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Chapter

64 Chapter 2

Chemical BondingIn caves we can sometimes see large structures known as stalagmites (Figure 1).These formations are made of crystals of calcium carbonate, CaCO3(s), alsoknown as limestone. Calcium carbonate, as its name and formula suggest, is acompound made up of three different elements. In addition to its crystallinestructure, calcium carbonate has high melting and boiling points and dissolvesto some extent in water. Many other compounds have similar physical and chem-ical properties, such as ordinary table salt, sodium chloride. Other compounds,such as water, H2O, and carbon dioxide, CO2, have significantly different prop-erties. How can we explain these similarities and differences? The answer lies inan understanding of the special forces of attraction, or bonds, that hold atomstogether in compounds. These forces form the foundation of chemical proper-ties and reactions.

As you know, we often develop models to help us understand abstract con-cepts. You are already familiar with several models: Bohr’s model of the atom; thewater cycle; the collision model. In this chapter we will develop theories andmodels of chemical bonding (including the character of atomic bonds) toexplain the nature and behaviour of matter and to classify compounds.

In this chapter, you will be

able to

• describe the role of electrons inionic and covalent bonding;

• relate the physical and chemicalproperties of compounds to thenature of their chemical bonds;

• predict the nature of bonds bycomparing electronegativityvalues;

• use a variety of models to repre-sent the formation and structureof compounds and molecular ele-ments;

• name a variety of compoundsusing common names, classicalnames, and IUPAC chemicalnomenclature.

2

1. Why do atoms form compounds? Use examples of compounds you arefamiliar with in your explanation.

2. Is there more than one type of force present between the atoms in com-pounds and, if so, how do these forces compare in strength and otherproperties?

3. How could the forces that hold atoms together in a compound determinethe chemical properties of that compound? Again, use examples offamiliar compounds and their properties in your speculations.

Throughout this chapter, note any changes in your ideas as you learn new con-cepts and develop your skills.

Reflect Learningonyour

Chemical Bonding 65

Figure 1

Stalagmites are shown extending upwardfrom this cavern floor. As you can see, thesecalcium carbonate formations have thepotential to join floor to ceiling in a continuous pillar.

In Chapter 1 you studied various representations, or models, of thestructure of atoms. In this activity, you will build models of a variety ofmolecules. This exercise will help you to see that different chemicalshave different molecular structures.

Materials: molecular models kit

• In your molecular model kit, identify the pieces representing each ofthe following elements: chlorine, bromine, carbon, nitrogen, oxygen,iodine, and hydrogen.

• Group the different pieces according to the number of holes present.(a) To what chemical family or families do these groups of pieces

correspond on the periodic table?(b) What do the holes in the pieces represent?(c) What do the sticks represent in your molecular model kit?

• Construct as many different compounds as possible using thepieces in the kit.(d) As best you can, name the compounds you have modelled

(later in this chapter you will learn a system for naming com-pounds).

(e) How do the connections differ among your models?(f) Classify the models by dividing them into groups. Provide a

rationale for your classification.(g) Organize the models into two more classification schemes.

Provide a rationale for each classification.

Try ThisActivity

Making Models of Compounds

66 Chapter 2

2.1 Classifying CompoundsThere are many ways to classify substances. You are already familiar with severalof them. For example, you know that iron is a solid at SATP (not a liquid or agas), a metal (not a nonmetal), and an element (not a compound). Similarly,water is a liquid and a compound. In this section, we will look at ways of furtherclassifying compounds.

One category of compounds includes table salt (sodium chloride), NaCl(s),bluestone (copper(II) sulfate), CuSO4(s), and baking soda (sodium bicarbonate),NaHCO3(s). If you think about their chemical formulas, you might notice thateach one is made up of a metal joined to a nonmetal. These compounds arecalled ionic compounds.

A second category includes sulfur dioxide, SO2, carbon dioxide, CO2, andammonia, NH3. Look at their formulas. What do you notice? They are all non-metals combined with nonmetals, and are called molecular compounds.

Many ionic and molecular compounds can be found within your own home.For example, window cleaners, household bleach, antacid tablets, and milk ofmagnesia contain ionic compounds. Vegetable oil, plastics, and sugar containmolecular compounds.

We can use electrical conductivity—the ability to conduct electricity—todistinguish between ionic and molecular compounds: Ionic compounds (manyof which dissolve readily in water) form solutions that conduct electricity.However, molecular compounds (some of which dissolve in water) form solu-tions that generally do not conduct electricity. Of course, there are exceptions tothis generalization.

Substances that form solutions that conduct electricity are called elec-

trolytes. Their solutions are called electrolytic solutions.Substances that form solutions that do not conduct electricity are called

nonelectrolytes. Their solutions are called nonelectrolytic solutions (Figure 1).

Just as a group of elements can share physical and chemical properties, so too cana class of compounds. For instance, solid ionic compounds generally have veryhigh melting points and are brittle—they can often be crushed fairly easily into apowder. Molecular compounds in solid form tend to have a softer or waxy texture,and many have melting and boiling points so low that they are gases or liquids atroom temperature. If we can classify a substance as either ionic or molecular, weshould be able to predict some of its properties. Conversely, if we know some of theproperties of a substance, we can classify it as ionic or molecular.

There are several properties that we could take into account. Generally, thesimplest to investigate are physical properties, such as state at SATP, solubility inwater, and ability to conduct electricity in solution.

The purpose of this investigation is to test the generalizations you haveencountered about ionic and molecular compounds. Complete theHypothesis/Prediction, Analysis, Evaluation, and Synthesis sections of the labreport.

ionic compound: a pure substanceformed from a metal and a nonmetal

molecular compound: a pure sub-stance formed from two or more differentnonmetals

electrical conductivity: the ability of amaterial to allow electricity to flow through it

electrolyte: a substance that forms a solu-tion that conducts electricity

Figure 1

Conductivity is used to distinguish betweenaqueous solutions of ionic and molecularcompounds.

I N Q U I R Y S K I L L S

QuestioningHypothesizingPredictingPlanningConducting

RecordingAnalyzingEvaluatingCommunicating

Investigation 2.1.1

Comparing Ionic and Molecular Compounds

Chemical Bonding 67

2.1

Question

Which of the following substances are ionic, and which are molecular: sodiumnitrate, NaNO3(s); sucrose, C12H22O11(s); sodium chloride, NaCl(s); potassiumsulfate, K2SO4(s); ethanol, C2H5OH(l); sodium bicarbonate, NaHCO3(s); and cal-cium sulfate, CaSO4(s)?

Hypothesis/Prediction

(a) Answer the Question, considering the state of matter at room temperatureand the chemical formula of each substance. Provide your reasoning foreach decision.

(b) Based on your classification, predict whether each substance will dissolve inwater and whether any solutions formed will conduct electricity.

Experimental Design

The ionic or molecular nature of several compounds will be determined byapplying the tests of solubility and conductivity.

Materials

lab apron sodium nitrate, NaNO3(s)eye protection sucrose, C12H22O11(s)8 50-mL beakers sodium chloride, NaCl(s)distilled water potassium sulfate, K2SO4(s)wax pencil ethanol, C2H5OH(l)scoopula sodium bicarbonate, NaHCO3(s)medicine dropper calcium sulfate, CaSO4(s)stirring rod low-voltage conductivity apparatus

Procedure

Part 1: Solubility1. Obtain a small amount of NaCl(s). Observe and, in a table, record its state

at the ambient temperature.

2. Pour about 10 mL of distilled water into a 50-mL beaker. Add a smallquantity of the chemical to the water (Figure 2).

3. Use a stirring rod to stir the mixture. Note whether the chemical dissolves.Record your observations. Label the beaker with the name of the chemicaland put it aside for the conductivity test in Part 2.

4. Repeat steps 1 through 3 for each of the other compounds.

Part 2: Conductivity5. Obtain a small sample of distilled water in a beaker. Test the electrical conduc-

tivity of the sample. A reading of zero should be indicated by the apparatus.

6. Test the conductivity of the mixture of sodium chloride and water that wasset aside in Part 1. Record your observations and rinse the probes in dis-tilled water.

7. Repeat step 6 for each of the remaining mixtures.

8. Dispose of the mixtures in the appropriate disposal containers provided byyour teacher. Rinse the beakers.

Ethanol is flammable; do notuse near an open flame. Useonly low-voltage conductivityapparatus.

Wear eye protection and anapron, and wash hands thor-oughly at the end of theinvestigation.

liquidsolid

Figure 2

For a solid chemical, use a small quantity in ascoopula as shown. If the chemical is aliquid, use a dropper full of liquid.

68 Chapter 2

Analysis

(c) Use the Evidence you gathered in this experiment to answer the Question(classify each of the substances as ionic or molecular).

Evaluation

(d) Do you have confidence in your observations? Do you feel that they can beused to accurately classify the substances? Explain.

(e) Compare your answer from the Analysis with the answer in yourHypothesis/Prediction (question (a)). How do you account for any differences?

Synthesis

(f) Were you able to accurately predict the properties of the substances basedon your initial classification? Why or why not?

(g) What assumptions are being made in this investigation?

Electrolytes and Classification

When table salt (sodium chloride), NaCl(s), dissolves in water it forms a solutionthat conducts electricity: It is an electrolyte, and dissolving it forms an elec-trolytic solution. Sodium chloride is a typical ionic compound: Its solution con-ducts electricity. Sugar will also dissolve in water. However, the sugar solutionwill not conduct electricity, so sugar is a nonelectrolyte. It is a typical molecularcompound: Its solution with water does not conduct electricity. In Unit 3, youwill learn more about conductivity and solutions. You will also find out moreabout molecular acids: important exceptions to the rule that molecular solutionsdo not conduct electricity.

Practice

Understanding Concepts

1. What types of elements combine to form(a) an ionic compound?(b) a molecular compound?

2. Briefly describe a diagnostic test for an ionic compound, and give atheoretical explanation for that test.

Applying Inquiry Skills

3. A student hypothesizes that an unknown substance is composed ofpositive and negative ions held together by the attraction of theiropposite charges. Design an experiment that would allow the studentto test this hypothesis.

4. Use the evidence in Table 1 to classify each of the five compounds asionic or molecular. Provide your reasoning for each classification.

2.2 Ionic BondingWhen sodium (a metal) is put in a vessel containing chlorine (a nonmetal), thetwo elements combine enthusiastically to form the compound sodium chloride,

Table 1: Observations of Five Unknown Compounds

Compound State at ConductivitySATP of solution

A solid yes

B liquid no

C gas yes

D solid yes

E solid no

Chemical Bonding 69

2.2

a substance that you have classified as ionic, as it is an electrolyte. Like some otherionic compounds that you are familiar with, for example, baking soda (sodiumhydrogen carbonate) and chalk (calcium carbonate), it is also brittle and has ahigh melting temperature. How do we explain the formation and properties ofthis compound?

You will recall, from Chapter 1, that atomic theory describes electronsmoving about the nucleus of the atom in energy levels, and that the electrons inthe outermost energy level are called the valence electrons. It is the valence elec-trons of an atom that form chemical bonds.

According to atomic theory, ionic compounds are formed when one or morevalence electrons are transferred from a metal atom to a nonmetal atom. Thisleaves the metal atom as a positive ion, or cation, and the nonmetal atom as anegative ion, or anion. The two oppositely charged ions are attracted to eachother by a force called an ionic bond (Figure 1).

Explaining the Properties of Ionic Compounds

Ionic compounds have similar properties: They are solids at SATP with highmelting points, and they are electrolytes. We can hypothesize that these proper-ties might be the result of the bonds formed between the ions, holding themfirmly in a rigid structure.

Although they are composed of ions, pure ionic compounds are electricallyneutral. Therefore, the sum total of the electrical charges on all the ions must bezero. Ionic compounds are made up of a fixed proportion of positive and negativeions. Consequently, ionic compounds can only be identified in terms of thesmallest unit of the compound, known as a formula unit, that would still have theproperties of the compound. In the case of sodium chloride, the sodium and chlo-ride ions are present in a 1:1 ratio, as indicated by its chemical formula, NaCl.

The anions and cations in an ionic compound are locked in a regular struc-ture, held by the balance of attractive bonds and electrical repulsion. The mostcommon model of ions shows them as spheres arranged in a regular three-dimensional pattern called a crystal lattice (Figure 2, page 70). We can actuallysee the shape of sodium chloride crystals—an observation that supports ourcrystal lattice model. We often see similar structures in bridges or scaffoldingwhere struts are joined one to the other in a repeating pattern.

Not all crystal lattices are square, like that of sodium chloride. Depending onthe sizes and charges of the ions that make up the substance, the crystal latticevaries in structure. All lattices are arranged so that each ion has the greatest pos-sible number of oppositely charged ions close, while keeping ions with the samecharge as far away as possible. In all cases, each ion is surrounded by ions ofopposite charge. In theory, this arrangement of ions creates strong attractions.This theory is supported by empirical evidence such as the hard surfaces andhigh melting and boiling points of ionic solids.

