounds. However, they are named
according to the rules for ionic compounds. For example, pure HCl is hydrogen
chloride. Hydrogen is named as though it was the positively charged ion and its name
is not changed. Chloride is named as though it was the negatively charged ion. The
rules for writing formulas for acids are the same as the rules for writing formulas for
other ionic compounds.
When an acid is dissolved in water, the name is changed. The current naming
system recommended by IUPAC is relatively new and the older, classical naming
system is used so frequently that it is helpful to learn both systems. In the IUPAC
naming system, the name of the pure acid is simply preceded by the term “aqueous.”
For example, when hydrogen chloride is dissolved in water, it becomes aqueous
hydrogen chloride. The classical names, however, are not quite as simple. To learn the
classical names, it is convenient to separate acids into two categories, those that contain
oxygen and those that do not.
Acids That Do Not Contain Oxygen
The classical name for acids that do not contain oxygen is formed by omitting the word
hydrogen, adding the prefix hydro- and the suffix -ic and acid to the root name. For
example, hydrogen chloride becomes hydrochloric acid. Table 2.7 lists some examples.
Table 2.7 Names of Some Common Acids without Oxygen
Pure Substance
(name)
Formula
H(negative ion)(aq)
Classical Name
hydro(root)ic acid
IUPAC Name
aqueous hydrogen
(negative ion)
hydrogen fluoride
HF(aq)
hydrofluoric acid
aqueous hydrogen fluoride
hydrogen cyanide
HCN(aq)
hydrocyanic acid
aqueous hydrogen cyanide
hydrogen sulfide
H2S(aq)
hydrosulfuric acid
aqueous hydrogen sulfide
Notice, in Table 2.7, that all of the examples except HCN are binary acids. That is, they
contain only hydrogen and a non-metal.
Acids That Contain Oxygen
Acids that contain oxygen are called oxoacids. They are composed of hydrogen,
oxygen, and atoms of at least one other element, which is usually, but not always, a
non-metal. The combination of oxygen and an atom of another element is essentially,
a negatively charged polyatomic ion. In fact, almost any of the negatively charged
polyatomic ions in Table 2.4 can be found in acids. However, notice what would happen
if you combined a hydrogen ion with a hydroxide ion, which is a polyatomic ion. You
would have HOH, which is water.
oxoacid an acid
composed of hydrogen,
oxygen, and atoms of at
least one other element
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Naming Oxoacids
The rules for determining IUPAC names for oxoacids are the same as the rules for
naming acids with no oxygen atoms. To learn the classical naming system, you need to
refer to the system for naming polyatomic ions with varying numbers of oxygen atoms
in Table 2.5. Just as there are families of ions with varying numbers of oxygen atoms,
there are families of acids with varying numbers of oxygen atoms. Table 2.8 relates the
prefixes and suffixes of the polyatomic ions to those of the corresponding acids. Notice
that the prefixes remain the same while the suffix -ite changes to -ous acid and the suffix
-ate changes to -ic acid.
Table 2.8 Classical Naming System for Families of Oxoacids
Examples
Name of Ion
hypo(root)ite
Name of Acid
(dissolved in water)
hypo(root)ous acid
Name of Ion
hypochlorite, ClO-
(root)ite
(root)ous acid
chlorite, ClO2
(root)ate
(root)ic acid
chlorate, ClO3-
per(root)ate
per(root)ic acid
Name of Acid
(dissolved in water)
hypochlorous acid, HClO
chlorous acid, HClO2
chloric acid, HClO3
perchlorate, ClO4
-
perchloric acid, HClO4
To name other oxoacids, look for the prefix (if any) and the suffix and match them to
the prefix (if any) and suffix of the acid in Table 2.8. For example, the ion nitrate has
no prefix and the suffix is -ate. The acid would then have no prefix and would have the
suffix -ic acid. The name of the pure substance, hydrogen nitrate, when dissolved in
water would be nitric acid. When the ion already includes one hydrogen atom, such as
hydrogen carbonate, HCO3-, simply add another hydrogen, H2CO3(aq). The name of
the ion would be carbonate, and the acid would be carbonic acid.
Learning Check
7. What is a binary ionic compound?
8. Write the names and chemical formulas for the
compounds containing the following.
a. potassium and sulfur
b. oxygen and magnesium
c. chlorine and iron
d. magnesium and nitrogen
e. hydrogen and iodine
f. calcium and hydroxide ion
9. Write the name of each compound.
a. CrBr2
c. HgCl
e. HNO3(aq)
b. Na2S
d. PbI2
f. KOH
10. Write the chemical formula for each compound.
a. zinc bromide
d. magnesium chloride
b. aluminum sulfide
e. hydrogen nitride
c. copper(II) nitride
f. copper(II) hydroxide
11. The root of the names of the following ions is fluor.
Name each ion, and explain how you decided on the
name.
a. FO−
b. FO2−
c. FO3−
d. FO4−
12. Write the chemical formula for each compound.
a. iron(II) sulfate
d. magnesium phosphate
b. sodium nitrate
e. hydrogen carbonate
c. copper(II) chromate
f. aluminum hydroxide
Writing Names and Formulas for Binary Molecular Compounds
The names of molecular compounds include more details than the names of ionic
compounds, because non-metals can combine in a variety of ratios. For example,
nitrogen and oxygen can combine to form six different molecular compounds:
NO, NO2, N2O, N2O3, N2O4, and N2O5. Clearly, the name nitrogen oxide could mean
any of these compounds. The rules for naming binary molecular compounds make
it possible for each compound to have its own name, which clearly describes the
numbers of atoms in the compound. These rules are listed on the following page.
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Naming Binary Molecular Compounds
The rules listed below explain how to name binary molecular compounds. The prefixes
that are used for naming these compounds are listed in Table 2.9. Three examples
follow, in Table 2.10.
Table 2.9 Prefixes
for Binary Molecular
Compounds
Number
Prefix
1
mono-
1. Name the element with the lower group number first. Name the element with the
higher group number second.
2
di-
3
tri-
2. The one exception to the first rule occurs when oxygen is combined with a
halogen. In this situation, the halogen is named first.
4
tetra-
5
penta-
3. If both elements are in the same group, name the element with the higher period
number first.
6
hexa-
7
hepta-
4. The name of the first element is unchanged.
8
octa-
9
nona-
10
deca-
Rules for Naming Binary Molecular Compounds
5. To name the second element, use the root name of the element and add the suffix
-ide.
6. If there are two or more atoms of the first element, add a prefix to indicate the
number of atoms.
7. Always add a prefix to the name of the second element to indicate the number of
atoms of this element in the compound. (If the second element is oxygen, an “o” or
“a” at the end of the prefix is usually omitted.)
Table 2.10 Examples of Naming Molecular Compounds
Steps
Atoms in Compound two nitrogen atoms and one
oxygen atom
five iodine atoms and one
phosphorus atom
two chlorine atoms and
seven oxygen atoms
1. Name the element with the lower
group number (to the left in the
periodic table) first. Name the
element with the higher group
number (to the right in the periodic
table) second.
2. The one exception to the first rule
occurs when oxygen is combined
with a halogen. In this situation, the
halogen is named first.
3. If both elements are in the same
group, name the element with the
higher period number first.
Nitrogen is in Group 15
and oxygen is in Group 16,
so nitrogen comes first.
_nitrogen _oxygen
Iodine is in Group 17 and
phosphorus is in Group 15,
so phosphorus comes first.
_phosphorus _iodine
Chlorine is in Group 17 and
oxygen is in Group 16, so
oxygen should be first and
chlorine should be second.
However, when oxygen is
combined with a halogen,
the halogen is named first.
_chlorine _oxygen
4. The name of the first element is
unchanged.
5. To name the second element, use the
root name of the element and add
the suffix -ide.
The name nitrogen is
unchanged, but oxygen
is changed to oxide.
_nitrogen _oxide
The name phosphorus is
unchanged, but iodine is
changed to iodide.
_phosphorus _iodide
The name chlorine is
unchanged, but oxygen is
changed to oxide.
_chlorine _oxide
6. If there are two or more atoms of
the first element, add a prefix to
indicate the number of atoms.
7. Always add a prefix to the name of
the second element to indicate the
number of atoms of this element
in the compound. (If the second
element is oxygen, an “o” or “a”
at the end of the prefix is usually
omitted.)
