CHAPTER ONE Chemical Bonding PROBLEM 1.3 Which of the following ions possess a noble gas electron config- uration? (a)K (c)H (f)Ca2 SAMPLE SOLUTION (a)Potassium has atomic number 19, and so a potassium atom has 19 electrons. The ion k therefore. has 18 electrons the same as the noble gas argon. The electron configurations of k and Ar are the same 252p63523p Transfer of an electron from a sodium atom to a chlorine atom yields a sodium cation and a chloride anion, both of which have a noble gas electron configuration Na(g) ci(g) Sodium atom Chlorine atom Were we to simply add the ionization energy of sodium (496 kJ/ mol) and the electron affinity of chlorine(-349 kJ/mol), we would conclude that the overall process is endothermic with AH =+147 kJ/mol. The energy liberated by adding an electron to by the German physicist Wal. chlorine is insufficient to override the energy required to remove an electron from ter Kossel in 1916, in order sodium. This analysis, however, fails to consider the force of attraction between the to explain the ability of sub- oppositely charged ions Na and CI, which exceeds 500 kJ/mol and is more than suf- de to conduct an electric ficient to make the overall process exothermic. Attractive forces between oppositely charged particles are termed electrostatic, or coulombic, attractions and are what we mean by an ionic bond between two atoms PROBLEM 1. 4 What is the electron configuration of c of c Does either one of these ions have a noble gas(closed-shell)electron configuration? onic bonds are very common in inorganic compounds, but rare in organic ones. The ionization energy of carbon is too large and the electron affinity too small for car- bon to realistically form a C4+ or C4- ion. What kinds of bonds, then, link carbon to other elements in millions of organic compounds? Instead of losing or gaining electrons, carbon shares electrons with other elements (including other carbon atoms) to give wha are called covalent bonds 1.3 COVALENT BONDS The covalent, or shared electron pair, model of chemical bonding was first suggested Gilbert Newton Lewis (born by G. N. Lewis of the University of California in 1916. Lewis proposed that a sharing of two electrons by two hydrogen atoms permits each one to have a stable closed-shell eley, Califor- electron configuration analogous to helium nia, 1946)has been called he greatest American 984is H: H ue of the journal of chemi. cal Education contains fi Two hydrogen atoms. Hydrogen molecule: articles describing Lewis'life each with a single alent bonding by way of and contributions to chem. electron a shared electron pair Back Forward Main MenuToc Study Guide ToC Student o MHHE Website
PROBLEM 1.3 Which of the following ions possess a noble gas electron configuration? (a) K (d) O (b) He (e) F (c) H (f) Ca2 SAMPLE SOLUTION (a) Potassium has atomic number 19, and so a potassium atom has 19 electrons. The ion K, therefore, has 18 electrons, the same as the noble gas argon. The electron configurations of K and Ar are the same: 1s 2 2s 2 2p6 3s 2 3p6 . Transfer of an electron from a sodium atom to a chlorine atom yields a sodium cation and a chloride anion, both of which have a noble gas electron configuration: Were we to simply add the ionization energy of sodium (496 kJ/mol) and the electron affinity of chlorine (349 kJ/mol), we would conclude that the overall process is endothermic with H° 147 kJ/mol. The energy liberated by adding an electron to chlorine is insufficient to override the energy required to remove an electron from sodium. This analysis, however, fails to consider the force of attraction between the oppositely charged ions Na and Cl– , which exceeds 500 kJ/mol and is more than suf- ficient to make the overall process exothermic. Attractive forces between oppositely charged particles are termed electrostatic, or coulombic, attractions and are what we mean by an ionic bond between two atoms. PROBLEM 1.4 What is the electron configuration of C? Of C? Does either one of these ions have a noble gas (closed-shell) electron configuration? Ionic bonds are very common in inorganic compounds, but rare in organic ones. The ionization energy of carbon is too large and the electron affinity too small for carbon to realistically form a C4 or C4 ion. What kinds of bonds, then, link carbon to other elements in millions of organic compounds? Instead of losing or gaining electrons, carbon shares electrons with other elements (including other carbon atoms) to give what are called covalent bonds. 