Na Cl ClNa

+ –

Figure 1

An electron is transferred from sodium tochlorine in the formation of an ionic bond.

chemical bond: the forces of attractionholding atoms or ions together

ionic bond: the electrostatic attractionbetween positive and negative ions in a com-pound; a type of chemical bond

formula unit: the simplest whole-numberratio of atoms or ions of the elements in anionic compound

crystal lattice: a regular, ordered arrange-ment of atoms, ions, or molecules

70 Chapter 2

It requires a great deal of energy to break the strong electrostatic attractionswithin a crystal lattice. The ions resist any movement, as even a slight shift wouldcause positive ions to move closer to other positive ions, and negative ions closerto other negative ions, resulting in strong repulsion. We can use this model toexplain why ionic substances are hard (the ions resist movement), and also whythey are often brittle. Once the lattice is broken, repulsions between ions of thesame charge will cause the substance to split into two crystals.

Finally, our model of the structure of ionic compounds can also explain theelectrical conductivity of their solutions. When ionic compounds are dissolvedin water, the positive and negative ions dissociate:

NaCl(s) → Na+(aq) + Cl–(aq)

The ions are responsible for carrying current when charged electrodes areplaced in the ionic solution.

The Formation of Ionic Compounds

We classify most simple compounds containing metallic elements as ionic.Elements within a chemical family (group) tend to participate in similar chem-ical reactions, producing ionic compounds with the same general formula. Forexample, the metals in Groups 13 to 15, except mercury, will form ionic oxides(compounds composed of a metal and oxygen) when burned in air. In the sameway, elements in Groups 1 and 2 form ionic compounds with oxygen: The oxidesformed by Group 1 elements have the general formula M2O while those ofGroup 2 elements (e.g., magnesium oxide, MgO) have the general formula MO,where M represents a metal ion. Because of its high melting point, magnesiumoxide is used to make objects that are exposed to very high temperatures, such ascrucibles, furnace linings, and thermal insulation.

From our discussion of trends in electronegativity, it makes sense that Group 1metals readily react with the elements in Group 17 to form ionic compoundswith the general formula MX. These compounds, which are composed of a metaland a halogen, are collectively referred to as ionic halides. Sodium chloride, anexample of an ionic halide, is found in large underground deposits in variousparts of the world. It is mined from these deposits and used as road salt, table salt,and as a reactant in many industrial processes.

The elements in Group 2 show a similar trend, as they also react with thehalogens to produce ionic halides with the general formula MX2.

In general, the addition of a metal from Group 1 or Group 2 to water willproduce hydrogen gas and a basic ionic compound. Calcium hydroxide,Ca(OH)2, is an example of an ionic compound that can be produced in this way.However, the reaction of calcium and water is quite vigorous. A safer means ofproducing this ionic compound is to react calcium oxide with water. Calciumhydroxide is also referred to as slaked lime and is used to make mortar andplaster for buildings.

Predicting Common Ions of Atoms

Noble gases are stable and virtually inert. Similarly, ions with eight valence elec-trons appear to have a special stability. This arrangement of electrons is knownas a stable octet. To reach this stable state, metal atoms of elements in Groups 1,2, and 3 will lose electrons to form cations, while elements in Groups 15, 16, and17 will gain electrons to form anions. By looking at an element’s position in theperiodic table, we can predict the charge on that element’s most stable ion. For

Figure 2

(a) The cubic structure of table salt crystalsprovides a clue about the internal struc-ture of sodium chloride.

(b) The arrangement of sodium and chlorineions in a crystal of sodium chloride. Thesodium and chloride ions occupy positionsin the crystal lattice known as latticepoints.

stable octet: a full shell of eight electronsin the outer energy level of an atom

(a)

(b)

Chemical Bonding 71

2.2

example, we can predict that a Group 1 element will tend to lose one electron,becoming a cation with a charge of 1+. A Group 16 element, on the other hand,will tend to gain two electrons to complete a stable octet, so it will form an anionwith a charge of 2–.

Hydrogen is a special case in that, theoretically, it can either give up an elec-tron to form H+, which is equivalent to a proton, or gain an electron to form H�,which has a filled shell like the noble gas helium.

Practice


1. What properties of ionic compounds suggest that ionic bonds arestrong?

2. What types of elements form ionic bonds with each other?

3. Which of the representative elements tend to form positive ions?Which tend to form negative ions?

4. What is the minimum number of different ions in the formula of anionic compound? Explain.

5. Predict the charge on the most stable ion formed by each of the fol-lowing elements. Indicate the ion by writing the symbol completewith charge.(a) sulfur(b) barium(c) bromine(d) chlorine(e) calcium(f) potassium(g) phosphorus(h) rubidium(i) beryllium


6. You have already discovered that solutions of ionic compounds inwater conduct electricity. You might wonder about the conductivity ofpure ionic compounds.(a) Design an experiment to investigate the conductivity of an ionic

solid. With your teacher’s approval, conduct your investigation.(b) Research the conductivity of molten (liquid) ionic compounds.(c) Assemble your findings into a report on the conductivity of ionic

compounds in various states.(d) Propose a hypothesis for the properties you observe.

Ions and the Human Body

Among the 12 elements that make up more than 99% of the human body (Table 1)are five metals: calcium, potassium, sodium, magnesium, and iron. These fivemetals, which form positively charged ions in solution, are essential for main-taining good health. For example, Mg2+, Na+, and K+ are major constituents ofblood plasma. Ca2+ is vital in the formation of bones and teeth. In addition topositively charged ions, some elements that form negatively charged ions are alsoessential for life. Chloride ions, Cl�, are another component of blood, and iodideions, I�, are required to prevent a condition called goitre, which results in theenlargement of the thyroid gland.

Table 1: 12 Most Common Elements in the Human Body

Element Percentage by mass

oxygen 65

carbon 19

hydrogen 9.5

nitrogen 3.2

calcium 1.5

phosphorus 1.0

sulfur 0.3

chlorine 0.2

sodium 0.2

magnesium 0.1

iodine <0.1

iron <0.1

72 Chapter 2

Practice

Making Connections

7. Research and report upon the importance of one of the ions thatmake up the human body. Your report should include: a description ofits biological importance; recommended daily minimum require-ments; the effects of deficiency or excess on the human body; andsome of the naturally occurring sources. Conclude your report with adiscussion on whether, and under what circumstances, you wouldrecommend that someone should artificially supplement his/herintake of this ion.

Follow the links for Nelson Chemistry 11, 2.2.

Representing Ionic Bonds

We generally find it easier to grasp a new concept if we have a model or mentalimage. An American chemist, G. N. Lewis, obviously thought the same thing, anddeveloped a model of the valence (bonding) electrons of single atoms ormonatomic ions. He represented his models on paper as electron dot diagrams,or Lewis symbols. These symbols consist of the chemical symbol for the elementplus dots representing the number of valence electrons. For example, a sodiumatom would be represented as

The convention for indicating more electrons is as follows:

These diagrams can help us to represent the process of ion formation.Electron dot diagrams also illustrate the theory that ionic bonds tend to pro-

duce full outer orbits of electrons: a configuration exactly the same as that of thenoble gases. Sodium has one valence electron. By transferring this electron toanother entity that has a stronger attraction for the electron, the resultingsodium ion will have the same electron configuration as neon.

A chlorine atom has seven valence electrons. By attracting an electron fromanother entity, the resulting chloride ion will have eight electrons in its valenceshell and the same number of electrons as its nearest noble gas, argon.

Representative elements around Groups 13 and 14 of the periodic table canalso achieve this special stability without losing or gaining electrons, as you willdiscover later in the chapter.

Another way to describe the process of forming ionic compounds is to saythat electron transfers result from the large difference in electronegativity betweenmetals and nonmetals. Nonmetals have a strong tendency to gain electrons andmetals do not. When an atom of low electronegativity, such as a metal atom, andan atom of high electronegativity, such as a nonmetal atom, are in proximity, oneor more electrons are transferred from the atom with low electronegativity, trans-forming both atoms into ions. The metal ion will be positive and the nonmetalion will be negative. The resulting oppositely charged ions attract each other andother ions, forming ionic bonds and resulting in a crystal lattice.

www.science.nelson.comGO TO

electron dot diagram or Lewis

symbol: a representation of an atom orion, made up of the chemical symbol anddots indicating the number of electrons in thevalence energy level

Na

Na Na + e

OB FNCBe NeLi

Cl + e Cl Ar )(compare with

Chemical Bonding 73

2.2

Using electron dot diagrams, we can show the formation of an ionic bondbetween sodium and chlorine.

When illustrating the formation of an ionic bond, we place square bracketsaround the ion to indicate that the charge is not associated with any particularelectron and that all the electrons in the valence shell are equivalent.

We can use the periodic table and electron dot diagrams to predict the for-mulas of other ionic compounds. By finding out how many electrons they tendto lose or gain to reach stable octets, we can figure out what ratio of ions willmake an electrically neutral compound.

As an example, suppose we were asked to draw electron dot diagrams toillustrate the formation of calcium fluoride, state the ratio of ions in the com-pound, and give the formula for the compound. Calcium is in Group 2, so willform an ion with a charge of 2+. Fluorine is in Group 17, so will form an ion witha charge of 1�.

To form an electrically neutral ionic compound, the ratio of calcium to fluorideions is 1:2. The formula is CaF2.

Sample Problem

Draw electron dot diagrams to illustrate the formation of magnesium oxide.Write the ion ratio and the chemical formula.

Solution

The two elements will combine in a ratio of 1:1. The formula is MgO.

Practice


8. (a) How do the electron dot diagrams of metal ions differ from thoseof nonmetal ions?

(b) How are the electron dot diagrams of metal ions similar to thoseof nonmetal ions?

9. Use electron dot diagrams to illustrate the formation of(a) lithium iodide(b) barium chloride(c) potassium oxide(d) calcium fluoride

10. Represent each of the following elements using electron dot dia-grams:(a) nitrogen(b) sulfur(c) argon(d) iodine(e) lithium(f) cesium(g) calcium(h) sodium

Na Cl ClNa+

Ca 2 2F Ca FF

Mg 2 2MgO O

74 Chapter 2

11. Use electron dot diagrams to determine the ratio in which oxygenwill combine with each of the following elements to form an ioniccompound. Label each diagram with the chemical formula of thecompound.(a) calcium(b) rubidium(c) strontium(d) aluminum

12. Represent the five halogens, using electron dot diagrams. How arethese diagrams consistent with the concept of a chemical family?

13. Explain, referring to stable octets, how the following ionic com-pounds are formed from pairs of elements; illustrate the formation ofeach compound with electron dot diagrams; and predict the formulaof each compound.(a) magnesium chloride(b) sodium sulfide(c) aluminum oxide (bauxite)(d) barium chloride(e) calcium fluoride (fluorite)(f) sodium iodide(g) potassium chloride (a substitute for table salt)

14. Give the common names for the following chemicals:(a) sodium bicarbonate(b) NaCl(c) calcium carbonate(d) Ca(OH)2


1. Use the concepts of periodic trends and electronegativity toexplain why ionic compounds are abundant in nature.

2. Give a theoretical reason why lithium and oxygen combine in theratio 2:1.

3. How are the topics of ion formation and periodic trends related?

4. Give the correct chemical (IUPAC) names for the following chemi-cals:(a) chalk(b) slaked lime(c) road salt(d) baking soda

5. A Group 1 metal (atomic number 55) is reacted with the mostreactive of the halogens. A very vigorous reaction results in theformation of a solid, white compound.(a) Represent the formation of the compound with electron dot

diagrams.(b) Write the formula of the compound formed.(c) What type of compound is formed?(d) Predict the physical properties of the resulting white com-

pound.(e) Explain the properties of the compound in terms of the

bonds formed.(f) Provide a theoretical explanation for the vigorous reaction.

Sections 2.1–2.2 Questions

Chemical Bonding 75

2.3

2.3 Covalent BondingIonic compounds, as you have just learned, contain many ions arranged in athree-dimensional structure. But not all compounds are brittle ionic solids withhigh melting points. Some, such as paraffin wax, are solids at SATP but have rel-atively low melting points. Others are liquids or gases, such as water and carbondioxide. In what way do these substances differ from ionic compounds? Can weexplain these differences with atomic theory? The answer is yes, but first weshould look at their formulas.

We have written the formulas of ionic compounds as simplest formularatios. The compound NaCl includes sodium and chloride ions in a 1:1 ratio. Wecould build a crystal from 8 ions of each element, or 8 million of each and itwould form the same structure and would have the same properties. There isonly one ionic compound of sodium and chloride ions, and that compoundalways contains a 1:1 ratio of the two ions.

