Covalent Bonding - HELP PDF

Section 8.1 Objectives ◗ Apply the octet rule to atoms that The Covalent Bond form covalent bonds. -!). )DEA Atoms gain stability when they share electrons ◗ Describe the formation of single, and form covalent bon...

Section 8.1 Objectives ◗ Apply the octet rule to atoms that The Covalent Bond form covalent bonds. -!). )DEA Atoms gain stability when they share electrons ◗ Describe the formation of single, and form covalent bonds. double, and triple covalent bonds. Real-World Reading Link Have you ever run in a three-legged race? ◗ Contrast sigma and pi bonds. Each person in the race shares one of their legs with a teammate to form a ◗ Relate the strength of a covalent single three-legged team. In some ways, a three-legged race mirrors how atoms bond to its bond length and bond share electrons and join together as a unit. dissociation energy. Review Vocabulary Why do atoms bond? chemical bond: the force Understanding the bonding in compounds is essential to developing that holds two atoms together new chemicals and technologies. To understand why new compounds form, recall what you know about elements that do not tend to form New Vocabulary new compounds—the noble gases. You read in Chapter 6 that all noble covalent bond gases have stable electron arrangements. This stable arrangement con- molecule sists of a full outer energy level and has lower potential energy than Lewis structure other electron arrangements. Because of their stable configurations, sigma bond noble gases seldom form compounds. pi bond endothermic reaction Gaining stability The stability of an atom, ion, or compound is exothermic reaction related to its energy; that is, lower energy states are more stable. In Chapter 7, you read that metals and nonmetals gain stability by trans- ferring (gaining or losing) electrons to form ions. The resulting ions have stable noble-gas electron configurations. From the octet rule in Chapter 6, you know that atoms with a complete octet, a configuration of eight valence electrons, are stable. In this chapter, you will learn that the sharing of valence electrons is another way atoms can acquire the stable electron configuration of noble gases. The water droplets shown in Figure 8.1 consist of water molecules formed when hydrogen and oxygen atoms share electrons. Figure 8.1 Each water droplet is made up of water molecules. Each water molecule is made up of two hydrogen atoms and one oxygen atom that have bonded by sharing electrons. The shapes of the drops are due to intermolecular forces acting on the water molecules. 240 Chapter 8 Covalent Bonding ©Charles Krebs/Getty Images  Force of repulsion Force of attraction The atoms are too Each nucleus attracts the other The distance is right for the attrac- If the atoms are forced far apart to have atom’s electron cloud. Repulsion tion between one atom’s protons closer together, the noticeable attraction occurs between nuclei and and the other atom’s electrons to nuclei and electrons or repulsion. between electron clouds. make the bond stable. repel each other. Figure 8.2 The arrows in this diagram What is a covalent bond? show the net forces of attraction and repulsion acting on two fluorine atoms as they move You just read that atoms can share electrons to form stable electron con- toward each other. The overall force between figurations. How does this occur? Are there different ways in which two atoms is the result of electron-electron electrons can be shared? How are the properties of these compounds repulsion, nucleus-nucleus repulsion, and different from those formed by ions? Read on to answer these questions. nucleus-electron attraction. At the position of maximum net attraction, a covalent bond forms. Shared electrons Atoms in nonionic compounds share electrons. Relate How is the stability of the bond The chemical bond that results from sharing valence electrons is a related to the forces acting on the atoms? covalent bond. A molecule is formed when two or more atoms bond covalently. In a covalent bond, the shared electrons are considered to be part of the outer energy levels of both atoms involved. Covalent bond- ing generally can occur between elements that are near each other on the periodic table. The majority of covalent bonds form between atoms of nonmetallic elements. Covalent bond formation Diatomic molecules, such as hydrogen (H 2), nitrogen (N 2), oxygen (O 2), fluorine (F 2), chlorine (Cl 2), bromine (Br 2), and iodine (I 2), form when two atoms of each element share elec- trons. They exist this way because the two-atom molecules are more sta- ble than the individual atoms. Consider fluorine, which has an electron configuration of 1s 22s 22p 5. Each fluorine atom has seven valence electrons and needs another elec- tron to form an octet. As two fluorine atoms approach each other, sever- al forces act, as shown in Figure 8.2. Two repulsive forces act on the atoms, one from each atom’s like-charged electrons and one from each Figure 8.3 Two fluorine atoms share atom’s like-charged protons. A force of attraction also acts, as one atom’s a pair of electrons to form a covalent bond. Note that the shared electron pair gives protons attract the other atom’s electrons. As the fluorine atoms move each atom a complete octet. closer, the attraction of the protons in each nucleus for the other atom’s electrons increases until a point of maximum net attraction is achieved. F + F At that point, the two atoms bond covalently and a molecule forms. If Fluorine Fluorine the two nuclei move closer, the repulsion forces increase and exceed the atom atom attractive forces. The most stable arrangement of atoms in a covalent bond exists at some optimal distance between nuclei. At this point, the net attraction is greater than the net repulsion. Fluorine exists as a diatomic molecule Bonding pair Complete because the sharing of one pair of electrons gives each fluorine atom a octets of electrons stable noble-gas configuration. As shown in Figure 8.3, each fluorine F F Lone atom in the fluorine molecule has one pair of electrons that are cova- pairs Fluorine lently bonded (shared) and three pairs of electrons that are unbonded molecule (not shared). Unbonded pairs are also known as lone pairs. Section 8.1 The Covalent Bond 241  Compare Melting Points 7. Turn the temperature knob on the hot plate to the highest setting. You will heat the How can you determine the relationship compounds for 5 min. Assign someone to time between bond type and melting point? The the heating of the compounds. properties of a compound depend on whether the 8. Observe the compounds during the 5-min bonds in the compound are ionic or covalent. period. Record which compounds melt and the order in which they melt. Procedure 9. After 5 min, turn off the hot plate and remove 1. Read and complete the lab safety form. the pie pan using a hot mitt or tongs. 2. Create a data table for the experiment. 