Chapter 8 Basic Concepts of Chemical Bonding

1. CHEMISTRYThe Central Science 9th Edition Chapter 8 Basic Concepts of Chemical Bonding David P. White Chapter 8

2. Chemical Bonds, Lewis Symbols, and the Octet Rule • Chemical bond: attractive force holding two or more atoms together. • Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals. • Ionic bond results from the transfer of electrons from a metal to a nonmetal. • Metallic bond: attractive force holding pure metals together. Chapter 8

3. Chemical Bonds, Lewis Symbols, and the Octet Rule • Lewis Symbols • As a pictorial understanding of where the electrons are in an atom, we represent the electrons as dots around the symbol for the element. • The number of electrons available for bonding are indicated by unpaired dots. • These symbols are called Lewis symbols. • We generally place the electrons one four sides of a square around the element symbol. Chapter 8

4. Chemical Bonds, Lewis Symbols, and the Octet Rule Lewis Symbols Chapter 8

5. Chemical Bonds, Lewis Symbols, and the Octet Rule • The Octet Rule • All noble gases except He has an s2p6 configuration. • Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). • Caution: there are many exceptions to the octet rule. Chapter 8

6. Ionic Bonding Consider the reaction between sodium and chlorine: Na(s) + ½Cl2(g)  NaCl(s) DHºf = -410.9 kJ Chapter 8

7. Ionic Bonding • The reaction is violently exothermic. • We infer that the NaCl is more stable than its constituent elements. Why? • Na has lost an electron to become Na+ and chlorine has gained the electron to become Cl-. Note: Na+ has an Ne electron configuration and Cl- has an Ar configuration. • That is, both Na+ and Cl- have an octet of electrons surrounding the central ion. Chapter 8

8. Ionic Bonding • NaCl forms a very regular structure in which each Na+ ion is surrounded by 6 Cl- ions. • Similarly, each Cl- ion is surrounded by six Na+ ions. • There is a regular arrangement of Na+ and Cl- in 3D. • Note that the ions are packed as closely as possible. • Note that it is not easy to find a molecular formula to describe the ionic lattice. Chapter 8

9. Ionic Bonding Chapter 8

10. Ionic Bonding • Energetics of Ionic Bond Formation • The formation of Na+(g) and Cl-(g) from Na(g) and Cl(g) is endothermic. • Why is the formation of Na(s) exothermic? • The reaction NaCl(s)  Na+(g) + Cl-(g) is endothermic (H = +788 kJ/mol). • The formation of a crystal lattice from the ions in the gas phase is exothermic: • Na+(g) + Cl-(g)  NaCl(s) H = -788 kJ/mol Chapter 8

11. Ionic Bonding • Energetics of Ionic Bond Formation • Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions. • Lattice energy depends on the charges on the ions and the sizes of the ions: • k is a constant (8.99 x 10 9 J·m/C2), Q1 and Q2 are the charges on the ions, and d is the distance between ions. Chapter 8

12. Ionic Bonding • Energetics of Ionic Bond Formation • Lattice energy increases as • The charges on the ions increase • The distance between the ions decreases. Chapter 8

14. Ionic Bonding • Electron Configurations of Ions of the Representative Elements • These are derived from the electron configuration of elements with the required number of electrons added or removed from the most accessible orbital. • Electron configurations can predict stable ion formation: • Mg: [Ne]3s2 • Mg+: [Ne]3s1not stable • Mg2+: [Ne] stable • Cl: [Ne]3s23p5 • Cl-: [Ne]3s23p6 = [Ar] stable Chapter 8

15. Ionic Bonding • Transition Metal Ions • Lattice energies compensate for the loss of up to three electrons. • In general, electrons are removed from orbitals in order of decreasing n (i.e. electrons are removed from 4s before the 3d). • Polyatomic Ions • Polyatomic ions are formed when there is an overall charge on a compound containing covalent bonds. E.g. SO42-, NO3-. Chapter 8