This is not true of the bonding of nonmetals with each other. For example,consider the simplest ratio formula CH. If carbon and hydrogen formed an ioniccompound with this ratio, we would expect that any structure in which these ele-ments were in this ratio would have the same properties. However, this is not thecase. Using a mass spectrometer and combustion analysis we can demonstratethat there are several compounds that have this simplest ratio formula. The gasacetylene (ethyne), C2H2(g), and the liquid benzene, C6H6(l), both have the sim-plest ratio formula CH, but they are otherwise much different in their physicaland chemical properties. Simplest ratio formulas indicate only the relative num-bers of atoms in a molecular compound; they give no information about theactual number of atoms or the arrangement of those atoms in a molecule. To dis-tinguish among molecular compounds, we need to represent them with formulasthat describe the molecules that make them up.

Molecules can be classified by the number of atoms that they contain.Molecules that contain only two atoms, such as carbon monoxide, CO, are calleddiatomic molecules. If they contain more than two atoms, such as ammonia,NH3, they are called polyatomic molecules.

Some elements also exist as molecules. Hydrogen and oxygen are examplesof elements composed of diatomic molecules. Sulfur, S8, and phosphorus, P4, arepolyatomic molecules.

Formation of Covalent Bonds

You already know that hydrogen can form a cation (H+) by losing a valence elec-tron, or it can form an anion (H–) by gaining an electron and filling its electronshell. However, two hydrogen atoms can each obtain a stable filled energy levelby sharing a pair of electrons. (Remember that the first energy level can only con-tain two electrons.) The covalent bond that results arises from the simultaneousattraction of two nuclei for a shared pair of electrons (Figure 1).

We can use the model of the Lewis symbol, or electron dot diagram, to com-municate the theory of covalent bonding. When an electron dot diagram is usedto represent covalent bonding, we adapt it slightly and call it a Lewis structure

(because it illustrates the structure of the molecule). A Lewis structure shows thevalence electrons surrounding each of the component atoms as dots, with theexception of the electrons that are shared: These shared electrons are representedby a dash. In effect, this dash represents a covalent bond.

diatomic molecule: a molecule consistingof two atoms of the same or different elements

polyatomic molecule: a molecule con-sisting of more than two atoms of the sameor different elements

covalent bond: the attractive forcebetween two atoms that results when electrons are shared by the atoms; a type of chemical bond

Lewis structure: a representation ofcovalent bonding based on Lewis symbols;shared electron pairs are shown as lines and lone pairs as dots

Figure 1

The pair of shared electrons between thenuclei of two hydrogen atoms results in acovalent bond.

Cl Cl Cl Cl

76 Chapter 2

According to the Lewis structure for a chlorine molecule, each Cl atom hasthree pairs of electrons that are not involved in the formation of a covalent bond.Each pair is referred to as a lone pair. There is one shared pair of electronsbetween the atoms.

Many elements will form bonds that result in a full valence shell for eachatom, so that each atom has the electron structure of an atom of a noble gaswhen the shared electrons are included. In other words, it has a stable octet. Thisgeneralization is referred to as the octet rule. There are many exceptions to thisrule, but you will not be learning about these in this course.

Sample Problem 1

Draw Lewis symbols for the reaction of two bromine atoms and a Lewis struc-ture for the resulting bromine molecule.

Solution

This method of representing molecules can be further simplified by notindicating lone pairs. This representation is referred to as a structural formula.The structural formula for chlorine would be Cl�Cl, and for bromine Br�Br. Asyou can see, they are easier to write, and are quite similar to the chemical for-mulas, Cl2 and Br2.

Each pair of shared electrons results in a single bond. Elements that needonly one more electron to complete their outer shell or energy level tend to formsingle bonds. Hydrogen and chlorine are typical examples.

Elements in Group 16 are two electrons short of a full outer shell. How, then,can they form covalent bonds with each other and still achieve a stable state? Theanswer is simple, though not necessarily obvious: They form a double bond, witheach pair of atoms sharing two pairs of electrons between them. Oxygen andcarbon dioxide are examples of molecules that include double bonds. The doublebond is represented, in a Lewis structure and structural formula, as a double dash.

Sample Problem 2

(a) Draw the Lewis structure for a molecule of oxygen.(b) Give the structural formula for the molecule.

Solution

(a)

(b)

Can you predict how atoms of nitrogen, which require three electrons eachto achieve a stable octet, might form N2? The two atoms form a triple bond bysharing three electron pairs.

The structural formula for a molecule of nitrogen is

The number of covalent bonds (shared electron pairs) that an atom can formis known as its bonding capacity (Table 1). Each atom of nitrogen, we have just

Br Br Br Br

lone pair: a pair of valence electrons notinvolved in bonding

octet rule: a generalization stating that,when atoms combine, the covalent bonds areformed between them in such a way thateach atom achieves eight valence electrons(two in the case of hydrogen)

structural formula: a representation ofthe number, types, and arrangement of atomsin a molecule, with dashes representingcovalent bonds

O + O O O

O O

N N N N

N Nbonding capacity: the number of elec-trons lost, gained, or shared by an atomwhen it bonds chemically

Chemical Bonding 77

2.3

learned, shares three electron pairs, so it has a bonding capacity of three. What isthe bonding capacity of oxygen? It always shares two electron pairs, so has abonding capacity of two. It is easy to find the bonding capacity of any element bylooking at the Lewis structure of a molecule containing that element: The numberof dashes associated with the element is the same as the bonding capacity.

So far, we have looked at molecules of elements. Because the atoms are thesame, each has the same bonding capacity and each contributes the samenumber of electrons to the covalent bond. The molecules of compounds, how-ever, consist of atoms of two or more different elements, often with differentbonding capacities. How do we decide on their arrangement, when we drawstructural formulas? The central position in the arrangement is often occupiedby the element with the highest bonding capacity. Carbon and nitrogen, forinstance, are commonly at the centre of a structural formula. Electronegativity isanother means of deciding upon the central atom. When there is a choice ofatoms for the central position in the molecule, choose the least electronegativeelement. Hydrogen is never the central atom since it can only form a single cova-lent bond. Halides and oxygen are also usually not the central atom.

There are exceptions to these generalizations, but they meet our needs inmost cases.

Sample Problem 3

(a) Draw a Lewis structure for a molecule of carbon dioxide.(b) Give the structural formula for the molecule.

Solution

(a)

(b)

Practice


1. Draw a Lewis structure and write the molecular formula for each ofthe following:(a) F2(g)(b) H2O(l)(c) CH4(g)(d) PCl3(s)(e) H2S(g)(f) SiO2(s)

Table 1: Bonding Capacities of Some Common Atoms

Atom Number of Number of Bonding capacityvalence electrons bonding electrons

carbon 4 4 4

nitrogen 5 3 3

oxygen 6 2 2

halogens 7 1 1

hydrogen 1 1 1

O + +C O O C O

O C O

78 Chapter 2

Coordinate Covalent Bonds

Many substances contain a combination of covalent and ionic bonding. Considerthe compound ammonium chloride, NH4Cl. This white, crystalline solid dis-solves rapidly in water and is an electrolyte—it dissociates to form a cation andan anion. It has many of the properties of an ionic compound, but it is composedonly of nonmetals. We explain the properties of ammonium chloride bydescribing it as an ionic compound composed of a chloride ion, Cl–, and a poly-

atomic ion, ammonium, NH4+. The bond holding the chloride and ammonium

ions together is ionic, but the bonds within the polyatomic ammonium ion arecovalent. There are several polyatomic ions, including NH4

+, SO42– (sulphate),

and CO32–(carbonate), all of which are covalently bonded groups of atoms car-

rying an overall charge.How does this arrangement fit with our description of covalent bonds?

Molecules that are composed of two or more different elements can sometimesform covalent bonds where both of the electrons making up the bond are pro-vided by the same atom. This type of bond is called a coordinate covalent bond.Consider the formation of the ammonium ion from the regular covalent mole-cule ammonia, NH3, and a hydrogen ion, H+. The hydrogen ion does not bringany electrons with it. To achieve a complete outer shell it can borrow two elec-trons from the atom with which it bonds.

To explain this bond, we can draw Lewis structures. First, we must establishthe Lewis structure for ammonia. We can arrange the atoms with nitrogen(which has the highest bonding capacity) at the centre, showing the five valenceelectrons of nitrogen and the one valence electron of each of the three hydrogenatoms.

We can show the pairs of shared electrons (covalent bonds) between adjacentnitrogen and hydrogen atoms as dashes.

Notice the lone pair of electrons in this structure. A hydrogen ion, which hasno electrons of its own, can bond to the ammonia molecule by sharing thisunbonded pair of electrons. This is the coordinate covalent bond.

Once the hydrogen ion is bonded, there is no way to tell which of the hydro-gens was the ion. Each of the four hydrogen atoms in the structure are equiva-lent: The positive charge is not really associated with any particular hydrogenatom. To indicate this, square brackets are placed around the entire ammoniumion and the positive charge is written outside the bracket.

polyatomic ion: a covalently bondedgroup of atoms with an overall charge

coordinate covalent bond: a covalentbond in which both of the shared electronscome from the same atom

NH H

H

NH H

H

H H

H HN H N H

HH

Chemical Bonding 79

2.3

The Strength of Covalent Bonds

Covalent bonds are strong. A large amount of energy is needed to separate theatoms that make up molecules. For this reason, molecules tend to be stable at rel-atively high temperatures: They do not easily decompose upon heating.

The stronger the bonds within the molecule, the greater the energy requiredto separate them. The strength of a bond between two atoms increases as thenumber of electron pairs in the bond increases. Therefore, triple bonds arestronger than double bonds, which are stronger than single bonds between thesame two atoms.

Drawing Lewis Structures and StructuralFormulas for Molecular Compounds

1. Arrange the symbols of the elements of the compound as you wouldexpect the atoms to be arranged in the compound. The element with thehighest bonding capacity is generally written in the central position(Figure 2(a)).

2. Add up the number of valence electrons available in each of the atoms(Figure 2(b)). If the structure is a polyatomic ion, add one electron foreach unit of negative charge, or subtract one for each unit of positivecharge.

3. Place one pair of electrons between each adjacent pair of elements(forming single covalent bonds) (Figure 2(c)).

4. Place pairs of the remaining valence electrons as lone pairs on the periph-eral atoms (not the central atom) (Figure 2(d)).

5. If octets are not complete, move lone pairs into bonding position betweenthose atoms and the central atom until all octets are complete(Figure 2(e)).

6. If the peripheral atoms all have complete octets and there are pairs of elec-trons remaining, place these electrons as lone pairs on the central atom.

7. Count the number of bonds between the central atom and the peripheralatoms. If this number exceeds the bonding capacity of the central atom,one or more of the bonds is coordinate covalent. To identify which ones,try removing the peripheral atoms one at a time. If you can do this andleave the central structure with complete octets, you have identified coordi-nate covalent bonds (Figure 2(f)).

8. To give the structural formula, remove the dots representing the lone pairsand replace bond dots with dashes (Figure 2(g)).

9. If you are representing a polyatomic ion, place brackets around the entirestructure and write the charge outside the brackets.

Practice


2. Draw Lewis structures and structural formulas for each of the fol-lowing molecules:(a) H2(g) (d) NF3(g)(b) O3(g) (e) N2H2(g)(c) OF2(g) (f) P2H4(g)

SUMMARY

SO

O

O

Figure 2

SO

O

O

3(6) + 6 = 24SO

SO

O

O

SO

O

O

Sulfur atom has an incomplete octet.

SO

O

OS

O

O

O

S

O

O

SO3 includes two coordinate covalent bonds.

O

(g)

(f)

(e)

(d)

(c)

(b)

(a)

80 Chapter 2

3. Draw Lewis structures and structural formulas for each of the fol-lowing polyatomic ions:(a) PO 3–

4(aq) (c) BrO –3(aq)

(b) OH–(aq) (d) ClO –

4(aq)

4. Which of the Lewis structures in questions 2 and 3 include coordinatecovalent bonds?

Explaining the Properties of Molecular Compounds

You are familiar with many molecular compounds: propane, C3H8(g), for the bar-becue, water, H2O(l), in your bathtub, and sugar, C12H22O11(s), for your coffee. Asyou can see, their physical properties vary greatly. In contrast to ionic com-pounds, which are all solids at SATP, molecular compounds may be solids, liq-uids, or gases (Table 2).

We have discussed the forces that bond atoms and ions together within acompound, the intramolecular forces (“intra” means within). These are suffi-cient to explain the existence of ionic and molecular compounds, and to explainmany of the properties of ionic compounds, but they aren’t sufficient to explainthe physical state of molecular compounds. Why is water a liquid at SATP and asolid at STP? Why isn’t it a gas at all temperatures? Something, some force, musthold the molecules together in the solid and liquid states.

