7 Chemical Bonding Chapter Goals 1 Lewis Dot

7 Chemical Bonding 化學鍵結

Chapter Goals 1. Lewis Dot Formulas of Atoms 路易士電子點結構式 Ionic Bonding 離子鍵結 2. Formation of Ionic Compounds 形成離子化合物 Covalent Bonding 共價鍵結 3. Formation of Covalent Bonds 形成共價鍵 4. Bond Lengths and Bond Energies 鍵長及鍵能 5. Lewis Formulas for Molecules and Polyatomic Ions 分子極多元 子離子之路易士結構式 6. Writing Lewis Formulas: The Octet Rule 八隅體法則 7. Formal Charges 形式電荷 8. Writing Lewis Formulas: Limitations of the Octet Rule 八隅體法 則之限制 9. Resonance 共振 10. Polar and Nonpolar Covalent Bonds 極性及非極性共價鍵結 11. Dipole Moments 偶極距 12. The Continuous Range of Bonding Types 鍵結形式之連續範圍 2

Introduction • Attractive forces that hold atoms together in compounds are called chemical bonds (化合物中將原 子拉進的吸引力稱之為化學鍵結) • The electrons involved in bonding are usually those in the outermost (valence) shell (與鍵結有關的電子通 常是指最外層的電子) 3

Introduction • Chemical bonds are classified into two types: Ø Ionic bonding 離子鍵結 results from electrostatic attractions among ions, which are formed by the transfer of one or more electrons from one atom to another. Ø 離子鍵為陽離子與陰離子間之庫倫靜電力，陽離子由金屬失去電子形成，陰離 子由非金屬得到電子形成。兩正負電荷之離子吸引力，稱為離子鍵。 ü Ionic compound 離子化合物 Na. Cl Ø Covalent bonding 共價鍵結 results from sharing one or more electron pairs between two atoms. Ø 價鍵為兩原子共用電子對之化學鍵，其結合原子皆為非金屬原子 ü Covalent compound H 2, Cl 2 4

Ionic Compound 離子化合物 Covalent Compound 共價化合物 Solids Gases, liquid, or solids High melting points (>400 o. C) Low melting points (<300 o. C) Soluble in polar solvents, such as water insoluble in polar solvents insoluble in nonpolar solvents, such as C 6 H 14 and CCl 4 Most are soluble in nonpolar solvents, such as C 6 H 14 , CCl 4 Molten compounds conduct electricity well because they contain ions Liquid and molten compounds do not conduct electricity Aqueous solutions conduct electricity well because they contain ions Aqueous solutions are usually poor conductors Two elements with different electronegativities, usually a metal and a nonmetal Two elements with similar electronegativities, usually nonmetals 5

Lewis Dot Formulas of Atoms 原子的路易 士電子點結構式 • Lewis dot formulas or Lewis dot representations are a convenient bookkeeping method for tracking valence electrons 價電子. – The electrons in the outermost occupied shells (最外層的 電子數) • s and p orbitals – Valence electrons are those electrons that are transferred or involved in chemical bonding (價電子是指與化學鍵結 有關或轉移的電子) • They are chemically important 6

Lewis Dot Formulas 路易士結構 of Atoms Elements that are in the same periodic group have the same Lewis dot structures (同族元素具相同路易士結構) 1 electron in valence shell 5 electrons in valence shell Not as useful for the transition and inner transition elements 7 (不適用於過渡元素)

Ionic Bonding 離子鍵結 Formation of Ionic Compounds 形成離子化 合物 • An ion is an atom or a group of atoms possessing a net electrical charge. • Ions come in two basic types: – positive (+) ions or cations • These atoms have lost 1 or more electrons. – negative (-) ions or anions • These atoms have gained 1 or more electrons. 8