16. Covalent Bonding • When two similar atoms bond, none of them wants to lose or gain an electron to form an octet. • When similar atoms bond, they share pairs of electrons to each obtain an octet. • Each pair of shared electrons constitutes one chemical bond. • Example: H + H  H2 has electrons on a line connecting the two H nuclei. Chapter 8

17. Covalent Bonding Chapter 8

18. Covalent Bonding • Lewis Structures • Covalent bonds can be represented by the Lewis symbols of the elements: • In Lewis structures, each pair of electrons in a bond is represented by a single line: Chapter 8

19. Covalent Bonding • Multiple Bonds • It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): • One shared pair of electrons = single bond (e.g. H2); • Two shared pairs of electrons = double bond (e.g. O2); • Three shared pairs of electrons = triple bond (e.g. N2). • Generally, bond distances decrease as we move from single through double to triple bonds. Chapter 8

20. Bond Polarity and Electronegativity • In a covalent bond, electrons are shared. • Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons. • There are some covalent bonds in which the electrons are located closer to one atom than the other. • Unequal sharing of electrons results in polar bonds. Chapter 8

21. Bond Polarity and Electronegativity • Electronegativity • Electronegativity: The ability of one atoms in a molecule to attract electrons to itself. • Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). • Electronegativity increases • across a period and • down a group. Chapter 8

22. Bond Polarity and Electronegativity Electronegativity

23. Bond Polarity and Electronegativity • Electronegativity and Bond Polarity • Difference in electronegativity is a gauge of bond polarity: • electronegativity differences around 0 result in non-polar covalent bonds (equal or almost equal sharing of electrons); • electronegativity differences around 2 result in polar covalent bonds (unequal sharing of electrons); • electronegativity differences around 3 result in ionic bonds (transfer of electrons). Chapter 8

24. Bond Polarity and Electronegativity • Electronegativity and Bond Polarity • There is no sharp distinction between bonding types. • The positive end (or pole) in a polar bond is represented + and the negative pole -. Chapter 8

25. Bond Polarity and Electronegativity • Dipole Moments • Consider HF: • The difference in electronegativity leads to a polar bond. • There is more electron density on F than on H. • Since there are two different “ends” of the molecule, we call HF a dipole. • Dipole moment, m, is the magnitude of the dipole: • where Q is the magnitude of the charges. • Dipole moments are measured in debyes, D. Chapter 8

26. Bond Polarity and Electronegativity • Bond Types and Nomenclature • In general, the least electronegative element is named first. • The name of the more electronegative element ends in –ide. • Ionic compounds are named according to their ions, including the charge on the cation if it is variable. • Molecular compounds are named with prefixes. Chapter 8

27. Chapter 8

28. Bond Polarity and Electronegativity Bond Types and Nomenclature Chapter 8

29. Drawing Lewis Structures • Add the valence electrons. • Write symbols for the atoms and show which atoms are connected to which. • Complete the octet for the central atom the complete the octets of the other atoms. • Place leftover electrons on the central atom. • If there are not enough electrons to give the central atom an octet, try multiple bonds. Chapter 8

30. Drawing Lewis Structures • Formal Charge • It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms. • To determine which structure is most reasonable, we use formal charge. • Formal charge is the charge on an atom that it would have if all the atoms had the same electronegativity. Chapter 8

31. Drawing Lewis Structures • Formal Charge • To calculate formal charge: • All nonbonding electrons are assigned to the atom on which they are found. • Half the bonding electrons are assigned to each atom in a bond. • Formal charge is: • valence electrons - number of bonds - lone pair electrons Chapter 8

32. Drawing Lewis Structures • Formal Charge • Consider: • For C: • There are 4 valence electrons (from periodic table). • In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. • Formal charge: 4 - 5 = -1. Chapter 8

33. Drawing Lewis Structures • Formal Charge • Consider: • For N: • There are 5 valence electrons. • In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. • Formal charge = 5 - 5 = 0. • We write: Chapter 8