Forces between molecules are called intermolecular forces (“inter” meansbetween). The evidence indicates that these forces are strong enough to causemolecules to arrange themselves in an orderly fashion to form a lattice structure(similar to that of ionic solids). Both solid water and solid carbon dioxide (dryice) show such structures when examined by crystallography (Figure 3).However, intermolecular forces must be weak compared to covalent bonds. Wecan deduce this from the observation that it is much easier to melt a molecularsolid than it is to cause the same substance to decompose. When water is heatedfrom –4°C to 104°C it changes state from a solid, to a liquid, and then to a gas,but it does not decompose to oxygen and hydrogen. The energy added in theform of heat is sufficient to overcome the intermolecular forces between the mol-ecules, but not the covalent bonds between the atoms. Adding a relatively smallamount of heat will cause a solid molecular compound to change state from asolid to a liquid, and then to a gas, but it takes much more energy to break thecovalent bonds between the atoms in the compound.

Later in this chapter, you will learn more about intermolecular forces andhow they affect the properties of substances.

Table 2: Comparison of Ionic and Molecular Solids

Properties Ionic Molecular

melting point high low

electrical conductivity in the solid state no conductivity no conductivityin the liquid state conductivity no conductivity

consistency of solid hard, brittle soft, waxy or flexible

Examples sodium chloride iodinecopper(II) sulfate phosphorus

intramolecular force: the attractiveforce between atoms and ions within a com-pound

intermolecular force: the attractiveforce between molecules

Chemical Bonding 81

2.3

Practice


5. Distinguish between bonding electrons and lone pairs.

6. Are the following pairs of atoms more likely to form ionic or covalentbonds?(a) sulfur and oxygen(b) iodine and iodine(c) calcium and chlorine(d) potassium and bromine

7. (a) List six examples of molecular elements and compounds and sixexamples of ionic compounds.

(b) Compare the two lists, referring to bond types to explain the con-trasting physical properties.

8. (a) How does the bonding capacity of nitrogen differ from that ofchlorine?

(b) Give a theoretical explanation for your answer to (a).

9. How are coordinate covalent bonds similar to covalent bonds? Howare they different?

10. (a) Use an electron dot diagram to explain the formula for nitrogen, N2.(b) Draw the Lewis structure for nitrogen.(c) Nitrogen is a fairly inert (unreactive) gas. Explain this, referring to

the bonds involved.

11. Illustrate the formation of each of the following molecular com-pounds, using electron dot diagrams and Lewis structures:(a) HCl (c) H2S(b) NH3 (d) CO2

12. Use the octet rule to develop a table that lists the bonding capacitiesfor carbon, nitrogen, oxygen, hydrogen, fluorine, chlorine, bromine,and iodine.

13. Use bonding capacities (Table 1, page 77) to draw the structural for-mula of each of the following molecules.(a) O2(b) Br2(c) H2O2(d) C2H4(e) HCN(f) C2H5OH(g) CH3OCH3(h) CH3NH2

14. (a) Illustrate the structure of a hydronium ion (H3O+) by drawing its Lewis structure.

(b) Name the bonds within the hydronium ion.(c) What kinds of bonds is this ion likely to form with other entities?

15. Is it correct for the structural formula of H2S to be written as H—H—S?Explain, using a diagram.

Figure 3

The molecules in solid carbon dioxide, CO2,form a crystal lattice.

82 Chapter 2

2.4 Electronegativity, Polar Bonds,and Polar Molecules

Why are some molecular substances solid, some liquid, and some gaseous atSATP? Why do different liquids have different boiling points? Why can waterstriders walk across the surface of a pond without falling in? Believe it or not, allthese phenomena depend on the bonds and forces between molecules.

Polar Covalent Bonds

So far, we have discussed models for two types of chemical bonding: ionic andcovalent. However, when a chemical bond is formed, it is not always exclusivelyone or the other.

When electrons are shared between two atoms, a covalent bond results.When the atoms are identical, such as in a chlorine molecule, the electrons areshared equally (Figure 1(a)). However, this is not the case for a compound likehydrogen chloride, where electrons are shared between two different elements. In


1. Compound A is formed when the element with atomic number 3combines with the element of atomic number 9. Compound B isformed when the element with atomic number 7 combines withthe element of atomic number 9.(a) Compare the properties of compounds A and B.(b) What types of compounds are A and B? Give reasons for

your answer.(c) Clearly show the structure of each compound formed, using

electron dot diagrams and Lewis structures.

2. Oxygen forms ionic bonds with aluminum to form bauxite (oraluminum oxide). However, oxygen forms covalent bonds withcarbon to form carbon dioxide. Use the concepts of electronega-tivity and periodic trends to explain these differences in bonding.

3. An alkali metal M reacts with a halogen X to form a compoundwith the formula MX.(a) Would this compound have ionic or covalent bonds? Explain.(b) Predict the physical properties of the compound MX.

4. The compound NaCl and the element Cl2 are both held togetherby chemical bonds.(a) Classify each substance as either an ionic or a covalent com-

pound.(b) Explain, in terms of chemical bonds and attractive forces,

why NaCl is a solid at SATP while Cl2 is a gas at SATP.

5. Using only oxygen and sulfur atoms, create as many compoundsas you can and draw Lewis diagrams for each. Which of the com-pounds contain only multiple bonds?


6. Design an experiment to test your classification of compounds Aand B in question 1.

7. Design an experiment to test your predictions of the physicalproperties of the compound MX in question 3.

Section 2.3 Questions

Figure 1

Electron densities of the bonding electrons intwo covalent molecules(a) Cl2(b) HCl

(a)

(b)

Chemical Bonding 83

2.4

this situation, the sharing is unequal, as the bonding electrons spend more timenear one atom than near the other (Figure 1(b)). The electrons in the H–Clbond in a hydrogen chloride molecule spend more time near the chlorine atomthan near the hydrogen atom. This is because of chlorine’s greater attraction forelectrons. Due to this unequal sharing of electrons, the hydrogen atom is, onaverage, slightly positively charged while the chlorine atom is slightly negativelycharged (Figure 2). We indicate these slight charges by d� and d�. (The Greekletter delta, d, indicates “a small difference.”) While chlorine shows the greaterattraction for the bonding electrons, the attraction is not strong enough to actu-ally bring about an electron transfer, as in an ionic compound. The bond issomewhere between an ionic bond and a covalent bond and is called a polar

covalent bond.We can predict which parts of the molecule will have d� and d� charges by

comparing the electronegativities of the atoms (given in the periodic table at theend of this text). The most electronegative atoms will attract the electron pairstrongly, and so will tend to have a d� charge in the covalent compounds theyform.

Sample Problem

(a) Draw the structural formula for methane (CH4, the main constituentof natural gas).

(b) Indicate which atoms are slightly positive and which are slightly nega-tive.

Solution

(a)

(b)

Each hydrogen atom carries a partial positive charge while the carbon atom car-ries a partial negative charge.

To predict whether a chemical bond between two atoms will be ionic, polarcovalent, or covalent, we must consider the electronegativities of the elementsinvolved. The absolute value of the difference in electronegativities of twobonded atoms provides a measure of the polarity in the bond: the greater the dif-ference, the more polar the bond (Table 1). According to the periodic trends inelectronegativity described in Chapter 1, fluorine atoms have the greatest ten-dency to gain or attract electrons and cesium atoms have the least. The differencebetween their electronegativities is 3.3 (cesium 0.7 and fluorine 4.0). Since thedifference in electronegativity for cesium and fluorine is large, they should (anddo) form an ionic bond. In other words, the fluorine atom is strong enough, rel-ative to the cesium atom, to “steal” a cesium electron.

�+ �–

Figure 2

Slight positive and negative charges are indi-cated by the Greek symbol d, delta.

polar covalent bond: a covalent bondformed between atoms with significantly dif-ferent electronegativities; a bond with someionic characteristics

H C

H

H

H

H�+

H�+

H�+ C�- H�+

Table 1: Electronegativity Differences of a Selection of Bonds

Bond Elecronegativity difference

Cl—Cl 0.0

C—Cl 0.3

C—H 0.4

Be—H 0.6

N—H 0.9

C—O 1.0

O—H 1.4

Mg—Cl 1.8

Na—Cl 2.1

Mg—O 2.3

Ba—O 2.6

Ca—F 3.0

84 Chapter 2

The difference in electronegativities of hydrogen (2.1) and chlorine (3.0) ismuch less. They instead form a polar covalent bond. The greater the differencein electronegativity, the more polar, and then more ionic, the bond becomes.This relationship describes a continuum, a range from covalent to ionic, ratherthan three different kinds of bonds (Figure 3). How do we know when to classifya chemical bond as fully ionic? By convention, a difference in electronegativitygreater than 1.7 indicates an ionic bond. However, there are many binary com-pounds in which the two atoms have an electronegativity difference of less than1.7, but the compound still has ionic properties. For example, MgI2 (electroneg-ativity difference 1.3) is an electrolyte, i.e., is an ionic compound.

Practice


1. Describe how we can use electronegativity values to predict the typesof bonds that will form within a compound.

2. Which type of bond—ionic or covalent—will form between each ofthe following pairs of atoms? Of the covalent bonds, indicate whichwould be the most polar.(a) H and Cl(b) Si and O(c) Mg and Cl(d) Li and O(e) N and O(f) O and O(g) I and Cl(h) Cr and O(i) C and Cl

3. Identify the more polar bond in each of the following pairs.(a) H—F; H—Cl(b) N—O; C—O(c) S—H; O—H(d) P—Cl; S—Cl(e) C—H; N—H(f) S—O; P—O(g) C—N; N—N

4. Draw Lewis structures for the following substances. In each case, useappropriate notation to indicate the atoms that are slightly positiveand those that are slightly negative.(a) H2O(b) Br2(c) HBr(d) PCl3(e) OF2

ionic covalentpolar covalent

1.7 03.3Electronegativity Difference

Figure 3

In this model of a bonding continuum, thecolour change from right to left indicates theincreasing difference in electronegativity andbond polarity.

Chemical Bonding 85

2.4

Bonding Characteristics

Polar Molecules

Molecules of water, ammonia (a reactant in the production of nitrogen fertil-izers), and sulfur dioxide (an industrial pollutant contributing to acid rain for-mation) all have polar covalent bonds holding their atoms together. If a moleculecontains polar covalent bonds, the entire molecule may have a positive end anda negative end, in which case it would be classified as a polar molecule. Not allmolecules containing polar covalent bonds are polar molecules, however.Carbon tetrachloride, CCl4(l), contains four polar covalent bonds and HCl(g) hasone. However, HCl(g) is a polar molecule (Figure 4(a)), while CCl4(l) is not(Figure 4(b)).

With practice you will be able to predict which molecules are likely to bepolar by looking at their chemical or structural formulas. Table 3 shows whichtypes of compounds tend to be polar molecules.

Practice


5. Some molecules contain polar covalent bonds but are not themselvespolar. Explain, with diagrams, how this is possible.

6. Use Lewis structures and electronegativity values to explain whymethane, CH4(g), is not a polar substance.

SUMMARY

polar molecule: a molecule that isslightly positively charged at one end andslightly negatively charged at the otherbecause of electronegativity differences

�–

�+

�–

�–

�–�+

�–

Figure 4

(a) A molecule of HCl acquires oppositelycharged ends because of the polar cova-lent nature of the H-Cl bond.

(b) Since a molecule of CCl4 is quite symmet-rical, it lacks oppositely charged ends andis not polar.

Table 2: Summary of Bonding

Intramolecular force Bonding model

ionic bond • involves an electron transfer, resulting in the formation of cations and anions

• cations and anions attract each other

polar covalent bond • involves unequal sharing of pairs of electrons by atoms of two different elements

• bonds can involve 1, 2, or 3 pairs of electrons, i.e., single (weakest), double, or triple (strongest) bonds

covalent bond • involves equal sharing of pairs of electrons• bonds can involve 1, 2, or 3 pairs of electrons, i.e., single (weakest),

double, or triple (strongest) bonds

(a)

(b)

Table 3: Guidelines for Predicting Polar and Nonpolar Molecules

Type Description Examples

Polar AB diatomic compounds CO(g)HAx any molecule with a single H HCl(g)AxOH any molecule with an OH at one end C2H5OH(l)OxAy any molecule with an O at one end H20(l), OCl2(g)NxAy any molecule with an N at one end NH3(g), NF3(g)

Nonpolar Ax all elements Cl2(g), N2(g)CxAy most carbon compounds (including organic CO2(g), CH4(g)

solvents, fats, and oils)

86 Chapter 2

Intermolecular Forces

As mentioned earlier, the properties of molecular compounds cannot beexplained simply by covalent bonds. If covalent bonds were the only forces atwork, most molecular compounds would be gases, as there would be no attrac-tion between the molecules strong enough to order the molecules into solids orliquids. The concepts of the polar molecule and small charges on atoms helpexplain why these molecular compounds are not all gases at SATP.

The existence of intermolecular forces was first suggested by Johannes vander Waals, toward the end of the 19th century, to explain why gases liquefy whenthey are cooled. Many different observations, such as surface tension and heat ofvaporization, provide evidence that there are three kinds of intermolecularforces, each with different strengths. Two of these are classified as van der Waals

forces (in honour of Johannes van der Waals): dipole–dipole forces and London

dispersion forces. The third intermolecular force—hydrogen bonding—is gen-erally not grouped with the other two.