Formation of Ionic Compounds • Monatomic ions consist of one atom. – Examples: • Na+ sodium ion, Ca 2+, Al 3+ -- cations • Cl- chloride ion, O 2 -, N 3 - -- anions • Polyatomic ions contain more than one atom. – Examples: • NH 4+ ammonium ion -- cation • NO 2 -, CO 32 -, SO 42 - sulfate ion – anions – The atoms of a polyatomic ion are held together by covalent bonds (多原子離子之原子間以共價鍵結) 9

Ionic Bonding is the attraction of oppositely charged ions (cations and anions) in large numbers to form a solid. Such a solid compound is called an ionic solid. • Elements that have low electronegativities and low ionization enengy — metals (oxidozed; lose electrons to form cations) React • Elements that have high electronegativities and very negative electron affinities — nonmetals (reduced; gain electrons to form anions) Ionic compound 10

Formation of Ionic Compounds • Reaction of Group IA Metals with Group VIIA Nonmetals 1 A metal 7 A nonmetal 2 Li(s) + F 2(g) 2 Li. F(s) Silver yellow white solid Solid gas with an 842 o. C melting point 11

Formation of Ionic Compounds • The underlying reason for the formation of Li. F lies in the electron configurations of Li and F. 1 s 2 s 2 p Li F These atoms form ions with these configurations. Li+ same configuration as [He] Li+ isoelectronic with He loss one electron F- same configuration as [Ne] F+ isoelectronic with Ne gained one electron 等電子離子 12

Formation of Ionic Compounds • We can also use Lewis dot formulas to represent the neutral atoms and the ions they form. 13

Formation of Ionic Compounds • The Li+ ion contains two electrons, same as the helium atom. – Li+ ions are isoelectronic 等電子離子 with helium. • The F- ion contains ten electrons, same as the neon atom. – F- ions are isoelectronic with neon. • Isoelectronic species contain the same number of electrons. • Most ionic compounds formed by reactions between representative metals 典型金屬and representative nonmetals 典型非金屬 14

Formation of Ionic Compounds • The reaction of potassium with bromine is a second example of a group IA metal with a Group VIIA non metal. – Write the reaction equation. 1 A metal 2 K(s) + 7 A nonmetal Br 2(g) 2 KBr(s) ionic solid 15

Formation of Ionic Compounds • We look at the electronic structures of K and Br. 4 s 4 p K [Ar] Br [Ar] and the d electrons The atoms form ions with these electronic structures. 4 s K+ Br- 4 p same configuration as [Ar] same configuration as [Kr] 16

Formation of Ionic Compounds • Write the Lewis dot formula representation for the reaction of K and Br. 17

Formation of Ionic Compounds • There is a general trend evident in the formation of these ions. • Cations become isoelectronic with the preceding noble gas. 之前的鈍氣 • Anions become isoelectronic with the following noble gas. 之後的鈍氣 18

Formation of Ionic Compounds • In general for the reaction of IA metals and VIIA nonmetals, the reaction equation is: 2 M(s) + X 2 2 M+ X-(s) – where M is the metals Li to Cs – and X is the nonmetals F to I. Electronically this is occurring. ns np M X ns np M+ X- 19

Formation of Ionic Compounds • Next we examine the reaction of IIA metals with VIIA nonmetals. • This reaction forms mostly ionic compounds. – Notable exceptions are Be. Cl 2, Be. Br 2, and Be. I 2 which are covalent compounds. • One example is the reaction of Be and F 2. Be(s) + F 2(g) Be. F 2(g) 20

Formation of Ionic Compounds • The valence electrons in these two elements are reacting in this fashion. 2 s 2 p Be [He] Be 2+ F [He] F- Next, draw the Lewis dot formula representation of this reaction. 21

Formation of Ionic Compounds • The remainder of the IIA metals and VIIA nonmetals react similarly. • Symbolically this can be represented as: M(s) + X 2 M 2+ X 2 M can be any of the metals Be to Ba. X can be any of the nonmetals F to Cl. 22

Formation of Ionic Compounds • For the reaction of IA metals with VIA nonmetals, a good example is the reaction of lithium with oxygen. • The reaction equation is: 23

