8aeierieo released energy is very often harnessed to do work,such as the burning of gasoline to mical reactions carried ou se of Ground state common The lowest energy state of a system. bon.the in accordance with the quantum chemist lowest energy form of carbon.If we place the electrons in a differe xamp ouldmanner (as an Excited state enerystate:when the electrons are rearranged back to this ground state,energy is released. ud in the low enengy orial 15 reghte rgest amount of en T First ionization caled the otential.The s electrons.therefore.have the highest ioniza- tion potential;however,the electrons in the 2p levels of carbon are the farthest from he cleus and are hel and ther d the weakest.They are e n the his is C.Lewis Dot Structures Chemists often focus on the electrons in the outermost shell of the atom because sare involved in the formation of chemical bonds and in che has fou or the ground- state ele ron cled va and the energy level in which they are found is called the valence shell.To illus- trate the outermost electrons of an atom,chemists commonly use a representation ure,named after the American chemist Gilb rt N.Lewi Lewis dot structure who devis al to th ow tel the of dots of an atom of that element.In Lewis dot structures.the atomic resents the core (i.e.,the nucleus and all inner shell electrons).Table 1.4 shows Lewis dot struc- e valence shell of 2 of th Clin period3 of the Periodic Table.the valence electrons belong to the third shell This shell is only partially filled with eight electrons;the 3s and 3p orbitals are fully occu- pied,but the five 3d orbitals can accommodate an additional ten electrons. Table 1.4 Lewis Dot Structures for Elements 1-18 1A 2A 3A 4A 5A 6A 7A 8A Gilbert N.Lewis (1875-1946) H He: introduced the theory of the electron pair that extended .Be:B:C:N:::Ne: our und standing of covalent aMgt:·:p::S::CI::: honor that we often refer to an "electron dot"structure as a Lewis structure.Bettmann/CORBIS 6
Unless otherwise noted all art on this page © Cengage Learning 2013 Chapter 1 Covalent Bonding and Shapes of Molecules 6 released energy is very often harnessed to do work, such as the burning of gasoline to drive the pistons in an internal combustion engine. In chemical reactions carried out in the laboratory, the release of energy commonly just heats up the reaction vessel. Let’s return to the energy-level diagram of carbon. In the ground state of carbon, the electrons are placed in accordance with the quantum chemistry principles (e.g., Aufbau principle, Hund’s rule, and Pauli exclusion principle) that dictate the lowest energy form of carbon. If we place the electrons in a different manner (as an example, only one electron in 2s and three electrons in 2p), we would have a higher energy state of carbon, referred to as an excited state. All of nature seeks its lowest energy state; when the electrons are rearranged back to this ground state, energy is released. Note that the electrons in the lowest energy orbital, 1s, are held tightest to the nucleus. It would take the largest amount of energy to remove these electrons relative to the others. The energy it takes to remove an electron from an atom or a molecule is called the ionization potential. The 1s electrons, therefore, have the highest ionization potential; however, the electrons in the 2p levels of carbon are the farthest from the nucleus and are held the weakest. They are the easiest to remove from the atom and therefore have the lowest ionization potential. This is analogous to it being easier to remove an object from the earth the farther it is from the surface. C. Lewis Dot Structures Chemists often focus on the electrons in the outermost shell of the atom because these electrons are involved in the formation of chemical bonds and in chemical reactions. Carbon, for example, with the ground-state electron configuration 1s2 2s2 2p2, has four outer-shell electrons. Outer-shell electrons are called valence electrons, and the energy level in which they are found is called the valence shell. To illustrate the outermost electrons of an atom, chemists commonly use a representation called a Lewis dot structure, named after the American chemist Gilbert N. Lewis (1875–1946), who devised it. A Lewis dot structure shows the symbol of the element surrounded by the number of dots equal to the number of electrons in the outer shell of an atom of that element. In Lewis dot structures, the atomic symbol represents the core (i.e., the nucleus and all inner shell electrons). Table 1.4 shows Lewis dot structures for the first 18 elements of the Periodic Table. The noble gases helium and neon have filled valence shells. The valence