10. Allow the pie pan to cool,and then place it in 3. Using a permanent marker, draw three lines the proper waste container. on the inside bottom of a disposable, 9-inch aluminum pie pan to create three, equal Analysis wedges. Label the wedges, A, B, and C. 1. State Which solid melted first? Which solid 4. Set the pie pan on a hot plate. did not melt? WARNING: Hot plate and metal pie pan will burn 2. Apply Based on your observations and data, skin—handle with care. describe the melting point of each solid as low, 5. Obtain samples of the following from your medium, high, or very high. teacher and deposit them onto the labeled 3. Infer Which compounds are bonded with ionic wedges as follows: sugar crystals (C 12H 22O 11), bonds? Which are bonded with covalent bonds? A; salt crystals (NaCl) B; paraffin (C 23H 48), C. 4. Summarize how the type of bonding affects 6. Predict the order in which the compounds the melting points of compounds. will melt. Single Covalent Bonds When only one pair of electrons is shared, such as in a hydrogen molecule, it is a single covalent bond. The shared electron pair is often referred to as the bonding pair. For a hydrogen molecule, shown in Figure 8.4, each covalently bonded atom equally attracts the pair of shared electrons. Thus, the two shared elec- trons belong to each atom simultaneously, which gives each hydrogen atom the noble-gas configuration of helium (1s 2) and lower energy. The hydrogen molecule is more stable than either hydrogen atom is by itself. Recall from chapter 5 that electron-dot diagrams can be used to show valence electrons of atoms. In a Lewis structure, they can represent the arrangement of electrons a molecule. A line or a pair of vertical dots between the symbols of elements represents a single covalent bond in a Lewis structure. For example, a hydrogen molecule is written as H—H or H:H. Figure 8.4 When two hydrogen atoms share a pair of electrons, each hydrogen atom is stable because it has a full outer-energy level. + → H + H → HH Hydrogen atom Hydrogen atom Hydrogen molecule 242 Chapter 8 Covalent Bonding Group 17 and single bonds The halogens—the Water group 17 elements—such as fluorine have seven valence electrons. To form an octet, one more electron a is needed. Therefore, atoms of group 17 elements form single covalent bonds with atoms of other nonmetals, such as carbon. You have already read that the atoms of some group 17 elements form covalent bonds with identical atoms. For example, fluorine exists as F 2 and 2H + O → H—O — chlorine exists as Cl 2. H Group 16 and single bonds An atom of a group Two Single Covalent Bonds 16 element can share two electrons and can form two covalent bonds. Oxygen is a group 16 element with an Ammonia electron configuration of 1s 22s 22p 4. Water is com- b posed of two hydrogen atoms and one oxygen atom. Each hydrogen atom has the noble-gas configuration of helium when it shares one electron with oxygen. Oxygen, in turn, has the noble-gas configuration of H neon when it shares one electron with each hydrogen — atom. Figure 8.5a shows the Lewis structure for a 3H + N → H—N — molecule of water. Notice that the oxygen atom has two single covalent bonds and two unshared pairs H of electrons. Three Single Covalent Bonds Group 15 and single bonds Group 15 elements Methane form three covalent bonds with atoms of nonmetals. c Nitrogen is a group 15 element with the electron con- figuration of 1s 22s 22p 3. Ammonia (NH 3) has three single covalent bonds. Three nitrogen electrons bond with the three hydrogen atoms leaving one pair of unshared electrons on the nitrogen atom. Figure 8.5b H shows the Lewis structure for an ammonia molecule. — Nitrogen also forms similar compounds with atoms of 4H + C → H—C—H — group 17 elements, such as nitrogen trifluoride (NF 3), H nitrogen trichloride (NCl 3), and nitrogen tribromide (NBr 3). Each atom of these group 17 elements and the Four Single Covalent Bonds nitrogen atom share an electron pair. Figure 8.5 These chemical equations show how atoms Group 14 and single bonds Atoms of group share electrons and become stable. As shown by the Lewis 14 elements form four covalent bonds. A methane structure for each molecule, all atoms in each molecule achieve a molecule (CH 4) forms when one carbon atom bonds full outer energy level. with four hydrogen atoms. Carbon, a group 14 ele- Describe For the central atom in each molecule, describe how the octet rule is met. ment, has an electron configuration of 1s 22s 22p 2. With four valence electrons, carbon needs four more elec- trons for a noble gas configuration. Therefore, when carbon bonds with other atoms, it forms four bonds. Because a hydrogen atom, a group 1 element, has one valence electron, it takes four hydrogen atoms to pro- vide the four electrons needed by a carbon atom. The Lewis structure for methane is shown in Figure 8.5c. Carbon also formts single covalent bonds with other nonmetal atoms, including those in group 17. Reading Check Describe how a Lewis structure shows a covalent bond. Section 8.1 The Covalent Bond 243  EXAMPLE Problem 8.1 Lewis Structure of a Molecule The pattern on the glass shown in Figure 8.6 was made by chemically etching its surface with hydrogen fluoride (HF). Draw the Lewis structure for a molecule of hydrogen fluoride. 1 Analyze the Problem You are given the information that hydrogen and fluorine form the molecule hydrogen fluoride. An atom of hydrogen, a group 1 element, has only one valence electron. It can bond with any nonmetal atom when they share one pair of electrons. An atom of fluorine, a group 17 element, needs one electron to complete its octet. Therefore, a single covalent bond forms when atoms of Figure 8.6 The frosted-looking hydrogen and fluorine bond. portions of this glass were chemically etched using hydrogen fluoride (HF), a weak acid. Hydrogen fluoride reacts with 2 Solve for the Unknown silica, the major component of glass, and To draw a Lewis structure, first draw the electron-dot diagram for forms gaseous silicon tetrafluoride (SiF 4) each of the atoms. Then, rewrite the chemical symbols and draw a and water. line between them to show the shared pair of electrons. Finally, add dots to show the unshared electron pairs. H + F → H—F Hydrogen Fluorine Hydrogen fluoride atom atom molecule 3 Evaluate the Answer Each atom in the new molecule now has a noble-gas configuration and is stable. PRACTICE Problems Extra Practice Page 979 and glencoe.com Draw the Lewis structure for each molecule. 1. PH 3 4. CCl 4 2. H 2S 5. SiH 4 3. HCl 6. Challenge Draw a generic Lewis structure for a molecule formed between atoms of Group 1 and Group 16 elements. The sigma bond Single covalent bonds are also called sigma bonds, represented by the Greek letter sigma (σ). A sigma bond occurs when the pair of shared electrons is in an area centered between the two atoms. When two atoms share electrons, their valence atomic orbitals VOCABULARY overlap end to end, concentrating the electrons in a bonding orbital ACADEMIC VOCABULARY between the two atoms. A bonding orbital is a localized region where Overlap bonding electrons will most likely be found. Sigma bonds can form to occupy the same area in part when an s orbital overlaps with another s orbital or a p orbital, or two The two driveways overlap at the street p orbitals overlap. Water (H 2O), ammonia (NH 3), and methane (CH 4) forming a common entrance. have sigma bonds, as shown in Figure 8.7. Reading Check List the orbitals that can form sigma bonds in a covalent compound. 