CHAPTER 2 Bonding and Molecular Structure 21 2.4 Electronegativity and Polarity for ed hy atoms of dissimilar ele ativities is called nolar a no olarcovalent bond exists betweer atoms having a very small or zero difference in electronegativity.A few relative electronegativities are F4.0)>03.5)>C1.N3.0)>Br2.8)>S.C.12.5)>H(2.1) electro ie elemenofnt bond is relaiivelyn hreeles partial charges should not be confused with ionic charges.Polar bonds are indicated bythe head points vector sum o ond moments gives the net dipole moment of the molecule Problem 2.13 What do the molecular dipole momentsu=0 for CO,and u=1.84 D for H,O tell you about the shapes of these molecules? InCO,: 6=&-8 ad each individul bond moments cancel: arges about the 00 H.O also has polar bonds.However.since there is a net dipole moment.the individual bond moments do not cancel,and the molecule must have a bent shape: H resultant moment 2.5 Oxidation Number The oxidation numbe when t ing electrons are assigned to the more equal harge on the specie Problem 2.14 Determine the oxidation number of each C.(ON).in:(a)CH.(b)CH,OH,(c)CH,NH (d)H,C=CH Use the data(ON)=-3;(ON)=1;(ON)=- All examples are molecules;therefore.the sum of all (ON)values is (a(ON0+4OND.=0: (ON0.+(4X1)=0:(ON)=-4 (b)(ONc+(ONo+4ONH=0:(ONc+(-2)+4=0:(ONc=-2 (C)(ONc+(ON、+5ONa=0:(ONc+(-3)+5=0:(ONc=-2 (d)Since both Catoms areequivalent. 2(ONe+4ON)4=0: 2(ONe+4=0 (ON)e=-22.4 Electronegativity and Polarity The relative tendency of a bonded atom in a molecule to attract electrons is expressed by the term electroneg￾ativity. The higher the electronegativity, the more effectively does the atom attract and hold electrons. A bond formed by atoms of dissimilar electronegativities is called polar. A nonpolar covalent bond exists between atoms having a very small or zero difference in electronegativity. A few relative electronegativities are F(4.0) > O(3.5) > Cl, N(3.0) > Br(2.8) > S, C, I(2.5) > H(2.1) The more electronegative element of a covalent bond is relatively negative in charge, while the less electroneg￾ative element is relatively positive. The symbols δ and δ represent partial charges (bond polarity). These partial charges should not be confused with ionic charges. Polar bonds are indicated by ; the head points toward the more electronegative atom. The vector sum of all individual bond moments gives the net dipole moment of the molecule. Problem 2.13 What do the molecular dipole moments μ 0 for CO2 and μ 1.84 D for H2 O tell you about the shapes of these molecules? In CO2: CHAPTER 2 Bonding and Molecular Structure 21 O is more electronegative than C, and each C—O bond is polar as shown. A zero dipole moment indicates a symmetrical distribution of δ charges about the δ carbon. The geometry must be linear; in this way, individual bond moments cancel: H2O also has polar bonds. However, since there is a net dipole moment, the individual bond moments do not cancel, and the molecule must have a bent shape: 2.5 Oxidation Number The oxidation number (ON) is a value assigned to an atom based on relative electronegativities. It equals the group number minus the number of assigned electrons, when the bonding electrons are assigned to the more electronegative atom. The sum of all (ON)’s equals the charge on the species. Problem 2.14 Determine the oxidation number of each C, (ON)C, in: (a) CH4, (b) CH3OH, (c) CH3NH2, (d) H2C=CH2. Use the data (ON)N 3; (ON)H 1; (ON)O 2. All examples are molecules; therefore, the sum of all (ON) values is 0. (a) (ON)C  4(ON)H 0; (ON)C  (4  1) 0; (ON)C 4 (b) (ON)C  (ON)O  4(ON)H 0; (ON)C  (2)  4 0; (ON)C 2 (c) (ON)C  (ON)N  5(ON)H 0; (ON)C  (3)  5 0; (ON)C 2 (d) Since both C atoms are equivalent, 2(ON)C  4(ON)H 0; 2(ON)C  4 0; (ON)C 2