There are two nitrogen
atoms, so the prefix is di-.
There is one oxygen atom, so
the prefix is mono-. Because
the second element is
oxygen, use mon-.
There is only one
phosphorus atom, so no
prefix is added. There are five
iodine atoms, so the prefix is
penta-.
There are two chlorine atoms,
so the prefix di- is added.
There are seven oxygen atoms
so the prefix should be hepta-.
However, the second element
The name of the compound is is oxygen so the “a” on heptais omitted. The prefix is hept-.
The name of the compound is phosphorus pentaiodide.
Th
e name of the compound is
dinitrogen monoxide.
dichlorine heptoxide.
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Writing Chemical Formulas for Binary Molecular Compounds
An important exception to all of the rules for naming and writing formulas for binary
molecular compounds occurs when the two elements in a compound are carbon
and hydrogen. Combinations of carbon and hydrogen constitute a large group of
compounds called hydrocarbons, which are a subgroup of a larger group of compounds
called organic compounds. Organic compounds consist of all compounds that contain
carbon atoms, other than carbon monoxide (CO), carbon dioxide (CO2), carbonates
(CO32-), cyanides (CN-), and carbides (several forms). Organic compounds have a
unique naming system. You will study organic chemistry in more advanced chemistry
courses. The following rules apply to inorganic compounds, which are all compounds
other than organic compounds. Table 2.11 provides three examples of naming binary
molecular compounds.
Rules for Writing Chemical Formulas for Binary Molecular Compounds
1. Write the symbol for the element with the lowest group number first.
2. Write the symbol for the element with the highest group number second.
3. The one exception to the first two rules occurs when oxygen is combined with
a halogen. In this case, the symbol for the halogen is written first.
4. If both elements are in the same group, write the symbol for the one with the
higher period number first.
5. If the number of atoms of either or both elements is greater than one, write the
number as a subscript beside the symbol. The absence of a subscript is understood
to mean one.
Table 2.11 Examples of Writing Chemical Formulas for Binary Molecular Compounds
Atoms in Compound two nitrogen atoms and
one oxygen atom
Steps
two chlorine atoms and one four bromine atoms and one
oxygen atom
silicon atom
1. Write the symbol for the element
with the lowest group number first.
2. Write the symbol for the element
with the highest group number
second.
Nitrogen is in Group 15
and oxygen is in Group 16,
so the symbol for nitrogen
is written first.
N_O_
Chlorine is in Group 17
and oxygen is in Group 16,
so the symbol for oxygen
should be written first.
O_Cl_
Bromine is in Group 17 and
silicon is in Group 14, so the
symbol for silicon is written
first.
Si_Br_
3. The one exception to the first
two rules occurs when oxygen is
combined with a halogen. In this
case, the symbol for the halogen is
written first.
Oxygen is not combined
with a halogen.
Oxygen and nitrogen are
not in the same group.
No changes are needed.
N_O_
Oxygen is combined with
the halogen chlorine, so the
symbol for chlorine comes
first.
Cl_O_
Oxygen is not in the
compound. Silicon and
bromine are not in the same
group.
No changes are needed.
Si_Br_
There are two nitrogen
atoms, so the subscript 2
is written beside N. There
is only one oxygen atom,
so there is no subscript
beside O.
There are two chlorine
atoms, so the subscript 2 is
written beside Cl. There is
only one atom of oxygen,
so there is no subscript
beside O.
There is one silicon atom, so
there is no subscript beside Si.
There are four bromine atoms,
so the subscript 4 is written
beside Br.
The formula is
The formula is
SiBr4.
N2O.
Cl2O.
4. If both elements are in the same
group, write the symbol for the one
with the higher period number first.
5. If the number of atoms of either or
both elements is greater than one,
write the number as a subscript
beside the symbol. The absence of a
subscript is understood to mean one.
The formula is
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Sample Problem
Names from Formulas and Formulas from Names
Problem
Write the names of the compounds for parts a and b. Write the formulas for parts c and d.
a. SF6
b. Cu(NO3)2
c. aluminum sulfide
d. sulfur trioxide
What Is Required?
You need to determine the names of SF6 and Cu(NO3)2
You also need to determine the formulas for aluminum sulfide and sulfur trioxide.
What Is Given?
You are given the formulas for two compounds: SF6 and Cu(NO3)2
You are given the names of two compounds: aluminum sulfide and sulfur trioxide
Plan Your Strategy
Act on Your Strategy
a. SF6 S is sulfur and F is fluorine. They are both non-metals so the compound _sulfur _fluoride
is molecular and you need prefixes.
Fluorine is the second element so you change the ending to –ide.
There is one sulfur atom so the prefix would be mono. However, it is the
first element so no prefix is needed.
There are six fluorine atoms so the prefix is hexa-.
b. Cu(NO3)2 Cu is copper and it is a metal. NO3 is a polyatomic ion and the
name is nitrate. The compound is ionic so you do not need prefixes.
sulfur hexafluoride
copper_ nitrate
A copper ion can have a charge of 1+ or 2+. Nitrate has a charge of 1- and copper(II) nitrate
there are two of the ions. Therefore the copper ion must have a charge of
2+ to make the compound neutral. Add (II) to the name of copper.
c. aluminum sulfide Aluminum is a metal. Its symbol is Al. Its charge is 3+.
Sulfur is a non-metal. Its symbol is S. Its charge is 2-.
The charges are not the same so you need subscripts. The subscript
(number of atoms) of each element is the same as the magnitude of the
charge of the other ion.
d. sulfur trioxide Sulfur and oxygen are both non-metals, so the compound is
molecular. You need subscripts. The symbol for sulfur is S. The symbol for
oxygen is O.
Sulfur has no prefix, so the number of atoms is assumed to be one and
therefore no subscript is needed.
Oxygen has the prefix tri-, meaning there are three oxygen atoms. Its
subscript is 3.
Al3+S2‒
Al2S3
S_O_
SO3
Check Your Solution
When you add up the charges on the ionic compounds, they add up to zero. The names and
symbols for the molecular compounds describe the same number of atoms of the same elements.
Practice Problems
1. Write the name of P4S7.
6. Write the formula for iron(III) oxide.
2. Write the name of Pb(NO3)2.
7. Write the formula for silicon dioxide.
3. Write the formula for manganese(IV) chloride.
8. Write the name of SeF6.
4. Write the formula for nitrogen triiodide.
9. Write the name of CaO.
5. Write the name of CuBr.
10. Write the formula for cobalt(III) nitrate.
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Drawing Structural Formulas for Molecular Compounds
structural formula
a diagram that has
the chemical symbols
connected by lines to
show the connections
among atoms in a
chemical compound
A chemical name or formula tells you how many atoms of each element are in a
molecule. However, it does not provide information about how the atoms are bonded
to one another. A Lewis structure shows you how the atoms are connected to each
other, but it is cumbersome to draw. Because chemists need an easier way to show
the connections between atoms, they developed structural formulas. You can draw a
structural formula from a Lewis structure by drawing a single line to represent a pair
of bonding electrons and omitting the lone pairs. Thus, one straight line represents one
bond. Figure 2.22 shows some Lewis structures you have seen before, as well as some
new Lewis structures, along with their structural formulas.
A
B
C
D
Cl
Cl
C
Cl
O
H
H
C
O
O
N
N
Cl
Cl
O
H
Cl
C
O
C
O
N
N
H
Cl
Cl
Figure 2.22 Compound (A) is carbon tetrachloride. Instead of drawing 32 dots, as you would
for a Lewis structure, you need to draw only four lines for a structural formula. Compound (B) is
dihydrogen monoxide. Compound (C) is carbon dioxide. Notice that the double bonds are drawn
as two lines. Compound (D) is nitrogen, which has a triple bond.
State the common name for dihydrogen monoxide.
SuggestedInvestigation
Inquiry Investigation 2-B,
Building Molecular Models
Although structural formulas provide more information than the simpler chemical
formulas, they are still only two-dimensional, while real molecules are three-dimensional.
Thus, to visualize a molecule completely, you need to build a model. You can use
something as simple as toothpicks and Styrofoam® balls to build three-dimensional
models. You can also use kits that have different sizes and colours of balls to represent
atoms of different elements. Regardless of the materials, these models can help you
visualize and analyze molecules.