244 Chapter 8 Covalent Bonding ©Visual Arts Library (London)/Alamy  H O N H C H H H H H H H Water (H2O) Ammonia (NH3) Methane (CH4) Figure 8.7 Sigma bonds formed in each Multiple Covalent Bonds of these molecules when the atomic orbital of each hydrogen atom overlapped end to end In some molecules, atoms have noble-gas configurations when they with the orbital of the central atom. share more than one pair of electrons with one or more atoms. Sharing Interpret Identify the types of orbitals multiple pairs of electrons forms multiple covalent bonds. A double that overlap to form the sigma bonds in covalent bond and a triple covalent bond are examples of multiple methane. bonds. Carbon, nitrogen, oxygen, and sulfur atoms often form multiple bonds with other nonmetals. How do you know if two atoms will form a multiple bond? In general, the number of valence electrons needed to form an octet equals the number of covalent bonds that can form. Double bonds A double covalent bond forms when two pairs of electrons are shared between two atoms. For example, atoms of the ele- ment oxygen only exist as diatomic molecules. Each oxygen atom has six valence electrons and must obtain two additional electrons for a noble- gas configuration, as shown in Figure 8.8a. A double covalent bond forms when each oxygen atom shares two electrons; a total of two pairs &/,$!",%3 of electrons are shared between the two atoms. Incorporate information from this section into Triple bonds A triple covalent bond forms when three pairs of elec- your Foldable. trons are shared between two atoms. Diatomic nitrogen (N 2) molecules contain a triple covalent bond. Each nitrogen atom shares three electron pairs, forming a triple bond with the other nitrogen atom as shown in Figure 8.8b. The pi bond A multiple covalent bond consists of one sigma bond and at least one pi bond. A pi bond, represented by the Greek letter pi (π), forms when parallel orbitals overlap and share electrons. The shared electron pair of a pi bond occupies the space above and below the line that represents where the two atoms are joined together. Figure 8.8 Multiple covalent bonds Two shared pairs form when two atoms share more than one of electrons pair of electrons. a. Two oxygen atoms form a O + O → O—O a double bond. b. A triple bond forms between two nitrogen atoms. b N + N → N— —N Three shared pairs of electrons Personal Tutor For an online tutorial on multiple covalent bonds, visit glencoe.com. Section 8.1 The Covalent Bond 245  Figure 8.9 Notice how the multiple bond between the two carbon atoms in p overlap ethene (C 2H 4) consists of a sigma bond and a σ bond pi bond. The carbon atoms are close enough that the side-by-side p orbitals overlap and H H forms the pi bond. This results in a doughnut- σ bond H H shaped cloud around the sigma bond. σ bond C σ bond C C—C σ bond H H H H p overlap π bond Interactive Figure To see an animation of sigma and pi bonding, visit glencoe.com. Ethene It is important to note that molecules having multiple covalent bonds contain both sigma and pi bonds. A double covalent bond, as shown in Figure 8.9, consists of one pi bond and one sigma bond. A triple covalent bond consists of two pi bonds and one sigma bond. The Strength of Covalent Bonds Recall that a covalent bond involves attractive and repulsive forces. In a molecule, nuclei and electrons attract each other, but nuclei repel other nuclei, and electrons repel other electrons. When this balance of forces is upset, a covalent bond can be broken. Because covalent bonds differ in strength, some bonds break more easily than others. Several factors influence the strength of covalent bonds. Bond length The strength of a covalent bond depends on the dis- tance between the bonded nuclei. The distance between the two bonded nuclei at the position of maximum attraction is called bond length, as shown in Figure 8.10. It is determined by the sizes of the two bonding atoms and how many electron pairs they share. Bond lengths for mole- cules of fluorine (F 2), oxygen (O 2), and nitrogen (N 2) are listed in Table 8.1. Notice that as the number of shared electron pairs increases, the bond length decreases. Bond length and bond strength are also related: the shorter the bond length, the stronger the bond. Therefore, a single bond, such as that in F 2, is weaker than a double bond, such as that in O 2. Likewise, the double bond in O 2 is weaker than the triple bond in N 2. Reading Check Relate covalent bond type to bond length. Figure 8.10 Bond length is the distance from the center of one nucleus to the center of the other nucleus of two bonded atoms. Covalent Bond Type and Table 8.1 Nuclei Bond Length Molecule Bond Type Bond Length F2 single covalent 1.43 × 10 -10 m O2 double covalent 1.21 × 10 -10 m Bond length N2 triple covalent 1.10 × 10 -10 m 246 Chapter 8 Covalent Bonding  Table 8.2 Bond-Dissociation Energy Molecule Bond-Dissociation Energy F2 159 kJ/mol O2 498 kJ/mol N2 945 kJ/mol Bonds and energy An energy change occurs when a bond between atoms in a molecule forms or breaks. Energy is released when a bond forms, but energy must be added to break a bond. The amount of energy required to break a specific covalent bond is called bond-dissociation energy and is always a positive value. The bond-dissociation energies for the covalent bonds in molecules of fluorine, oxygen, and nitrogen are listed in Table 8.2. Bond-dissociation energy also indicates the strength of a chemical Figure 8.11 Breaking the C–C bonds in charcoal and the O–O bonds in bond because of the inverse relationship between bond energy and the oxygen in air requires an input of bond length. As indicated in Table 8.1 and Table 8.2, the smaller bond energy. Energy is released as heat and length, the greater the bond-dissociation energy. The sum of the bond- light when bonds form producing CO 2. dissociation energy values for all of the bonds in a molecule is the Thus, the burning of charcoal is an amount of chemical potential energy in a molecule of that compound. exothermic reaction. The total energy change of a chemical reaction is determined from the energy of the bonds broken and formed. An endothermic reaction occurs when a greater amount of energy is required to break the exist- ing bonds in the reactants than is released when the new bonds form in the products. An exothermic reaction occurs when more energy is released during product bond formation than is required to break bonds in the reactants. See Figure 8.11. Section 8.1 Assessment Section Summary 7. -!). )DEA Identify the type of atom that generally forms covalent bonds. ◗ Covalent bonds form when atoms 8. Describe how the octet rule applies to covalent bonds. share one or more pairs of electrons. 9. Illustrate the formation of single, double, and triple covalent bonds using ◗ Sharing one pair, two pairs, and three Lewis structures. pairs of electrons forms single, double, 10. Compare and contrast ionic bonds and covalent bonds. and triple covalent bonds, respectively. 