shell of helium is filled with two electrons; that of neon is filled with eight electrons. Neon and argon have in common an electron configuration in which the s and p orbitals of their valence shells are filled with eight electrons. The valence shells of all other elements shown in Table 1.4 contain fewer than eight electrons. For C, N, O, and F in period 2 of the Periodic Table, the valence electrons belong to the second shell. With eight electrons, this shell is completely filled. For Si, P, S, and Cl in period 3 of the Periodic Table, the valence electrons belong to the third shell. This shell is only partially filled with eight electrons; the 3s and 3p orbitals are fully occupied, but the five 3d orbitals can accommodate an additional ten electrons. Table 1.4 Lewis Dot Structures for Elements 1–18* H He 1A 2A 3A 4A 5A 6A 7A 8A Be Mg B Al Ne Ar C Si Li Na N P O S F Cl *These dots represent electrons from the valence shell. They are arranged as pairs or single electrons in accordance with Hund’s rule. Ground state The lowest energy state of a system. Excited state A state of a system at higher energy than the ground state. First ionization potential The energy needed to remove the most loosely held electron from an atom or a molecule. Valence electrons Electrons in the valence (outermost) shell of an atom. Valence shell The outermost occupied electron shell of an atom. Lewis dot structure The symbol of an element surrounded by the number of dots equal to the number of electrons in the valence shell of the atom. Gilbert N. Lewis (1875–1946) introduced the theory of the electron pair that extended our understanding of covalent bonding and of the concept of acids and bases. It is in his honor that we often refer to an “electron dot” structure as a Lewis structure. © Bettmann/CORBIS Copyright 2013 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it
1.2 Lewis Model of Bonding 2 Lewis Model of Bonding ing and rea ou ns of thes valence shell of eight electrons (2p),and argon with a valence shell of eight electrons (3s3p).The tendency of atoms to react in ways that achieve an outer shell of eight va ectrons is partic Example 1.2 The Octet Rule valence electrons. Show how the loss of an electron from a sodium atom leads to a stable octet Solution The ground-state electron configurations for Na and Na are Na (11 electrons):123 Na*(10 electrons):1 Thus,Na has a complete octet of electrons in its outermost(valence)shell and has the same electron configuration as neon,the noble gas nearest it in atomic number. Problem 1.2 Show how each chemical change leads to a stable octet. (a)Sulfur formsS. (b)Magnesium forms Mg?+ A.Formation of Chemical Bonds According to Lewis's model,atoms interact in such a way that each participating atom acquires a completed outer-shell electron configuration resembling that of the noble gas nearest to it in atomic number.Atoms acquire completed valence shells in two ways. 1.An atom may become ionic(ie,lose or gain enough electrons to acquire a com- pletely filled valence shell).An atom that gains electrons becomes an anion Anior (a negatively charged ion),and an atom that (a positiv ely c arge n and ely char ach vice versa,in a definite geometric arrangement that depends on the crystal.When ionic in 2.An atom as an I lete its valence shell.A chemical bond formed by sharing electrons is called a 3. these bond Covalent bond s are called polar y in the next sectio B.Electronegativity and Chemical Bonds How do we estimate the degree of ionic or covalent character in a chemical bond?One way is to compare the electronegativities of the atoms involved.Electronegativity is Unless otherwise noted all art on this page Cengage Leaming 2013
Unless otherwise noted all art on this page © Cengage Learning 2013 7 1.2 Lewis Model of Bonding 1.2 Lewis Model of Bonding In 1916, Lewis devised a beautifully simple model that unified many of the observations about chemical bonding and reactions of the elements. He pointed out that the chemical inertness of the noble gases indicates a high degree of stability of the electron configurations of these elements: helium with a valence shell of two electrons (1s2 ), neon with a valence shell of eight electrons (2s2 2p6 ), and argon with a valence shell of eight electrons (3s2 3p6 ). The tendency of atoms to react in ways that achieve an outer shell of eight valence electrons is particularly common among second-row elements of Groups 1A–7A (the main-group elements) and is given the special name octet rule. Example 1.2 | The Octet Rule Show how the loss of an electron from a sodium atom leads to a stable octet. Solution The ground-state electron configurations for Na and