1.3 COVALENT BONDS The covalent, or shared electron pair, model of chemical bonding was first suggested by G. N. Lewis of the University of California in 1916. Lewis proposed that a sharing of two electrons by two hydrogen atoms permits each one to have a stable closed-shell electron configuration analogous to helium. H Two hydrogen atoms, each with a single electron H Hydrogen molecule: covalent bonding by way of a shared electron pair H H Na(g) ±£ Sodium atom NaCl(g) Sodium chloride Cl(g) Chlorine atom 12 CHAPTER ONE Chemical Bonding Ionic bonding was proposed by the German physicist Walter Kossel in 1916, in order to explain the ability of substances such as sodium chloride to conduct an electric current. Gilbert Newton Lewis (born Weymouth, Massachusetts, 1875; died Berkeley, California, 1946) has been called the greatest American chemist. The January 1984 issue of the Journal of Chemical Education contains five articles describing Lewis’ life and contributions to chemistry. Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website
1.3 Covalent Bonds Structural formulas of this type in which electrons are represented as dots are called Lewis structures The amount of energy required to dissociate a hydrogen molecule H2 to two sep arate hydrogen atoms is called its bond dissociation energy (or bond energy). For H it is quite large, being equal to 435 k/mol(104 kcal/mol). The main contributor to the strength of the covalent bond in H, is the increased binding force exerted on its two electrons. Each electron in H,"feels"the attractive force of two nuclei, rather than one as it would in an isolated hydrogen atom Covalent bonding in F2 gives each fluorine 8 electrons in its valence shell and a stable electron configuration equivalent to that of the noble gas neon: Two fluorine atoms. each with seven electrons in nding by way of ts valence shell electron pai PROBLEM 1.5 Hydrogen is bonded to fluorine in hydrogen fluoride by a cova- lent bond. Write a Lewis formula for hydrogen fluoride The Lewis model limits second-row elements(Li, Be, B, C, N, O, F, Ne) to a total of 8 electrons(shared plus unshared) in their valence shells. Hydrogen is limited to 2. Most of the elements that we'll encounter in this text obey the octet rule: in forming compounds they gain, lose, or share electrons to give a stable electron configuration characterized by eight valence electrons. When the octet rule is satisfied for carbon nitrogen, oxygen, and fluorine, they have an electron configuration analogous to the noble gas neon Now lets apply the Lewis model to the organic compounds methane and carbon tetrafluoride to write a Combine. C. and four h Lewis structure H: C: H for methane H Combine·C· and four Lewis structure F: C: F for carbon tetrafluoride Carbon has electrons in its valence shell in both methane and carbon tetrafluoride. By forming covalent bonds to four other atoms, carbon achieves a stable electron configu ration analogous to neon. Each covalent bond in methane and carbon tetrafluoride is quite strong--comparable to the bond between hydrogens in H2 in bond dissociation energy PROBLEM 1.6 Given the information that it has a carbon -carbon bond write a satisfactory Lewis structure for C2H(ethane) Representing a 2-electron covalent bond by a dash () the Lewis structures for hydrogen fluoride, fluorine, methane, and carbon tetrafluoride become Back Forward Main MenuToc Study Guide ToC Student o MHHE Website