Building your own models gives you an understanding of the three-dimensional
structure of molecules that you cannot attain any other way. However, modern
computers can now generate molecular models of very small to very large molecules.
The image in Figure 2.23 is a space-filling model. The relative sizes of the spheres and the
way they fit together is an excellent representation of the shape of the actual molecule.
Figure 2.23 This is a computer-generated model of ascorbic acid (vitamin C), C6 H8 O6.
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Section 2.2
RE V I E W
Section Summary
• The name of a binary ionic compound starts with the
name of the metal element and, if necessary, a roman
numeral indicating the charge on the ion. This is followed
by the name of the non-metal element with the ending
changed to -ide.
• The formula for a binary ionic compound starts with the
symbol for the metal element followed by the symbol for
the non-metal element. Subscripts indicate the numbers
of atoms of the two elements.
• Bases are named according to the rules for ionic
compounds.
• When acids are dissolved in water, they are named
according to different rules than when they are in their
pure form.
• The name of a binary molecular compound starts
with the name of the element that has the lower group
number. The name of the element that has the higher
group number is last, and the ending is changed to -ide.
Prefixes are used to indicate the numbers of atoms of the
two elements. However, a prefix is not used for the first
element if there is only one atom of this element.
• The formula for a binary molecular compound starts with
the symbol for the element with the lower group number,
followed by the symbol for the element with the higher
group number. Subscripts indicate the numbers of atoms
of the two elements in the compound.
• A structural formula shows how the atoms in a
compound are attached to each other.
Review Questions
1.
2.
3.
4.
T/I Turn to page 64 and read the “dinner table”
statements at the top of the page. Then answer the
following questions.
a. What do you think are the common names for
sodium chloride, dihydrogen monoxide, and
sodium hydrogen carbonate? Write the chemical
formulas for these compounds.
b. What do you think is the common name for sucrose?
c. Identify each compound as an ionic compound or a
molecular compound. Explain your reasoning.
K/U Explain why prefixes that indicate the numbers
of atoms of the different elements are not needed in the
names of ionic compounds.
K/U
What is a polyatomic ion?
What is the difference between a sulfate ion and
a sulfite ion? How would you be able to determine the
difference without looking up the names in a table?
8.
T/I The following six compounds contain nitrogen
and oxygen: NO, NO2, N2O, N2O3, N2O4, and N2O5.
Write the names of these compounds.
9.
Write the formula for each compound.
a. phosphorus pentachloride
b. difluorine monoxide
c. sulfur trioxide
d. silicon tetrabromide
e. cobalt(II) hydroxide
f. sulfur hexafluoride
10.
Write the name of each compound.
a. CO
c. CS2
e. SiO2
g. Ba(OH)2
b. BCl3
d. CCl4
f. PI3
h. H3BO3(s)
11.
K/U Explain why the name of C3H8 is not tricarbon
octahydride.
K/U
5.
T/I Write the name of each compound.
a. Al2O3
c. Na3P
e. NH4Cl
g. HNO3(aq)
b. HgI2
d. K3PO4
f. LiClO4
h. LiOH(aq)
6.
T/I Write the formula for each compound.
a. zinc oxide
d. magnesium iodide
b. iron(II) sulfide
e. cobalt(III) chloride
c. potassium hypochlorite f. sodium cyanide
7.
K/U Why must the name of a molecular compound
include prefixes to indicate the numbers of atoms of
the elements in the compound?
T/I
12.
T/I
Draw Lewis structures for these compounds.
From your Lewis structures draw structural formulas.
a. NF3
b. HCN
c. ClNO
13. C In a group, discuss the advantages and
disadvantages of using structural formulas.
14.
T/I
T/I First draw a Lewis structure for each compound.
Then, using your diagram, draw a structural formula.
a. two carbon atoms bonded to each other, and two
hydrogen atoms bonded to each carbon atom
b. two carbon atoms bonded to each other, with three
hydrogen atoms bonded to one carbon atom, and
one hydrogen atom and one oxygen atom bonded to
the second carbon atom
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SECTION
2.3
Key Terms
melting point
boiling point
dipole
dipole-dipole force
intermolecular forces
electrical conductivity
melting point the
temperature at which
a compound changes
from a solid to a liquid
Comparing the Properties of Ionic
and Molecular Compounds
It is not a coincidence that water melts at 0°C and boils at 100°C. The Celsius temperature
scale is based on the melting point and boiling point of water. All compounds have
melting points and boiling points, but these temperatures vary widely with the type of
substance. What factors determine the melting point and boiling point of a compound?
Do the same factors affect the other properties of a compound? This section will answer
these and other questions concerning the properties of ionic and molecular compounds.
Melting Points and Boiling Points of Compounds
As you read above, boiling points and melting points are unique to each pure compound.
Thus, they can provide important information about the characteristics of the compound.
For example, the melting point and boiling point of a compound reveal information about
the strength of the attractions that are holding the particles (ions or molecules) of the
compound together. Consider what is happening to a compound when it melts or boils.
Melting Point
The melting point of a compound is the temperature at which it changes from a solid
to a liquid at standard atmospheric pressure (the pressure exerted on the ground by
dry air at sea level, or 101.325 kPa). In a solid, the particles—ions or molecules—are so
strongly attracted to one another that they cannot pull apart. You can imagine a solid as
particles held together by springs, as shown in Figure 2.24 (A).
You might recall from previous science courses that, no matter how low the
temperature, all particles have some kinetic energy. So, although the particles in a
solid cannot pull away from the surrounding particles, they are always vibrating. You
have probably learned that temperature is directly related to the kinetic energy of the
particles in a substance. As energy in the form of heat enters a substance, the kinetic
energy, and thus the temperature of the substance, increases. When the kinetic energy of
the particles is great enough for the particles to pull away from one another, as shown in
Figure 2.24 (B), the temperature stops increasing and the compound melts. If the melting
point of a compound is very high, you know that a large amount of energy is needed for
the particles to pull away from one another. Therefore, the forces holding them together
must be very strong. A low melting point tells you that the particles are easily pulled
apart, and thus the forces attracting them to one another are relatively weak.
A
B
Figure 2.24 (A) Even in a solid, all particles are moving. (B) As the temperature of a substance
increases, more and more particles have enough energy to break away from their nearest
neighbour. Note that, in the case ofCHEM11_2.032A_33A.ai
molecular compounds, the spheres represent the entire
molecule and the springs represent attractive interactions between individual molecules.
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Boiling Point
The boiling point of a compound is the temperature at which it changes from a
liquid to a gas at standard atmospheric pressure. In a liquid, particles have enough
kinetic energy to pull away from one neighbouring particle, only to be attracted to
another neighbouring particle. The particles slide past one another. At the boiling
point, the particles have enough kinetic energy to completely break away from all the
other particles and the compound becomes a gas. Gas particles have enough energy
to bounce off one another when they collide rather than sticking together. Thus, the
boiling point of a compound, like the melting point, provides information about the
strength of the forces between the particles. A high boiling point indicates that the
attractive forces between the particles in a liquid are very strong. A low boiling point
tells you that these forces are relatively weak.
boiling point the
temperature at which
a compound changes
from a liquid to a gas
Forces between Particles in a Compound
As you read, a comparison of the melting and boiling points of a variety of substances
can provide information about the strength of the forces between ions in ionic
compounds and between molecules in molecular compounds. Note that when a
molecular compound melts or boils, the covalent bonds remain intact.
• A low melting point or boiling point means that particles with small amounts of
kinetic energy can break away from the adjacent particles. Thus the forces between
particles are weak.
• A very high melting point or boiling point means that the particles must have a very
large amount of kinetic energy to break away, and thus the forces between particles
are strong.
Keeping these relationships in mind, consider the data in Table 2.12.
Table 2.12 Melting Points and Boiling Points of Some Common Compounds
Compound
Melting Point (°C)
Boiling Point (°C)
ethanol (grain alcohol), C2H5OH
-114
+78.3
ammonia, NH3
-77.7
-33.3
cesium bromide, CsBr
+636
+1300
hydrogen, H2
-259
-253
hydrogen chloride, HCl
-114
-85
magnesium oxide, MgO
+2825
+3600
-182
-161
methane (natural gas), CH4
nitrogen, N2
-210
-196
sodium chloride, NaCl
+801
+1465
0
+100
water, H2O
If you analyze the data in Table 2.12 and classify the compounds into three categories,
you will get the results in Table 2.13. An analysis of the melting points would give the
same categories.