11. Contrast sigma bonds and pi bonds. ◗ Orbitals overlap directly in sigma 12. Apply Create a graph using the bond-dissociation energy data in Table 8.2 bonds. Parallel orbitals overlap in pi and the bond-length data in Table 8.1. Describe the relationship between bond bonds. A single covalent bond is a length and bond-dissociation energy. sigma bond but multiple covalent bonds are made of both sigma and 13. Predict the relative bond-dissociation energies needed to break the bonds in the pi bonds. structures below. ◗ Bond length is measured nucleus-to- a. H — C — — C—H b. H H — — — — nucleus. Bond-dissociation energy is needed to break a covalent bond. C —C H H Self-Check Quiz glencoe.com Section 8.1 The Covalent Bond 247 ©Charles O’Rear/CORBIS Section 8.2 Objectives ◗ Translate molecular formulas into Naming Molecules binary molecular compound names. MAIN Idea Specific rules are used when naming binary molecular ◗ Name acidic solutions. compounds, binary acids, and oxyacids. Review Vocabulary Real-World Reading Link You probably know that your mother’s mother is your grandmother, and that your grandmother’s sister is your great-aunt. oxyanion: a polyatomic ion in which an element (usually a nonmetal) is But what do you call your grandmother’s brother’s daughter? Naming molecules bonded to one or more oxygen atoms requires a set of rules, just as naming family relationships requires rules. New Vocabulary Naming Binary Molecular Compounds oxyacid Many molecular compounds have common names, but they also have scientific names that reveal their composition. To write the formulas and names of molecules, you will use processes similar to those described in Chapter 7 for ionic compounds. Start with a binary molecular compound. Note that a binary molecu- lar compound is composed only of two nonmetal atoms—not metal atoms or ions. An example is dinitrogen monoxide (N 2O), a gaseous anesthetic that is more commonly known as nitrous oxide or laughing gas. The naming of nitrous oxide is explained in the following rules. 1. The first element in the formula is always named first, using the entire element name. N is the symbol for nitrogen. 2. The second element in the formula is named using its root and adding the suffix -ide. O is the symbol for oxygen so the second word is oxide. 3. Prefixes are used to indicate the number of atoms of each element that are present in the compound. Table 8.3 lists the most common prefixes used. There are two atoms of nitrogen and one atom of oxygen, so the first word is dinitrogen and second word is monoxide. There are exceptions to using the prefixes shown in Table 8.3. The first element in the compound name never uses the mono- prefix. For example, CO is carbon monoxide, not monocarbon monoxide. Also, if using a prefix results in two consecutive vowels, one of the vowels is usually dropped to avoid an awkward pronunciation. For example, notice that the oxygen atom in CO is called monoxide, not monooxide. Interactive Table Explore Table 8.3 Prefixes in Covalent Compounds naming covalent compounds at glencoe.com. Number of Atoms Prefix Number of Atoms Prefix 1 mono- 6 hexa- 2 di- 7 hepta- 3 tri- 8 octa- 4 tetra- 9 nona- 5 penta- 10 deca- 248 Chapter 8 Covalent Bonding EXAMPLE Problem 8.2 Naming Binary Molecular Compounds Name the compound P 2O 5, which is used as a drying and dehydrating agent. 1 Analyze the Problem You are given the formula for a compound. The formula contains the elements and the number of atoms of each element in one molecule of the compound. Because only two different elements are present and both are nonmetals, the compound can be named using the rules for naming binary molecular compounds. 2 Solve for the Unknown First, name the elements involved in the compound. phosphorus The first element, represented by P, is phosphorus. oxide The second element, represented by O, is oxygen. Add the suffix –ide to the root of oxygen, ox-. phosphorus oxide Combine the names. Now modify the names to indicate the number of atoms present in a molecule. diphosphorus pentoxide From the formula P 2O 5, you know that two phosphorus atoms and five oxygen atoms make up a molecule of the compound. From Table 8.3, you know that di- is the prefix for two and penta- is the prefix for five. The a in penta- is not used because oxide begins with a vowel. 3 Evaluate the Answer The name diphosphorus pentoxide shows that a molecule of the compound contains two phosphorus atoms and five oxygen atoms, which agrees with the compound’s chemical formula, P 2O 5. PRACTICE Problems Extra Practice Page 979 and glencoe.com Name each of the binary covalent compounds listed below. 14. CO 2 15. SO 2 16. NF 3 17. CCl 4 18. Challenge What is the formula for diarsenic trioxide? Common names for some molecular compounds Have you ever enjoyed an icy, cold glass of dihydrogen monoxide on a hot day? You probably have but you most likely called it by its common name, water. Recall from Chapter 7 that many ionic compounds have common names in addition to their scientific ones. For example, baking soda is sodium hydrogen carbonate and common table salt is sodium chloride. Many binary molecular compounds, such as nitrous oxide and water, were discovered and given common names long before the present-day naming system was developed. Other binary covalent compounds that are generally known by their common names rather than their scientific names are ammonia (NH 3), hydrazine (N 4H 4), and nitric oxide (NO). Reading Check Apply What are the scientific names for ammonia, hydrazine, and nitric oxide? Section 8.2 Naming Molecules 249  Naming Acids Water solutions of some molecules are acidic and are named as acids. Acids are important compounds with specific properties and will be discussed at length in Chapter 18. If a compound produces hydrogen ions (H +) in solution, it is an acid. For example, HCl produces H + in solution and is an acid. Two common types of acids exist—binary acids and oxyacids. Naming binary acids A binary acid contains hydrogen and one other element. The naming of the common binary acid known as hydrochloric acid is explained in the following rules. 1. The first word has the prefix hydro- to name the hydrogen part of the compound. The rest of the first word consists of a form of the root of the second element plus the suffix -ic. HCl (hydrogen and chlorine) becomes hydrochloric. 2. The second word is always acid. Thus, HCl in a water solution is called hydrochloric acid. Although the term binary indicates exactly two elements, a few acids that contain more than two elements are named according to the rules for naming binary acids. If no oxygen is present in the formula for the acidic compound, the acid is named in the same way as a binary acid, except that the root of the second part of the name is the root of the polyatomic ion that the acid contains. For example, HCN, which is composed of hydrogen and the cyanide ion, is called hydrocyanic acid in solution. Naming oxyacids An acid that contains both a hydrogen atom and an oxyanion is referred to as an oxyacid. Recall from Chapter 7 that an oxyanion is a polyatomic ion containing one or more oxygen atoms. The following rules explain the naming of nitric acid (HNO 3), an oxyacid. 1. First, identify the oxyanion present. The first word of an oxyacid’s name consists of the root of the oxyanion and the prefix per- or hypo- if it is part of the name, and a suffix. If the oxyanion’s name ends with the suffix -ate, replace it with the suffix -ic. If the name of the oxyanion ends with the suffix -ite, replace it with the suffix -ous. NO 3, the nitrate ion, becomes nitric. 