Chapter 8. Basic Concepts of Chemical Bonding. Lecture Presentation. John D. Bookstaver St. Charles Community College Cottleville, MO

Transcription

1 Lecture Presentation Chapter 8 of Chemical John D. Bookstaver St. Charles Community College Cottleville, MO

2

3 Chemical Bonds Chemical bonds are the forces that hold the atoms together in substances. Three basic types of bonds Ionic Electrostatic attraction between ions. Covalent Sharing of electrons. Metallic Metal atoms bonded to several other atoms.

4 Lewis Symbols G.N. Lewis pioneered the use of chemical symbols surrounded with dots to symbolize the valence electrons around an atom.

5 A. Yes, all three are correct. B. No, the first two are different Lewis structures because of differing placement of unpaired electron. C. No, because the Lewis structure on the right has only five valence electrons and elemental Cl has seven valence electrons.

6 Lewis Symbol of Element (3-18)

7 The Octet Rule When forming compounds, atoms tend to add or subtract electrons until they are surrounded by eight valence electrons

8 Ionic

9 Let s start by exploring the energetics of ionic compound formation

10 Energetics of Ionic Formation of NaCl As we saw in the last chapter, it takes 495 kj/mol to remove one electron from Na.

11 Energetics of Ionic Formation of NaCl We get 349 kj/mol back by giving those electrons to chlorine.

12 Energetics of Ionic But these numbers don t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

13 Energetics of Ionic There must be a third piece to the puzzle. What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion.

14 The crystal Structure of Sodium Chloride

15 Lattice Energy Lattice energy is the energy required to completely separate a mole of a solid ionic compound into its gaseous ions.

16 Lattice Energy The energy associated with electrostatic interactions is governed by Coulomb s law: E el = κ Q 1Q 2 d where Q 1 and Q 2 are the ionic charges, and d is the distance between the ions in the solid.

17 Lattice Energy and Ionic Charge The lattice energy is strongly influenced by the charges of the ions because the charge on each ion can double or triple. Compound Cation charge Anion charge Lattice energy NaF kj/mol CaF kj/mol CaO kj/mol

18 Lattice Energy and Distance The effect of distance (the sum of ionic radii) is usually small compared to charge effects because size differences are usually small. Compound r (pm) Lattice energy (kj/mol) NaF NaCl NaBr

19 Sample Exercise Magnitudes of Lattice Energies Arrange the ionic compounds NaF, CsI, and CaO in order of increasing lattice energy.

20 Sample Exercise Charges on Ions Predict the ion generally formed by (a) Sr, (b) S, (c) Al.

21 Covalent

22 Covalent In covalent bonds, atoms share electrons. There are several electrostatic interactions in these bonds: Attractions between electrons and nuclei, Repulsions between electrons, Repulsions between nuclei.

23 Multiple Bonds Single Bond - sharing one pair of electrons F Double Bond - sharing two pairs of electrons Triple Bond - sharing three pairs of electrons F O C O N N

24 Bond Length N N N N N N 147 pm 125 pm 110 pm

25 Polar Covalent Bonds Though atoms often form compounds by sharing electrons, the electrons are not always shared equally. Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

26 Electronegativity Electronegativity is the ability of atoms in a molecule to attract electrons to themselves. On the periodic chart, electronegativity increases as you go from left to right across a row. from the bottom to the top of a column.

27

28 Dipole Moment When two atoms share electrons unequally, a bond dipole results. The dipole moment, µ, produced by two equal but opposite charges separated by a distance, r, is calculated: µ = Qr It is measured in debyes (D).

29 Polar Covalent Bonds \ The greater the difference in electronegativity, the more polar is the bond.

30 Sample Exercise 8.4 Bond Polarity In each case, which bond is more polar: (a) B Cl or C Cl, (b) P F or P Cl? Indicate in each case which atom has the partial negative charge.

31 Shared and localized electrons are associated with which type of bonding? A. Metallic B. Covalent C. Ionic

32 Which choice below properly ranks the ionic compounds in terms of increasing lattice energy? A. LiF < LiCl < NaCl B. NaCl < LiF < NaCl C. LiCl < NaCl < LiF D. NaCl < LiCl < LiF

33 Lewis Structures Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

34 Multiple Bonds Single Bond - sharing one pair of electrons F Double Bond - sharing two pairs of electrons Triple Bond - sharing three pairs of electrons F O C O N N

35 Definitions The bond order is the number of electron pairs shared between two atoms. The skeleton structure shows which atoms are bonded to each other. A central atom is bonded to two or more other atoms. A terminal atom is bonded to only one other atom.