Table 2.13 Categories of Compounds Based on Boiling Point
High Boiling Point
Intermediate Boiling Point
Low Boiling Point
cesium bromide, CsBr
ethanol, C2H5OH
hydrogen, H2
magnesium oxide, MgO
ammonia, NH3
nitrogen, N2
sodium chloride, NaCl
hydrogen chloride, HCl
methane, CH4
water, H2O
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SuggestedInvestigation
Inquiry Investigation
2-A, Modelling Ionic
Compounds
Compounds with High Melting Points and Boiling Points
Consider the compounds with high boiling points in Table 2.13. These compounds
are all ionic. Their high boiling points are explained by the fact that the attractive
electrostatic forces between oppositely charged particles create very strong bonds.
An examination of the structure of ionic compounds will reveal why so much energy
is needed to break these bonds. Figure 2.25 shows the arrangement of sodium and
chloride ions in a crystal of sodium chloride. The same structure continues throughout
an entire crystal. Notice that each chloride ion is attracted to six adjacent sodium ions,
and each sodium ion is attracted to six adjacent chloride ions. Because the attractive
forces are all the same, there are no specific pairs of sodium and chloride ions that you
could identify as “molecules.” Each ion is strongly attracted to all the adjacent ions of
the opposite charge. There are continuous chains of ions that are attracted to each other
throughout the entire crystal, making the structure very stable. The formula, NaCl,
simply means that there is a 1:1 ratio of sodium to chloride ions in the entire crystal.
One sodium ion and one chloride ion are referred to as a formula unit of sodium
chloride, never as a molecule of sodium chloride.
A
Cl-
B
C
Na+
Figure 2.25 (A) The yellow spheres represent chloride ions and the blue spheres represent
sodium ions. (B) This model is called a “ball and stick” model. The balls represent the ions, and
the sticks represent the bonds. (C) This model is called a “space-filling model.” It shows that, in
the actual crystal, the ions are packed tightly together.
Compounds with Intermediate Melting Points and Boiling Points
Now consider the compounds with intermediate boiling points in the second
column of Table 2.13. If you look for a similarity among these compounds, you will
find that they are all molecular compounds. As well, they all have one or more polar
bonds. Depending on the overall structure of a molecule that has polar bonds, the
entire molecule can be polar. Figure 2.26 shows models of water and ammonia to
illustrate why they are polar. One end of each molecule is slightly negative, while the
other end is slightly positive.
A
B
δ-
δδ+
δ+
δ+
water
δ+
δ+
ammonia
Figure 2.26 The white spheres represent hydrogen atoms, the red sphere represents an oxygen
atom, and the blue sphere represents a nitrogen atom. You can think of the polarity as being
caused by the electrons spending more time around the oxygen atom in water (A) and the
nitrogen atom in ammonia (B) and less time around the hydrogen atoms.
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Representing Polar Molecules
δ+
dipole a molecule
with a slightly positively
charged end (positive
pole) and a slightly
negatively charged end
(negative pole)
dipole-dipole force
the attractive force
between the positive
end of one molecule
and the negative end of
another molecule
δ-
δ+
δ+
δ-
intermolecular forces
attractive forces that act
between molecules
δ+
δ-
δ-
δ+
δ+
δ-
δ-
δ-
δ+
δ+
A polar molecule is often represented as an oval shape with a slightly positively charged
end (positive pole) and a slightly negatively charged end (negative pole). Because a polar
molecule has one slightly positive end and one slightly negative end, it is often called a
dipole. Figure 2.27 shows how the positive ends of polar molecules attract the negative
ends of other polar molecules. This attractive force, called a dipole-dipole force, is much
smaller than the forces between ions. The dipole-dipole force is the main attractive force
that acts between polar molecules. The intermediate strength of this force results in the
intermediate boiling points of compounds that are composed of polar molecules.
δ-
Figure 2.27 Each oval represents a polar molecule. As the positively charged end of one
molecule is attracted to the negatively charged end of another, the molecules form a
continuous network.
Compounds with Low Melting Points and Boiling Points
Finally, consider the compounds in Table 2.13 that have low boiling points. Notice that
their molecules are all non-polar. The bonds between the carbon and hydrogen atoms
in methane are slightly polar, but the molecule, as shown in Figure 2.28, is symmetrical.
Therefore, the polarities of the bonds cancel one another in the whole molecule.
Nevertheless, some attractive forces exist between the molecules. Although non-polar
molecules have no distinct separation of charge, it is still possible for the positive nuclei
of atoms in one molecule to attract the electrons in a neighbouring molecule. These
attractions are very weak. As a consequence of these weak forces, compounds that are
composed of non-polar molecules have much lower boiling points than compounds
that are composed of polar molecules of a similar size.
In summary, these three interactions (strong attractive forces between ions, weaker
dipole-dipole attractive forces between polar molecules, and very weak attractive forces
between non-polar molecules), determine the boiling points and melting points of pure
substances. Because the dipole-dipole forces and the weak attractive forces act between
molecules, they are called intermolecular forces. This distinguishes them from the
covalent bonds that act within molecules. The intermolecular forces determine the
melting points and boiling points of molecular compounds.
δ+
δδ+
δ+
δ+
methane
Figure 2.28 The black
sphere represents a carbon
atom and the white spheres
represent hydrogen atoms in
this methane molecule. Each
bond is slightly polar, but the
symmetry of the molecule
makes it non-polar.
Learning Check
13. Explain what is happening, on the level of ions and
molecules, when a substance is melting.
15. Why is it incorrect to refer to a “molecule” of a
compound such as potassium iodide?
14. One compound has a melting point of 714°C.
Another compound, which is similar in size and
appearance, has a melting point of 146°C. How
would you classify these compounds based on their
melting points?
16. What is a dipole-dipole force?
17. Why do non-polar molecules have very low melting
and boiling points?
18. What forces are included within the category of
intermolecular forces?
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Other Properties of Ionic and Molecular Compounds
The strength and the type of bonds and intermolecular forces that exist among ions and
molecules affect several properties, in addition to melting points and boiling points.
Among these properties are solubility in water and electrical conductivity.
Solubility in Water
Whether or not a substance dissolves in water is an important property. For example,
many vitamins and nutrients in food (Figure 2.29) move through your bloodstream
from your digestive system to all of the tissues in your body because they are soluble
in water. Similarly, waste materials that are water soluble are carried to your kidneys
where they are eliminated from your body. Many chemical processes can take place
only when the compounds are dissolved in water. It is not always possible to predict
whether a compound will dissolve in water. However, differing trends in solubility can
be clearly seen when considering the polarities of substances.
Figure 2.29 All of the nutrients in these foods are critical to good health. The nutrients that are
soluble in water reach your bloodstream and are carried to your tissues quickly.
For a substance to dissolve in water, the water molecules must be more strongly
attracted to particles of that substance than to other water molecules. As you know,
water molecules are polar, having a slightly positive end and a slightly negative end.
The positive end will attract a negative ion or the negatively charged end of another
polar molecule. Likewise, the negative end of a water molecule will attract a positive
ion or the positively charged end of another polar molecule. Consequently, water
will dissolve many ionic compounds and polar compounds. For example, table sugar
(sucrose) is a polar molecular compound, and table salt (sodium chloride) is an ionic
compound. Both are soluble in water.
Water molecules are much more strongly attracted to each other than to non-polar
molecules. Therefore, most non-polar compounds do not dissolve in water. For
example, fats and oils are mixtures of non-polar compounds and they do not dissolve
in water.
Note that not all ionic compounds and polar molecular compounds are soluble
in water. In Unit 4, you will read more about solubility and learn how to determine
whether a particular compound is soluble in water.
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Electrical Conductivity
Electrical conductivity is the ability of a substance to allow an electric current to exist
within it. A substance can conduct an electric current only if charges (electrons or ions)
can move independently of one another. In a pure metal, electrons can move somewhat
freely because they are not tightly bound to the metal atoms. When a pure metal is
conducting an electric current, electrons are moving with ease from one metal atom to
the next.