Structural formulas of this type in which electrons are represented as dots are called Lewis structures. The amount of energy required to dissociate a hydrogen molecule H2 to two separate hydrogen atoms is called its bond dissociation energy (or bond energy). For H2 it is quite large, being equal to 435 kJ/mol (104 kcal/mol). The main contributor to the strength of the covalent bond in H2 is the increased binding force exerted on its two electrons. Each electron in H2 “feels” the attractive force of two nuclei, rather than one as it would in an isolated hydrogen atom. Covalent bonding in F2 gives each fluorine 8 electrons in its valence shell and a stable electron configuration equivalent to that of the noble gas neon: PROBLEM 1.5 Hydrogen is bonded to fluorine in hydrogen fluoride by a covalent bond. Write a Lewis formula for hydrogen fluoride. The Lewis model limits second-row elements (Li, Be, B, C, N, O, F, Ne) to a total of 8 electrons (shared plus unshared) in their valence shells. Hydrogen is limited to 2. Most of the elements that we’ll encounter in this text obey the octet rule: in forming compounds they gain, lose, or share electrons to give a stable electron configuration characterized by eight valence electrons. When the octet rule is satisfied for carbon, nitrogen, oxygen, and fluorine, they have an electron configuration analogous to the noble gas neon. Now let’s apply the Lewis model to the organic compounds methane and carbon tetrafluoride. Carbon has 8 electrons in its valence shell in both methane and carbon tetrafluoride. By forming covalent bonds to four other atoms, carbon achieves a stable electron configuration analogous to neon. Each covalent bond in methane and carbon tetrafluoride is quite strong—comparable to the bond between hydrogens in H2 in bond dissociation energy. PROBLEM 1.6 Given the information that it has a carbon–carbon bond, write a satisfactory Lewis structure for C2H6 (ethane). Representing a 2-electron covalent bond by a dash (—), the Lewis structures for hydrogen fluoride, fluorine, methane, and carbon tetrafluoride become: Combine to write a Lewis structure for methane and fourC H CH H H H Combine to write a Lewis structure for carbon tetrafluoride and fourC F F F F F C Fluorine molecule: covalent bonding by way of a shared electron pair F F Two fluorine atoms, each with seven electrons in its valence shell F F 1.3 Covalent Bonds 13 Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website
CHAPTER ONE Chemical Bonding F:H-C一H:F一C一F: H Methane Carbon tetrafluoride 1.4 DOUBLE BONDS AND TRIPLE BONDS Lewis's concept of shared electron pair bonds allows for 4-electron double bonds and 6-electron triple bonds. Carbon dioxide(CO2)has two carbon-oxygen double bonds and the octet rule is satisfied for both carbon and oxygen. Similarly, the most stable Lewis structure for hydrogen cyanide(HCN) has a carbon-nitrogen triple bor Hy drogen cyanide H:C∷:N H一 Multiple bonds are very common in organic chemistry. Ethylene( C2H4)contains a carbon-carbon double bond in its most stable lewis structure. and each carbon has a ipleted octet. The most stable Lewis structure for acetylene( C2H2) contains a car- bon-carbon triple bond. Here again, the octet rule is satisfied. H H H: C: C: H H一C≡C-H PROBLEM 1.7 Write the most stable Lewis structure for each of the following (a)Formaldehyde, CH,O. Both hydrogens are bonded to carbon. (A solution of hyde in water is sometimes used to Biological specimens. (b)Tetrafluoroethylene, C2F4. (The starting material for the preparation of Teflon. (c) Acrylonitrile, C3HaN. The atoms are connected in the order CCCN, and all hydrogens are bonded to carbon. (The starting material for the preparation of acrylic fibers such as Orlon and Acrilan) SAMPLE SoLUTION (a)Each hydrogen contributes 1 valence carbon contributes 4, and oxygen 6 for a total of 12 valence electrons told that both hydrogens are bonded to carbon. Since carbon forms four n its sta- ble compounds, join carbon and oxygen by a double bond. The partial structure so generated accounts for 8 of the 12 electrons. Add the remaining four electrons to oxygen as unshared pairs to complete the structure of formaldehyd H Partial structure showing Complete Lewis structure Back Forward Main MenuToc Study Guide ToC Student o MHHE Website