Na1 are: Na (11 electrons): 1s2 2s2 2p6 3s1 Na1 (10 electrons): 1s2 2s2 2p6 Thus, Na1 has a complete octet of electrons in its outermost (valence) shell and has the same electron configuration as neon, the noble gas nearest it in atomic number. Problem 1.2 Show how each chemical change leads to a stable octet. (a) Sulfur forms S2–. (b) Magnesium forms Mg21. A. Formation of Chemical Bonds According to Lewis’s model, atoms interact in such a way that each participating atom acquires a completed outer-shell electron configuration resembling that of the noble gas nearest to it in atomic number. Atoms acquire completed valence shells in two ways. 1. An atom may become ionic (i.e., lose or gain enough electrons to acquire a completely filled valence shell). An atom that gains electrons becomes an anion (a negatively charged ion), and an atom that loses electrons becomes a cation (a positively charged ion). A positively charged ion and a negatively charged ion attract each other. This attraction can lead to the formation of ionic crystals such as sodium chloride, in which each positive ion is surrounded by negative ions and vice versa, in a definite geometric arrangement that depends on the crystal. When atoms are held together primarily by attraction of oppositely charged ions, we say that an ionic interaction exists between them. (This ionic interaction is often referred to as an ionic bond.) 2. An atom may share electrons with one or more other atoms to complete its valence shell. A chemical bond formed by sharing electrons is called a covalent bond. 3. Bonds may be partially ionic and partially covalent; these bonds are called polar covalent bonds. Polar covalent bonds are defined more precisely in the next section. B. Electronegativity and Chemical Bonds How do we estimate the degree of ionic or covalent character in a chemical bond? One way is to compare the electronegativities of the atoms involved. Electronegativity is Octet rule Group 1A–7A elements react to achieve an outer shell of eight valence electrons. Anion An atom or a group of atoms bearing a negative charge. Cation An atom or a group of atoms bearing a positive charge. Ionic interaction Attraction between oppositely charged ions. Covalent bond A chemical bond formed between two atoms by sharing one or more pairs of electrons. Electronegativity A measure of the force of an atom’s attraction for electrons. Copyright 2013 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it
Chapter t Bonding and Shapes of Molecules Table 1.5 Electronegativity Values for Some Atoms(Pauling Scale) 1A 2A 西 3A 4A 5A 6A 7A 3B 4B 5B 6B 7B 8B 吕04吕8吕88 ☐15-1g25-20 us Pauling(19 received the 1954 Nobel prize in Chemistry for his contributions to our u rstanding he1962 sure of an atom's at action for electro rize fo on behalf of international control theratpmg of nuclear weapons testing. On the Pauling scale,fluorine,the most electronegative element,is assigned D Bettian/CORBIS to fuorinesu thenote that they As you stud values in thi note that they hin ap the Pe a the increasing p cha on the nucleus results in a greater force of attraction for the atom's valence electrons.Electronegativity decreases from top to bottom be- cause the increasing distance of the valer ce electrons from the nucleus results in a tween th e nucle us ana ssed As proceed from left to right in a row of the Periodic Table the contraction occurs because as you go across a row,the electrons are placed in the same shell,but the arge or nuc ei is increasing thereby pulling the ectrons in means tha ower m energ in the tabl the electro ons from atoms as you move toward the right in the Periodic Table (with some exceptions),meaning that these atoms have a higher first ionization potential. wtonsid r ad on to the atoms. Fo atio The ad. dition of an electron is called the electron affinity.which becomes more favorable as tials The a e nu electronegativity of an atom reflects its tendency to hold on to and to acquire elec- trons,the phenomenon arises from a combination of ionization potentials and elec- tron a from left to right and up to down in the can chide that f et have t v held elec trons of any atom that can make bonds.Further,fluorine most tightly holds any elec tron that it gains during ion formation or covalent bond formation.Hence,fluorine has the highest electronegativity of any atom. Unless other e noted all art on this pageCengage Learning 2013