In any type of compound, electrons are held tightly by the atoms. In an ionic
compound, electrons have moved from a metal atom to a non-metal atom. Once
they are bound to the non-metal atom, however, they are held tightly. A pure ionic
compound can only conduct an electric current under conditions in which entire ions
can move independently of one another.
As you know, ionic compounds are solid at room temperature. In a solid, the
oppositely charged ions are held rigidly together. Therefore, in their solid form, ionic
compounds cannot conduct an electric current. When an ionic compound is in the
liquid state, however, the ions are free to move independently of one another. This
occurs only at very high temperatures, but, at these temperatures, ionic compounds can
conduct an electric current.
Ionic compounds can also conduct an electric current when in an aqueous state, as
shown in Figure 2.30. When an ionic compound is dissolved in water, the ions are free
of other ions because they are surrounded by water molecules. Thus, ionic solutions
can also conduct an electric current.
e-
electrical conductivity
the ability of a substance
or an object to allow an
electric current to exist
within it
ee-
positive
electrode
negative
electrode
e-
power
source
Figure 2.30 Positive ions are attracted to a negative electrode, and negative ions are attracted
to a positive electrode. The ions create an electric current as they move around each other in
opposite directions.
When atoms are bound together in a molecular compound, they are sharing electrons. The
electrons never leave one atom completely. Therefore, there are no positive and negative
charges that are independent of one another. This means that molecular compounds
cannot conduct an electric current regardless of whether they are non-polar or polar.
If a polar compound is dissolved in water and electrodes are placed in the water, the
molecules will orient themselves so that their positively charged end is directed toward
the negative electrode and their negatively charged end is directed toward the positive
electrode. However, the charges never leave the molecules. Thus, even in a water solution,
molecular compounds cannot conduct an electric current. You might recall that acids
are molecular compounds when in a pure form but come apart and become ionic when
dissolved in water. Therefore, aqueous solutions of acids do conduct an electric current.
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Section 2.3
RE V I E W
Section Summary
• The strength of the attractive forces acting between ions
or molecules determines the melting point and boiling
point of a compound.
• Ionic compounds usually have the highest melting points
and boiling points. Polar molecules have intermediate
melting points and boiling points, and non-polar
molecules have the lowest melting points and boiling
points for molecules of similar sizes.
• Ionic and polar compounds are likely to be soluble in
water. Non-polar compounds are insoluble in water.
• For a substance to conduct an electric current, oppositely
charged particles must be free to move independently of
one another.
Review Questions
1.
K/U
Explain the basis of the Celsius temperature
13.
K/U Can polar molecular compounds conduct
electric current under either of the conditions that you
described in question 12? Explain why or why not.
14.
A
To be transported throughout the body in the
bloodstream, fat molecules must be bound to protein
molecules, as shown in the following figure. Explain
why you think this is necessary.
scale.
2.
K/U Describe, on the level of individual particles,
what happens to a substance when it is heated.
3.
K/U What property of particles determines whether
they will pull away from adjacent particles?
4.
How would you classify a compound that has a
boiling point of -182°C? Explain your answer.
K/U
5.
K/U Explain why compounds consisting of polar
molecules are likely to have a higher melting point
than compounds consisting of non-polar molecules.
6.
T/I What would you predict about the melting point
of a compound that will not dissolve in water? Explain
your thinking.
7.
Explain how an attractive force can exist
between non-polar molecules.
8.
If a compound has very high melting and
boiling points, is the compound likely to be soluble in
water? Explain the relationship between these two
properties of a compound.
9.
two types of fat bound to protein
red blood cell
K/U
15.
A
You might have heard the saying, “Like dissolves
like.” From what you have learned about solubility,
comment on the validity of this statement.
16.
Two molecular compounds, X and Y, have
similar masses. Compound X is solid at room
temperature, has a melting point of 146°C, and is
soluble in water. Compound Y is liquid at room
temperature, has a melting point of -10°C, and is
not soluble in water.
a. What would you predict about the polarities of
compound X and compound Y?
b Based on your predictions, explain the differences
in their melting points and solubilities.
T/I
C
Use sketches to show how a non-polar molecule
can have polar bonds.
10.
K/U Describe what must happen, on a particle level,
for a substance to dissolve in water.
11.
K/U Glycerol is a compound that dissolves readily in
water. The water solution of glycerol, however, will not
conduct an electric current. What would you predict
about the properties of glycerol?
12.
K/U Under what two conditions can an ionic
compound conduct an electric current?
T/I
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Inquiry
INVESTIGATION
2-A
Skill Check
Initiating and Planning
✓
Performing and Recording
✓
Analyzing and Interpreting
✓
Communicating
Materials
• crystal structure model kit
or
• polystyrene balls of two
different sizes
• toothpicks
Modelling Ionic Compounds
In this investigation, you will build models of crystals of three different ionic
compounds to better visualize their structures.
Pre-Lab Questions
1. What information is provided by a formula for an ionic compound?
2. Describe the bonding that occurs between ions in an ionic compound.
3. What do the “balls” and “sticks” represent in a ball-and-stick model
of a compound?
Question
What can you predict about the structures of crystals by building models?
A
S2Zn2+
Procedure
1. Your teacher will give you a crystal structure model kit, or polystyrene
balls and toothpicks.
2. Carefully study the arrangement of the ions in the sodium chloride
model in Figure 2.25.
3. Choose which size or colour of balls you will use to represent the sodium
and chloride ions. Discuss, with your partner, how the ions are arranged and
how you will connect the “ions.” Build a model of a sodium chloride crystal.
zinc sulfide
B
4. Study the illustrations of the zinc sulfide and calcium fluoride crystals
shown here. Each illustration shows one “repeating unit” for a crystal of
each of the two compounds.
5. Repeat step 3 for zinc sulfide and for calcium fluoride. Build models of at
least two “repeating units” for each compound.
Ca2+
F-
6. Compare your models with another group’s models. If your models are
not the same, discuss the differences and decide which, if any, are the
correct models.
Analyze and Interpret
1. What is the ratio of metal ion to non-metal ion in each of your models?
calcium fluoride
These diagrams represent “repeating
units” for each of the ionic compounds.
Note that the ions on the outer edges
(sulfur in the zinc sulfide and calcium
in the calcium fluoride) are bonded to
more oppositely charged ions in the
adjacent “repeating units.” This is the
reason that some sulfur ions in (A) do
not appear connected to other ions.
2. Provide a possible explanation as to why the ratio of metal to non-metal
ions can be the same and the structures of the crystals can be different.
Conclude and Communicate
3. How well do you think your models represent real crystals? Describe ways
in which your models are similar to real crystals and ways in which they
are different.
Extend Further
4. RESEARCH Using print and Internet resources, research the technique first
used by chemists to determine the crystal structure of ionic compounds.
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Inquiry
INVESTIGATION
2-B
Skill Check
Initiating and Planning
✓
Performing and Recording
✓
Analyzing and Interpreting
✓
Communicating
Materials
• molecular model kit
Building Molecular Models
Models are very important for chemists. You cannot see detailed features
of a molecule, even with a microscope. However, you can build a model
that shows some of the properties that chemists have determined through
experimentation. In this investigation, you will use a molecular model kit
to assemble models of a few molecules.
• pen
• paper
James Watson and Francis Crick won a Nobel prize for their discovery
of the structure of DNA (deoxyribonucleic acid). In this photograph, they
are discussing their model.
Pre-Lab Questions
1. When drawing a Lewis structure, what basic rule tells you where the
electrons must be?
2. What characteristics of a Lewis structure tell you whether a bond
between two atoms is a single bond, a double bond, or a triple bond?
3. What is the difference between a two-dimensional image and a
three-dimensional image?
Question
What can you predict about the structures of molecules by building models?
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Procedure
1. Your teacher will provide a molecular model kit for
you to use.
2. Copy the table shown below into your notebook. Use
your table for drawings in Procedure steps 3, 4, and 5.
Each cell in the last two columns must be large enough
for you to draw Lewis structures and sketches of
your molecular models.
Data Table for Model Building
Name
Formula
Lewis Structure
Sketch of Shape
Analyze and Interpret
1. Compare your models and sketches with those of your
classmates. Discuss any differences.