2. The second word of the name is always acid. HNO 3 (hydrogen and the nitrate ion) becomes nitric acid. Table 8.4 shows how the names of several oxyacids follow these rules. Notice that the hydrogen in an oxyacid is not part of the name. Table 8.4 Naming Oxyacids Compound Oxyanion Acid Suffix Acid Name HClO 3 chlorate -ic chloric acid HClO 2 chlorite -ous chlorous acid HNO 3 nitrate -ic nitric acid HNO 2 nitrite -ous nitrous acid 250 Chapter 8 Covalent Bonding  Interactive Table Explore Formulas and Names of naming covalent compounds Table 8.5 Some Covalent Compounds glencoe.com. Formula Common Name Molecular Compound Name H 2O water dihydrogen monoxide NH 3 ammonia nitrogen trihydride N 2H 4 hydrazine dinitrogen tetrahydride HCl muriatic acid hydrochloric acid C 9H 8O 4 aspirin 2-(acetyloxy)benzoic acid You have learned that naming covalent compounds follows different sets of rules depending on the composition of the compound. Table 8.5 summarizes the formulas and names of several covalent compounds. Note that an acid, whether a binary acid or an oxyacid, can have a com- mon name in addition to its compound name. PRACTICE Problems Extra Practice Page 979 and glencoe.com Name the following acids. Assume each compound is dissolved in water. 19. HI 20. HClO 3 21. HClO 2 22. H 2SO 4 23. H 2S 24. Challenge What is the formula for periodic acid? Writing Formulas from Names The name of a molecular compound reveals its composition and is important in communicating the nature of the compound. Given the name of any binary molecule, you should be able to write the correct chemical formula. The prefixes used in a name indicate the exact num- ber of each atom present in the molecule and determine the subscripts used in the formula. If you are having trouble writing formulas from the names for binary compounds, you might want to review the naming rules listed on pages at the beginning of this section. The formula for an acid can also be derived from the name. It is helpful to remember that all binary acids contain hydrogen and one other element. For oxyacids—acids containing oxyanions—you will need to know the names of the common oxyanions. If you need to review oxyanion names, see Table 7.9 in the previous chapter. PRACTICE Problems Extra Practice Page 979 and glencoe.com Give the formula for each compound. 25. silver chloride 26. dihydrogen oxide 27. chlorine trifluoride 28. diphosphorus trioxide 29. strontium acetate 30. Challenge What is the formula for carbonic acid? Section 8.2 Naming Molecules 251  Look at the Examples: formula of HBr, H2SO3, the molecule. and N02 Does the compound form an acidic aqueous solution? Yes No (H2SO3 and HBr) (NO2) Name the first element in the molecule. Use a prefix if the number Name as an acid. of atoms is greater than one. To Is there an oxygen name the second element, indicate the present in the number present by using a prefix + compound? root of second element + -ide. NO2 is nitrogen dioxide. No Yes (HBr) (H2SO3) Hydro + root of second Root of oxyanion present + -ic element + -ic, then acid. if the anion ends in -ate, or + -ous if the anion ends in -ite, then acid. HBr (aq) is hydrobromic acid. H2SO3 is sulfurous acid. Figure 8.12 Use this flowchart to name molecular compounds when their formulas are known. Apply Which compound above is an oxyacid? Which is a binary acid? The flowchart in Figure 8.12 can help you determine the name of a molecular covalent compound. To use the chart, start at the top and work downward by reading the text contained in the colored boxes and applying it to the formula of the compound you wish to name. Section 8.2 Assessment Section Summary 31. MAIN Idea Summarize the rules for naming binary molecular compounds. ◗ Names of covalent molecular 32. Define a binary molecular compound. compounds include prefixes for the 33. Describe the difference between a binary acid and an oxyacid. number of each atom present. The final letter of the prefix is dropped if 34. Apply Using the system of rules for naming binary molecular compounds, the element name begins with a describe how you would name the molecule N 2O 4. vowel. 35. Apply Write the molecular formula for each of these compounds: iodic acid, ◗ Molecules that produce H+ in solu- disulfur trioxide, dinitrogen monoxide, and hydrofluoric acid. tion are acids. Binary acids contain 36. State the molecular formula for each compound listed below. hydrogen and one other element. a. dinitrogen trioxide d. chloric acid Oxyacids contain hydrogen and an b. nitrogen monoxide e. sulfuric acid oxyanion. c. hydrochloric acid f. sulfurous acid 252 Chapter 8 Covalent Bonding Self-Check Quiz glencoe.com Section 8.3 Objectives ◗ List the basic steps used to draw Molecular Structures Lewis structures. MAIN Idea Structural formulas show the relative positions of ◗ Explain why resonance occurs, and atoms within a molecule. identify resonance structures. Real-World Reading Link As a child, you might have played with plastic ◗ Identify three exceptions to the building blocks that connected only in certain ways. If so, you probably noticed octet rule, and name molecules in which these exceptions occur. that the shape of the object you built depended on the limited ways the blocks interconnected. Building molecules out of atoms works in a similar way. Review Vocabulary ionic bond: the electrostatic force Structural Formulas that holds oppositely charged particles In Chapter 7, you learned about the structure of ionic compounds— together in an ionic compound substances formed from ionic bonds. The covalent molecules you have read about in this chapter have structures that are different from those New Vocabulary of ionic compounds. In studying the molecular structures of covalent structural formula compounds, models are used as representations of the molecule. resonance The molecular formula, which shows the element symbols and coordinate covalent bond numerical subscripts, tells you the type and number of each atom in a molecule. As shown in Figure 8.13, there are several different models that can be used to represent a molecule. Note that in the ball-and-stick and space-filling molecular models, atoms of each specific element are represented by spheres of a representative color, as shown in Table R-1 on page 968. These colors are used for identifying the atoms if the chemical symbol of the element is not present. One of the most useful molecular models is the structural formula, which uses letter symbols and bonds to show relative positions of atoms. You can predict the structural formula for many molecules by drawing the Lewis structure. You have already seen some simple examples of Lewis structures, but more involved structures are needed to help you determine the shapes of molecules. Figure 8.13 All of these models can be used to show the relative locations of atoms and electrons in the phosphorus trihydride (phosphine) molecule. Compare and contrast the types of information contained in each model. PH3 H—P—H Molecular formula H Lewis structure Space-filling molecular model H—P—H H Structural formula Ball-and-stick molecular model Section 8.3 Molecular Structures 253  Lewis structures Although it is fairly easy to draw Lewis structures for most compounds formed by nonmetals, it is a good idea to follow a regular procedure. Whenever you need to draw a Lewis structure, follow the steps outlined in this Problem-Solving Strategy. Problem-Solving Strategy Drawing Lewis Structures 1. Predict the location of certain atoms. The atom that has the least attraction for shared electrons will be the central atom in the molecule. This element is usually the one closer to the left side of the periodic table. The central atom is located in the center of the molecule; all other atoms become terminal atoms. Hydrogen is always a terminal, or end, atom. Because it can share only one pair of electrons, hydrogen can be connected to only one other atom. 2. Determine the number of electrons available for bonding. This number is equal to the total number of valence electrons in the atoms that make up the molecule. 