Dipole–dipole forces are the forces of attraction between oppositely chargedends of polar molecules (e.g., HCl). The positive end of each molecule attractsthe negative ends of neighbouring molecules—rather like a weak version of ionicbonds. London dispersion forces, by contrast, exist between all molecules—bothpolar and nonpolar—so are the only intermolecular forces acting between non-polar molecules. Chemists believe that the weak attractive forces are the result oftemporary displacements of the electron “cloud” around the atoms in a mole-cule, resulting in extremely short-lived dipoles. Because the dipoles last for onlytiny fractions of a second, the attraction is continually being lost, so the forces arevery weak. Dipole–dipole forces, when they exist, tend to be much stronger thanLondon dispersion forces. You will learn more about these forces later in thiscourse.

Hydrogen Bonds

Water, a polar molecule, consists of one atom of oxygen bound by single cova-lent bonds to two hydrogen atoms. Its structure is simple, but water exhibitssome rather unusual properties: higher than expected melting and boilingpoints, high vapour pressure, high surface tension, and the ability to dissolve alarge number of substances. To explain these properties, we must consider theintermolecular forces that exist between water molecules. As a result of the largedifference in electronegativity between hydrogen and oxygen, the O–H bonds ina molecule of water are highly polar covalent. The oxygen atom in a molecule ofwater carries a slight negative charge while the hydrogen atoms carry a small pos-itive charge (Figure 5). As a result, the hydrogen atoms of one water moleculeexert a strong force of attraction on the oxygen atom of neighbouring water mol-ecules. This kind of intermolecular force is referred to as a hydrogen bond.Hydrogen bonds occur among highly polar molecules containing F–H, O–H,and N–H bonds. Although a hydrogen bond is similar to a dipole–dipole force,it is stronger than any of the van der Waals forces.

Water molecules have a tendency to “stick together” because of hydrogenbonding. This is one of the reasons why water appears to “climb” the sides of anarrow tube (Figure 6). We can also explain water’s high melting and boilingpoints using the concept of the hydrogen bond: Large amounts of energy arerequired to break the hydrogen bonds in the solid and liquid states.

van der Waals forces: weak intermolec-ular attractions, including London dispersionforces and dipole–dipole forces

dipole–dipole force: an attractive forceacting between polar molecules

London dispersion force: an attractiveforce acting between all molecules, includingnonpolar molecules

hydrogen bond: a relatively strongdipole–dipole force between a positivehydrogen atom of one molecule and a highlyelectronegative atom (F, O, or N) in anothermolecule

�–

�+�+

Figure 5

(a) The oxygen atom in a water molecule hastwo single polar covalent bonds and twolone pairs of electrons. This results in aV-shaped molecule.

(b) A water molecule has a slight negativecharge on the oxygen atom and slightpositive charges on the hydrogens.

H

O

H

(b)

(a)

Chemical Bonding 87

2.4

Most substances are more dense in the solid state than in the liquid state.However, you have probably seen solid ice cubes floating in a glass of water. Ifthey float, they must be less dense than the liquid. Hydrogen bonds enable us toexplain this unusual behaviour: ice is less dense than liquid water because itforms an open lattice structure when it freezes, with a great deal of empty spacebetween the molecules (Figure 7). This lower density of water in the solid stateenables aquatic life to survive the winter. Instead of freezing from the bottom up,lakes freeze from the top down, creating an insulating layer of ice at the top. Thisis one of the many reasons why scientists consider water to be indispensable forthe existence of life.

Molecular models provide a way of representing molecules in three dimen-sions. In this activity, you will practise various ways of modelling molecules.

Materials: molecular model kits(a) Draw Lewis structures to represent the following molecules:

water, hydrogen sulfide, hydrogen fluoride, methane, carbondioxide, ammonia, and nitrogen.

(b) Write the structural formula for each compound.

• Assemble models to represent each of the substances.

(c) Which of the molecules required double or triple covalentbonds? Explain why.

(d) Which of the molecules would you expect to be polar?(e) Give examples of other compounds that you would expect to

have similar shapes to water, ammonia, and hydrogen fluoride.Explain your answers using diagrams and models.

Try ThisActivity

Molecular Models

Figure 6

The curved surface of water in a container iscalled the meniscus.

Figure 7

In ice, hydrogen bonds between the mole-cules result in a regular hexagonal crystalstructure. The • • • H– represents a hydrogen(proton) being shared between two pairs ofelectrons. In liquid water these bonds con-stantly break and reform as molecules movepast each other.

H H

H

HH

H

H H

H

H H

O

O

O

H

O

O

O

88 Chapter 2

Practice


7. Water is known for its many anomalous properties. Use your knowl-edge of intermolecular forces and intramolecular bonding to explaintheoretically why lakes freeze from top to bottom.

8. Using Table 3 (page 85), predict whether each of the following mole-cules would be polar or nonpolar.(a) CH3OH(l) (d) PCl3(s)(b) I2(s) (e) HC2H3O2(aq)(c) HBr(g) (f) CCl4(l)

9. Which compound(s) in question 8 were classified as nonpolar butcontain polar covalent bonds? Explain how this is possible.

10. How are hydrogen bonds different from other dipole–dipole forces?

11. Predict the type of intermolecular force that would exist betweenmolecules of the following:(a) I2(s)(b) H2O(l)(c) NH3(g)

12. Explain each of your predictions in question 11 theoretically.


13. A student was provided with seven sample liquids and asked toinvestigate whether each liquid is affected by a positively or nega-tively charged object. Complete the Prediction, Evidence, Analysis,and Synthesis sections of the investigation report.

Question

Does a charged object have an effect on a thin stream of the fol-lowing liquids: NCl3, H2O, Br2, CCl4, CH3OH, H2O2, and vegetable oil?

Prediction

(a) Using the generalizations of polar and nonpolar molecules, pre-dict an answer to the Question.

Experimental Design

A thin stream of each liquid was tested by holding a positivelycharged object near the liquid stream. The procedure was repeatedwith a negatively charged object.

Evidence

(b) Which of the seven substances could be samples 1 to 3? Whichcould be 4 to 7?

Analysis

(c) Use the Evidence to answer the Question.

Synthesis

(d) Provide a theoretical explanation to justify your answer in (b).(e) Speculate as to why the liquids were affected by both positive

and negative charges.

Table 4: Effects of Charged Objects on Seven Liquids

Samples Positive charge Negative charge

1 - 3 no effect no effect

4 - 7 stream moved toward charged object stream moved toward charged object

Chemical Bonding 89

2.5

2.5 The Names and Formulas ofCompounds

Prior to the late 1700s, there was no systematic method of naming compounds.Substances were named in a variety of ways. In some cases, the name referred tothe use of the compound; in other cases, it incorporated an obvious property, orperhaps referred to the sources of the substance. These common names gavelittle, if any, information about the composition of the compound.

The classical system of nomenclature based on the Latin names of elementswas devised in 1787 by French chemist Guyton de Morveau (1737–1816). Asmore and more compounds were discovered, however, it became apparent that amore comprehensive and consistent system of nomenclature was needed.

Today, we are aware of a very large number of compounds, and there is everypossibility that more will be discovered in the future. We need a system that iseasy to use and provides information on the composition of every compound. Asa result, scientists now use the system of chemical nomenclature chosen by theInternational Union of Pure and Applied Chemistry (IUPAC) (sometimes justcalled the IUPAC system) when naming compounds. Just as many immigrants toCanada have learned a new language, you will soon become familiar with the lan-guage of chemical nomenclature. Chemical nomenclature provides us with a sys-tematic means of both naming and identifying compounds.

In describing a chemical compound, chemists may use its name or its for-mula. The formula of an ionic compound tells us the ratio of elements present.The formula of a molecular compound indicates the number of atoms of eachelement present in a molecule of the compound. Every chemical compound hasa unique name. Knowing the names of common chemicals and being able towrite their formulas correctly are useful skills in chemistry. The common andIUPAC names for some familiar compounds are shown in Table 1.


1. Explain why the bonding in compounds cannot be described onlyin terms of covalent bonds and ionic bonds.

2. How does the sharing of the bonding electrons in a molecule ofCl2 compare to that in a molecule of HCl? Explain your answer,including diagrams.

3. Both boron and nitrogen form compounds with chlorine. In eachcase, the formula has the general form AX3.(a) Classify each of these compounds as ionic or molecular.

Justify your answer.(b) How are the B—Cl bonds and N—Cl bonds similar? How are

they different?(c) What other properties can you predict for the two com-

pounds? Use the concepts of electronegativity and Lewisstructures to justify your answers.

4. Carbon tetrachloride, CCl4(l), is a nonpolar substance, althoughcarbon and chlorine have different electronegativities. Explainthis lack of polarity.


chemical nomenclature: a system,such as the one approved by IUPAC, ofnames used in chemistry

Salt of the Gods

Ammonium chloride was at one time referredto as sal ammoniac. The Greeks built atemple dedicated to the God Ammon afterparts of Egypt had been conquered byAlexander the Great (356–323 B.C.). Driedcamel manure was used as fuel for fires inthis temple. Over many years of manureburning, a white, crystalline, saltlike materialwas gradually deposited on the walls of thetemple. This deposit became known as “saltof Ammon” or sal ammoniac.

DID YOU KNOW ?

Table 1: Common and IUPAC Names for Some Compounds

Common name IUPAC name

quicklime calcium oxide

laughing gas dinitrogen monoxide

saltpetre sodium nitrate

potash potassium carbonate

muriatic acid hydrochloric acid

baking soda sodium hydrogen carbonate(sodium bicarbonate)

cream of tartar potassium hydrogen tartrate

grain alcohol ethanol or ethyl alcohol

sal ammoniac ammonium chloride

90 Chapter 2

Binary Ionic Compounds

The simplest compounds are called binary compounds. Binary ionic compoundsconsist of two types of monatomic ions (ions consisting of one charged atom).In the formula of a binary ionic compound, the metal cation is always writtenfirst, followed by the nonmetal anion. (This reflects the periodic table: metals tothe left, nonmetals to the right.) The name of the metal is stated in full and thename of the nonmetal ion has an -ide suffix; for example, NaCl(s) is sodium chlo-ride and LiBr(s) is lithium bromide. Binary ionic compounds can be made up ofmore than two ions, providing they are of only two kinds: aluminum oxide isAl2O3(s).

If we know what ions make up a compound, we can often predict the com-pound’s formula. First, we determine the charges on each type of ion making upthe compound. The charge on an ion is sometimes called the valence. We thenbalance the charges to determine the simplest ratio in which they combine. Wecan predict the charge on the most common ion formed by each representativeelement by counting the number of electrons that would have to be gained or lostto obtain a stable octet. (Ionic charges for the most common ions of elements areshown in the periodic table at the back of this text.) For example, the compoundmagnesium chloride, a food additive used to control the colour of canned peas,is composed of the ions Mg2+ and Cl–. For a net charge of zero, the ratio of mag-nesium to chlorine ions must be 1:2. The formula for magnesium chloride istherefore MgCl2(s).

Sample Problem 1

Predict the formula for magnesium oxide—a source of dietary magnesium whenused as a food additive.

Solution

Charge on magnesium ion: 2+Charge on oxygen ion: 2–The ratio of magnesium ions to oxide ions that produces a net charge of zero is 1:1.The formula is therefore MgO.

You may be familiar with a method known as the crisscross rule for pre-dicting the formula of an ionic compound. The crisscross rule works as follows:

1. Write the symbol of each of the elements in the order in which they appearin the name of the compound.

2. Write the valence number (electrons lost or gained in forming that ele-ment’s most stable ion) above the symbol of each of the elements.

3. Crisscross the numbers written above the symbols such that the valencenumber of one element becomes a subscript on the other.

4. Divide each subscript by the highest common factor. The resulting sub-scripts indicate the ratio of ions present in the compound.

5. Omit any subscript equal to 1 from the formula.

For example, if we were to use the crisscross rule to predict the formula ofmagnesium chloride, we would go through these steps.

1. Write symbols in order:

Mg Cl

binary compound: a compound com-posed of two kinds of atoms or two kinds ofmonatomic ions

valence: the charge on an ion

Chemical Bonding 91

2.5

2. Write valences above symbols:

2 1

Mg Cl

3. Crisscross the valences, making them subscripts:

Mg1 Cl24. Divide subscripts by highest common factor (= 1):

Mg1 Cl25. Remove 1 subscripts:

MgCl2

Sample Problem 2

Use the crisscross rule to predict the formula of barium sulfide.

Solution

2 2

Ba S

Ba2 S2The formula of barium sulfide is BaS.

Most transition metals and some representative metals can form more thanone kind of ion. Metals that can have more than one valence, or charge, are clas-sified as multivalent. For example, iron can form an Fe2+ ion or an Fe3+ ion,although Fe3+ is more common. The periodic table at the end of this text showsthe most common ion of each element first, with one alternative ion chargebelow. It does not list all of the possible ions of the element.

