CHAPTER 7

1. CHAPTER 7 • Covalent Bonds and Molecular Structure

2. Introduction • Attractive forces that hold atoms together in compounds are called chemical bonds. • The electrons involved in bonding are those in the outermost (valence) shell. • Electrons repel each other so the bonds try to get as far apart as possible (especially in the smaller molecules. • Hybridization allows this separation to occur! • This establishes electronic and molecular geometry.

3. Introduction • Chemical bonds are classified into two types: Ionic bonding results from electrostatic attractions among ions, which are formed by the transfer of one or more electrons from one atom to another. Ionic bonds are common between metals and nonmetals. Covalent bonding results from sharing one or more electron pairs between two atoms. Covalent bonds occur between nonmetals. • We will be concentrating on the bonding and geometry of covalent molecules.

4. The Octet Rule • Representative elements achieve noble gas configurations in most of their compounds. • Metals tend to lose electrons to drop back to a noble gas configuration. • Nonmetals tend to gain or share electrons to attain a noble gas configuration. • With the exception of He, all noble gases have 8 valence electrons. • This tendency to achieve a valence “noble gas” configuration is called the octet rule. • The octet rule was considered absolute until the late 1960s.

5. DEN • Bonds between metals and non-metals tend to be ionic(DEN>1.7) • Bonds between non-metals tend to be covalent (DEN <1.7) • The magnitude ofDEN establishes the direction and strength of the dipole moment in a compound or ion.

6. Lewis Dot Representations of Atoms • The Lewis dot representation or Lewis dot formula is used to show bonding between atoms. • Only valence electrons, only outer s and p electrons,are involved in bonding. • You may represent a shared pair of electrons as a dash or bond in these drawings, which makes these Lewis Dot-Dash representations. • Hybridization of orbitals is • VERY important!

7. Hybridization • Hybridization is the mixing of 2 or more orbitals to form a new type of atomic orbital. • This allows most of the nonmetals and B and Al to share 8 electrons (or more) with other atoms. • This would not be possible otherwise. • Hybridization also allows for the maximum separation of orbitals within the molecule (or ion).

8. Hybridization • Beginning with Be, all metalloids and nonmetals of Row 2 and 3 hybridize. • This allows Be to share its electrons and form up to 2 bonds with different atoms or groups. • Technically this is 4 outer electrons and violates the octet rule. • Beginning with B, the Row 2 elements can form up to 4 bonds to different atoms or groups. • Al can also have a maximum of 4 bonds. • B and Al compounds generally have 3 bonds (6 outer electrons which violates the octet rule).

9. Hybridization • Row 3 elements beginning with P can form more than 4 bonds. • This violates the octet “rule.” • These elements use their empty d orbitals in the formation of these hybrids. • FYI. These hybrids were not recognized until the early 1970s with the routine use of x-ray crystallography.

10. Hybridization • Orbitals may either be used to share electrons (form bonds – up to one pair/bond per orbital overlap) or to house a pair of nonbonded electrons. • The elements hybridize to produce to maximum separation of bonds and nonbonded electron pairs. • Nonbonded electrons require more ‘elbow room’ than bonds. • The electronic and molecular geometry are determined by hybridization.

11. Lewis Dot Representations of Hybridized Atoms Elements in the same group have the same Lewis Dot configurations.

12. Ionic Bonding • Cations or positive (+) ions are atoms or groups of atoms that have lost 1 or more electrons. Anions or negative (–) ions are atoms or groups of atoms that have gained 1 or more electrons. Ions may contain covalent bonds.

13. Ionic Bonding • Metals react with nonmetals to form ionic compounds. IA VIIA 2 Li + F2 2 Li+ F– silver metal yellow gas white solid mp = 842 °C

14. Ionic Bonding • We can also use Lewis formulas to represent the neutral atoms and the ions they form. Li+ ions contain two e–, same as He, 1s2. F- ions contain ten electrons same as Ne, 1s22s22p6. Li+ ions are isoelectronic with He. F- ions are isoelectronic with Ne. Isoelectronic species contain the same number of valence electrons.