3. Determine the number of bonding pairs. To do this, divide the number of electrons available for bonding by two. 4. Place the bonding pairs. Place one bonding pair (single bond) between the central atom and each of the terminal atoms. 5. Determine the number of bonding pairs remaining. To do this, subtract the number of pairs used in Step 4 from the total number of bonding pairs determined in Step 3. These remaining pairs include lone pairs as well as pairs used in double and triple bonds. Place lone pairs around each terminal atom (except H atoms) bonded to the central atom to satisfy the octet rule. Any remaining pairs will be assigned to the central atom. 6. Determine whether the central atom satisfies the octet rule. Is the central atom surrounded by four electron pairs? If not, it does not satisfy the octet rule. To satisfy the octet rule, convert one or two of the lone pairs on the terminal atoms into a double bond or a triple bond between the terminal atom and the central atom. These pairs are still associated with the terminal atom as well as with the central atom. Remember that carbon, nitrogen, oxygen, and sulfur often form double and triple bonds. Apply the Strategy Study Example Problems 8.3 through 8.5 to see how the steps in the Problem-Solving Strategy are applied. 254 Chapter 8 Covalent Bonding EXAMPLE Problem 8.3 Lewis Structure for a Covalent Compound with Single Bonds Ammonia is a raw material used in the manufacture of many materials, including fertilizers, cleaning products, and explosives. Draw the Lewis structure for ammonia (NH 3). 1 Analyze the Problem Math Handbook Ammonia molecules consist of one nitrogen Dimensional Analysis atom and three hydrogen atoms. Because page 956 hydrogen must be a terminal atom, nitrogen is the central atom. 2 Solve for the Unknown Find the total number of valence electrons available for bonding. 1 N atom × __ + 3 H atoms × __ 5 valence electrons 1 valence electron 1 N atom 1 H atom = 8 valence electrons There are 8 valence electrons available for bonding. __8 electrons = 4 pairs Determine the total number of 2 electrons/pair bonding pairs. To do this, divide the number of available electrons by two. Four pairs of electrons are available for bonding. H—N—H Place a bonding pair (a single bond) — between the central nitrogen atom H and each terminal hydrogen atom. Determine the number of bonding pairs remaining. 4 pairs total - 3 pairs used Subtract the number of pairs used in = 1 pair available these bonds from the total number of pairs of electrons available. The remaining pair—a lone pair—must be added to either the terminal atoms or the central atom. Because hydrogen atoms can have only one bond, they have no lone pairs. H—N—H Place the remaining lone pair on the — central nitrogen atom. H 3 Evaluate the Answer Each hydrogen atom shares one pair of electrons, as required, and the central nitrogen atom shares three pairs of electrons and has one lone pair, providing a stable octet. PRACTICE Problems Extra Practice Page 980 and glencoe.com 37. Draw the Lewis structure for BH 3. 38. Challenge A nitrogen trifluoride molecule contains numerous lone pairs. Draw its Lewis structure. Section 8.3 Molecular Structures 255  EXAMPLE Problem 8.4 Lewis Structure for a Covalent Compound with Multiple Bonds Carbon dioxide is a product of all cellular respiration. Draw the Lewis structure for carbon dioxide (CO 2). 1 Analyze the Problem The carbon dioxide molecule consists of one carbon atom and two oxygen atoms. Because carbon has less attraction for shared electrons, carbon is the central atom, and the two oxygen atoms are terminal. 2 Solve for the Unknown Find the total number of valence electrons available for bonding. 1 C atom × __ + 2 O atoms × __ 4 valence electrons 6 valence electrons 1C atom 1O atom = 16 valence electrons There are 16 valence electrons available for bonding. __16 electrons = 8 pairs Determine the total number of Personal Tutor For an online tutorial on 2 electrons/pair bonding pairs by dividing the number greatest common factors, visit glencoe.com. of available electrons by two. Eight pairs of electrons are available for bonding. Place a bonding pair (a single bond) O—C—O between the central carbon atom and each terminal oxygen atom. Determine the number of bonding pairs remaining. Subtract the number of pairs used in these bonds from the total number of pairs of electrons available. 8 pairs total - 2 pairs used Subtract the number of pairs used in these bonds from the total number of = 6 pairs available pairs of electrons available. O—C—O Add three lone pairs to each terminal oxygen atom. Determine the number of bonding pairs remaining. 6 pairs available - 6 pairs used Subtract the lone pairs from the = 0 pairs available pairs available. Examine the incomplete structure above (showing the placement of the lone pairs). Note that the carbon atom does not have an octet and that there are no more electron pairs available. To give the carbon atom an octet, the molecule must form double bonds. O— —C — —O Use a lone pair from each O atom to form a double bond with the C atom. 3 Evaluate the Answer Both carbon and oxygen now have an octet, which satisfies the octet rule. PRACTICE Problems Extra Practice Page 980 and glencoe.com 39. Draw the Lewis structure for ethylene, C 2H 4. 40. Challenge A molecule of carbon disulfide contains both lone pairs and multiple-covalent bonds. Draw its Lewis structure. 256 Chapter 8 Covalent Bonding Lewis structures for polyatomic ions Although the unit acts as an ion, the atoms within a polyatomic ion are covalently bonded. The procedure for drawing Lewis structures for polyatomic ions is similar to drawing them for covalent compounds. The main difference is in find- ing the total number of electrons available for bonding. Compared to the number of valence electrons present in the atoms that make up the ion, more electrons are present if the ion is negatively charged and fewer are present if the ion is positive. To find the total number of electrons available for bonding, first find the number available in the atoms pres- ent in the ion. Then, subtract the ion charge if the ion is positive, and add the ion charge if the ion is negative. EXAMPLE Problem 8.5 Lewis Structure for a Polyatomic Ion Draw the correct Lewis structure for the polyatomic ion phosphate (PO 4 3-). 1 Analyze the Problem You are given that the phosphate ion consists of one phosphorus atom and four oxygen atoms and has a charge of 3-. Because phosphorus has less attraction for shared electrons than oxygen, phosphorus is the central atom and the four oxygen atoms are terminal atoms. 