1.4 DOUBLE BONDS AND TRIPLE BONDS Lewis’s concept of shared electron pair bonds allows for 4-electron double bonds and 6-electron triple bonds. Carbon dioxide (CO2) has two carbon–oxygen double bonds, and the octet rule is satisfied for both carbon and oxygen. Similarly, the most stable Lewis structure for hydrogen cyanide (HCN) has a carbon–nitrogen triple bond. Multiple bonds are very common in organic chemistry. Ethylene (C2H4) contains a carbon–carbon double bond in its most stable Lewis structure, and each carbon has a completed octet. The most stable Lewis structure for acetylene (C2H2) contains a carbon–carbon triple bond. Here again, the octet rule is satisfied. PROBLEM 1.7 Write the most stable Lewis structure for each of the following compounds: (a) Formaldehyde, CH2O. Both hydrogens are bonded to carbon. (A solution of formaldehyde in water is sometimes used to preserve biological specimens.) (b) Tetrafluoroethylene, C2F4. (The starting material for the preparation of Teflon.) (c) Acrylonitrile, C3H3N. The atoms are connected in the order CCCN, and all hydrogens are bonded to carbon. (The starting material for the preparation of acrylic fibers such as Orlon and Acrilan.) SAMPLE SOLUTION (a) Each hydrogen contributes 1 valence electron, carbon contributes 4, and oxygen 6 for a total of 12 valence electrons. We are told that both hydrogens are bonded to carbon. Since carbon forms four bonds in its stable compounds, join carbon and oxygen by a double bond. The partial structure so generated accounts for 8 of the 12 electrons. Add the remaining four electrons to oxygen as unshared pairs to complete the structure of formaldehyde. Partial structure showing covalent bonds O X C O X C H H ± ± Complete Lewis structure of formaldehyde H H ± ± orEthylene: C H H H H C CœC H H H H ± ± ± ± H H C orAcetylene: H±CPC±HC O orCarbon dioxide: O C OœCœO H N H orHydrogen cyanide: C ±CPN H±C±H H W W H Methane Carbon tetrafluoride F±C±F F W W F Hydrogen fluoride H±F Fluorine F±F 14 CHAPTER ONE Chemical Bonding Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website
1.5 Polar Covalent Bonds and Electronegativity 1.5 POLAR COVALENT BONDS AND ELECTRONEGATIVITY Electrons in covalent bonds are not necessarily shared equally by the two atoms that they connect. If one atom has a greater tendency to attract electrons toward itself than the other we say the electron distribution is polarized, and the bond is referred to as a polar cova lent bond. Hydrogen fluoride, for example, has a polar covalent bond. Because fluorine attracts electrons more strongly than hydrogen, the electrons in the H-F bond are pulled toward fluorine, giving it a partial negative charge, and away from hydrogen giving it a partial positive charge. This polarization of electron density is represented in various ways 6H—F6- H一F CThe symbols and (The symbol +represents indicate partial positive the direction of polarization and partial negativ of electrons in the h-F bond) y of an atom to draw the electrons in a covalent bond toward itself is referred to electronegativity. An electronegative element attracts electrons: an electropositive one donates them. Electronegativity increases across a row in the peri- odic table. The most electronegative of the second-row elements is fluorine; the most electropositive is lithium Electronegativity decreases in going down a column. Fluorine is more electronegative than chlorine. The most commonly cited electronegativity scale was devised by Linus Pauling and is presented in Table 1. 2. PROBLEM 1.8 Examples of carbon-containing compounds include methane( CHa) chloromethane(CHaCi), and methyllithium(CH3 Li). In which one does carbon bear the greatest partial positive charge? The greatest partial negative charge? Technology, where he Centers of positive and negative charge that are separated from each other consti- in 1925. In addition to re tute a dipole. The dipole moment u of a molecule is equal to the charge e(either the pauling studied the structure between the centers of charge: the Nobel Prize in chemistry for that work in 1954. Paul efforts to limit the testing of uclear weapons. He was TABLE 1.2 Selected Values from the Pauling Electronegativity Scale ive won two Nobel prizes a woman. Can you name Group number Period H2u B C N 3.0 18 0.8 1.0 2.8 2.5 Back Forward Main MenuToc Study Guide ToC Student o MHHE Website
