강의개요 Chapter 7 Covalent Bonding Lewis 구조 : 옥테트규칙 Lewis구조그리기 공명구조 분자의기하학적구조 극성 결합의극성 분자의극성 원자궤도함수와혼성궤도함수 Copyright 2005 연세대학교이학계열일반화학및실험 (1) 강의노트 7.1 Lewis Structures; The ctet Rule A Lewis structure shows the distribution of outer (valence) electrons in an atom, molecule, or polyatomic ion. 7.1 Lewis Structures; The ctet Rule (cont.) + F F 2 + Unshared electrons are shown as dots, bonds as straight lines. In 2 and F, as in most molecules and polyatomic ions, nonmetal atoms, except, are surrounded by 8 electrons, an octet. In this sense, each atom has a noble gas structure. ctet Rule
1) Writing Lewis Structures Step 1. 원자배치 Step 1. 원자배치 Central atom : 전기음성도가낮은원자를중심에배치 Step 2. 원자가전자의수계산 Step 3. 단일결합골격그리기 Terminal atom :,, alogen 은끝에놓이는원자 Step 4. 나머지전자의배치 Step 5. 다중결합그리기 Step 2. 원자가전자의수계산 Step 3. 단일결합골격그리기 원자가전자 (valence electrons) 의수 = 단일결합으로이루어진골격그리기 음전하수더하고, 양전하수뺀다. Cl - ion: valence e - C 3 molecule: valence e - S 2-3 ion: valence e - Cl C S
Step 4. 나머지전자의배치 남은전자수계산 : 하나의단일결합당두개의전자를줄임 Cl - ion: valence e - left C 3 molecule: valence e - left S 2-3 ion: valence e - left Step 5. 다중결합그리기 중심원자가 ctet Rule을만족하도록 중심원자에전자가부족하면 Terminal atom에서한쌍의전자를없애고중심원자와다중결합을만듦 C, N,, P, S 남은전자배치 : Terminal atom 에비공유전자쌍배치 (ctet rule), 남은전자는중심원자에배치 : ( Cl ) - C ( S ) 2- ex1) N 3 - ex1) N 3- (cont.) S1. 원자배치 N S2. # valence e - = S3. 단일결합골격 N S4. # of valence e - left: S4. Terminal atom 에비공유전자쌍배치 N N 원자는단지 6 개의전자에의해둘러싸여있음. 부족한전자 2 개를보충하려면이중결합하나가필요함 : N ex2) N 2 의구조? 10 valence e - N N
2) Resonance Forms To explain the fact that all three bonds in the nitrate ion are the same length, invoke the concept of resonance: = N N N = The true structure is a hybrid of the three forms. resonance forms are obtained by moving electrons, not atoms resonance can be expected when it is possible to draw more than one structure that follows the octet rule. 3) Formal Charge ex) Methyl alcohol (C 4 )? C Formal charge = number of valence electrons in the free atom C f = X (Y+Z/2) X : # of valence e - in the free atom (last digit of the group number) Y : # of unshared e - owned by the atom in the Lewis structure Z : # of bonding e - charged by the atom in the Lewis structure C 3) Formal Charge (cont.) ex) C 4? C C 4) Exceptions to the ctet Rule: Electron-Deficient Molecules In a few molecules, there are less than 8 electrons around the central atom: The formal charges are as close to zero as possible. Any negative formal charge is located on the most strongly electronegative atom.
5) Exceptions to the ctet Rule: Expanded ctets 7.2 Molecular Geometry Give central atom expanded octet: Consider XeF 4 : 36 valence e -. The octet structure only uses 32, therefore the remaining 4 are added to the central atom, giving it 12. VESPR principle: electron pairs around a central atom tend to be oriented so as to be as far apart as possible. 1) Ideal Geometries with Two to Six Electronpairs on the Central Atom 2 to 6 atoms around central atom, no unshared pairs: BeF 2 BeF 3 CF 4 PF 5 SF 6 2) Effect of Unshared Pairs on Molecular Geometry Unshared pairs (lone pairs) : the molecular geometry is quite different when one or more unshared pairs are present AX 2 E (GeF 2 ): AX 3 E (N 3 ): AX 2 E 2 ( 2 ): 3) Multiple Bonds as no effect upon geometry: multiple bond behaves like a single bond BF 3 and S 3 : BeF 2 and C 2 :
7.3 Polarity of Molecules 1) Polar and Nonpolar Covalent Bonds Bond Polarity : between unlike atoms, a chemical bond that has positive and negative ends. Not for same atoms. Molecular Polarity : polar molecule contains positive and negative poles. Non-polar molecule has no poles. Diatomic molecules Polyatomic molecules All bonds are polar unless the two atoms joined are identical ( ). Extent of polarity depends upon difference in electronegativity. E.N. = 0 C E.N. = 0.3 F E.N. = 1.8 2) Polar and Nonpolar Molecules Diatomic molecules: polar if atoms differ Cl Cl Cl Cl molecules line up in an electric field, Cl 2 molecules don t. Polyatomic molecules even though bonds are polar, molecule may be nonpolar if bonds are symmetrically arranged: F Be F C 4 C 3 Cl 7.4 Atomic rbitals; ybridization Bonding Model Valance Bond Theory Molecular rbital Theory Molecular rbital Theory In molecules, the orbitals occupied by electron pairs are seldom pure s or p orbitals. Instead, they are hybrid orbitals, formed by combining s, p, and d orbitals.
1) ybrid rbitals: sp, sp 2, sp 3, sp 3 d, sp 3 d 2 sp: s orbital + p orbital two sp hybrids 2s 2p Be in BeF 2 ( ) ( )( )( ) sp 2 : s orbital + two p orbitals three sp 2 hybrids 2s 2p B in BF 3 ( ) ( )( )( ) ybridization with 5 or 6 Electron Pairs sp 3 d, sp 3 d 2 Note that expanded octets do not occur with atoms in the second period (e.g.. N,, F) since there are no 2d orbitals. sp 3 : s orbital + 3 p orbitals four sp 3 hybrids 2s 2p C in C 4 ( ) ( ) ( ) ( ) 2) Multiple Bonds 3) Sigma and Pi Bonds Unshared pairs can be hybridized ( 2, N 3 ). nly one of the electron pairs in a multiple bond is hybridized. S = = C = sp 2 hybridization for sulfur sp hybridization for carbon When a bond consists of an electron pair in a hybrid orbital, the electron density is concentrated along the bond axis and is symmetrical about it. Such a bond is called a sigma bond. The extra electron pairs in a multiple bond are located in unhybridized orbitals which are not concentrated along the bond axis. Instead, they are concentrated in lobes north and south of the axis. Such bonds are called pi bonds.