2 Solve for the Unknown Find the total number of valence electrons available for bonding. Real-World Chemistry Phosphorus and Nitrogen 1 P atom × __ + 4 O atoms × __ 5 valence electrons 6 valence electrons P atom O atom + 3 electrons from the negative charge = 32 valence electrons __ 32 electrons = 16 pair Determine the total number of 2 electrons/pair bonding pairs. O — Draw single bonds from each O—P—O terminal oxygen atom to the central phosphorus atom. — O 16 pairs total - 4 pairs used Subtract the number of pairs = 12 pairs available used from the total number of pairs of electrons available. Add three lone pairs to each terminal oxygen atom. 12 pairs available - 12 lone pairs used = 0 Algal blooms Phosphorus and 3- Subtracting the lone pairs used from nitrogen are nutrients required for O the pairs available verifies that there algae growth. Both can enter lakes — O—P—O are no electron pairs available for the and streams from discharges of phosphorus atom. The Lewis structure — sewage and industrial waste, and in O for the phosphate ion is shown. fertilizer runoff. If these substances build up in a body of water, a rapid 3 Evaluate the Answer growth of algae, known as an algal All of the atoms have an octet, and the group has a net charge of 3-. bloom, can occur, forming a thick layer of green slime over the water’s surface. When the algae use up the PRACTICE Problems Extra Practice Page 980 and glencoe.com supply of nutrients, they die and decompose. This process reduces the 41. Draw the Lewis structure for the NH 4 + ion. amount of dissolved oxygen in the water that is available to other 42. Challenge The ClO 4 - ion contains numerous lone pairs. aquatic organisms. Draw its Lewis structure. Section 8.3 Molecular Structures 257 ©Suzanne Long/Alamy  Figure 8.14 The nitrate ion (NO 3 -) exhibits resonance. a. These resonance a O - b structures differ only in the location of the dou- ble bond. The locations of the nitrogen and oxy- N gen atoms stay the same. b. The actual nitrate O O O - ion is like an average of the three resonance structures in a. The dotted lines indicate possi- N ble locations of the double bond. O - O - O O N N O O O O Resonance Structures Using the same sequence of atoms, it is possible to have more than one correct Lewis structure when a molecule or polyatomic ion has both a double bond and a single bond. Consider the polyatomic ion nitrate (NO 3 -), shown in Figure 8.14a. Three equivalent structures can be used to represent the nitrate ion. Resonance is a condition that occurs when more than one valid VOCABULARY Lewis structure can be written for a molecule or ion. The two or more SCIENCE USAGE V. COMMON USAGE correct Lewis structures that represent a single molecule or ion are Resonance referred to as resonance structures. Resonance structures differ only in Science usage: a phenomenon related the position of the electron pairs, never the atom positions. The location to the stability of a molecule; a large of the lone pairs and bonding pairs differs in resonance structures. The vibration in a mechanical system molecule O 3 and the polyatomic ions NO 3 -, NO 2 -, SO 3 2-, and CO 3 2- caused by a small periodic stimulus commonly form resonance structures. The new molecule had several resonance structures. It is important to note that each molecule or ion that undergoes resonance behaves as if it has only one structure. Refer to Figure 8.14b. Common usage: a quality of Experimentally measured bond lengths show that the bonds are identi- richness or variety cal to each other. They are shorter than single bonds but longer than The sound of the orchestra had double bonds. The actual bond length is an average of the bonds in the resonance. resonance structures. PRACTICE Problems Extra Practice Page 980 and glencoe.com Draw the Lewis resonance structures for the following molecules. 43. NO 2 - 44. SO 2 45. O 3 46. Challenge Draw the Lewis resonance structure for the ion SO 3 2-. Exceptions to the Octet Rule Generally, atoms attain an octet when they bond with other atoms. Figure 8.15 The central nitrogen Some molecules and ions, however, do not obey the octet rule. There atom in this NO 2 molecule does not satisfy the octet rule; the nitrogen atom has only are several reasons for these exceptions. seven electrons in its outer energy level. Odd number of valence electrons First, a small group of mole- Incomplete octet cules might have an odd number of valence electrons and be unable to form an octet around each atom. For example, NO 2 has five valence N electrons from nitrogen and 12 from oxygen, totaling 17, which cannot O O form an exact number of electron pairs. See Figure 8.15. ClO 2 and NO are other examples of molecules with odd numbers of valence electrons. 258 Chapter 8 Covalent Bonding  H H H H Figure 8.16 In this reaction between boron trihydride (BH 3) and ammonia (NH 3), — — — — H—B + N—H → H—B—N—H the nitrogen atom donates both electrons — — — — that are shared by boron and ammonia, H H H H forming a coordinate covalent bond. Interpret Does the coordinate The boron atom has no electrons The nitrogen atom shares covalent bond in the product molecule to share, whereas the nitrogen both electrons to form the satisfy the octet rule? atom has two electrons to share. coordinate covalent bond. Suboctets and coordinate covalent bonds Another exception to the octet rule is due to a few compounds that form suboctets—stable configurations with fewer than eight electrons present around an atom. This group is relatively rare, and BH 3 is an example. Boron, a group 3 nonmetal, forms three covalent bonds with other nonmetallic atoms. H—B—H — H The boron atom shares only six electrons, to few to form an octet. Such compounds tend to be reactive and can share an entire pair of electrons donated by another atom. A coordinate covalent bond forms when one atom donates both of the electrons to be shared with an atom or ion that needs two electrons to form a stable electron arrangement with lower potential energy. Refer to Figure 8.16. Atoms or ions with lone pairs often form coordinate covalent bonds with atoms or ions that need two more electrons. Expanded octets The third group of compounds that does not fol- low the octet rule has central atoms that contain more than eight valence electrons. This electron arrangement is referred to as an expanded octet. An expanded octet can be explained by considering the d orbital that occurs in the energy levels of elements in period three or higher. An example of an expanded octet, shown in Figure 8.17, is the bond forma- tion in the molecule PCl 5. Five bonds are formed with ten electrons shared in one s orbital, three p orbitals, and one d orbital. Another exam- ple is the molecule SF 6, which has six bonds sharing 12 electrons in an s orbital, three p orbitals, and two d orbitals. When you draw the Lewis structure for these compounds, extra lone pairs are added to the central atom or more than four bonding atoms are present in the molecule. Reading Check Summarize three reasons why some molecules do not conform to the octet rule. Figure 8.17 Prior to the reaction of PCl 3 and Cl 2, every reactant atom follows the octet rule. After the reaction, the product, PCl 5, has an expanded octet containing ten electrons. Cl Cl Cl Cl P + Cl Cl P Cl Cl Cl Cl Expanded octet Section 8.3 Molecular Structures 259  EXAMPLE Problem 8.6 Lewis Structure: Exception to the Octet Rule Xenon is a noble gas that will form a few compounds with nonmetals that strongly attract electrons. Draw the correct Lewis structure for xenon tetrafluoride (XeF 4). 1 Analyze the Problem You are given that a molecule of xenon tetrafluoride consists of one xenon atom and four fluorine atoms. Xenon has less attraction for electrons, so it is the central atom. 2 Solve for the Unknown First, find the total number of valence electrons. 