Unless otherwise noted all art on this page © Cengage Learning 2013 Chapter 1 Covalent Bonding and Shapes of Molecules 8 a measure of an atom’s attraction for electrons that it shares in a chemical bond with another atom. The most widely used scale of electronegativities (Table 1.5) was devised by Linus Pauling in the 1930s. On the Pauling scale, fluorine, the most electronegative element, is assigned an electronegativity of 4.0 and all other elements are assigned values in relation to fluorine. As you study the electronegativity values in this table, note that they generally increase from left to right within a period of the Periodic Table and generally decrease from top to bottom within a group. Values increase from left to right because the increasing positive charge on the nucleus results in a greater force of attraction for the atom’s valence electrons. Electronegativity decreases from top to bottom because the increasing distance of the valence electrons from the nucleus results in a lower attraction between the nucleus and these electrons. Let’s further analyze the trends in the Periodic Table we just discussed. As you proceed from left to right in a row of the Periodic Table, the atoms get smaller. This contraction occurs because as you go across a row, the electrons are placed in the same shell, but the charge on the nuclei is increasing, thereby pulling the electrons in closer. This means that the orbitals get lower in energy as you move from left to right in the table and that the atoms hold their electrons tighter. It therefore takes more energy to remove the electrons from atoms as you move toward the right in the Periodic Table (with some exceptions), meaning that these atoms have a higher first ionization potential. In contrast, consider adding rather than removing an electron to the atoms. For example, when an electron is added to the halogens (Group 7A), energy is released because these atoms achieve a noble gas configuration. The energy released upon addition of an electron is called the electron affinity, which becomes more favorable as you move from left to right in a row of the Periodic Table. In contrast, as you proceed down a column in the Periodic Table, the principal quantum levels increase and the outermost electrons are farther from the nuclei, are held less tightly, and have lower ionization potentials. The atoms also have decreasing electron affinities. Because the electronegativity of an atom reflects its tendency to hold on to and to acquire electrons, the phenomenon arises from a combination of ionization potentials and electron affinities. When combining the trends of moving from left to right and up to down in the Periodic Table, you can conclude that fluorine must have the most tightly held electrons of any atom that can make bonds. Further, fluorine most tightly holds any electron that it gains during ion formation or covalent bond formation. Hence, fluorine has the highest electronegativity of any atom. Electron affinity Energy added or released when an electron is added to an atom or a molecule. Table 1.5 Electronegativity Values for Some Atoms (Pauling Scale) <1.0 1.0 – 1.4 1.5 – 1.9 2.0 – 2.4 2.5 – 2.9 3.0 – 4.0 1A 3B 4B 5B 6B 7B 8B 1B 2B 2A 3A 4A 5A 6A 7A Li 1.0 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 Mg 1.2 Ca 1.0 Sr 1.0 Ba 0.9 Sc 1.3 Y 1.2 La 1.1 Ti 1.5 Zr 1.4 Hf 1.3 V 1.6 Nb 1.6 Ta 1.5 Cr 1.6 Mo 1.8 W 1.7 Mn 1.5 Tc 1.9 Re 1.9 Fe 1.8 Ru 2.2 Os 2.2 Co 1.8 Rh 2.2 Ir 2.2 Ni 1.8 Pd 2.2 Pt 2.2 Cu 1.9 Ag 1.9 Au 2.4 Zn 1.6 Cd 1.7 Hg 1.9 B 2.0 Al 1.5 Ga 1.6 In 1.7 Tl 1.8 C 2.5 Si 1.8 Ge 1.8 Sn 1.8 Pb 1.8 N 3.0 P 2.1 As 2.0 Sb 1.9 Bi 1.9 O 3.5 S 2.5 Se 2.4 Te 2.1 Po 2.0 F 4.0 Cl 3.0 Br 2.8 I 2.5 At 2.2 Be 1.5 H 2.1 Linus Pauling (1901–1994) was the first person to receive two unshared Nobel prizes. He received the 1954 Nobel Prize in Chemistry for his contributions to our understanding of chemical bonding. He received the 1962 Nobel Peace Prize for his efforts on behalf of international control of nuclear weapons testing. © Bettmann/CORBIS Copyright 2013 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it
2 Example 1.3 Electronegativity Lewis Model of Bonding Judging from their relative positions in the Periodic Table,which element in each set is more electronegative (a)Lithium or carbon (b)Nitrogen or oxyger (c)Carbon or oxyger Solution (a)C>Li b)0>N (d)O>C Problem 1.3 Judging from their reaive the Periodi Table,which eem et is more electronegativ (a)Lithium or potassium (b)Nitrogen or phosphorus (c)Carbon or silicon Formation of lons eswaereertronie3atiityto e v nce s ity between interacting atoms is 1.9 or greater.As an example,ions are formed from sodium(electronegativity 0.9)and fluorine(electronegativity 4.0).In the following -headed (barbed)curved arrow to show the transfer of one NaTF:一Na:P时 ns品cfedgredkatonofoienstastnadone a(1223)+F122)→a*(1222)+F-(1222 m and fluorine form ions tha uration as neon,the noble gas nearest each in atomic Covalent Bonds A covalent bond is formed het s of ele hydrogen molecule.When two hydrogen atoms bond,the single electrons from each combine to form an ele on pair.This hared pair completes the valence shell of each ord ng to e Le s mo bolize the covalent bond formed by the sharing of a pair of electrons. H.