hw02sol ch 13

hw02sol ch 13 - CHAPTER]3 BONDING: GENERAL CONCEPTS...

Info iconThis preview shows pages 1–3. Sign up to view the full content.

View Full Document Right Arrow Icon

Info iconThis preview has intentionally blurred sections. Sign up to view the full version.

View Full DocumentRight Arrow Icon
This is the end of the preview. Sign up to access the rest of the document.

Unformatted text preview: CHAPTER]3 BONDING: GENERAL CONCEPTS Chemical Bonds and Electronegativity II. Electronegativity is the ability of an atom in a molecule to attract electrons to itself. Electronegativity is a bonding term. Electron affinity is the energy change when an electron is added to a substance. Electron affinity deals with isolated atoms in the gas phase. A covalent bond is a sharing of electron pair(s) in a bond between two atoms. An ionic bond is a complete transfer of electrons from one atom to another to form ions. The electrostatic attraction of the oppositely charged ions is the ionic bond. A pure covalent bond is an equal sharing of shared electron pair(s) in a bond. A polar covalent bond is an unequal sharing. Ionic bonds form when there is a large difference in electronegativity between the two atoms bonding together. This usually occurs when a metal with a small electronegativity is bonded to a nonmetal having a large electronegativity. A pure covalent bond forms between atoms having identical or nearly identical eletronegativities. A polar covalent bond forms when there is an intermediate electronegativity difference. In general, nonmetals bond together by forming covalent bonds, either pure covalent or polar covalent. Ionic bonds form due to the strong electrostatic attraction between two oppositely charged ions. Covalent bonds form because the shared electrons in the bond are attracted to two different nuclei, unlike the isolated atoms where electrons are only attracted to one nuclei. The attraction to another nuclei overrides the added electron-electron repulsions. . (+1)(-1) 10 12. a. There are two attractions of the form , where r = I x 10- m = 0.1 nm. r v = 2 x (2.31 x 10- 19 J nm)[(+l)(-l)] = -4.62 x 10- 18 J = -5 X 10- 18 J O.lnm b. There are 4 attractions of + I and -1 charges at a distance of 0.1 nm from each other. The two negative charges and the two positive charges repel each other across the diagonal of the square. This is at a distance of fi x 0.1 nm. 490 491 CHAPTER 13 BONDING: GENERAL CONCEPTS V=4 x (2.31 x 10- 19 )[ (+1)(- 1)] + 2.31 x 10- 19 [(+1)(+1)] 0.1 .fi (0.1) + 2.31 x 10-19 [(-1)(-1)] .fi(0.1) v = -9.24 x 10- 18 J + 1.63 x 10- 18 J + 1.63 x 10- 18 J = -5.98 x 10- 18 J =-6 x 10- 18 J Note: There is a greater net attraction in arrangement b than in a. 13. Using the periodic table, we expect the general trend for electronegativity to be: 1. Increase as we go from left to right across a period 2. Decrease as we go down a group a. C <N<O b. Se < S < Cl c. Sn<Ge<Si d. Tl <Ge<S e. Rb <K <Na f. Ga<B<O 14. The most polar bond will have the greatest difference in electronegativity between the two atoms. From positions in the periodic table, we would predict: a. Ge-F b. P-Cl c. S-F d. Ti-Cl e. Sn-H f. Tl-Br 15. The general trends in electronegativity used in Exercises 13.13 and 13.14 are only rules of thumb. In this exercise we use experimental values of electronegativities and can begin to see several exceptions. The order of EN using Figure 13.3 is: see several exceptions....
View Full Document

Page1 / 10

hw02sol ch 13 - CHAPTER]3 BONDING: GENERAL CONCEPTS...

This preview shows document pages 1 - 3. Sign up to view the full document.

View Full Document Right Arrow Icon
Ask a homework question - tutors are online