15. Ionic Bonding • Cations become isoelectronic with the preceding noble gas. • Anions become isoelectronicthe with following noble gas.

16. Driving Force for Reactions • Elements will gain or lose electrons in order to achieve a stable outer orbitals. • Metals alwayslose electrons. • Nonmetals usuallygain electrons unless they are reacting with O or F (then they lose electrons). • The driving force for ALL reactions is the octet, or more specifically, the stable outer orbital. • Remember Li, Be, and B form ions that are isoelectronic with He.

17. Ionic Bonding • Hydrogen • H forms ionic bonds with IA + IIA metals. • Hydrogen exists as H– or H: – in these compounds. • Transition metal hydrides do exist but are unstable in an oxygen atmosphere. • ALL other hydrogen compounds are covalent! • This specifically includes the IIIA and post-transition metal hydrides. • These compounds do react with water to liberate H2 just like the other metal hydrides. • Unlike the standard metal hydrides, these can covalently transfer H to other covalent molecules.

18. Covalent Bonding • Covalent bonds are formed when two atoms share electrons. • Bonding between 2 atoms • sharing 2 electrons = single covalent bond • sharing 4 electrons = double covalent bond • sharing 6 electrons = triple covalent bond • single covalent bond = 1 s bond • double covalent bond = 1 s and 1 p bond • triple covalent bond = 1 s and 2 p bonds

19. Types of Covalent Bonds • Maximum overlap bonds are called sigma bonds,s • Minimum overlap bonds are called pibonds, p • Only p orbitals can form p bonds • Only one (1) sigma bond can form between any two atoms, all additional bonds formed between the two atoms are pi bonds.

20. Covalent Bonding • Lewis dot representation • H molecule formation This is a sigma bond and can be designated as H—H This is the Lewis dash formula

21. Covalent Bonding • Extremes in bonding • Pure covalent bonds have the electrons equally shared by the atoms (dipole moment = 0). • Pure ionic bonds exist between ions, the cation has completely lost one (or more electrons) and the anion has gained those electrons. • Most compounds fall somewhere between these two extremes with unequal sharing of electron pairs.

22. Lewis Dot Formulas for Molecules and Polyatomic Ions • Homonuclear Diatomic Molecules • Hydrogen, H2 Fluorine, F2 Nitrogen, N2

23. Lewis Dot Formulas for Molecules and Polyatomic Ions • Heteronuclear Diatomic Molecules • Hydrogen Halides • Hydrogen fluoride, HF Hydrogen chloride, HCl Hydrogen bromide, HBr

24. Lewis Dot Formulas for Molecules and Polyatomic Ions Ammonia molecule, NH3 Ammonium ion, NH4+ • Water, H2O

25. Counting Electrons Around Atoms in Structures • General Rules • Each bond counts as 2 electrons. • Each multiple bond counts electrons as equal to 2 times the number of bonds. • Each nonbonded electron pair counts as 2 electrons. • A single electron counts as a single electron.

26. Counting Electrons Around Atoms in Structures • For example 8 electrons 8 electrons 8 electrons 8 electrons Octet satisfied for each atom

27. Assigning Formal Charges to Atoms • General Rules • Each bond counts as 1 electron for each atom. • Each nonbonded electron pair counts as 2 electrons for that atom. • Each non-paired electron counts as one electron for that atom. • Total the number of electrons around the atom and subtract this from the group number of the atom. • The difference = the charge • too many = – charge too few = + charge

28. Assigning Formal Charges to Atoms • For example 6 electrons: VIA – 6 = 0 no charge -1 -1 +2 7 electrons 7 electrons: VIA – 7 = -1 4 electrons: VIA – 4 = +2

29. Resonance • Charges and p bonds can move. • This “rearrangement” is called resonance. • For example, using SO3 This is used to explain the properties shown by the compound

30. Resonance • These resonance structures do not give an accurate picture of SO3. • SO3 is a gas. Gases are not expected to be charged. • A better representation is probably

31. Limitations of the Octet Rule for Lewis Formulas • Species That Violate the Octet Rule covalent compounds of Be covalent compounds of IIIA Group species containing an odd number of electrons species in which the central element must have a share of more than 8 valence electrons to accommodate all substituents compounds of the d- and f-transition metals

32. Limitations of the Octet Rule for Lewis Formulas • In cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations. • The central atom does not.