+H→H:H Symbolized H--H AH=-435 k](-104 kcal)/mol In this pairing a large amount of energy is released,meaning that two hydro have cia Later in the chapter,we see that electrons have both wave and particle character (Section 1.6).When bonds are formed by the sharing of two electrons between adjacent Unless otherwise noted all art on this page Cengage Leaming 2013 9
Unless otherwise noted all art on this page © Cengage Learning 2013 9 1.2 Example 1.3 Lewis Model of Bonding | Electronegativity Judging from their relative positions in the Periodic Table, which element in each set is more electronegative? (a) Lithium or carbon (b) Nitrogen or oxygen (c) Carbon or oxygen Solution All of the elements in these sets are in the second period of the Periodic Table. Electronegativity in this period increases from left to right. (a) C . Li (b) O . N (c) O . C Problem 1.3 Judging from their relative positions in the Periodic Table, which element in each set is more electronegative? (a) Lithium or potassium (b) Nitrogen or phosphorus (c) Carbon or silicon Formation of Ions Ions are formed by the transfer of electrons from the valence shell of an atom of lower electronegativity to the valence shell of an atom of higher electronegativity. As a rough guideline, we say that ions will form if the difference in electronegativity between interacting atoms is 1.9 or greater. As an example, ions are formed from sodium (electronegativity 0.9) and fluorine (electronegativity 4.0). In the following equation, we use a single-headed (barbed) curved arrow to show the transfer of one electron from sodium to fluorine. Na Na 1 1 F F – In forming Na1F2, the single 3s valence electron of sodium is transferred to the partially filled valence shell of fluorine: Na(1s22s22p63s1) 1 F(1s22s22p5) 4 Na1(1s22s22p6) 1 F2(1s22s22p6) As a result of this transfer of one electron, both sodium and fluorine form ions that have the same electron configuration as neon, the noble gas nearest each in atomic number. The attraction between ions is what permits ionic salts such as sodium fluoride to form a strong crystal lattice and gives them a high melting point. Covalent Bonds A covalent bond is formed between atoms that share one or more pairs of electrons to give a noble gas configuration to each atom. The simplest example occurs in the hydrogen molecule. When two hydrogen atoms bond, the single electrons from each combine to form an electron pair. This shared pair completes the valence shell of each hydrogen. According to the Lewis model, a pair of electrons in a covalent bond functions in two ways simultaneously: it is shared by two atoms and at the same time fills the outer (valence) shell of each. We use a line between the two hydrogens to symbolize the covalent bond formed by the sharing of a pair of electrons. H∙ 1 ∙H 4 H:H Symbolized H!H H0 5 2435 kJ (2104 kcal)/mol In this pairing, a large amount of energy is released, meaning that two hydrogen atoms are unstable relative to H2. The same amount of energy, called the bond dissociation enthalpy (BDE, also known as the bond dissociation energy) would have to be absorbed to break the bond. Later in the chapter, we see that electrons have both wave and particle character (Section 1.6). When bonds are formed by the sharing of two electrons between adjacent Copyright 2013 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it
Chapter t Bonding and atoms,the system becomes more stable because the wave character of the electrons is Shapes of Molecules zed re veto two separate atoms.The wave th and shield the repulsion between the two positively charged nuclei.The lowering of the he ectrons The dista ngtled the Bond length The dis em)or A WesSI units of picometers in this book many chemists still use 100pm. mum of four bor nds ca vith pair （called a lone pair),one fewer bond is po In many situations,filled valence shells can be satisfied only when bonded atoms ectrons.In these c between two atoms form a A good way to distinguish covalent bonds from ionic attraction is the fact that covalent bonds have defined geometries and connectivities resulting from the sharing ons of atoms t eRitonoftheatomsasoc part in cova rough ioni nd on Polar Covalent Bonds Although all covalent bonds involve the sharing of electrons,they differ widely in the degree of sharing.Homonuclear diatomics such