Chapter 10. CHEMICAL BONDING

1. + Na Cl - Na+ + Cl Chapter 10. CHEMICAL BONDING Lewis Electron-Dot Symbols He H Li Be B C N O F Ne Na Mg Al Si P S Cl Ar The Ionic Bonding Model • electrons are transferred from one element to another and bonds are formed between the two oppositely charged ions

2. Cl + Mg Cl - + 2 Mg2+ Cl

3. Energy Considerations in Ionic Bonding Consider the following reactions, Now we can determine that, • the reaction in the gas phase is endothermic.

4. the reaction of the gas phase ions, to form the solid is quite exothermic and is called the lattice energy The lattice energy is the energy required to break a solid ionic compound into a mole of gaseous ions. (See Chapter 12, Section 12.7 and 12.9) • difficult to determine but is important because it gives us a way to compare or quantify the strength of the ionic bond.

5. Lattice Energy (Chapters 12.7 and 12.9, Do problems 58, 58, 73-76 from Ch. 12) • the lattice energy indicates the strength of the ionic interaction which influences the melting point, hardness and solubility of ionic compounds. • the lattice energy cannot be measured directly but we can use Hess’ Law to determine it. • lattice energies are determined by means of a Born-Haber cycle which is a series of chosen steps, from elements to ionic compounds, for which all the enthalpies are known except the lattice energy. • because the enthalpy is a state function we choose hypothetical steps whose enthalpy changes we can measure which take us from reactants to LiF formation. These steps are not the actual steps that occur when lithium reacts with fluorine.

7. We begin with the elements in their standard states Li(s) and F2(g). Step 1. Vaporize lithium metal. BDE Step 2. Break F2 up into atoms, Step 3. Ionize Li. Step 4. Ionize F. Step 5. Form crystalline solid from the gaseous reactant ions. The reaction of solid lithium and fluorine gas to form the salt, LiF is,

8. or

10. eg. Determine lattice energy for Na2S(s) given: 1st electron affinity of S is 203 kJ mol-1 2nd electron affinity of S is -694 kJ mol-1 ionization energy of Na is 496 kJ mol-1 the heat of sublimation of solid S is 277 kJ mol-1 the heat of sublimation of solid Na is 107 kJ mol-1 and the heat of formation of Na2S is -366 kJ mol-1

11. Periodic Trends in Lattice Energy From Coulomb’s Law, and since then cations and anions lie as close to each other as possible so that the distance between them is the sum of their radii

12. as we move down a group of metals or non-metals the ionic size increases so the lattice energy decreases

13. as the charge on the ions increases, the lattice energy increases Mg2+ O2- Li+ F- r 76 pm 133 pm 72 pm 140 pm DlatticeHo 1050 kJ mol-1 3923 kJ mol-1

14. Using the Model to Explain the Properties of Ionic Solids • ionic solids are hard, brittle, and rigid • ionic solids don’t conduct electricity but the liquids do as well as solutions of the ionic compound

15. the melting and boiling points of ionic compounds are extremely high because it requires enormous amounts of energy to disrupt the forces between the ions and form ion pairs. • it should be noted that the vapor exists as ion pairs not individual ions and the condensed phases do not exist as separate molecules.

16. boiling point melting point NaF NaCl NaBr NaI

17. H H H H Covalent Bonding • sharing electrons is the principle way that atoms interact chemically, based on the number of known covalent compounds compared to the number of known ionic compounds or H-H

18. or H F H F H F H H C H H Basic Lewis Diagrams Systematic method for drawing Lewis diagrams eg. Draw the Lewis diagram for CH4 1. Count up the number of valence electrons C - 4e 4H - 4e 8e 2. Decide which atoms are bonded together and draw a bond, which is equivalent to two electrons between the bonded atoms. 3. Determine how many electrons are left and add them in as lone pairs about the appropriate nucleus trying to obey the “octet rule” or ensuring a full outer shell. For the present example we are finished.

19. eg. Draw the Lewis structure for NH3 eg. Draw the Lewis structure for N2

20. eg. Draw the Lewis structure for ethyne, C2H2. eg. Draw the Lewis structure for ethanol, C2H5OH.

21. Bond Strengths and Bond Lengths

23. C-X bond lengths and bond strengths, X=C, N, O, F C-X bond lengths and bond strengths, X=F, Cl, Br, I

24. Bond lengths and bond strengths for different CC bonds

25. C O C O C O • eg. Using the periodic table to rank the bonds in each set in order of decreasing bond length and strength, • S-F, S-Br, S-Cl , ,

26. Using the Model to Explain the Properties of Covalent Solids • the model proposes that electron sharing between pairs of atoms leads to strong bonds between those atoms within a molecule • there are two types of compounds with covalent bonds that have very different properties Molecular solids • low mpt (large range) • soft

27. Network Covalent Solids Diamond SiO2 3550 oC • very high mpt 1550 oC • hard • compounds held together by covalent bonds are poor electrical conductors in all phases or when dissolved in water

28. Infrared (IR) Spectroscopy • characterization of substances

30. 2917 cm-1 O F 1033 cm-1 1049 cm-1 Cl O 732 cm-1 1746 cm-1 Br 611 cm-1 S I 820 cm-1 533 cm-1

31. +

32. d- d+ H F Polar Covalent Bonds • the bond energies of H-H and F-F are 432 and 159 kJ mol-1, respectively. • we might consider, then, that the bond strength of H-F is intermediate between the two, ~296 kJ mol-1. • the bond strength of H-F is, however, 565 kJ mol-1. Linus Pauling reasoned that the difference is due to an electrostatic (charge) contribution to the HF bond energy. fluorine holds the shared electrons closer to it than does H, thereby giving the bond some “ionic character” Electronegativity: the relative ability of a bonded atom to attract shared electrons.

33. Pauling Electronegativity Scale.

34. in general the smaller the atom the higher the electronegativity

35. d- d+ H F Bond Polarity • the HF bond is polar. F has a higher electronegativity so the electrons in the bond are “attracted” more to F than H. the arrow points to the negative end of the polar bond or to the more electronegative bond. Bond Character • it was remarked earlier that HF had “partial ionic character” • due to differences in electronegativities, covalent compounds, other than homonuclear diatomics, have some partial ionic character

36. “non-zero covalent character” “0 % ionic character” : always some sharing of electrons : H2, N2, Cl2, etc.

37. The difference in electronegativity, then, can give a measure of the polarity of the bond. eg. Use the polar arrow to indicate the polarity of each bond, N-H, F-N, I-Cl. eg. Rank the following bonds in order if increasing polarity; H-N, H-O, H-C

38. Properties of period 3 chlorides

39. Metallic Bonding: in a piece of metal what holds the atoms together? • the “electron-sea model” proposes that all the metal atoms in the sample contribute their valence electrons to form an electron sea which is delocalized throughout the metal. The positive cores are submerged in this electron sea in an orderly array. • the valence electrons are “shared” among all atoms in the substance • the solid is held together by the mutual attraction of the metal cations for the mobile, highly delocalized electrons

40. Using the Model to Explain the Properties of Metallic Solids • the picture of the metallic solid is a regular array of ionic cores in a mobile sea of electrons Conductivity • the sea of electrons is mobile so electrons are easily transferred through the solid making metals excellent conductors of electricity • metals are also good conductors of heat because the delocalized electrons can transfer heat better than localized electrons in a covalent bond or in ionic solids Mechanical Properties

41. Melting Point • typically melting points (and boiling points) are quite high since the cationic core and its electron(s) must break away from the others • the alkali earth metals boil at a higher temperature because they form 2+ ions and twice as many electrons meaning stronger bonding