Formation of Ionic Compounds • Draw the electronic configurations for Li, O, and their appropriate ions. 2 s 2 p Li [He] O [He] 2 s 2 p Li 1+ O 2 - Draw the Lewis dot formula representation of this reaction. 24

Formation of Ionic Compounds • The remainder of the IA metals and VIA nonmetals behave similarly. • Symbolically this can be represented as: 2 M (s) + X M 21+ X- M can be any of the metals Li to Cs. X can be any of the nonmetals O to Te. 25

Formation of Ionic Compounds • The reaction of IIA metals and VA nonmetals also follows the trends that we have established in this chapter. • The reaction of calcium with nitrogen is a good example. • The reaction equation is: 3 Ca(s) + N 2(g) Ca 3 N 2 (s) 26

Formation of Ionic Compounds • Draw the electronic representation of Ca, N, and their ions. 4 s 4 p Ca [Ar] Ca 2+ 2 s 2 p N [He] N 3 - • Draw the Lewis dot representation of this reaction. 27

Formation of Ionic Compounds • Other IIA and VA elements behave similarly. • Symbolically, this reaction can be represented as: 3 M(s) + 2 X(g) M 32+ X 23 M can be the IIA elements Be to Ba. X can be the VA elements N to As. 28

Formation of Ionic Compounds • d-transition Metal Ions 鈧 – The outermost s electrons and energy level lower d electrons are always the first ones lost when transition metals form simple ions. 3 d 4 s Sc[Ar] Sc 3+[Ar] 3 e- lost Zn[Ar] Zn 2+[Ar] 2 e- lost – Most other 3 d-transition metals can form at least two cations in their compounds. 3 d 4 s Co[Ar] Co 2+[Ar] 2 e- lost Co[Ar] Co 3+[Ar] 3 e- lost 29

Formation of Ionic Compounds 3 A H, a nonmetal, forms ionic compounds with IA and IIA metals for example, Li. H, KH, Ca. H 2, and Ba. H 2. Other hydrogen compounds are covalent. Group IA and IIA can form peroxide (contain O 22 - ion ) or superoxide (contain O 2 - ion). The peroxide and superoxide ions contain atoms that are covalently 30 bonded to one another

Group IA and IIA can form peroxide (contain O 22 - ion ) or superoxide (contain O 2 - ion). The peroxide and superoxide ions contain atoms that are covalently bonded to one another 31

Formation of Ionic Compounds • Ionic compounds form extended three dimensional arrays of oppositely charged ions. • Ionic compounds have high melting points because the coulomb force庫侖力, which holds ionic compounds together, is strong. 32

Formation of Ionic Compounds • Coulomb’s Law describes the attraction of positive ions for negative ions due to the opposite charges. 33

Formation of Ionic Compounds • Small ions with high ionic charges have large Coulombic forces of attraction. • Large ions with small ionic charges have small Coulombic forces of attraction. • Use this information, plus the periodicity rules from Chapter 6, to arrange these compounds in order of increasing attractions among ions KCl, Al 2 O 3, Ca. O 34

Covalent Bonding 共價鍵結 • Covalent bonds are formed when atoms share electrons. It Occurs when the electronegativity difference between elements (atoms) is zero or relativity small (電負度幾乎沒差) • The bonds between atoms within a molecule (intramolecular bonds 分子內鍵結) are relatively strong, but the force of attraction between molecules (intermolecular forces 分子間 鍵結) are relatively weak lower melting and boiling points than ionic compound (較離 子化合物的熔點及沸點低) – If the atoms share 2 electrons a single covalent bond is formed (若原子共享 2個電子則形成單一共價鍵) – If the atoms share 4 electrons a double covalent bond is formed (若原子共享 4個電子則形成二個共價鍵) – If the atoms share 6 electrons a triple covalent bond is formed (若原子共享 6個電子則形成三個共價鍵) • The attraction between the electrons is electrostatic in nature – The atoms have a lower potential energy when bound. 35