36 Writing Lewis Structures PCl 3 Keep track of the electrons: 5 + 3(7) = Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge.

37 Writing Lewis Structures 2. The central atom is the least electronegative element that isn t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26 6 = 20

38 Writing Lewis Structures 3. Fill the octets of the outer atoms. Keep track of the electrons: 26 6 = 20; = 2

39 Writing Lewis Structures 4. Fill the octet of the central atom. Keep track of the electrons: 26 6 = 20; = 2; 2 2 = 0

40 Writing Lewis Structures 5. If you run out of electrons before the central atom has an octet form multiple bonds until it does.

41 Sample Exercise Lewis Structure with a Multiple Bond Draw the Lewis structure for HCN.

42 Sample Exercise Lewis Structure for a Polyatomic Ion Draw the Lewis structure for the BrO 3 ion.

43 Writing Lewis Structures Then assign formal charges. For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. Subtract that from the number of valence electrons for that atom: the difference is its formal charge.

44 Writing Lewis Structures The best Lewis structure is the one in which the atoms bear formal charge closest to 0. puts a negative charge on the most electronegative atom.

45 iclicker Question The formal charges on O 1, O 2 and O 3 in the Lewis structure depicted here for O 3 are: A. -1, +1, B. 0, -1, +1 C. +1, -1, 0 D. -1, +1, 0 1 3

46 Lewis Structure of Ozone This is the Lewis structure we would draw for ozone, O 3.

47 Molecular Model of Ozone But this is at odds with the true, observed structure of ozone, in which both O O bonds are the same length. both outer oxygens have a charge of 1/2.

48 Resonance Structure One Lewis structure cannot accurately depict a molecule like ozone. We use multiple structures, resonance structures, to describe the molecule.

49 Resonance Just as green is a synthesis of blue and yellow ozone is a synthesis of these two resonance structures.

50 Resonance In truth, the electrons that form the second C O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. They are not localized; they are delocalized.

51 Resonance The organic compound benzene, C 6 H 6, has two resonance structures. It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

52 Breakdown of the Octet Rule Case 1: ions or molecules with an odd number of electrons. Case 2: ions or molecules with less than an octet. Case 3: ions or molecules with more than eight valence electrons (an expanded octet).

53 Case 1: Odd Number of Electrons Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons. Example: NO

54 Case 2 : Fewer Than Eight Valence Electrons Example: H, He, Be, B

55 Case 3: More Than Eight Electrons

56 Covalent Bond Strength Bond enthalpy: energy is required to break the bond in one mole of a gaseous substance. H 0 = D(Cl Cl) = +242 kj/mol.

57 Bond Enthalpies for C-H Bonds CH 4 CH 3 + H H 0 = +438 kj/mol CHF 3 CF 3 + H H 0 = +429 kj/mol CHCl 3 CCl 3 + H H 0 = +380 kj/mol the average value (413 kj/mol) for C-H bonds. H 0 is positive endothermic

58 Average Bond Enthalpies

59 Importance of Bond Enthalpies The bond enthalpies can be used to calculate approximate enthalpies of reaction. H reaction = Σ(bond energies of bonds broken) - Σ(bond energies of bonds formed) Example: CH 4 (g) + Cl 2 (g) CH 3 Cl(g) + HCl(g) H 0 = -104 kj/mol H 0 is negative - exothermic

60 Example : How to Estimate Enthalpies of Reaction? Calculate an approximate enthalpy change for the reaction: CH 4 (g) + Cl 2 (g) CH 3 Cl(g) + HCl(g) Lewis structures need to be written to determine the types and numbers of bonds broken and formed.

61 Enthalpies of Reaction So, H = [4D(C H) + D(Cl Cl)] [D(C Cl) + 3D(C H) + D(H Cl)] = [4(413 kj) + (242 kj)] [(328 kj) + 3(413kJ) + (431 kj)] = 104 kj

62 Sample Exercise 8.12 Using Average Bond Enthalpies Using data from Table 8.4, estimate H for the reaction

63 Bond Enthalpy and Bond Length We can also measure an average bond length for different bond types. As the number of bonds between two atoms increases, the bond length decreases.