Copper is a multivalent element. It is capable of bonding with chlorine intwo different ratios to form two different chloride compounds: CuCl and CuCl2.How are the names of these compounds different? The IUPAC system of namingcompounds containing multivalent ions is very simple. The name of the metalion includes the charge on the ion, indicated by Roman numerals in brackets.Consequently, CuCl(s) (in which copper has a charge of 1+) is copper(I) chloride,and CuCl2(s) (in which copper has a charge of 2+) is copper(II) chloride. Thissystem of naming is sometimes referred to as the Stock system.

To determine the chemical name of a compound containing a multivalent metalion, we have to figure out the necessary charge on that ion to yield a net charge of zero.If we are given the formula, we simply have to calculate the equivalent negative charge.The metal’s charge is then written, in Roman numerals, after the name of the metal.For example, if you were asked to name the compound MnO2, using IUPAC nomen-clature, you would first look at the charge on the nonmetal ions. In this case, the chargeon each O is 2–, so the total negative charge is 4 –. The charge on the Mn ion must be4+. Consequently, the IUPAC name for MnO2 is manganese(IV) oxide.

If the ion of a multivalent metal is not specified in a name, it is assumed thatthe charge on the ion is the most common one.

multivalent: the property of having morethan one possible valence

92 Chapter 2

Sample Problem 3

The formula of a compound is found to be SnCl2. What is its IUPAC name?

Solution

Charge on each Cl ion: 1–Total negative charge is 2–Charge on Sn ion: 2+The IUPAC name for SnCl2 is tin(II) chloride.

Sample Problem 4

The formula of a compound is found to be Fe2O3. What is its IUPAC name?

Solution

Total negative charge: 6–Total positive charge: 6+Charge on each Fe ion: 3+The IUPAC name of Fe2O3 is iron(III) oxide.

The classical nomenclature system has, in the past, been used for namingcompounds containing multivalent metals with no more than two possiblecharges. In this system, the Latin name for the element along with the suffix -icwas applied to the larger charge, and the suffix -ous was applied to the smallercharge. The compounds formed by copper and chlorine were therefore known ascuprous chloride (CuCl(s)) and cupric chloride (CuCl2(s)).

In many industries, the classical system is still used extensively. Table 2 showsa comparison of the classical and IUPAC names of multivalent metal ions.

Table 2: Classical and IUPAC Names of Common Multivalent Metal Ions

Metal Ion Classical name IUPAC name

iron Fe2+ ferrous iron(II)Fe3+ ferric iron(III)

copper Cu+ cuprous copper(I)Cu2+ cupric copper(II)

tin Sn2+ stannous tin(II)Sn4+ stannic tin(IV)

lead Pb2+ plumbous lead(II)Pb4+ plumbic lead(IV)

antimony Sb3+ stibnous antimony(III)Sb5+ stibnic antimony(V)

cobalt Co2+ cobaltous cobalt(II)Co3+ cobaltic cobalt(III)

gold Au+ aurous gold(I)Au2+ auric gold(II)

mercury Hg+ mercurous mercury(I)Hg2+ mercuric mercury(II)

Chemical Bonding 93

2.5

Practice


1. How were compounds named before the advent of systematic meansof chemical nomenclature?

2. State the common names of each of the following chemicals:(a) hydrochloric acid(b) sodium hydrogen carbonate(c) dinitrogen monoxide(d) ethanol

3. How many elements are there in each binary compound?

4. Describe the IUPAC system of naming a binary ionic compound.

5. If an element is described as“multivalent,” what characteristic does ithave?

6. Write the formula for each of the following compounds:(a) calcium fluoride (l) mercury(I) oxide(b) sodium sulfide (m) nickel(II) bromide(c) aluminum nitride (n) zinc oxide(d) aluminum chloride (o) cobalt(III) chloride(e) potassium oxide (p) strontium bromide(f) calcium chloride (q) gold(I) fluoride(g) copper(II) sulfide (r) lithium chloride(h) lead(II) bromide (s) strontium nitride(i) silver iodide (t) barium bromide(j) barium nitride (u) tin(IV) iodide(k) iron(II) fluoride

7. Write the names of each of the following binary ionic compounds,using the IUPAC system of chemical nomenclature:(a) table salt, NaCl(s) (h) chalcocite, Cu2S(s)(b) lime, CaO(s) (i) galena, PbS2(s)(c) road salt, CaCl2(s) (j) hematite, Fe2O3(s)(d) magnesia, MgO(s) (k) molybdite, MoO3(s)(e) bauxite, Al2O3(s) (l) argentite, Ag2S(s)(f) zinc ore, ZnS(s) (m) zincite, ZnO(s)(g) cassiterite, SnO2(s)

8. Write the IUPAC names for the following binary ionic compounds:(a) Na2O(s) (g) Ni2O3(s)(b) SnCl4(s) (h) Ag2S(s)(c) ZnI2(s) (i) FeCl2(s)(d) SrCl2(s) (j) KBr(s)(e) AlBr3(s) (k) CuI2(s)(f) PbCl4(s) (l) NiS(s)

9. Write the chemical formulas and names (using IUPAC chemicalnomenclature) for the binary ionic compounds formed by each of thefollowing pairs of elements:(a) strontium and oxygen(b) sodium and sulfur(c) silver and iodine(d) barium and fluorine(e) calcium and bromine(f) lithium and chlorine

10. Write the chemical formulas for the following ionic compounds:(a) mercury(II) sulfide, cinnabar ore(b) molybdenum(IV) sulfide, molybdenite ore(c) manganese(IV) oxide, pyrolusite ore(d) nickel(II) bromide

94 Chapter 2

(e) copper(II) chloride (f) iron(III) iodide

11. Rename each of the following compounds, using the IUPAC systemof nomenclature:(a) ferrous sulfide(b) plumbic bromide(c) stannous chloride

Reflecting

12. Why is it sometimes helpful to understand a system of nomenclaturethat has been replaced by another, more comprehensive, system?

Compounds with Polyatomic Ions

Many familiar compounds (such as sulfuric acid, H2SO4(aq), used in car batteries,and sodium phosphate, Na3PO4(aq), a food additive typically found in processedcheese) are composed of three different elements. Compounds of this type areclassified as tertiary compounds. (Many compounds containing polyatomic ionsconsist of more than three different elements, but we will not be dealing withthese at this stage.) Tertiary ionic compounds are composed of a metal ion anda polyatomic ion (a covalently bonded group of atoms, possessing a net charge).We treat polyatomic ions much like regular monatomic ions when we write themin formulas or chemical equations.

Polyatomic ions that include oxygen are called oxyanions. One example isthe nitrate ion, NO3

–. (Compounds involving the nitrate ion are often used in theprocessing of foods, particularly cured meats, where they are often used to con-trol colour. Potassium nitrate and sodium nitrate are added to foods to controlthe growth of microorganisms.) In determining the name of a compound con-taining an oxyanion, the first part of the name is easy: It is the name of the metalcation. The second part requires more thought: We have to consider the threeparts of the ion indicated in Figure 1.

There are four polyatomic ions formed from combinations of chlorine andoxygen. Note that all of these oxyanions have the same charge, despite the factthat their formulas are different.

• ClO� is the hypochlorite ion;• ClO2

– is the chlorite ion;• ClO3

– is the chlorate ion; and• ClO4

– is the perchlorate ion.

Note that in each name the stem is –chlor–. The suffixes and prefixes varyaccording to the number of oxygen atoms in the ion, as described below.

• The per–ate oxyanion has one more oxygen atom than does the -ateoxyanion.

• The -ite oxyanion has one fewer oxygen atom than does the -ate oxyanion.• The hypo–ite oxyanion has one fewer oxygen atom than the -ite oxyanion.

Table 3 indicates some polyatomic ions commonly found in compounds.You do not need to memorize each entry in this table. If you become familiarwith the -ate oxyanions, the most stable and common combination of nonmetaland oxygen, then you need only remember the simple relationship between thenames of the ions.

The crisscross method can be applied to predict the formula of ionic com-pounds involving polyatomic ions. From the ion charges (whether for a single

tertiary compound: a compound com-posed of three different elements

oxyanion: a polyatomic ion containingoxygen

ClO –3

ion charge

determines the stem of theion name

determines the suffix and prefix to the stem of the ion name

Figure 1

The chlorate anion

Chemical Bonding 95

2.5

ion or for a polyatomic ion), determine the number of each ion necessary to yielda net charge of zero.

Sample Problem 5

Write the formula of copper(II) nitrate. (Use Table 3 to find the charge of anitrate ion.)

Solution

Charge on each Cu ion: 2+Charge on each NO3 ion: 1�

2 1

Cu NO3

Cu1 (NO3)2The formula of copper(II) nitrate is Cu(NO3)2.

Notice how brackets are used in the formula if there is more than one of thepolyatomic ions. Brackets are not required with one polyatomic ion or withsimple compounds.

Sample Problem 6

Use IUPAC nomenclature to name CaCO3, which is commonly added to break-fast cereals as a source of calcium.

Solution

CaCO3 is calcium carbonate.

Hydrates

Consumers are sometimes surprised to find a tiny white pouch when they openthe box containing a newly purchased pair of boots. What is it? What purposedoes it serve? The pouch contains a white, crystalline powdered desiccant (asubstance that absorbs water) called silica gel (SiO2(s)). The pouch keeps the airinside the box dry so mildew and other moulds will not grow. Similarlyabsorbent compounds are included in such diverse products as powdered foods,talcum powder, and cat litter. Many tertiary ionic compounds form crystals thatcontain molecules of water within the crystal structure. Such compounds arereferred to as hydrates. When heat is applied to a hydrate, it will decompose toproduce water vapour and an associated ionic compound, indicating that thewater is loosely held to the ionic compound. The water molecules are assumedto be electrically neutral in the compound. When this water, called water ofhydration, is removed, the product is referred to as anhydrous. Bluestone,hydrated copper(II) sulfate, is an example of a hydrate. Its formula is writtenCuSO4·5 H2O(s). Notice how the chemical formula includes both the formula ofthe compound and the formula for water. This is true of the chemical formulasfor all hydrated compounds. The formula for bluestone indicates the associationof five water molecules with each unit of copper(II) sulfate (Figure 2, page 96).The IUPAC names for ionic hydrates indicate the number of water molecules bya Greek prefix (Table 4, page 96), so bluestone, or CuSO4·5H2O(s), is calledcopper(II) sulfate pentahydrate.

Table 3: IUPAC Names and Formulas of Some Common Polyatomic Ions

Name Formula

acetate C2H3O2�

bromate BrO3�

carbonate CO32�

hydrogen carbonate(bicarbonate) HCO3

�

hypochlorite ClO–

chlorite ClO2�

chlorate ClO3�

perchlorate ClO4�

chromate CrO42�

dichromate Cr2O72�

cyanide CN–

hydroxide OH–

iodate IO3�

permanganate MnO4�

nitrite NO2�

nitrate NO3�

phosphate PO43�

hydrogen phosphite HPO32�

hydrogen phosphate HPO42�

dihydrogen phosphite H2PO3�

dihydrogen phosphate H2PO4�

sulfite SO32�

sulfate SO42�

hydrogen sulfide(bisulfide) HS–

hydrogen sulfite(bisulfite) HSO3

�

hydrogen sulfate(bisulfate) HSO4

�

thiosulfate S2O3�

ammonium NH +4

hydrate: a compound that contains wateras part of its ionic crystal structure (theoret-ical definition); a compound that decomposesto an ionic compound and water vapour whenheated (empirical definition)

Naming Hydrates

When researching scientific references, youmay encounter an alternative naming systemfor hydrates, in which CuSO4·5H2O(s) is calledcopper(II) sulfate-5-water.

DID YOU KNOW ?