Conclude and Communicate
2. What can you learn from models that you cannot learn
from Lewis structures?
3. Summarize the strengths and limitations of creating
molecular models using kits. What can you infer from
the models? What features of the molecules cannot be
inferred from the models?
Extend Further
3. In your table, write the name and formula, and
draw a Lewis structure of each molecule below.
a. hydrogen bonded to hydrogen
b. chlorine bonded to chlorine
c. oxygen bonded to two hydrogen atoms
d. carbon bonded to two oxygen atoms
e. nitrogen bonded to three hydrogen atoms
f. carbon bonded to four chlorine atoms
g. nitrogen bonded to three fluorine atoms
4. INQUIRY Describe the difference between ball-andstick models and space-filling models. Discuss the
advantages and disadvantages of using each type of
model.
5. RESEARCH Using print and Internet resources,
research a discovery of a structure in chemistry
or biochemistry that depended heavily on
model building.
4. Look through your textbook, and choose three molecules
that are not in the list above. Record the names and
formulas for these compounds in your data table.
Draw Lewis structures of these molecules in your table.
5. Based on your Lewis structures, build models of all
the molecules. Make a sketch of each of your models.
Consult the directions that came with the kit for
information about assembling the models.
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STSE
Case Study
Feminization of Male Fish
Monitoring the Effects of Environmental Estrogens
Scenario
To meet your community involvement requirement for
your Grade 12 diploma, you have begun volunteering
at a conservation area in your region. The Conservation
Authority has been actively developing a new watershed
protection plan. This plan is especially important because a
new wastewater treatment plant was built along a river in
the middle of the watershed several years ago. Local citizens
are concerned about how this wastewater treatment plant
might impact the quality of drinking water in the area. The
Conservation Authority has been collecting data related to
the health and status of organisms in local stream and river
ecosystems. You have recently been helping to collect data
by sampling fish populations in the watershed.
The sampling data that you have collected in this year
show some startling changes in the ecosystem. In particular,
approximately 50% of the male fish sampled also have some
female anatomy. As well, there are about five times as many
females compared to males in the population. These results
suggest that some type of environmental estrogen has been
entering the water from some, as yet unknown, source.
Natural and Synthetic Estrogens
There are many natural estrogens. These include estrogens
produced in living plants as well as those that are produced
in animals. However, there are also many environmental
estrogens. These environmental estrogens are synthetic
compounds that mimic estrogen activity and that are
released into the environment as by-products of industry.
More than 60 chemical substances, including dioxin and
DDT, have been identified as environmental estrogens.
These compounds are produced for use in many different
sectors, including the pharmaceutical, plastics, and
detergent-manufacturing industries.
Nonylphenol ethoxylates (NPEs)
Data previously collected by the Conservation
Authority in your region indicate that
one particular class of environmental
estrogens, called nonylphenol ethoxylates
(NPEs), is the likely cause of the changes in
the fish you have observed. These chemicals
have been used for more than 40 years as
detergents, emulsifiers (which keep oils
dispersed in a liquid to prevent clumping)
wetting agents (that lower the surface tension
of a liquid so that liquids mix together more
easily) and dispersing agents (that are used to
keep particles in a suspension from clumping and
coming out of suspension). NPEs are not produced
naturally–their presence in the environment is
entirely the result of human activity. The two most
likely sources of environmental NPEs are wastewater
from industrial operations and the water output from
municipal wastewater treatment plants.
Waterways and marshlands are the breeding grounds for many
species of fish. Environmental estrogens that mimic the action
of the female hormone, estrogen, can get into these waters from
waste treatment plants or industrial waste. When developing male
fish are exposed to these estrogens, they can become feminized.
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Research and Analyze
1. As you read, more than 60 chemical substances,
including NPEs, dioxin, and DDT have been
identified as environmental estrogens. Research
sources of environmental estrogens other than
NPEs, including products manufactured by the
pharmaceutical industry. What properties of some
common pharmaceutical products allow them to
persist in water systems and influence the growth
and development of organisms?
United States, in 2010.
European Union, and the
the
,
ada
Can
in
Es
NP
of
Status
in the Canadian
red to be “toxic” as defined
Canada NPEs are conside
eless, their use is not
Act, (CEPA), 1999. Neverth
Environmental Protection
the future, research is
t that they are banned in
es of
legally banned. In the even
that can serve the purpos
test alternative chemicals
underway to develop and
Es.
NP
of
production
rently responsible for the
the compounds that are cur
sewage treatment practices
te
qua
ade
re
wn that, whe
Canadian studies have sho
e a significant risk to
osure to NPEs does not pos
are employed, normal exp
not generally pose a risk to
current use of NPEs does
t
human health. As well, the
to the aquatic environmen
However, discharge directly
the aquatic environment.
ms.
anis
aquatic org
ted waste is likely to harm
of untreated or partially trea
e
European Union. They hav
are legally banned in the
European Union NPEs
are
ts
duc
pro
ts. The new
expensive but safer produc
been replaced with more
.
alcohol-based compounds
ition that NPEs and
Union have taken the pos
an
ope
Eur
the
Offi cials in
th in humans with
se serious illness or even dea
their by-products can cau
even at very low exposures.
t make them susceptible,
pre-existing conditions tha
factor to breast cancer,
t NPEs are a contributing
They have established tha
among other diseases.
States. In fact, they are
not banned in the United
United States NPEs are
available compounds for
ured as cheap and readily
widely used. NPEs are favo
eve that NPEs and their
s in the United States beli
are
industrial processes. Official
a sewage treatment plant
in water after treatment at
by-products that remain
c limits for toxicity.
within established scientifi
y are aware of the
the United States is that the
The position of officials in
ironment, and believe
its of NPEs have on the env
als
impact that established lim
they feel that other chemic
able. On the other hand,
and
,
these impacts to be accept
ood
erst
und
l
are not wel
as replacements for NPEs
that might be considered
m.
the
for
its
lim
sible to set safe
thus, it would not be pos
2. Research some effects of estrogen-mimicking
agents (environmental estrogens) on human health.
What are the accepted limits and concentrations of
some of the common environmental estrogens in
ecosystems? How consistently and effectively are
ecosystems monitored in Canada for the presence
of environmental estrogens?
3. Environmental estrogens appear to have a
significant impact on biological organisms and
ecosystems. Proven environmental estrogens, for
example, pesticides such as DDT, have been banned
from use in many parts of the world. However, these
synthetic estrogens persist in the environment
and organisms, including humans, continue to be
exposed to them due to natural cycling of air and
water. Analyze the impact of banning a substance
such as an environmental estrogen in certain parts
of the world but not in others.
Take Action
1. Plan In a group, discuss the controversy
surrounding the use of environmental estrogens.
What are some key issues to consider when
analyzing how to reduce or eliminate the impacts of
these compounds? Share the results of the research
and analysis you conducted in questions 1 to 3
above.
2. Act Prepare an informational brochure that could
be handed out by the Conservation Authority at
their public information session on the watershed
protection plan. Ensure that you create a brochure
that will enable your audience to understand
the properties of environmental estrogens, the
potential risks associated with their presence in
the environment, and possible ways to reduce the
impact of environmental estrogens. Support your
position with information from credible sources.
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Chapter 2
Section 2.1
SUMMARY
The Formation of Ionic and Covalent Bonds
An ionic bond forms when oppositely charged ions
attract each other. A covalent bond forms when two
atoms share one or more pairs of electrons.
Key Terms
bonding pair
covalent bond
double bond
electronegativity difference
ionic bond
ionic compound
Lewis structure
Section 2.2
lone pair
molecular compound
octet rule
polar covalent bond
polyatomic ion
single bond
triple bond
• An ionic bond forms when a negatively charged ion and a
positively charged ion are attracted to each other.
• A covalent bond forms when two atoms share one or more
pairs of electrons.
• A polyatomic ion consists of two or more atoms that
are covalently bonded together and carry a charge. A
polyatomic ion can form an ionic compound with a simple
ion or another polyatomic ion of the opposite charge.
• A chemical bond can be non-polar covalent, slightly polar
covalent, polar covalent, or mostly ionic, depending on the
electronegativity difference between the two atoms that
are bonded together.
Writing Names and Formulas for Ionic and Molecular Compounds
The name and chemical formula for a compound
describe exactly how many atoms of each element
are present in one particle of the compound.