1.5 POLAR COVALENT BONDS AND ELECTRONEGATIVITY Electrons in covalent bonds are not necessarily shared equally by the two atoms that they connect. If one atom has a greater tendency to attract electrons toward itself than the other, we say the electron distribution is polarized, and the bond is referred to as a polar covalent bond. Hydrogen fluoride, for example, has a polar covalent bond. Because fluorine attracts electrons more strongly than hydrogen, the electrons in the H±F bond are pulled toward fluorine, giving it a partial negative charge, and away from hydrogen giving it a partial positive charge. This polarization of electron density is represented in various ways. The tendency of an atom to draw the electrons in a covalent bond toward itself is referred to as its electronegativity. An electronegative element attracts electrons; an electropositive one donates them. Electronegativity increases across a row in the periodic table. The most electronegative of the second-row elements is fluorine; the most electropositive is lithium. Electronegativity decreases in going down a column. Fluorine is more electronegative than chlorine. The most commonly cited electronegativity scale was devised by Linus Pauling and is presented in Table 1.2. PROBLEM 1.8 Examples of carbon-containing compounds include methane (CH4), chloromethane (CH3Cl), and methyllithium (CH3Li). In which one does carbon bear the greatest partial positive charge? The greatest partial negative charge? Centers of positive and negative charge that are separated from each other constitute a dipole. The dipole moment of a molecule is equal to the charge e (either the positive or the negative charge, since they must be equal) multiplied by the distance between the centers of charge: e d (The symbols and indicate partial positive and partial negative charge, respectively) H±F H±F (The symbol represents the direction of polarization of electrons in the H±F bond) 1.5 Polar Covalent Bonds and Electronegativity 15 TABLE 1.2 Selected Values from the Pauling Electronegativity Scale Group number Period 1 2 3 4 5 I H 2.1 Li 1.0 Na 0.9 K 0.8 II Be 1.5 Mg 1.2 Ca 1.0 III B 2.0 Al 1.5 IV C 2.5 Si 1.8 V N 3.0 P 2.1 VI O 3.5 S 2.5 VII F 4.0 Cl 3.0 Br 2.8 I 2.5 Linus Pauling (1901–1994) was born in Portland, Oregon and was educated at Oregon State University and at the California Institute of Technology, where he earned a Ph.D. in chemistry in 1925. In addition to research in bonding theory, Pauling studied the structure of proteins and was awarded the Nobel Prize in chemistry for that work in 1954. Pauling won a second Nobel Prize (the Peace Prize) for his efforts to limit the testing of nuclear weapons. He was one of only four scientists to have won two Nobel Prizes. The first double winner was a woman. Can you name her? Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website
CHAPTER ONE Chemical Bonding Because the charge on an electron is 4.80X 10 electrostatic units(esu) and tances within a molecule typically fall in the 10 cm range, molecular dipole are on the order of 10 8 esu- cm. In order to simplify the reporting of dipole this value of 10 esu' cm is defined as a debye, D. Thus the experimentally determined honor of Peter Debye, a dipole moment of hydrogen fluoride, 1.7 X 10esucm is stated as 1.7 Dutch scientist who did Table 1.3 lists the dipole moments of various bond types. For H-F, H-Cl, ortant work in many H-Br, and H-I these"bond dipoles"are really molecular dipole moments. A polar was awarded the Nobel Prize molecule has a dipole moment, a nonpolar one does not. Thus, all of the hydrogen halides are polar molecules. In order to be polar, a molecule must have polar bonds, bu can't have a shape that causes all the individual bond dipoles to cancel. We will have more to say about this in Section 1. ll after we have developed a feeling for the three- dimensional shapes