Formation of Covalent Bonds • Representation of the formation of an H 2 molecule from H atoms. The electron of each atom is attracted by the positively charged nucleus of the other atom, so the electron density begins to shift (blue arrows) (電子受 到帶正電的原子核的吸引, 電子團開始變化 ) The electron clouds of the two atoms repel one another, and so do the nuclei of the two atoms (Red arrows) (兩個原子亦會有排斥力) The two 1 s orbitals overlap 36

Formation of Covalent Bonds • This figure shows the potential energy of an H 2 molecule as a function of the distance between the two H atoms. 太靠近, 產生斥力 太遠, 引力太小, 無鍵結 斥力與引力達成 平衡, 穩定的排列 For any covalent bond there is an internuclear distance where the attractive and repulsive forces balance. This distance is the bond length (彼此間的距離稱鍵長). The energy difference is the bond energy (能量的 差異稱為鍵能) 37

Bond dissociation energy 鍵離解能 Bond energy 鍵能 38

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Formation of Covalent Bonds We can use Lewis dot formulas to show covalent bond formation. 1. H molecule formation representation. H + H H: H or H-H 2. HF molecule formation H + F : H: F: or HF 3. F 2 molecule formation : F: F: or F : F + F: 2 40

Dot formula • CO 2 • NH 4+ N: 5 electrons H: 1 electrons Total 9 electrons H O or H O H H • H 2 O Dash formula線結構式

Writing Lewis Formulas: The Octet Rule八隅體法則 • N - A = S rule Simple mathematical relationship to help us write Lewis dot formulas. – N = number of electrons needed to achieve a noble gas configuration. (要達到鈍氣組態的電子數目, 通常為 8) • N usually has a value of 8 for representative elements. • N has a value of 2 for H atoms. – A = number of electrons available in valence shells of the atoms. (原子的價電子數) • A is equal to the periodic group number for each element. • A is equal to 8 for the noble gases. – S = number of electrons shared in bonds. (形成鍵結可共享的電 子數) – A-S = number of electrons in unshared, lone, pairs. (不共享的電 子數, 又稱孤電子對) 42

Lewis Formulas for Molecules and Polyatomic Ions • First, we explore Lewis dot formulas of homonuclear diatomic molecules. –Two atoms of the same element. 1. Hydrogen molecule, H 2. N=2 x 2=4 e- needed A=2 x 1=2 e- available S=N-A=2 e- shared H: H or H−H F F− 2. Fluorine, F 2 F or F 3. Nitrogen, N 2 N or N N N N=2 x 8=16 e- needed A=2 x 7=14 e- available S=N-A=2 e- shared N=2 x 8=16 e- needed A=2 x 5=10 e- available S=N-A=6 e- shared 2 molecule with 6 electrons 3 covalent bond 43

Lewis Formulas for Molecules and Polyatomic Ions • Next, look at heteronuclear diatomic molecules. –Two atoms of different elements. • Hydrogen halides 鹵化氫 are good examples. 1. hydrogen fluoride, HF F H− F or H 2. hydrogen chloride, HCl Cl or H − Cl H 3. hydrogen bromide, HBr Br H Br or H − N=1 x 2+1 x 8=10 e- needed A=1 x 1+1 x 7=8 e- available S=N-A=2 e- shared 44

Lewis Formulas for Molecules and Polyatomic Ions • Now we will look at a series of slightly more complicated heteronuclear molecules. • Water, H 2 O H N=2 x 2+1 x 8=12 e- needed A=2 x 1+1 x 6=8 e- available S=N-A=4 e- shared H O or H • Ammonia molecule , NH 3 H N H H or H H N=3 x 2+1 x 8=14 e- needed A=3 x 1+1 x 5=8 e- available S=N-A=6 e- shared 45