as HN O and F2share th e ele ns equally be atoms C are said to e n nt b the hond inc es with increasing difference in electronegativity between the bonded atoms(Table 1.6). toms onalent bond bettem is classifed as ference in electronegativity between these two at- n ole of a polar covalent bond is that of H-C1 The difference ir electronegativity between chlorine and hydrogen is 3.0-2.1=0.9.An ir portant re electro m gains a grea ons and fraction of the shared electrons and acquires a partial positive charge,indicated by the symbol +Alternatively,we show the direction of bond H_&二d Table 1.6 Classification of Chemical Bonds Type of Bond Less than 0.5 Nonpolar covalent 05to1.9 Polar covalent Greater than 1.9 Ions formed 10
Unless otherwise noted all art on this page © Cengage Learning 2013 Chapter 1 Covalent Bonding and Shapes of Molecules 10 atoms, the system becomes more stable because the wave character of the electrons is stabilized relative to two separate atoms. The wave that represents the electrons in a bond is partially concentrated in the space between the two nuclei, leading to repulsion between these electrons. In contrast, the electrons are attracted to each nucleus and shield the repulsion between the two positively charged nuclei. The lowering of the energy of the wave character of the electrons along with their added attraction to each nucleus is balanced with the repulsion between the nuclei and between the electrons. This balance results in an optimum internuclear distance called the bond length. The distance between nuclei participating in a chemical bond is called the bond length. Every covalent bond has a characteristic bond length. In H!H, it is 74 pm (picometer; 1 pm 5 10212 m). We use SI units of picometers in this book; many chemists still use Å (Ångstroms); 1 Å 5 100 pm. Because each bond requires two electrons, a maximum of four bonds can form with second-row atoms. For each unshared pair of electrons on an atom (called a lone pair), one fewer bond is possible. In many situations, filled valence shells can be satisfied only when bonded atoms share more than two electrons. In these cases, multiple covalent bonds form between the same two atoms. For example, four electrons shared between two atoms form a double bond. Six shared electrons form a triple bond. A good way to distinguish covalent bonds from ionic attraction is the fact that covalent bonds have defined geometries and connectivities resulting from the sharing of electrons. In other words, the number and positions of atoms taking part in covalent bonds are defined. In crystals, the positions of the atoms associated through ionic attraction depend on the particular crystal lattice. Polar Covalent Bonds Although all covalent bonds involve the sharing of electrons, they differ widely in the degree of sharing. Homonuclear diatomics such as H2, N2, O2, and F2 share the electrons equally between the two atoms and are said to have nonpolar covalent bonds. Many compounds such as HCl and H2O share the electrons in the bond unequally and are said to contain polar covalent bonds. The polarity in the bond increases with increasing difference in electronegativity between the bonded atoms (Table 1.6). A covalent bond between carbon and hydrogen, for example, is classified as nonpolar covalent because the difference in electronegativity between these two atoms is 2.5 2 2.1 5 0.4. An example of a polar covalent bond is that of H!Cl. The difference in electronegativity between chlorine and hydrogen is 3.0 2 2.1 5 0.9. An important consequence of the unequal sharing of electrons in a polar covalent bond is that the more electronegative atom gains a greater fraction of the shared electrons and acquires a partial negative charge, indicated by the symbol d2. The less electronegative atom has a smaller fraction of the shared electrons and acquires a partial positive charge, indicated by the symbol d1. Alternatively, we show the direction of bond polarity using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end. d1 d2 H Cl H Cl Bond length The distance between nuclei in a covalent bond in picometers (pm; 1 pm 5 10212 m) or Å (1 Å 5 10210 m). Nonpolar covalent bond A covalent bond between atoms whose difference in electronegativity is less than approximately 0.5. Polar covalent bond A covalent bond between atoms whose difference in electronegativity is between approximately 0.5 and 1.9. Table 1.6 Classification of Chemical Bonds Difference in Electronegativity Between Bonded Atoms Type of Bond Less than 0.5 Nonpolar covalent 0.5 to 1.9 Polar covalent Greater than 1.9 Ions formed Copyright 2013 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it