of molecules The bond dipoles in Table 1.3 depend on the difference in electronegativity of the bonded atoms and on the bond distance. The polarity of a C-H bond is relatively low; substantially less than a C-O bond, for example. Don't lose sight of an even more important difference between a C-H bond and a C-O bond, and that is the direction of the dipole moment. In a C-H bond the electrons are drawn away from H, toward C In a C-o bond, electrons are drawn from C toward O. As we'll see in later chap ters, the kinds of reactions that a substance undergoes can often be related to the size and direction of key bond dipoles 1.6 FORMAL CHARGE Lewis structures frequently contain atoms that bear a positive or negative charge. molecule as a whole is neutral, the sum of its positive charges must equal the sum of its negative charges. An example is nitric acid, HNO As written, the structural formula for nitric acid depicts different bonding patterns for its three oxygens. One oxygen is doubly bonded to nitrogen, another is singly bonded TABLE 1.3 Selected Bond Dipole Moments Bond* Dipole moment, D Bond* Dipole moment, D H—F C-F H—C H—C C≡N 3.6 irection of the dipole moment is toward the more electronegative atom In the listed examples en and carbon are the positive ends of the dipoles. Carbon is the negative end of the dipole ted with the C-H bond Back Forward Main MenuToc Study Guide ToC Student o MHHE Website
Because the charge on an electron is 4.80 1010 electrostatic units (esu) and the distances within a molecule typically fall in the 108 cm range, molecular dipole moments are on the order of 1018 esu·cm. In order to simplify the reporting of dipole moments this value of 1018 esu cm is defined as a debye, D. Thus the experimentally determined dipole moment of hydrogen fluoride, 1.7 1018 esu cm is stated as 1.7 D. Table 1.3 lists the dipole moments of various bond types. For H±F, H±Cl, H±Br, and H±I these “bond dipoles” are really molecular dipole moments. A polar molecule has a dipole moment, a nonpolar one does not. Thus, all of the hydrogen halides are polar molecules. In order to be polar, a molecule must have polar bonds, but can’t have a shape that causes all the individual bond dipoles to cancel. We will have more to say about this in Section 1.11 after we have developed a feeling for the threedimensional shapes of molecules. The bond dipoles in Table 1.3 depend on the difference in electronegativity of the bonded atoms and on the bond distance. The polarity of a C±H bond is relatively low; substantially less than a C±O bond, for example. Don’t lose sight of an even more important difference between a C±H bond and a C±O bond, and that is the direction of the dipole moment. In a C±H bond the electrons are drawn away from H, toward C. In a C±O bond, electrons are drawn from C toward O. As we’ll see in later chapters, the kinds of reactions that a substance undergoes can often be related to the size and direction of key bond dipoles. 1.6 FORMAL CHARGE Lewis structures frequently contain atoms that bear a positive or negative charge. If the molecule as a whole is neutral, the sum of its positive charges must equal the sum of its negative charges. An example is nitric acid, HNO3: As written, the structural formula for nitric acid depicts different bonding patterns for its three oxygens. One oxygen is doubly bonded to nitrogen, another is singly bonded H±O±N O O œ ± 16 CHAPTER ONE Chemical Bonding TABLE 1.3 Selected Bond Dipole Moments Bond* H±F H±Cl H±Br H±I H±C H±N H±O Dipole moment, D 1.7 1.1 0.8 0.4 0.3 1.3 1.5 Bond* C±F C±O C±N CœO CœN CPN Dipole moment, D 1.4 0.7 0.4 2.4 1.4 3.6 *The direction of the dipole moment is toward the more electronegative atom. In the listed examples hydrogen and carbon are the positive ends of the dipoles. Carbon is the negative end of the dipole associated with the C±H bond. The debye unit is named in honor of Peter Debye, a Dutch scientist who did important work in many areas of chemistry and physics and was awarded the Nobel Prize in chemistry in 1936. Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website