Lewis Formulas for Molecules and Polyatomic Ions • Lewis formulas can also be drawn for molecular ions. • One example is the ammonium ion , NH 4+. N=1 x 8+4 x 2=16 e- need 1 C atom 4 H atom A=1 x 5+4 x 1 -1=8 e- available 1 C atom 4 H atom (+1 charge) S=N-A=16 -8=8 e- share • Notice that the atoms other than H in these molecules have eight electrons around them. 46

A Guild to Writing Lewis Formulas 1. Select a reasonable (symmetrical) “skeleton” for the molecule or polyatomic ion 選擇最合理分子當作骨架 – The least electronegative element is usually the central element, except the H 電負度最小者通常為中心分子, 氫除外 – Carbon bonds to two, three or four atoms, but never more than four 碳可與2, 3及4個原子鍵結 – Oxygen atoms do not bond to each other except in (氧原子不 會鍵結在一起, 除非) • O 2 and O 3 • hydrogen peroxide, H 2 O 2, and the peroxide contain the O 22 group • Superoxide, which contain the O 2 - group – In ternary oxoacids, hydrogen usually bond to an O atom, not to the central atom, HNO 2 三元含氧酸中, 氫通常與氧鍵結 – For those have more than one central atom, the most symmetrical skeletons possible are used, such as C 2 H 4, P 2 O 74 -

A Guild to Writing Lewis Formulas 2. Calculate N, the number of valence shell electrons needed by all atoms in the molecule or iron to achieve noble gas configurations PF 3 N=1 x 8(P atom)+3 x 8(F atoms) = 32 e- need CH 3 OH N=1 x 8(C atom)+4 x 2(H atoms)+1 x 8 (O atom) = 24 e- need NO 3 - N=1 x 8(N atom)+3 x 8(O atoms) = 32 e- need Calculate A, the number of electrons available in the outer shells of all the atoms PF 3 A=1 x 5(P atom)+3 x 7(F atoms) = 26 e- available CH 3 OH A=1 x 4(C atom)+4 x 1(H atoms)+1 x 6 (O atom) = 14 e- available NO 3 - A=1 x 5(N atom)+3 x 6(O atoms)+1 (for 1 -charge) = 24 e- available

A Guild to Writing Lewis Formulas Calculate S, total number of electrons shared in the molecule or ion, using the relationship S=N-A PF 3 S=N-A = 32 -26=6 e- shared (3 pairs of e-shared) CH 3 OH S=24 -14= 10 e- shared (5 pairs of e-shared) NO 3 - S= 32 -24=8 e- hared (4 pairs of e-shared)

3. Place the S electron in to the skeleton as shared pairs. Use double and triple bonds only when necessary. S=6 (3 pairs of e-shared) S=10 (5 pairs of e-shared) S=8 (4 pairs of e-shared)

4. Check the additional electrons into the skeleton as unshared pairs to fill the octet of every A group element (except H, shared only 2 e-) PF 3 CH 3 OH NO 3 -

Writing Lewis Formulas: The Octet Rule • The octet rule states that representative elements usually attain stable noble gas electron configurations in most of their compounds. • Lewis dot formulas are based on the octet rule. • We need to distinguish between bonding (or shared) electrons and nonbonding (or unshared or lone pairs) of electrons. 52

Writing Lewis Formulas: The Octet Rule • Example 7 -2: Write Lewis dot and dash formulas for hydrogen cyanide, HCN. N = 2 (H) + 8 (C) + 8 (N) = 18 A = 1 (H) + 4 (C) + 5 (N) = 10 S= 8 A-S = 2 • This molecule has 8 electrons in shared pairs and 2 electrons in lone pairs. or C C N N H C 53

Writing Lewis Formulas: The Octet Rule • Example 7 -3: Write Lewis dot and dash formulas for the sulfite ion, SO 32 -. N = 8 (S)+3 x 8 (O) = 32 A = 6 (S)+3 x 6 (O)+ 2 (- charge)= 26 S = 6 A-S = 20 • Thus this polyatomic ion has 6 electrons in shared pairs and 20 electrons in lone pairs. • Which atom is the central atom in this ion? 54