33. Limitations of the Octet Rule for Lewis Formulas • Write dot and dash formulas for BBr3. This gives The bromines each have an octet, boron is surrounded by six electrons

34. Polar and Nonpolar Covalent Bonds • Nonpolar covalent bonds occur when electrons are shared equally. • These have a symmetrical charge distribution. • This seldom occurs with different elements. • H2 N2

35. Polar and Nonpolar Covalent Bonds • Polar covalent bonds result from unequally shared electrons. • These molecules have asymmetrical charge distribution and have dipole moments as a result of this asymmetry. • The greater the electronegativity difference between atoms, the greater the dipole moment.

36. Polar and Nonpolar Covalent Bonds Consider H–F and H–I

37. Continuous Range of Bonding Types • All bonds have some ionic and some covalent character (unless DEN = 0). • Even with a very smallDEN = 0.4, HI is about 17% ionic. • The greater the electronegativity differences the more polar the bond.

38. Continuous Range of Bonding Types • Molecules whose centers of positive and negative charge do not coincide are polar or are said to have a dipole moment. • Dipole moment has the symbol m. • m is the product of the distance, d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q

39. Dipole Moments • Molecules that have a small separation of charge have a small m. • Molecules that have a large separation of charge have a large m.

40. Dipole Moments For a compound to be polar, there must be at least one polar bond present or one lone pair of electrons. The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments donot cancel one another. Symmetrical molecules, even if they have very polar bonds, are nonpolar.

41. Two Simple Theories of Covalent Bonding • Valence Shell Electron Pair Repulsion Theory • VSEPR was developed by R. J. Gillespie in the 1950’s as an improvement over the existing Valence Bond Theory • Valence Bond Theory • VB (L. Pauling – 1930’s & 40’s)was the first theory to use hybridized orbitals to explain observations of molecular bonding and geometry.

42. VSEPR and VB • VSEPR and VB Theories are virtually identical • The two theories vary slightly in their approach. • You may use whichever Theory you feel more comfortable with. • The important thing is the final molecular and electronic geometry of the molecule.

43. Stereochemistry • Stereochemistry is the study of the 3 dimensional shapes of molecules. • Some questions to examine: • Why the interest in molecular shapes? • What role does molecular shape play in life? • How is molecular shape determined? • How can molecular shapes be predicted?

44. Electronic vs Molecular Geometry • Electronic geometry describes the spatial arrangement of orbitals around an atom. • This arrangement is determined by the hybridization of the central atom. • Molecular geometry describes the general shape of the atoms arranged around the central atom. • Remember the electronic geometry describes where the orbital and therefore bonds are. • The hybridization of the central atom determines the electronic and therefore the molecular geometry of the molecule.

45. VSEPR Theory • VSEPR dictates that the regions of high electron density around the central atom are as far apart as possible to minimize repulsions. • Key to this is that nonbonded electron pairs occupy more space than bonded electrons do. • There are five basic orbital configurations. • These shapes are based upon the number of regions of high electron density within the molecule. • Molecular shapes are based upon these basic orbital configurations.

46. VSEPR Theory • Two regions of high electron density linear bond angle = 180°

47. VSEPR Theory • Three regions of high electron density trigonal planar bond angles = 120°

48. VSEPR Theory • Four regions of high electron density dihedral bond angles = 109.5o tetrahedral

49. VSEPR Theory • Five regions of high electron density trigonal bipyrimidal bond angles = 90 °, 120 °, and 180°

50. VSEPR Theory • Six regions of high electron density octahedral bond angles = 90°, and 180°