34. Drawing Lewis Structures • Formal Charge • The most stable structure has: • the lowest formal charge on each atom, • the most negative formal charge on the most electronegative atoms. • Resonance Structures • Some molecules are not well described by Lewis Structures. • Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms Chapter 8

35. Drawing Lewis Structures • Resonance Structures • Example: experimentally, ozone has two identical bonds whereas the Lewis Structure requires one single (longer) and one double bond (shorter). Chapter 8

36. Drawing Lewis Structures Resonance Structures Chapter 8

37. Drawing Lewis Structures • Resonance Structures • Resonance structures are attempts to represent a real structure that is a mix between several extreme possibilities. Chapter 8

38. Drawing Lewis Structures • Resonance Structures • Example: in ozone the extreme possibilities have one double and one single bond. The resonance structure has two identical bonds of intermediate character. • Common examples: O3, NO3-, SO42-, NO2, and benzene. Chapter 8

39. Drawing Lewis Structures • Resonance in Benzene • Benzene consists of 6 carbon atoms in a hexagon. Each C atom is attached to two other C atoms and one hydrogen atom. • There are alternating double and single bonds between the C atoms. • Experimentally, the C-C bonds in benzene are all the same length. • Experimentally, benzene is planar. Chapter 8

40. Drawing Lewis Structures • Resonance in Benzene • We write resonance structures for benzene in which there are single bonds between each pair of C atoms and the 6 additional electrons are delocalized over the entire ring: • Benzene belongs to a category of organic molecules called aromatic compounds (due to their odor). Chapter 8

41. Exceptions to the Octet Rule • There are three classes of exceptions to the octet rule: • Molecules with an odd number of electrons; • Molecules in which one atom has less than an octet; • Molecules in which one atom has more than an octet. • Odd Number of Electrons • Few examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons. Chapter 8

42. Exceptions to the Octet Rule • Less than an Octet • Relatively rare. • Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. • Most typical example is BF3. • Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds. Chapter 8

43. Exceptions to the Octet Rule • More than an Octet • This is the largest class of exceptions. • Atoms from the 3rd period onwards can accommodate more than an octet. • Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density. Chapter 8

44. Strengths of Covalent Bonds • The energy required to break a covalent bond is called the bond dissociation enthalpy, D. That is, for the Cl2 molecule, D(Cl-Cl) is given by H for the reaction: • Cl2(g)  2Cl(g). • When more than one bond is broken: • CH4(g)  C(g) + 4H(g) H = 1660 kJ • the bond enthalpy is a fraction of H for the atomization reaction: • D(C-H) = ¼H = ¼(1660 kJ) = 415 kJ. • Bond enthalpies can either be positive or negative. Chapter 8

46. Strengths of Covalent Bonds • Bond Enthalpies and the Enthalpies of Reactions • We can use bond enthalpies to calculate the enthalpy for a chemical reaction. • We recognize that in any chemical reaction bonds need to be broken and then new bonds get formed. • The enthalpy of the reaction is given by the sum of bond enthalpies for bonds broken less the sum of bond enthalpies for bonds formed. Chapter 8

47. Strengths of Covalent Bonds • Bond Enthalpies and the Enthalpies of Reactions • Mathematically, if Hrxn is the enthalpy for a reaction, then • We illustrate the concept with the reaction between methane, CH4, and chlorine: • CH4(g) + Cl2(g)  CH3Cl(g) + HCl(g) Hrxn = ? Chapter 8

48. Strengths of Covalent Bonds

49. Strengths of Covalent Bonds • Bond Enthalpies and the Enthalpies of Reactions • In this reaction one C-H bond and one Cl-Cl bond gets broken while one C-Cl bond and one H-Cl bond gets formed. • The overall reaction is exothermic which means than the bonds formed are stronger than the bonds broken. • The above result is consistent with Hess’s law. Chapter 8

50. Strengths of Covalent Bonds • Bond Enthalpy and Bond Length • We know that multiple bonds are shorter than single bonds. • We can show that multiple bonds are stronger than single bonds. • As the number of bonds between atoms increases, the atoms are held closer and more tightly together. Chapter 8