Writing Lewis Formulas: The Octet Rule • What kind of covalent bonds, single, double, or triple, must this ion have so that the six shared electrons are used to attach the three O atoms to the S atom? or O S O 2 O 2 O S O O 55

Example 7 -1 Writing Lewis Formula Write the Lewis formula for the nitrogen molecule, N 2, carbon disulfide, CS 2 and the carbonate ion, CO 32 - O S =C= S –C = or or O O C 2 - O N N or C S 2 - N=1 x 8+3 x 8=32 e- needed A=1 x 4+3 x 6+2=24 e- available S=N-A=8 e- shared Four pairs are shared S (b) CO 3 N=1 x 8+2 x 8=24 e- needed A=1 x 4+2 x 6=16 e- available S=N-A=8 e- shared 3 molecule with 8 electrons 2 x 2 covalent bond N (b) CS 2 N (a) N 2 2 - – O N=2 x 8=16 e- needed A=2 x 5=10 e- available S=N-A=6 e- shared 2 molecule with 6 electrons 3 covalent bond Exercise 29, 30, 38

Formal Charge 形式電荷 • Calculation of a formal charge on a molecule is a mechanism for determining correct Lewis structures • The formal charge is the hypothetical charge on an atom in a molecule or polyatomic ion. • The best Lewis structures will have formal charges on the atoms that are zero or nearly zero. 形式電荷：基於共用電子對由兩鍵結原子平分共用的原則 下，此原子的價電子數與在自由原子狀態的價 電子數兩者的差值。 58

Formal Charge Rules for Assigning Formal Charge 1. Formal Charge = group number – (number of bonds + number of unshared e-) 2. An atom that has the same number of bonds as its periodic group number has a formal charge of 0. 3. a. The formal charges of all atoms must sum to 0 in molecules. b. The formal charges must sum to the ion’s charge for a polyatomic ion. ♦ 形式電荷的定義: (路易士結構中的) 價電子數 = 孤電子數 + ½鍵結電子數 形式電荷 =自由原子的價電子數 =自由原子的價電子數 孤電子數 ½鍵結電子數 59

In NH 3 N atom: group 5 A, has 3 bonds and 2 unshared e. FC =(group number) –[(number of bonds)+(number of unshared e-)] = 5 -(3+2)=0 H: FC= 1 -(1+0)=0 the formal charge of N and H are both zero in NH 3, so the sum of the formal charged is 0+3(0)=0 In NH 4+ the atom has four bonds and no shared efor N FC=5 -(4+0)=1 for H FC=1 -(1+0)=0 NH 4+=1+4(0)=1 60

Formal Charge Cl N O i Consider nitrosyl chloride, NOCl 7 – (2+4) = +1 5 – (3+2) = 0 6 – (1+6) = -1 ii Cl N O 7 – (1+6) = 0 5 – (3+2) = 0 6 – (2+4) = 0 (ii) is a preferable Lewis formula because it has smaller charges than (i) 61

Writing Lewis Formulas: Limitations of the Octet Rule • There are some molecules that violate違背the octet rule. –For these molecules the N - A = S rule does not apply: 1. The covalent compounds of Be. – Use 4 electrons as N number 2. The covalent compounds of the IIIA Group. – Use 6 electrons as N number 3. Species which contain an odd number of electrons, such as NO 2. – NO with 11 valence shell electrons and NO 2 with 17. 4. Species in which the central element must have a share of more than 8 valence electrons to accommodate all of the substituents. 5. Compounds of the d- and f-transition metals. 62

Writing Lewis Formulas: Limitations of the Octet Rule • In those cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations. • The central atom does not have a noble gas configuration but may have fewer than 8 (exceptions 1, 2, & 3) or more than 8 (exceptions 4 & 5). 63

