Chapter 8 Concepts of Chemical Bonding

1. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 8Concepts of Chemical Bonding John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

2. November 15Chapter 8 – Chemical Bonding Section 1 and 2 Bonding types Lewis Structures Ionic bonding HW: 1,3,4,7 to 19 odd ,24,25 There will be a quiz any day on hw questions

3. Chemical Bonds • Three basic types of bonds: • Ionic • Electrostatic attraction between ions • Covalent • Sharing of electrons • Metallic • Metal atoms bonded to several other atoms

4. Ionic BondingElectron Transfer

5. Energetics of Ionic Bonding As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium.

6. Energetics of Ionic Bonding We get 349 kJ/mol back by giving electrons to chlorine.

7. Energetics of Ionic Bonding • But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

8. 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.

9. Energetics of Ionic Bonding • 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.

10. Q1Q2 d Eel =  Lattice Energy • This third piece of the puzzle is the lattice energy: The energy required to completely separate a mole of a solid ionic compound into its gaseous ions. • The energy associated with electrostatic interactions is governed by Coulomb’s law:

11. Lattice energy depends on the charges of 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. • Lattice energy increases as • * The charges on the ions increase • * The distance between the ions decreases.

13. 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.

14. 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 NaCl(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

15. Energetics of Ionic Bonding By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.

16. Born-Haber CycleApplication of Hess’s Law • Vaporize the metal (enthalpy of vaporization) • Na (s)  Na (g) • Break diatomic nonmetal molecules (if applicable) (bond enthalpy) • ½ Cl2 (g)  Cl- • Remove electron(s) from metal (ionization energy) • Na (g)  Na+ (g) + e-

17. 4 Add electron(s) to nonmetal (electron affinity) Cl (g) + e- Cl- (g)

18. 5 Put ions together to form compound (lattice energy) • Na+ + Cl- NaCl (s)

19. Overall Reaction: • Na (s) + ½ Cl2 (g)  NaCl (s) • This is useful because all quantities are directly measurable except lattice energy. The Born-Haber cycle can be used to calculate lattice energy from the other values.

22. Lattice Energy • Lattice energy, then, increases with the charge on the ions. • It also increases with decreasing size of ions.

23. Examples – Which in each pair would have the greatest lattice energy? • LiF or LiCl • NaCl or MgCl2 • KBr or KI • MgCl2 or MgO • CaO or NaF

24. Examples – Which in each pair would have the greatest lattice energy? • LiF or LiCl • NaCl or MgCl2 • KBr or KI • MgCl2 or MgO • CaO or NaF

25. Energetics of Ionic Bonding • These phenomena also helps explain the “octet rule.” • Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.

26. November 17 • Properties of ionic compounds/ Naming ionic compounds • COVALENT BONDS • LEWIS STRUCTURES • RESONANCE STRUCTURES • EXCEPTIONS TO OCTECT RULE • STRENGTHS OF COVALENT BONDS

27. Properties of Ionic Substances • Crystalline structure, result from electrostatic forces that keep ions in a rigid well defined three dimensional arrangement. • hard and brittle. • Represented by empirical formulas • High melting points • Do not conduct electricity in the solid state but they conduct electricity when are dissolved in water or in the liquid state

28. Naming Ionic Compounds • Binary compounds – ending ide • When metal has several oxidation states indicate the o.n. with the stock system. Examples • Lead (IV) oxide • Manganese (III) Nitrate • Iron (II) Phosphate • Calcium Sulfate

29. COVALENT BONDING Sharing of electrons. Typical bond between non metals. Can be polar or non polar. Can be single, double or triple bonds.

30. General properties of substances containing covalent bonds • They are molecular substances. • Generally they have low melting points. • They exists in the 3 states. • In the solid state many are pliable (like plastics). • They are not conductors of heat or electricity.

31. Naming Covalent Compounds • Use IUPAC nomenclature with roman numerals for the oxidation numbers

32. Bond Types and Nomenclature • In general, the least electronegative element is named first. • The name of the more electronegative element ends in –ide. • Compounds are named according to their ions (or perceived ions), including the charge on the cation if it is variable. • Compounds of metals with high O.N. have properties similar to covalent compounds and sometimes are named using the convention for molecular compounds • Mn2O7 can be called Dimanganese heptoxide or Manganese(VII)oxide

33. Bond Types and Nomenclature

34. Bond Types and Nomenclature

35. Covalent Bonding • In these bonds atoms share electrons. Pairs of electrons become the glue that bind to atoms together. • There are several electrostatic interactions in these bonds: • Attractions between electrons and nuclei • Repulsions between electrons • Repulsions between nuclei

36. Polar Covalent Bonds • Although 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.

37. Electronegativity: • The ability of atoms in a molecule to attract electrons to itself. • 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.

38. Electronegativity and Bond Polarity • Difference in electronegativity is a gauge of bond polarity: • electronegativity differences between 0 and 0.5 result in non-polar covalent bonds (equal or almost equal sharing of electrons); • electronegativity differences around between 0.5 and 2 result in polar covalent bonds (unequal sharing of electrons); • electronegativity differences greater that 2 result in ionic bonds (transfer of electrons).

39. Non Polar Covalent bondsEqual electron sharing • Happens between atoms with same electronegativity. • All diatomic molecules have non polar covalent bonds!

40. Polar Covalent Bonds • When two atoms share electrons unequally, a bond dipole results. • The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:  = Qr • It is measured in debyes (D).

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

42. Lewis Structures(Lewis electron-dot structures) Lewis structures are representations of molecules showing all electrons, bonding and nonbonding. A line between 2 atoms represents a pair of electrons (2 electrons = 1 line )

43. 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.

44. 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, then 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.

45. 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. PCl3 Writing Lewis Structures 5 + 3(7) = 26

46. Writing Lewis Structures • 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

47. Writing Lewis Structures • Fill the octets of the outer atoms. Keep track of the electrons: 26  6 = 20  18 = 2

48. Writing Lewis Structures • Fill the octet of the central atom. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0

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

50. Examples – Draw Lewis Structures for each: • H2O • CO2 • NCl3 • SO2 • SO3