96 Chapter 2

Practice


13. How many elements are there in a tertiary compound?

14. Use each of the following terms correctly in a sentence about the for-mation of compounds:(a) polyatomic ion(b) oxyanion(c) hydrate

15. Write the IUPAC name for each of the following ionic compounds:(a) NaNO3(s) (found in tobacco)(b) NaNO2(s) (a meat preservative)(c) Cu(NO3)2(s) (forms a blue solution in water)(d) CuNO3(s) (forms a green solution in water)(e) Al2(SO3)3(s) (a food additive in pickles)(f) Ca(OH)2(s) (firming agent in fruit products)(g) PbCO3(s) (cerussite, a mineral popular with collectors)(h) Sn3(PO4)2(s) (use to fix paints to silk)(i) Fe2(SO4)3(s) (a mineral found on Mars)

16. Write the chemical formula for each of the following ionic com-pounds:(a) calcium carbonate (active ingedient in antacids)(b) sodium bicarbonate (a foaming agent added to foods)(c) sodium hypochlorite (a component of bleach)(d) calcium sulfate (plaster of Paris)(e) ammonium nitrate (used in fertilizers)(f) ammonium phosphate (a leavening agent added to foods)(g) copper(II) sulfate (used as a fungicide)(h) sodium hydroxide (a strong base used as a washing agent)(i) potassium permanganate (a traditional antiseptic)

17. Use IUPAC chemical nomenclature to name each of the followingionic compounds containing polyatomic ions:(a) LiClO3(s) (n) Ag2SO4(s)(b) BaSO4(s) (o) Hg(BrO3)2(s)(c) Hg2CO3(s) (p) Fe2(CO3)3(s)(d) Mg(NO3)2(s) (q) NH4ClO(s)(e) Fe(BrO3)3(s) (r) Au(NO3)3(s)(f) Na3PO4(s) (s) Mg(BrO3)2(s)(g) NH4IO3(s) (t) NaIO(s)(h) AuC2H3O2(s) (u) Zn(ClO2)2(s)(i) Zn3(PO4)2(s) (v) SnCO3(s)

2+Cu

2–4SO

Figure 2

In a model of the compound copper(II) sulfatepentahydrate, the copper(II) ions are sur-rounded by four water molecules. The fifthwater molecule is hydrogen-bonded to thesulfate ion.

Table 4: Prefixes Used When Naming Hydrated Compounds

Number of water Prefix in chemicalmolecules nomenclaturein chemical formula

1 mono

2 di

3 tri

4 tetra

5 penta

6 hexa

7 hepta

8 octa

9 nona

10 deca

Chemical Bonding 97

2.5

(j) Sb(ClO3)5(s) (w) SrSO3(s)(k) MnSO3(s) (x) NiPO4(s)(l) KBrO(s) (y) Cu(C2H3O2)2(s)(m) AlPO5(s) (z) Ba3(PO5)2(s)

18. Write the IUPAC name and chemical formula for each of the followingionic compounds containing polyatomic ions:(a) cuprous hypophosphite(b) stannic chlorite(c) ferrous bromate(d) ferric chlorite(e) plumbic sulfate

19. Name each of the following hydrated ionic compounds:(a) bluestone, CuSO4·5 H2O(s)(b) Na2SO4·10 H2O(s)(c) MgSO4·7 H2O(s)

20. Write the chemical formulas for the following ionic hydrates:(a) iron(III) oxide trihydrate (rust)(b) aluminum chloride hexahydrate (component of antiperspirant)(c) sodium thiosulfate pentahydrate (photographic “hypo”)(d) cadmium(II) nitrate tetrahydrate (photographic emulsion)(e) lithium chloride tetrahydrate (in fireworks)(f) calcium chloride dihydrate (deicer)

21. How would you convert a hydrate to an anhydrous compound?

Naming Molecular Compounds

So far, in this section, we have been looking at the names and formulas of ioniccompounds. We now turn to molecular compounds, for which a different systemis used.

If a binary compound is formed from two nonmetals, it is classified as a molec-ular compound. Even though there are fewer nonmetals than metals, there is a widevariety of compounds formed from the combination of two nonmetals, becausetwo nonmetals may combine to form more than one compound. For example, N2O,NO, and NO2 are three of the several binary compounds that can be formed fromnitrogen and oxygen. Each compound has different properties, and so different uses.

In naming compounds formed from two nonmetals, a Greek prefix isattached to the name of each element in the binary compound indicating thenumber of atoms of that element in the molecule. The common prefixes andtheir numerical equivalences are shown in Table 5. If there is only one of the firsttype of atom, we leave out the prefix “mono.”

Suppose you are asked to write the IUPAC name for the chemical compoundrepresented by the formula N2O. Looking at the first element, you can see thatthe subscript after the nitrogen is two, so the prefix for nitrogen is “di.” Lookingat the second element, you can see that there is only one oxygen atom, so theprefix for oxygen will be “mono.” Therefore, the formula’s IUPAC name is dini-trogen monoxide.

Sample Problem 7

What is the IUPAC name for the chemical compound CF4?

Solution

C: carbon (not monocarbon)F: tetrafluorideThe IUPAC name for CF4 is carbon tetrafluoride.

Table 5: Prefixes Used When Naming Binary Covalent Compounds

Subscript in Prefix in chemicalchemical formula nomenclature

1 mono

2 di

3 tri

4 tetra

5 penta

6 hexa

7 hepta

8 octa

9 nona

10 deca

98 Chapter 2

Once again, hydrogen is an exception to this rule: The common practice isnot to use the prefix system for hydrogen. For example, we do not call H2S dihy-drogen sulfide, but simply hydrogen sulfide.

Practice


22. Write the chemical formula for each of the following molecules:(a) nitrogen (n) sulfur tetrafluoride(b) carbon dioxide (o) phosphorus pentachloride(c) carbon monoxide (p) disulfur dichloride(d) nitrogen dioxide (q) carbon tetrachloride(e) nitrogen monoxide (r) sulfur trioxide(f) dinitrogen oxide (s) sulfur hexafluoride(g) dinitrogen tetroxide (t) chlorine dioxide(h) sulfur dioxide (u) dinitrogen pentoxide(i) diiodine pentoxide (v) phosphorus trichloride(j) silicon tetrafluoride (w) silicon tetrachloride(k) boron trifluoride (x) carbon disulfide(l) phosphorus triiodide (y) phosphorus pentabromide(m) diphosphorus pentoxide (z) carbon tetrafluoride

23. Name the compound indicated by each of the following formulas:(a) SF6(g) (f) IF7(g)(b) N2O3(g) (g) BF3(g)(c) NO2(g) (h) P2S5(s)(d) PCl3(l) (i) P2O5(s)(e) PCl5(s)

Naming Acids

You are already familiar with many acids and bases. You may have taken acetyl-salicylic acid (also known as ASA, or Aspirin) to treat a headache, poured a littleacetic acid solution (vinegar) on your fish and chips, or used an ammonia solu-tion to clean the smears from a mirror. You also know one of the tests for acidsand bases: acids turn blue litmus red, and bases turn red litmus blue. But youmight have been somewhat mystified about the naming of these substances.

Acids are well-known, long-established chemicals. They were originally nameddecades or even centuries ago, and the use of traditional names persists. TheInternational Union of Pure and Applied Chemistry suggests that names of acidsshould be derived from the IUPAC name for the compound. According to this rule,sulfuric acid would be called aqueous hydrogen sulfate. However, it is difficult to getpeople to change to new names when the old, familiar names are so widely used.

Let us take, as a simple example, HCl(g). It is a binary compound formedfrom a combination of hydrogen and a halogen. When a gas, it is namedhydrogen chloride. When it is dissolved in water, the resulting aqueous solutiondisplays a set of specific properties called acidic, and the name of the substancechanges. Binary acids are classically named by using the prefix hydro- with thestem of the name of the most electronegative element and the ending -ic. Thename “hydrogen” does not appear. Instead, the word “acid” is added after thehydro-stem-ic combination, as indicated in Table 6. Consequently, the classicalname for HCl(aq) is hydrochloric acid. Its IUPAC name is aqueous hydrogenchloride. Similarly, HBr(aq) is hydrobromic acid or aqueous hydrogen bromide.

Note that the difference between the solution and the pure binary compoundis indicated by the presence or absence of the subscript (aq) in the formula.

Chemical Bonding 99

2.5

A second group of acids is named (by the classical system) in the same way asbinary acids. In this group, the IUPAC names for the polyatomic ions end in -ide(e.g., the cyanide ion, CN–). Looking at Table 6, you can see that the classicalname for the acidic solution HCN(aq) will be hydrocyanic acid.

A third group of acids is formed from various combinations of oxyanions (neg-ative polyatomic ions consisting of a nonmetal plus oxygen) with hydrogen. Perhapsthe best-known example is H2SO4(aq), or sulfuric acid, which is one of the mostwidely produced industrial chemicals in the world. It is used to make pharmaceuti-cals, detergents, and dyes, and is a component of car batteries. Phosphoric acid,H3PO4(aq), is another example of an acid formed from an oxyanion and hydrogen.Phosphoric acid is an ingredient in soft drinks, is a reagent in the manufacture offertilizers, and can be used as a rust remover. These acids are classified as oxyacids

because they incorporate oxyanions. Table 7 compares the classical name to theIUPAC name for a series of acids derived from a chlorine-based oxyanion.

The classical names for oxyacids can be derived according to the simple rulesin Table 8.

When naming oxyacids, we omit the word hydrogen and add the word“acid.” For example, to name the acidic solution with the formula HNO2(aq), wewould first consider the IUPAC name: hydrogen nitrite (from Table 3, or fromthe Table of Polyatomic Ions in Appendix C). “Nitrite” changes to “nitrous,” wedrop the “hydrogen” from the front of the name, and add “acid” to the end. Thus,HNO2(aq) is called nitrous acid.

Table 6: Naming Systems for Binary Acids

Formula Classical name IUPAC name

HF(aq) hydrofluoric acid aqueous hydrogen fluoride

HCl(aq) hydrochloric acid aqueous hydrogen chloride

HBr(aq) hydrobromic acid aqueous hydrogen bromide

HI(aq) hydroiodic acid aqueous hydrogen iodide

H2S(aq) hydrosulfuric acid aqueous hydrogen sulfide

Table 7: Classical and IUPAC Nomenclature System for Chlorine-Based Oxyacids

Classical name IUPAC name Formula

perchloric acid aqueous hydrogen perchlorate HClO4(aq)

chloric acid aqueous hydrogen chlorate HClO3(aq)

chlorous acid aqueous hydrogen chlorite HClO2(aq)

hypochlorous acid aqueous hydrogen hypochlorite HClO(aq)

hydrochloric acid aqueous hydrogen chloride HCl(aq)

Table 8: Rules for Naming Acids and Oxyanions

Name of Example Formula Classical Exampleoxyanion name of acid

per–ate persulfate SO52� per–ic acid persulfuric acid

-ate sulfate SO42� -ic acid sulfuric acid

-ite sulfite SO32� -ous acid sulfurous acid

hypo–ite hyposulfite SO22� hypo–ous acid hyposulfurous acid

oxyacid: an acid containing oxygen,hydrogen, and a third element

100 Chapter 2

If you were asked to write the formula for an acid, you would first have tofigure out the names of the ions involved, then their symbols or formulas, thentheir ratio. For example, what is the formula for phosphoric acid? The -ic endingindicates the presence of the -ate oxyanion of phosphorus: phosphate. The phos-phate oxyanion is PO4

3– with a charge of 3–. The cation in oxyacids is alwayshydrogen, which has a charge of 1+. To find the ratio of the ions, use the criss-cross method:

1 3

H (PO4)This gives the subscripts for each ion.

H3 (PO4)1Divide each subscript by the highest common factor—in this case, 1:

H3 (PO4)1The hydrogen ions and the phosphate oxyanions will combine in a ratio of 3:1.Therefore, the correct formula is H3PO4(aq).

Sample Problem 8

What is the formula for the oxyacid sulfurous acid?

Solution

sulfurous indicates sulfite ion: SO32�

1 2

H2 (SO3)1The formula for sulfurous acid is H2SO3(aq).

Sample Problem 9

What is the formula for the oxyacid hypochlorous acid?

Solution

hypochlorous indicates one fewer oxygen atom than a chlorite ion: ClO–

1 1

H1 (ClO)1The formula for hypochlorous acid is HClO(aq).

Naming Bases

Chemists have discovered that all aqueous solutions of ionic hydroxides arebases. (You will learn more about bases in Chapter 8.) Other solutions have alsobeen classified as bases, but for the time being we will restrict our exploration ofbases to aqueous ionic hydroxides such as NaOH(aq) and Ba(OH)2(aq). Noticethat these bases are formed of a combination of a metal cation with one or morehydroxide anions. The name of the base is the name of the ionic hydroxide: inthis case, aqueous sodium hydroxide and aqueous barium hydroxide.

Chemical Bonding 101

2.5

Practice


24. Write the chemical formulas for the following acids:(a) aqueous hydrogen chloride (g) aqueous hydrogen nitrite(b) hydrochloric acid (h) nitric acid(c) aqueous hydrogen sulfate (i) hydrobromic acid(d) sulfuric acid (j) hyposulfurous acid(e) aqueous hydrogen acetate (k) hydroiodic acid(f) acetic acid (l) aqueous hydrogen perchlorate

25. Name each of the following compounds, using both the classical andthe IUPAC nomenclature systems:(a) H2SO3(aq) (d) H2CO3(aq) (g) HCN(aq)(b) H3PO4(aq) (e) H2S(aq) (h) H2SO4(aq)(c) HCN(aq) (f) HCl(aq) (i) H3PO4(aq)

26. Write the names of the following bases:(a) KOH(aq) (b) CaOH2(aq)

27. Write the formulas of the following bases:(a) aqueous magnesium hydroxide(b) aqueous sodium hydroxide(c) aqueous aluminum hydroxide


1. The atmosphere of Saturn contains traces of ammonia, whilehydrogen cyanide is a component of the atmosphere of one ofSaturn’s many moons.(a) Write formulas for each of these compounds.(b) Classify each substance as ionic or covalent.(c) Which compound contains both ionic and covalent bonds?

