Chapter 8. Covalent Bonding

Transcription

1 Chapter 8 Covalent Bonding

2 Two Classes of Compounds Usually solids with high melting points Many are soluble in polar solvents such as water. Most are insoluble in nonpolar solvents such as hexane. Molten compounds conduct electricity. Aqueous solutions conduct electricity well because the contain mobile charged particles (ions). Can be gases, liquids, or solids (solids usually have low m.p.) Many are insoluble in polar solvents. Most are soluble in nonpolar solvents such as hexane. Liquid and molten compounds do not conduct electricity. Aqueous solutions are usually poor conductors of electricity because most do not contain charged particles.

3 Two Types of Bonding Models Ionic Bonds and Covalent Bonds These models represent extreme cases. It is not likely that any molecule exist at either extreme.

4 Ionic Bonds Formed by electrostatic (charge) attraction. There must be cations and anions. Compounds that are formed by ionic bonds are known as salts.

5 Why do molecules form?

6 Covalent Bonds Covalent bonds are formed when two atoms share one or more pairs of electrons. Number of Electron Pairs Shared Type of Bonding single covalent double covalent triple covalent

7 When do we see covalent compounds? We usually observe covalent compounds when the difference between the electronegativity of two atoms involved in the bond is small. Thus, we can t give electrons to one atom or the other as we did with ionic salts.

8 The Octet Rule Derived from electronic configurations. The representative elements will attempt to achieve a noble gas configuration in most of their compounds. Remember, hydrogen achieves a noble gas configuration when it has two electrons, so it does not need an octet.

9 Valence Electrons The outer-shell electrons. The electrons used to explain bonding. Electron Configurations of neutral atoms. (we don t have ions) We will only be concerned with the Representative Elements. Look at the Periodic Table.

10 Lewis Dot Representations Alkali metal, M, has an outershell configuration of ns 1 Halogen,X, has an outershell configuration of ns 2 np 5

11 Common Lewis Dot Representations

12 Electrons In Molecules Electrons can be: shared electrons (bonding electrons) unshared electrons (lone pair electrons)

13 CS 2 CS 2 4 (for C) + 2*6 (for 2 S) = 16 valence e Central element is usually the first one written Always attempt to draw the most symmetric skeleton you can.. Have we placed our 16 e?

14 HNO 2 HNO 2 5 (for N) + 2*6 (for 2 O) + 1 (for H) = 18 valence e Add e for each negative charge, subtract e for each positive charge For oxo-acids, the hydrogen is attached to a terminal oxygen, not the central element. Don t place electron pairs on H. Have we placed our 18 e? Have we placed our 18 e?

15 PH 3 PH 3 5 (for P) +3*1 (for H) = 8 valence e Add e for each negative charge, subtract e for each positive charge Will hydrogen atoms accept any more electrons? Have we placed our 8 e?

16 ClF 4 + ClF (for Cl) + 4*7 (for F) 1(for + charge) = 34 valence e Add e for each negative charge, subtract e for each positive charge Have we placed our 34 e?

17 Exceptions Covalent compounds of Be. Most covalent compounds of Group IIIA, especially B. Compounds containing odd number of electrons. When the central element needs more than eight electrons in its valence shell.

18 C 3 H 8 C 3 H 8 3*4 (for 3 C) + 8*1 (for 8 H) = 20 valence e Have we placed our 20 e?

19 NO 3 - NO 3-5 (for N) + 3*6 (for 3 O) + 1 (for negative charge) = 24 valence e Add e for each negative charge, subtract e for each positive charge Central element is usually the first one written Always attempt to draw the most symmetric skeleton you can.. Have we placed our 24 e?

20 Resonance Write a Lewis Dot Structure for CO Only show bonding electrons. [ O O C O ]2- [ O O C O ]2- [ O O C O ]2- [ O C O O ]2- [ O O C O ]2- [ O O C O ]2-

21 Resonance Is used to describe or express the true chemical structure of compound which cannot be accurately represented by any one valence-bond structure

22 Practice with Resonance H Starting with the following structure, determine how many resonance structures there are for C 6 H 6. H H C C C C H C C H H Draw the Lewis Structures for SO 3 and SO Include any resonance structures needed.