1 Xe atom × __ + 4 F atoms × __ = 36 valence electrons 8 valence electrons 7 valence electrons 1Xe atom 1F atom __36 electrons = 18 pairs Determine the total number of bonding pairs. 2 electrons/pair F F Use four bonding pairs to bond the four F atoms to Xe the central Xe atom. F F 18 pairs available - 4 pairs used = 14 pairs available Determine the number of remaining pairs. 3 pairs 14 pairs - 4 F atoms × _ = 2 pairs unused Add three pairs to each F atom to obtain an octet. 1F atom Determine how many pairs remain. F F Xe Place the two remaining pairs on the central Xe atom. F F 3 Evaluate the Answer This structure gives xenon 12 total electrons—an expanded octet—for a total of six bond positions. Xenon compounds, such as the XeF 4 shown here, are toxic because they are highly reactive. PRACTICE Problems Extra Practice Page 980 and glencoe.com Draw the expanded octet Lewis structure for each molecule. 47. ClF 3 48. PCl 5 49. Challenge Draw the Lewis structure for the molecule formed when six fluorine atoms and one sulfur atom bond covalently. Section 8.3 Assessment Section Summary 50. MAIN Idea Describe the information contained in a structural formula. ◗ Different models can be used to 51. State the steps used to draw Lewis structures. represent molecules. 52. Summarize exceptions to the octet rule by correctly pairing these molecules ◗ Resonance occurs when more than and phrases: odd number of valence electrons, PCl 5, ClO 2, BH 3, expanded octet, one valid Lewis structure exists for less than an octet. the same molecule. 53. Evaluate A classmate states that a binary compound having only sigma bonds ◗ Exceptions to the octet rule occur in displays resonance. Could the classmate’s statement be true? some molecules. 54. Draw the resonance structures for the dinitrogen oxide (N 2O) molecule. 55. Draw the Lewis structures for CN -, SiF 4, HCO 3 -, and, AsF 6 -. 260 Chapter 8 Covalent Bonding Self-Check Quiz glencoe.com Section 8.4 Objectives ◗ Summarize the VSEPR Molecular Shapes bonding theory. MAIN Idea The VSEPR model is used to determine molecular ◗ Predict the shape of, and the bond shape. angles in, a molecule. Real-World Reading Link Have you ever rubbed two balloons in your hair ◗ Define hybridization. to create a static electric charge on them? If you brought the balloons together, Review Vocabulary their like charges would cause them to repel each other. Molecular shapes are also affected by the forces of electric repulsion. atomic orbital: the region around an atom’s nucleus that defines an electron’s probable location VSEPR Model The shape of a molecule determines many of its physical and chemical New Vocabulary properties. Often, shapes of reactant molecules determine whether or VSEPR model not they can get close enough to react. Electron densities created by the hybridization overlap of the orbitals of shared electrons determine molecular shape. Theories have been developed to explain the overlap of bonding orbitals and can be used to predict the shape of the molecule. The molecular geometry, or shape, of a molecule can be determined once a Lewis structure is drawn. The model used to determine the molecular shape is referred to as the Valence Shell Electron Pair Repulsion model, or VSEPR model. This model is based on an arrange- ment that minimizes the repulsion of shared and unshared electron pairs around the central atom. Bond angle To understand the VSEPR model better, imagine bal- loons that are inflated to similar sizes and tied together, as shown in Figure 8.18. Each balloon represents an electron-dense region. The repulsive force of this electron-dense region keeps other electrons from entering this space. When a set of balloons is connected at a central point, which represents a central atom, the balloons naturally form a shape that minimizes interactions between the balloons. The electron pairs in a molecule repel one another in a similar way. These forces cause the atoms in a molecule to be positioned at fixed angles relative to one another. The angle formed by two terminal atoms and the central atom is a bond angle. Bond angles predicted by VSEPR are supported by experimental evidence. Unshared pairs of electrons are also important in determining the shape of the molecule. These electrons occupy a slightly larger orbital than shared electrons. Therefore, shared bonding orbitals are pushed together by unshared pairs. Figure 8.18 Electron pairs in a mole- cule are located as far apart as they can be, just as these balloons are arranged. Two pairs form a linear shape. Three pairs form a trigonal planar shape. Four pairs form a tetrahedral shape. Linear Trigonal planar Tetrahedral Section 8.4 Molecular Shapes 261 Matt Meadows  Connection Biology The shape of food molecules is important to our sense of taste. The surface of your tongue is covered with taste buds, each of which contains from 50 to 100 taste receptor cells. Taste receptor cells can detect five distinct tastes—sweet, bitter, salty, sour, and umami (the taste of MSG, monosodium glutamate)—but each receptor cell responds best to only one taste. VOCABULARY The shapes of food molecules are determined by their chemical WORD ORIGIN structures. When a molecule enters a taste bud, it must have the correct Trigonal planar shape for the nerve in each receptor cell to respond and send a message comes from the Latin words to the brain. The brain then interprets the message as a certain taste. trigonum, which means triangular, When such molecules bind to sweet receptors, they are sensed as sweet. and plan-, which means flat The greater the number of food molecules that fit a sweet receptor cell, the sweeter the food tastes. Sugars and artificial sweeteners are not the only sweet molecules. Some proteins found in fruits are also sweet mol- ecules. Some common molecular shapes are illustrated in Table 8.5. Hybridization A hybrid occurs when two things are combined and the result has char- acteristics of both. For example, a hybrid automobile uses both gas and electricity as energy sources. During chemical bonding, different atomic Figure 8.19 A carbon atom’s 2s and 2p electrons occupy the hybrid sp 3 orbitals undergo hybridization. To understand this, consider the bond- orbitals. Notice that the hybrid orbitals ing involved in the methane molecule (CH 4). The carbon atom has four have an intermediate amount of potential valence electrons with the electron configuration[He]2s 22p 2. You might energy when compared with the energy of expect the two unpaired p electrons to bond with other atoms and the the original s and p orbitals. According to 2s electrons to remain an unshared pair. However, carbon atoms under- VSEPR theory, a tetrahedral shape mini- mizes repulsion between the hybrid orbit- go hybridization, a process in which atomic orbitals mix and form new, als in a CH 4 molecule. identical hybrid orbitals. Identify How many faces does the The hybrid orbitals in a carbon atom are shown in Figure 8.19. tetrahedral shape formed by the sp 3 Note that each hybrid orbital contains one electron that it can share orbitals have? with another atom. The hydrid orbital is called an sp 3 orbital because the four hybrid orbitals form from one s orbital and three p orbitals. H Carbon is the most common element that undergoes hybridization. sp3 The number of atomic orbitals that mix and form the hybrid orbital equals the total number of pairs of electrons, as shown in Table 8.5. In