How could you verify this prediction?(d) If hydrogen cyanide is added to water, what substance is

formed? If the substance were tested with litmus paper, whatwould you expect the results to be?

2. Caustic potash (KOH) is a component of oven cleaners. Muriaticacid (HCl) is used to clean mortar from bricks.(a) Classify each of these compounds. Provide your reasoning.(b) Create electron dot diagrams and structural formulas for

each compound.(c) Use the IUPAC system to name each compound.(d) What can you predict about the reactivity of these sub-

stances? How would this affect the way you handle the sub-stances? Explain your answer.

3. (a) Write the IUPAC names and formulas for as many com-pounds as you can, using only the following elements: K, C,H, F, Mg, O, Cl, and Na.

(b) State the common name, if there is one, for each compound.(c) Classify each compound as polar or nonpolar. In each case,

use Lewis structures and periodic trends to justify your classification.

(d) Classify the compounds as binary, tertiary, acidic, or basic.(e) Indicate which compounds, if any, contain both ionic and

covalent bonding.


102 Chapter 2

SummaryChapter 2

Key Expectations

Throughout this chapter, you have had the

opportunity to do the following:

• Identify chemical substances and reactions in everydayuse or of environmental significance. (all sections)

• Explain how different elements combine to form covalent and ionic bonds, using the octet rule. (2.2)

• Demonstrate an understanding of the formation ofionic and covalent bonds and explain the properties ofthe products. (2.2, 2.3)

• Predict the ionic character or polarity of a given bondusing electronegativity values, and represent the forma-tion of ionic and covalent bonds using diagrams. (2.2,2.3, 2.4)

• Draw Lewis structures, construct molecular models,and give the structural formulas for compounds con-taining single and multiple bonds. (2.3)

• Relate common names of substances to their systematic(IUPAC) names. (all sections)

• Write, using IUPAC or traditional (classical) systems,the formulas of binary and tertiary compounds,including those containing elements with multiplevalences, and recognize the formulas in various con-texts. (2.5)

• Use appropriate scientific vocabulary to communicateideas related to chemical reactions. (all sections)

Key Terms

binary compoundbonding capacitychemical bondchemical nomenclaturecoordinate covalent

bondcovalent bondcrystal latticediatomic moleculedipole–dipole forceelectrical conductivityelectrolyteelectron dot diagramformula unithydratehydrogen bondintermolecular forceintramolecular forceionic bondionic compound

Lewis structureLewis symbolLondon dispersion forcelone pairmolecular compoundmultivalentoctet ruleoxyacidoxyanionpolar covalent bondpolar moleculepolyatomic ionpolyatomic moleculestable octetstructural formulatertiary compoundvalencevan der Waals forces

Summarize the concepts presented in this chapter by constructing andcompleting a table like Table 1. Try to include as many examples of com-pounds as possible for each type of intramolecular bond.

Make aSummary

Table 1: Summarizing Bonds and Forces

Compound Properties Lewis Intramolecular Polarity Intermolecular structure bond type forces

Revisit your answers to the Reflect on Your Learning questionsat the beginning of this chapter.

• How has your thinking changed?• What new questions do you have?

Reflect Learningonyour

Review



1. Describe the relationship between the positions of tworeacting elements in the periodic table and the type ofcompound and bond they form.

2. How does an ionic bond differ from a covalent bond?

3. (a) Briefly summarize and explain the properties ofionic and molecular compounds.

(b) Explain why electrical conductivity is generally asuitable test for ionic compounds.

4. (a) What is intramolecular bonding? Describe the maintypes with examples.

(b) What is the relationship between the type of bondand the properties of a substance?

5. For each chemical bond or force listed below, indicatewhich types of entities are involved.(a) covalent bond(b) dipole–dipole force(c) hydrogen bonds(d) ionic bond

6. How does the information conveyed by the chemicalformulas for ionic compounds differ from the informa-tion conveyed by the chemical formulas for molecularcompounds? Include examples with your explanation.

7. Explain why halogens tend to form diatomic molecules.

8. Sketch the bonding continuum. Indicate where Na—Cland Cl—Cl bonds would fit along the continuum.Provide a theoretical explanation of your answer.

9. Draw Lewis symbols (electron dot diagrams) to repre-sent a single atom of each of the following elements:(a) Ca (e) S(b) Al (f) Br(c) K (g) Ne(d) N

10. For each of the following substances, predict whetherthe bonds present are ionic, covalent, or polar covalent:(a) I2 (s) (d) Fe2O3(s)(b) SO2(g) (e) KBr(s)(c) OCl2(g) (f) SrO(s)

11. The most common oxides of Period 3 elements are asfollows: Na2O(s), MgO(s), Al2O3(s), SiO2(s), P2O5(s),SO2(g), and Cl2O(g).(a) Classify the oxides as either ionic or molecular.(b) Use electron dot diagrams or Lewis structures to

show the formation of each compound.(c) What differences would you expect to observe in

the properties of each of the compounds?

(d) How is the difference in electronegativity of theconstituent elements related to the properties of thecompound?

12. Create a table like Table 1 and complete it for the fol-lowing compounds: HF(g), BCl3(g), SiH4(g), CCl4(l),NCl3(g), H2O2(l), CO2(g), HCN(g)

13. Use electronegativity values to predict the polarity ofthe bond formed by each of the following pairs of ele-ments. In each case, represent the formation of thebonds using diagrams.(a) carbon and chlorine(b) calcium and fluorine(c) aluminum and chlorine(d) silicon and oxygen(e) carbon and oxygen

14. Use Lewis structures and structural formulas to repre-sent the following compounds. Indicate which com-pounds, if any, involve coordinate covalent bonds.(a) NH4Cl(s) (d) H2O(l)(b) BF3(g) (e) NH3BF3(g)(c) NH3(g)

15. Write the chemical formula (plus state) for each of thefollowing substances:(a) sodium hydrogen sulfate (toilet bowl cleaner)(b) sodium hydroxide (lye, drain cleaner)(c) carbon dioxide (dry ice, soda pop)(d) acetic acid (vinegar)(e) sodium thiosulfate pentahydrate (photographic

“hypo”)(f) sodium hypochlorite (laundry bleach)(g) octasulfur (vulcanizing rubber)(h) potassium nitrate (meat preservative)(i) phosphoric acid (rust remover)(j) iodine (disinfectant)(k) aluminum oxide (alumina, aluminum ore)(l) potassium hydroxide (caustic potash)

(m) aqueous hydrogen carbonate (carbonated beverages)

16. Write the chemical formula for each of the followingsubstances:(a) magnesium bromide(b) carbon disulfide(c) mercury(II) nitrite

Chapter 2

Table 1: Structures of Covalent Compounds

Compound Lewis Structural Typesstructure formula of bonds

104 Chapter 2

(d) hydrochloric acid(e) lithium hydroxide(f) silver carbonate(g) aluminum perchlorate(h) copper(II) sulfate(i) sulfur trioxide(j) nickel(III) phosphate(k) magnesium oxide(l) dinitrogen monoxide(m) iron(II) persulfate(n) carbonic acid(o) calcium hydroxide(p) zinc hypochlorite(q) lead(IV) perchlorate(r) phosphorous pentabromide(s) arsenic(V) chloride(t) bismuth(III) nitrate(u) sodium hypochlorite(v) oxygen dichloride(w) tin(II) bromide(x) sulfuric acid(y) potassium hydroxide(z) barium carbonate

17. Write the chemical formula for each of the followingsubstances:(a) ammonium dihydrogen phosphite(b) lithium hydrogen sulfite(c) potassium hydrogen sulfate(d) barium chloride trihydrate(e) sodium dihydrogen phosphate(f) sodium hydrogen carbonate

18. Give the names of the following substances, usingIUPAC chemical nomenclature:(a) CaCO3(s) (marble, limestone, chalk)(b) P2O5(s) (fertilizer)(c) MgSO4

•7 H2O(s) (Epsom salts)(d) N2O(g) (laughing gas, an anesthetic)(e) Na2SiO3(s) (water glass)(f) Ca(HCO3)2(s) (hard-water chemical)(g) HCl(aq) (muriatic acid, gastric fluid)(h) CuSO4

•5 H2O(s) (copperplating, bluestone)(i) H2SO4(aq) (acid in car battery)(j) Ca(OH)2(s) (slaked lime)(k) SO3(g) (a cause of acid rain)(l) NaF(s) (toothpaste additive)

19. Give the IUPAC names of the following substances:(a) NaCl(s) (d) Pb(C2H3O2)2(s)(b) P2O3(s) (e) NH4OCl(s)(c) HNO3(aq) (f) Sn(BrO3)4(s)

(g) Sb2O3(s) (q) Ba(C2H3O2)2(s)(h) Zn(IO3)2(s) (r) ICl(s)(i) Fe(NO4)2(s) (s) AuCl3(s)(j) Ca(OH)2(s) (t) MgS(s)(k) KI(s) (u) N2F2(g)(l) SF2(g) (v) NiSO4(s)(m) HBr(aq) (w) H2S(aq)(n) CuCO3(s) (x) AgBrO3(s)(o) Al2(SO3)3(s) (y) LiClO4(s)(p) NH4OH(l)

20. Give the names of the following substances, usingIUPAC chemical nomenclature:(a) CaHPO4(s) (d) LiHCO3(s)(b) CuSO4

•7 H2O(s) (e) KHSO4(s)(c) Na2HPO4(s)

21. Write the name and formula (with state at SATP) forthe compound formed by each of the following pairs ofelements. Where a molecular compound is formed, givethe structural formula. For ionic compounds, assumethe most common ion charges for the ions.(a) potassium and bromine(b) silver and iodine(c) lead and oxygen(d) zinc and sulfur(e) copper and oxygen(f) lithium and nitrogen


22. A forensic chemist was given samples of four unknownsolutions, the identity of which could affect the out-come of a court case involving an electrocution. Thechemist had reason to believe that the four substanceswere KCl(aq), C2H5OH(aq), HCl(aq), and Ba(OH)2(aq).The investigation was designed to identify the chemi-cals. Complete the Analysis and Synthesis sections ofthe report.

Question

What is the identity of each of the four substances,labelled 1 to 4?

Experimental Design

Each of the samples was dissolved in water and testedwith a conductivity apparatus and litmus paper. Tastetests were not carried out.


24. Boron nitride (BN) is a compound that involves a largenumber of covalent bonds. Use the Internet to researchthe structure of boron nitride. Use your findings tomake predictions about the physical properties ofboron nitride. How do the properties of boron nitrideand diamond compare? Could boron nitride serve as asubstitute for diamond in some practical way? Provide atheoretical argument that supports your answer.

Follow the links for Nelson Chemistry 11, Chapter 2Review.

25. Glycerol is a sweet-tasting, colourless, viscous liquidthat is considered (by many national food safety organi-zations) to be quite safe to eat. The food industry addsglycerol (sometimes called glycerine) to a wide range ofprocessed cheese and meat, baked goods, and candies.(a) Look for glycerol on ingredients labels at home or

on the shelves of stores. Make a list of foods inwhich glycerol is used.

(b) Research the physical and chemical properties ofglycerol.

(c) Research the chemical formula and molecularshape of glycerol. Use this information, as well asyour knowledge of intra- and intermolecularbonding, to explain at least three of the propertiesof glycerol.

(d) Write a short report, outlining why glycerol is souseful to the food industry.


Exploring

26. London dispersion forces are named for Fritz London,who explained the origin of intermolecular forcesbetween nonpolar molecules in 1928. ResearchLondon’s work and write a brief report on his majorfindings.


Evidence

Analysis

(a) Answer the Question.

Synthesis

(b) Why was the water used to prepare the solutionsalso tested?

(c) Which of the solutions, 1, 2, 3, or 4, could havebeen involved in somebody getting electrocuted?Use electron dot diagrams and the concept of elec-tronegativity to explain your conclusion.

Making Connections

23. In historical dramas you may have seen someonebrought back to consciousness by having a bottle of“smelling salts” waved under their noses. Smelling saltswere made by mixing perfume and the active ingre-dient, which was once called sal volatile.(a) What is the IUPAC name and chemical formula for

sal volatile?(b) Based on the nature of the bonds in this substance,

what properties would you predict that it woulddemonstrate?

(c) Does your prediction fit with the evidence thatsmelling salts have a strong, sharp, ammonia smell?Explain.

(d) Using the Internet, find out how smelling salts usedto be administered, how they worked, and com-ment on the safety concerns that this use wouldraise today.






Table 2: Evidence for Identifying Unknown Solutions

Solution Conductivity Litmus

water none no change

1 high no change

2 high blue to red

3 none no change

4 high red to blue

2 chemical bonding - chute's web...

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