Models of Chemical Bonding

1. Models of Chemical Bonding

2. Models of Chemical Bonding 9.1 Atomic Properties and Chemical Bonds 9.2 The Ionic Bonding Model 9.3 The Covalent Bonding Model 9.4 Between the Extremes: Electronegativity and Bond Polarity 9.5 An Introduction to Metallic Bonding

3. Chemical Bonds Chemical Bonds • The attractive forces that hold atoms or ions together to form molecules or crystals Octet rule • atoms tend to gain, lose, or share valence electrons to get an octet. • Everything wants to be like a noble gas. • exceptions • near He obey duet rule • Transition metals • n = 3 and above

4. Figure 9.1 A general comparison of metals and nonmetals

5. Types of Chemical Bonding Typically: 1. Metal with nonmetal: electron transfer leads to ionic bonding 2. Nonmetal with nonmetal: electron sharing leads to covalent bonding 3. Metal with metal: electron pooling leads to metallic bonding

6. Figure 9.2 The three models of chemical bonding

7. The A group number gives the number of valence electrons. Place one dot per valence electron around the four sides of the element symbol. Do not pair dots until all four sides have an electron. . . . : . . . : N . N . . N N : . . : . Lewis Electron-Dot Symbols - A method for depicting valence electrons and interactions of atoms For main group elements - Example: Nitrogen, N, is in Group 5A and therefore has 5 valence electrons.

8. Figure 9.3 Lewis electron-dot symbols for elements in Periods 2 and 3 Nonmetals - The number of unpaired dots indicates the number of electrons it gains, or the number of covalent bonds it usually forms. Metals – The total number of dots is the maximum number of electrons it may lose when forming a cation.

9. Sec 9.2 Ionic Bonding • In ionic bonding, electrons are gained or lost, the resulting bonds are based on electrostatic attraction. • ex Na has 1 valence e, Cl has 7 • If Na could only get rid of 1, if Cl could only gain 1…. • When sodium metal is placed in Cl2 gas, they react by transferring 1 e- from Na to Cl to form Na+ and Cl-. Now each has an octet. Because both now have a charge they are attracted to each other to form NaCl.

10. PROBLEM: Use partial orbital diagrams and Lewis symbols to depict the formation of Na+ and O2- ions from the atoms, and determine the formula of the compound. O2- Na Na 2s 2p 3s 3p O 2s 2p . . Na : 3s 3p + O : : 2Na+ + O 2- : : . : . Na SAMPLE PROBLEM 9.1 Depicting Ion Formation PLAN: Draw orbital diagrams for the atoms and then move electrons to make filled outer levels. It can be seen that 2 sodiums are needed for each oxygen. SOLUTION: 2 Na+

11. + F- 1s22s22p6 Li+ Li 1s 2s 2p 1s 2s 2p F- + + F 1s 2s 2p 1s 2s 2p . : . Li : F : Li+ + : F - : + : : Three ways to represent the formation of Li+ and F- through electron transfer. Figure 9.4 Electron configurations Li 1s22s1 + F 1s22s22p5 Li+ 1s2 Orbital diagrams Lewis electron-dot symbols

12. Energy in Ionic Bonding • Li(g) Li+(g) + e- IE1 = 520 kJ • F(g) + e-  F-(g) EA = -328 kJ • So the process would appear to be endothermic • Li(g) + F(g)  Li+(g) + F-(g) E = 192kJ • Overall the process is very exothermic, this is because of the lattice energy. • the enthalpy change of gaseous ions coalescing into a crystalline solid. • Indicates the strength of the two ions attraction • Influences melting point, hardness, and solubility • Ionic solids exist only because the lattice energy drives the unfavorable electron transfer.

13. Calculating lattice energy • Lattice energy cannot be directly measured, so it is found by using Hess’s Law. • The enthalpy change for an overall reaction is the sum of the enthalpy changes of the reactions which make it up. • Lattice energies are calculated by using a Born-Haber Cycle • A series of chosen steps from elements to ionic compounds for which all the enthalpies are known. • The steps are hypothetical and not the actual steps of the process

14. Figure 9.6 The Born-Haber cycle for lithium fluoride

15. Periodic Trends in Lattice Energy Coulomb’s Law charge A X charge B electrostatic force a distance2 energy = force X distance therefore charge A X charge B electrostatic energy a distance cation charge X anion charge a DH0lattice electrostatic energy a cation radius + anion radius

16. Trends in lattice energy • Effect of ion size. • Increasing the size of the ions decreases lattice energy, therefore attraction between cations and anions decreases down in a group • Effect of ionic charge. • Increasing the charge of the ions increases the lattice energy.

17. Figure 9.7 Trends in lattice energy

18. Properties of ionic compounds • Ionic compounds are hard, rigid, and brittle • This is a result of ions being held in specific positions in a crystal. So a crystal retains it’s shape until enough energy is applied to shift positions and crack the crystal.

19. Figure 9.8 Electrostatic forces and the reason ionic compounds crack.

20. Properties of ionic compounds • Do not conduct electricity in the solid state • Ions in fixed positions • Do conduct when melted or dissolved • Ions can move independently

21. Solid ionic compound Molten ionic compound Ionic compound dissolved in water Figure 9.9 Electrical Conductance and Ion Mobility

22. Properties of ionic compounds • High melting and boiling points (all solid at RT) • Enough energy must be supplied to free ions from the attractions of the surrounding ions • Ionic compounds vaporize as ion-pairs even though no “molecules” exist in the crystal

23. Table 9.1 Melting and Boiling Points of Some Ionic Compounds Compound mp (0C) bp (0C) CsBr 636 1300 NaI 661 1304 MgCl2 714 1412 KBr 734 1435 CaCl2 782 >1600 NaCl 801 1413 LiF 845 1676 KF 858 1505 MgO 2852 3600

24. Figure 9.10 Vaporizing an ionic compound.

25. Sec 9.3 Covalent Bonding • Elements can also form octets by sharing e- between them, the bonds that result are called covalent bonds. • usually occurs in nonmetal/nonmetal compounds. • More compounds are covalent than ionic. • A single Cl atom has 7 valence electrons, in a sample of pure Cl one atom cannot steal an electron from another, so they share to form Cl2. • Molecule – a compound formed by 2 or more atoms joined by covalent bonds that behaves as a single particle.