Example 7 -1 Writing Lewis Formula Write the Lewis formula for the beryllium chloride, Be. Cl 2, boron trichloride, BCl 3 and the phosphorus pentafluoride, PF 5 or Cl Be Cl or Cl Be Cl Cl Cl B Cl F F P F N=5 x 8+1 x 8=48 e- needed A=5 x 7+1 x 5=40 e- available S=N-A=8 e- shared 5 F bond to P. This requires the sharing a minimum 10 e-. Increase S from 8 e- to 10 e-. F or F P F (b) PF 5 N=3 x 8+1 x 6=30 e- needed A=3 x 7+1 x 3=24 e- available S=N-A=6 e- shared (b) BCl 3 N=2 x 8+1 x 4=20 e- needed A=2 x 7+1 x 2=16 e- available S=N-A=4 e- shared (a) Be. Cl 2 64

Example 7 -1 Writing Lewis Formula Write the Lewis formula for the sulfir tetrafluoride, SF 4, the triiodide ion, I 3 - (a) SF 4 N=1 x 8+4 x 8=40 e- needed A=1 x 6+4 x 7=34 e- available S=N-A=6 e- shared Requires a minimum 8 e- (b) I 3 N=3 x 8=24 e- needed A=3 x 7+1 x 1=22 e- available S=N-A=2 e- shared Exercise 62 65

Resonance 共振 • Resonance is a flawed method of representing molecules. – There are no single or double bonds in SO 3. • In fact, all of the bonds in SO 3 are equivalent. • The best Lewis formula of SO 3 that can be drawn is: 共價鍵結構的混成體, 稱之為共振 66

The typical C-O single bond length is 1. 43Å The typical C=O double bond length is 1. 22Å The C-O bond in the CO 32 ion is at 1. 29Å Delocalization of bonding electron 非定域電子鍵 67

Polar and Nonpolar Covalent Bonds 極性 及非極性共價鍵 • Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds. – Nonpolar covalent bonds have a symmetrical charge distribution (對稱的電荷分佈) • To be nonpolar the two atoms involved in the bond must be the same element to share equally. • Some examples of nonpolar covalent bonds. – H 2 – N 2 68

Polar and Nonpolar Covalent Bonds • Covalent bonds in which the electrons are not shared equally are designated as polar covalent bonds – Polar covalent bonds have an asymmetrical charge distribution • To be a polar covalent bond the two atoms involved in the bond must have different electronegativities (具不同的電負 度) • Some examples of polar covalent bonds. • HF 69

Polar and Nonpolar Covalent Bonds • Shown below is an electron density map of HF – Blue areas indicate low electron density – Red areas indicate high electron density • Polar molecules have a separation of centers of negative and positive charge, an asymmetric charge distribution. 70

Polar and Nonpolar Covalent Bonds • Shown below are electron density maps of the hydrogen halides. – Notice that the charge separation decreases as we move from HF to HI. 71

Polar and Nonpolar Covalent Bonds • Polar molecules can be attracted by magnetic and electric fields. 72

Dipole Moments電偶極矩 • Molecules whose centers of positive and negative charge do not coincide同位, 重疊, have an asymmetric charge distribution, and are polar. – These molecules have a dipole moment. • The dipole moment has the symbol . • is the product of the distance, d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q. 73

Dipole Moments 電偶極矩 • Molecules that have a small separation of charge have a small . • Molecules that have a large separation of charge have a large . • For example, HF and HI: 74

Dipole Moments • There are some nonpolar molecules that have polar bonds. • There are two conditions that must be true for a molecule to be polar. – There must be at least one polar bond present or one lone pair of electrons. – The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another. 75

The Continuous Range of Bonding Types • Covalent and ionic bonding represent two extremes. 1. In pure covalent bonds electrons are equally shared by the atoms. 2. In pure ionic bonds electrons are completely lost or gained by one of the atoms. • Most compounds fall somewhere between these two extremes. 76

Continuous Range of Bonding Types • All bonds have some ionic and some covalent character. – For example, HI is about 17% ionic • The greater the electronegativity differences the more polar the bond. 77