Key Term
alkali
oxoacid
structural formula
Key Concepts
• The name of a binary ionic compound starts with the name
of the metal element and, if necessary, a roman numeral
indicating the charge on the ion. This is followed by the
name of the non-metal element with the ending changed
to -ide.
• The formula for a binary ionic compound starts with the
symbol for the metal element followed by the symbol for
the non-metal element. Subscripts indicate the numbers
of atoms of the two elements.
Section 2.3
Key Concepts
• The octet rule can be used to predict how bonds will form.
• Bases are named according to the rules for ionic
compounds.
• When acids are dissolved in water, they are named according
to different rules than when they are in their pure form.
• The name of a binary molecular compound starts with the
name of the element that has the lower group number. The
name of the element that has the higher group number is
last, and the ending is changed to -ide. Prefixes are used
to indicate the numbers of atoms of the two elements.
However, a prefix is not used for the first element if there
is only one atom of this element.
• The formula for a binary molecular compound starts with
the symbol for the element with the lower group number,
followed by the symbol for the element with the higher
group number. Subscripts indicate the numbers of atoms
of the two elements in the compound.
• A structural formula shows how the atoms in a compound
are attached to each other.
Comparing the Properties of Ionic and Molecular Compounds
The type of bonds and the shape of the particles
influence the properties of compounds, such as their
melting and boiling points, solubility in water, and
electrical conductivity.
Key Terms
boiling point
dipole
dipole-dipole force
electrical conductivity
intermolecular forces
melting point
Key Concepts
• The strength of the attractive forces acting between ions or
molecules determines the melting point and boiling point
of a compound.
• Ionic compounds usually have the highest melting points
and boiling points. Polar molecules have intermediate
melting points and boiling points, and non-polar molecules
have the lowest melting points and boiling points.
• Many ionic and polar compounds are soluble in water.
Non-polar compounds are insoluble in water.
• For a substance to conduct an electric current, oppositely
charged particles must be free to move independently of
one another.
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Chapter 2
REVIEW
Knowledge and Understanding
Select the letter of the best answer below.
1. Which statement about ionic compounds is false?
a. An ionic compound is comprised of ions held
together by an electrostatic force.
b. An ionic compound typically consists of a metal ion
and a non-metal ion.
c. An ionic compound contains the same number of
oppositely charged ions.
d. An ionic compound has a zero net charge.
e. The composition of an ionic compound can often be
predicted by the octet rule.
2. The circled electrons in this Lewis diagram are called
a.
b.
c.
d.
e.
unpaired electrons
free electrons
an electron pair
a bonding pair
an unbound pair
O
O
3. The electronegativity of magnesium is 1.3, and the
electronegativity of oxygen is 3.4. The bond that forms
between them is
a. mostly ionic
b. polar covalent
c. slightly polar covalent
d. non-polar covalent
e. none of the above
4. The chemical name of Mg(ClO3)2 is
a. magnesium chloride
b. magnesium dichlorite
c. magnesium chlorite
d. magnesium chlorate
e. magnesium hypochlorite
5. The element that comes second in the name of a binary
molecular compound is the element that
a. has the lower group number
b. has the higher group number
c. has the higher period number
d. is the non-metal
e. has the greater mass
6. The chemical name of SiBr4 is
a. monosilicon tetrabromide
b. silicon hexabromide
c. monosilicon pentabromide
d. silicon octabromide
e. silicon tetrabromide
7. Which statement about the properties of compounds
is true?
a. A compound that has a very high melting point is
a liquid at room temperature.
b. Ionic bonds are stronger than intermolecular forces.
c. Non-polar molecules experience no intermolecular
forces.
d. A compound that has a very low boiling point is a
liquid at room temperature.
e. Dipole-dipole forces are stronger than the force
between oppositely charged ions.
8. Which compound is most likely to be soluble in water?
a. a non-polar compound
b. a slightly polar compound
c. a polar compound
d. an ionic compound
e. all of the above
Answer the questions below.
9. In this chapter, you read that ores are metals combined
with non-metals. How would you classify the
compounds that are found in ores? Why?
10. Several different gaseous compounds that consist of
non-metals are found in the atmosphere. How would
you classify these gaseous compounds? Why?
11. Aluminum ions have a charge of 3+ and oxide ions
have a charge of 2-. How can aluminum ions and
oxide ions combine to form a compound with a net
charge of zero?
12. Copy the following diagram and complete a Lewis
structure for the compound. Draw a circle around each
atom and its electrons and describe how each atom
satisfies the octet rule.
H
H
H
C
C
Cl
H
F
13. Explain the meaning of the term “bond dipole.”
14. The boiling point of a compound depends on a balance
between two conditions. What are these conditions?
Explain.
15. Describe the two forces that make up intermolecular
forces.
16. State which type of compound, ionic or molecular,
can conduct electric current. What conditions are
necessary for this type of compound to conduct
electric current?
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Chapter 2
REVIEW
Thinking and Investigation
17. Use Lewis diagrams to predict the ratio of metal to
non-metal ions in a compound formed by each pair of
elements.
a. magnesium and fluorine
b. potassium and bromine
c. rubidium and chlorine
d. calcium and oxygen
18. Each of the following Lewis structures has an error
in it. State what the error is, and draw the correct
Lewis structure.
A
H
B
H
C
C
O
F
F
C
F
H
C
H
22. Name each compound.
a. SO2
b. N2O4
c. CO
d. Cl2O
23. Write the formula for each compound.
a. dihydrogen monoxide
b. sulfur trioxide
c. silicon tetrachloride
24. Identify the errors in each phrase or statement, and
rewrite it correctly.
a. four molecules of potassium bromide
b. The compound NaHSO4 is sodium sulfate.
c. The compound KNO2 is potassium nitrate.
Communication
25.
The type of chemical bond in a compound
determines the physical and chemical
properties of that compound. Name and sketch two
different types of chemical bonds. For each bond type,
describe two ways in which it influences the properties
of the compound.
26.
It is important to use chemicals properly to
minimize the risks to human health and the
environment. You read that when sodium, a highly
reactive metal, is combined with chlorine, a toxic gas,
the product, sodium chloride, is very safe. Using print
and Internet resources, research another element or
compound that can be made safe by reacting it with
another element or compound. Share your findings in
the format of your choosing.
H
C
O
H
19. Name each compound.
a. MgCl2
b. Na2O
c. FeCl3
d. CuO
e. Ba(ClO)2
f. NH4NO3
g. H2CrO4(aq)
h. H3PO4(s)
i. KOH
j. Cd(OH)2
20. Write the formula for each compound.
a. gold(III) chloride
g. aqueous hydrogen
b. magnesium oxide
chloride
c. lithium nitrite
h. sulfuric acid
d. calcium phosphide
i. cobalt(II) hydroxide
e. manganese(II) sulfide j. lithium hydroxide
f. calcium hypochlorite
21. Draw a Lewis structure of each molecule consisting
of the following combinations of atoms.
a. one carbon atom bonded to three hydrogen atoms
and one chlorine atom
b. one carbon atom bonded to two sulfur atoms
c. two iodine atoms bonded together
d. three carbon atoms bonded together in a chain;
three hydrogen atoms bonded to each of the carbon
atoms on the ends; an oxygen atom bonded to the
central carbon atom
27. In some Lewis diagrams, one of the chemical symbols
might have no dots. Draw an example of this, and
explain why one of the symbols has no dots.
28. Identify the chemical bonds in the following
compounds as mostly ionic, polar covalent, slightly
polar covalent, or non-polar covalent. Show and
explain the calculations you used to identify the bonds.
a. calcium chloride
b. carbon dioxide
c. nitrogen
d. silicon tetrachloride
29. “If there were no intermolecular forces, all molecular
compounds would be gases.” Do you agree or disagree
with this statement? Explain your reasoning, as if
you were explaining it to a classmate who did not
understand intermolecular forces.
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30. The boiling points of argon (−186°C) and fluorine
(−188°C) are quite similar. Write a paragraph that
you could read to help a Grade 10 student understand
why these boiling points are similar, based on
intermolecular forces.
31. Molecules of methane, CH4, and water, H2O, have
similar masses. However, their boiling points are very
different. The boiling point of methane is −161°C, and
the boiling point of water is +100°C. Draw sketches of
these molecules, and use your sketches to explain why
their boiling points are so different.