26. Covalent Bonding • atoms share electrons by overlapping orbitals so electrons can exist in the orbitals of both atoms at once. • These shared or bonding pairs of electrons are represented by lines in structures. • Other valence electrons that are not involved in bonding are called unshared or lone pairs. • Every pair of electrons shared between atoms is a bond • 1 pair – single bond • 2 pairs – double bond , stronger • 3 pairs – triple bond, strongest • aka bond order • Lewis structures of Cl2, O2, N2

27. Figure 9.11 Covalent bond formation in H2.

28. Figure 9.12 The attractive and repulsive forces in covalent bonding.

29. Bond Energy • Bond energy or Bond Strength - the energy required to overcome the attraction of covalently bonded atoms. • It is defined as energy required to break bonds in 1 mole of gaseous atoms. • Bond energy depends on the specific elements involved. • It can vary from molecule to molecule so table values are averages.

32. Internuclear distance (bond length) Covalent radius Internuclear distance (bond length) Covalent radius Internuclear distance (bond length) Covalent radius Internuclear distance (bond length) Covalent radius Figure 9.13 Bond length and covalent radius.

33. Bond length, bond energy, and bond order are closely related • Higher bond order is shorter, and stronger for a given set of atoms • With a constant bond order, longer bonds are usually weaker.

35. (b) C = O, C - O, C O Bond length: C - O > C = O > C O Bond strength: C O > C = O > C - O SAMPLE PROBLEM 9.2 Comparing Bond Length and Bond Strength PROBLEM: Using the periodic table, but not Tables 9.2 and 9.3, rank the bonds in each set in order of decreasing bond length and bond strength: (a) S - F, S - Br, S - Cl PLAN: (a) The bond order is one for all and sulfur is bonded to halogens; bond length should increase and bond strength should decrease with increasing atomic radius. (b) The same two atoms are bonded but the bond order changes; bond length decreases as bond order increases while bond strength increases as bond order increases. SOLUTION: (a) Atomic size increases going down a group. (b) Using bond orders we get Bond length: S - Br > S - Cl > S - F Bond strength: S - F > S - Cl > S - Br

36. Properties of covalent cmpds • The physical properties of molecular compounds, are not related to the strength of their covalent bonds. • Most covalent compounds have low m.p. and b.p. because the strong covalent bonding is typically isolated within molecules. The attractions between separate molecules, called intermolecular forces, are what must be overcome to melt or boil these covalent substances. • The physical properties of network covalent solids, are related to the strength of their covalent bonds. • In these substances there are no individual molecules, the covalent bonding extends in 3-D throughout the substance. • Ex Quartz (SiO2) very hard, mp 1550°C • Diamond (C) hardest known substance, mp 3550°C

37. Strong covalent bonding forces within molecules Weak intermolecular forces between molecules Strong forces within molecules and weak forces between them. Figure 9.14

38. Figure 9.15 Covalent bonds of network covalent solids.

39. Properties of covalent cmpds • Most covalent substance are poor electrical conductors, when solid, liquid, or dissolved.

40. Sec. 9.4 Between the extremes • Most real bonds fall somewhere between the ideal of ionic or covalent bonding theory. • The type of bond atoms form depends on electronegativity • some atoms attract e- more strongly than others, we say these are more electronegative • electronegativity increases going right and up the table • Covalent bonds in which e- are not shared equally because of electronegativity differences are called polar covalent bonds

41. Figure 9.16 The Pauling electronegativity (EN) scale.

42. Figure 9.17 Electronegativity and atomic size.

43. Representation of Polar Bonds

44. SAMPLE PROBLEM 9.3 Determining Bond Polarity from EN Values PROBLEM: (a) Use a polar arrow to indicate the polarity of each bond: N-H, F-N, I-Cl. (b) Rank the following bonds in order of increasing polarity: H-N, H-O, H-C. PLAN: (a) Use Figure 9.16(button at right) to find EN values; the arrow should point toward the negative end. (b) Polarity increases across a period. SOLUTION: (a) The EN of N = 3.0, H = 2.1; F = 4.0; I = 2.5, Cl = 3.0 N - H F - N I - Cl (b) The order of increasing EN is C < N < O; all have an EN larger than that of H. H-C < H-N < H-O

45. Nonpolar, Polar, or Ionic • In general if the electronegativity difference between two bonded atoms is: • 0, usually between identical nonmetal atoms, called nonpolar covalent • < .4 , mostly covalent • .4 to 1.7, 2 different nonmetals called polar covalent • > 1.7 , usually nonmetals and reactive metals, is mostly ionic • Note: there is no perfect ionic bond.

46. 3.0 2.0 DEN 0.0 Figure 9.18 Boundary ranges for classifying ionic character of chemical bonds.

47. Figure 9.19 Percent ionic character of electronegativity difference (DEN).

48. Figure 9.20 Li F

49. Figure 9.21 Properties of the Period 3 chlorides.

50. Sec. 9.5 Metallic Bonding • Solid Metals and metal alloys have metallic bonding • Electron Sea Model • All the metal atoms contribute their valence electrons to a delocalized pool of electrons. The metal cations are held together by attraction to the delocalized electrons.