23 Formal Charge Construct used to help us decide which Lewis structure is most likely. Assumes that bonding electrons are shared equally Though we know they are not shared equally in polar bonds (later in this chapter)

24 Calculating Formal Charge F.C. = Valence electrons [lone pair electrons + ½ shared electrons] Type of electrons Valence Electrons Element - Lone Pair Electrons - ½ Shared Electrons = Formal Charge

25 Example Calculation Calculate the formal charge on each atom in SO 2. Type of electrons O S O Valence Electrons Lone Pair Electrons ½ Shared Electrons = Formal Charge

26 Using Formal Charge Prefer structures where formal charge is minimized. If you have a choice between the two structures: A B 2- or A - B - Which structure is more likely?

27 Problems Chapter 8 Problems 52 and 58

28 Hydrocarbons Alkanes and branched-alkanes discussed in Chapter 3. (C n H n+2 ) As we saw with C 3 H 8, carbon will form single covalent bonds with another carbon, or with another element (commonly H, N, O, X(halogens)) This ability makes carbon unique and special.

29 Alkenes Draw the Lewis Structure for C 2 H 4.

30 Isomers Are these molecules the same?

31 Rotation around a bond Single bonds are allowed to rotate around the axis of the bond. Double and Triple bond do not rotate around the axis of the bond.

32 Bond Length Bond length is measured from nucleus to nucleus. Estimate by considering atomic radii. Values are estimations because electronic effects of other elements to the binary pair considered affect the distance. (So the only way to know the bond length is to measure it.)

33 Bond Length and Multiple Bonds As you increase the electron density between two atoms, the bond length gets shorter. Can think: There are now more electrons on the carbon that are attracted to the oxygen nucleus (and vice versa since the electrons are shared).

34 Bond Energy Break Requires energy to break a bond. Endothermic What is the sign on H?

35 Bond Energy Make Releases energy (exothermic) What is the sign of H?

36 Bond Energies

37 NF 3 ClO 3 SO 3 2 CHClF 2

38 Figure 8.4

39 Calculate the enthalpy of combustion of ethanol, if the reaction is: C 2 H 5 OH(g) + 3O 2 (g) 2CO 2 (g) + 3H 2 O(g) To ensure that you count all the bonds, draw a picture. Use the balanced equation Use Lewis Structures

40 Calculate the enthalpy of combustion of ethanol, if the reaction is: C 2 H 5 OH(g) + 3O 2 (g) 2CO 2 (g) + 3H 2 O(g) How many C-H bonds? How many C-C bonds? How many C-O bonds? How many O-H bonds? How many O=O bonds? We use D for Dissociation. H diss = 5*D (C-H) +D (C-C) +D (C-O) +D (O-H) +3*D (O=O) H diss = 5(416 kj) kj kj kj + 3(498 kj) H diss = 4733 kj

41 Calculate the enthalpy of combustion of ethanol, if the reaction is: C 2 H 5 OH(g) + 3O 2 (g) 2CO 2 (g) + 3H 2 O(g) How many C=O bonds? How many O-H bonds? 4 6 H form = [4*D (C=O) + 6* D (OH) ] We make this negative because bond formation is exothermic. H form = [4(803 kj) + 6(467kJ)] H form = 6014 kj

42 Finishing the Math H diss = 4733 kj H form = 6014 kj All we have left to do is add these two values together! H rxn = 1281 kj Does this make sense? Does burning ethanol in air result in a release in heat?

43 Bond Polarity Covalent bonds may be either Polar Nonpolar

44 Nonpolar Covalent Bonds Electrons are shared equally by both nuclei. The covalent bonds in all homonuclear diatomic molecules must be nonpolar. Homonuclear-same element

45 Polar Covalent Bonds The electron pairs are shared unequally. Go back to electronegativity. δ+ partial positive charge δ- partial negative charge

46 Dipoles When we have a polar covalent bond, an electric dipole is created. The dipole is measurable as the dipole moment.

47 Dipoles Which atom would have a partial positive/negative charge for the binary combinations below? C-O C-N C-Li C-Cl

48 Bond Character as a Continuum EN Bond Character >1.8 Mostly ionic Polar covalent <0.4 Mostly covalent 0 Nonpolar covalent