32. Write the names of the following ions: I-, IO−, IO2−,
IO3−, IO4−. The last four ions are polyatomic ions.
Design a different naming system that you think would
be descriptive of the ions and easy to remember.
33. Draw a structural formula based on the Lewis structure
shown here. Explain, in detail, the relationship between
the two diagrams.
O
H
C
O
38. Pure sodium can be extracted from sodium chloride
using a process called electrolysis. Sodium ions
can pick up electrons from one electrode and form
sodium atoms. Chloride ions can give up electrons
to the other electrode and form chlorine atoms,
which then combine to form molecules of chlorine
gas. The diagram shown here is a simplified sketch
of the apparatus. Imagine that you were asked to
design the containers and other equipment for
this process. Review what you have learned about
compounds that carry an electric current and about
the properties of sodium metal and chlorine gas.
Describe the challenges you would have to overcome
when designing the equipment. Present some possible
solutions to these challenges.
e-
e+ power supply positive
electrode
negative
electrode
Cl2 gas
←ClNa+→
molten sodium
chloride
Na metal
H
34. Summarize your learning in this chapter using
a graphic organizer. To help you, the Chapter 2
Summary lists the Key Terms and Key Concepts. Refer
to Using Graphic Organizers in Appendix A to help
you decide which graphic organizer to use.
Application
35. You have two white crystalline solids. One is an ionic
compound, and the other is a molecular compound.
Design an investigation to determine which is which.
Assume that your investigation cannot involve
dissolving them in water.
36. Water and methanol, CH3OH (a type of alcohol), mix
together in any proportions. Find their boiling points.
Then, based on the boiling points you find, design a
method you could use to separate water and methanol
that are mixed together.
39. In 1906, the Nobel Prize in Chemistry was awarded
to French chemist, Henri Moissan, for isolating
fluorine in its pure elemental form. Why would
this achievement be deserving of such a prestigious
honour? Use your understanding of the properties of
the elements, as well as chemical bonds, to explain
your answer.
40. You might have heard advertisements about detergents
that “break up grease.” Oil and grease consist of large
non-polar molecules, which are very insoluble in
water. Nevertheless, detergents, which seem to dissolve
in water, can remove oil and grease from clothing in
water. A space-filling model of a typical detergent
molecule is shown below. Study the model, and provide
a possible explanation for how detergents can remove
grease from clothing.
non-polar
polar
37. Suppose that you have two colourless solutions. One is
a solution of an ionic compound in water, and the other
is a solution of a molecular compound in water. Design
an investigation to determine which solution is which.
Describe the tests you would perform and the results
you would expect for each solution.
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Chapter 2
Self-Assessment
Select the letter of the best answer below.
1.
K/U Ionic bonds form between which two types of
elements?
a. metals and metalloids
b. metals and non-metals
c. metalloids and non-metals
d. two non-metals
e. two metals
2.
K/U What is the correct name of the compound
formed from Fe3+ and Cl−?
a. iron chloride
b. iron(I) chloride
c. iron chloride(I)
d. iron(II) chloride
e. iron(III) chloride
3.
K/U The electrons in a non-polar bond are
a. gained by one atom and lost by the other
b. shared equally
c. shared unequally
d. gained
e. lost
4.
Which compound is ionic?
a. KBr
b. H2O
c. HCl(g)
d. NH3
e. CH4
5.
K/U Which compound is molecular?
a. NaOH
b. PbCl2
c. MnO2
d. CaCrO4
e. SiO2
6.
7.
K/U What type of bonding occurs within a water
molecule and between water molecules?
a. non-polar covalent within a water molecule and
ionic between water molecules
b. polar covalent within a water molecule and dipoledipole between water molecules
c. polar covalent within a water molecule and ionic
between water molecules
d. non-polar covalent within a water molecule and
dipole-dipole between water molecules
e. ionic within a water molecule and dipole-dipole
between water molecules
8.
K/U Which statement about ionic compounds is
false?
a. They may be soluble in water.
b. They have very high melting points.
c. They cannot conduct electric current when melted.
d. They are held together with ionic bonds.
e. They are solid at room temperature.
9.
K/U Which statement about molecular compounds
is false?
a. They can be solid, liquid, or gas at room
temperature.
b. They conduct electric current when dissolved in
water.
c. They are held together with covalent bonds.
d. They have low to moderate boiling points.
e. They can be polar.
K/U
The name of the compound N2O4 is
nitrogen oxide
dinitrogen dioxide
nitrogen tetraoxide
dinitrogen tetroxide
tetranitrogen dioxide
K/U
a.
b.
c.
d.
e.
10.
Intermolecular forces consist of
a. ionic and covalent bonds
b. dipole-dipole forces and ionic bonds
c. dipole-dipole forces and weak attractive forces
d. covalent bonds and dipole-dipole forces
e. weak attractive forces and ionic bonds
K/U
Use sentences and diagrams, as appropriate, to answer the
questions below.
11.
A
State the octet rule, and give an example of how
you would apply it to determine the number of sodium
ions that would be needed to form an ionic compound
with sulfur.
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12.
A
Draw Lewis diagrams of the following elements
in the order given: hydrogen, carbon, carbon,
hydrogen. Connect the atoms to form a Lewis structure
of a molecule. (Make sure that you keep the atoms in
the order given.) Is the bond between the two carbon
atoms a single, double, or triple bond? Explain how
you know what type of bond it is.
13.
Draw the Lewis structure of a nitrogen
molecule. Circle and label the bonding pairs and the
lone pairs.
14.
Draw a diagram to show how you would use the
charges on ions to determine the subscripts for the
chemical formula for an ionic compound. Explain your
diagram. Under what circumstances would you use
this method?
15.
16.
C
Use a Venn diagram to compare Lewis
structures and structural formulas. Give an example
of each. Under what circumstances would you use a
structural formula instead of a Lewis structure?
21.
T/I The melting points of three compounds are
listed below. Predict the type of attractive forces
between the particles of each compound when the
compound is in its solid form.
Melting Points of Three Compounds
Compound
T/I
2489
nitrogen trichloride
−40
A classmate asks, “How could there possibly be
any intermolecular forces between non-polar
compounds?” Answer your classmate’s question, using
a diagram to support your explanation.
23.
T/I List the following compounds in the order of
their boiling points, from lowest to highest without
knowing any exact boiling points. Explain your
reasoning for your order, based on the structures given.
C
H
H
T/I
Name each compound.
Si3N4
PCl5
SF6
ClF3
C
+
H
O
Cl
Cl
K 2+ O
K
H
methanol
chlorine
potassium oxide
24.
C
Make a sketch to represent dipole-dipole forces.
Explain what is happening in your sketch. What effect
do dipole-dipole forces have on the properties of a
polar compound?
25.
Describe the circumstances that are necessary
for a compound to conduct an electric current.
T/I
a.
b.
c.
d.
−182.79
22.
Name each compound.
Mg3(PO4)2
NaIO3
AlPO4
NaHCO3
Write the chemical formula for each compound.
potassium thiocyanate
yttrium chloride
iron(III) sulfide
tin(II) fluoride
Melting Point (°C)
scandium oxide
ethane
T/I
a.
b.
c.
d.
18.
20.
C
For each pair of atoms below, predict whether
the bond between the atoms will be non-polar
covalent, slightly polar covalent, polar covalent, or
mostly ionic. For each slightly polar bond or polar
covalent bond, indicate which atom will be slightly
positive and which atom will be slightly negative.
a. carbon and fluorine
b. oxygen and nitrogen
c. chlorine and chlorine
d. manganese and oxygen
Write the formula for each compound.
sulfur trioxide
carbon monoxide
diselenium dibromide
nitrogen triiodide
T/I
a.
b.
c.
d.
K/U
a.
b.
c.
d.
17.
19.
K/U
Self-Check
If you
missed
question …
1
2
3
4
5
6
7
8
9
10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25
Review
2.1 2.2 2.1 2.1 2.1 2.2 2.1 2.3 2.1 2.3 2.1 2.1 2.1 2.2 2.1 2.2 2.2 2.2 2.2 2.2 2.3 2.3 2.3 2.3 2.3
section(s)…
1.3
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