1. Bonding Bringing the atoms together

2. Until now, we have been consumed with describing individual atoms of elements. However, isolating individual atoms in most elements is an arduous task given the scale. Most matter is composed of many, many atoms. Matter Pure substances* Elements Compounds Mixtures Homogeneous Heterogeneous *Pure is very hard to achieve in quantity More than one atom

3. Pure substances – typically more than one atom So what is holding the polyatomic elements together? Image by E.B. Watson

4. Covalent bonds – the sharing of electrons 1s 1 to 1s 2 ([He]) 1s 1 & [He] 2s 2 2p 5 to [Ne] [He] 2s 2 2p 5 to [Ne]

5. Two hydrogen atoms in close proximity can share their electrons so that each takes on an electronic structure similar to He – a noble gas. The diatomic H-H system: Covalent bonding Images by E.B. Watson

6. Electron Swapping Best with elements of nearly identical electronegativity. e.g. H 2 Hydrogen is 1s 1 Another proton and electron would make He, a much more stable configuration- so when you bring two hydrogens together they share their electrons so each has 2 (sort of). The Covalent Bond Model

7. This electron sharing is a very strong bond. This tends to by asymmetrical, and therefor difficult to pack into repeating lattices H 2 O 2 Cl 2 F 2 Image from Gill, 1996

8. Image from Pauling, 1970

9. Image modified from Zoltai and Stout, 1984 Diamond’s excited state

10. Sharing an electron in 2p Sharing an electron in two vacancies with 1s in 2p

11. Image after Zoltai and Stout, 1984 Electron distribution

12. Double bonding Two electron pairs are shared Triple bonding Three electron pairs are shared

13. Coordinate covalent bonds – the shared electron is donated by atom. Most carry an overall negative charge (CO 3 ) 2-, (OH) -, (HCO 3 ) -, (PO 4 ) 3-

14. lithium (Li) 1s 2 2s 1 The 2s electrons are delocalized – move around For atoms with appropriate electron structure, electron swapping is quite easy when matter is condensed Metallic bonding Image by E.B. Watson

16. Covalently bonded atoms have no charge, but most examples will orient themselves in an electric field Couloumb’s law So there must be an ionic character to the bond

18. Ionization - losing or gaining valence electrons Can only attract so far – “solid spheres” Ionic bonding

19. Electronegativity! Ionic and covalent bonding are endmembers there is a range of between these two extremes. How do we know if the bonding is dominated by one form or another. Measuring electronegativity E AB > ½ E(A 2 ) + ½E(B 2 ) E AB = ½ E(A 2 ) + ½E(B 2 ) +   is excess energy due to ionic attraction (+ and -)  = 23(X A - X B ) ---X is electronegativity --

20. Periodicity of electronegativity

22. Ionic-covalent character makes molecules dipolar F, X = 4 H, X = 2.1  = 23|X A - X B |

23. Silcates Si-OSulfides M-S Fe1.65Fe-S:|1.8-2.5| = 0.7 S2.5 Si1.8 Si-O: |1.8-3.5| = 1.7 O3.5 Silcates vs. Sulfides

24. A compound containing just two elements "-ide“. The metal is named first: NaClsodium chloride Multiple valence state Fe 3+ or Fe 2+. Higher valence state with the suffix "- ic" and the lower with "-ous": Fe 2 O 3 ferric oxide FeOferrous oxide Today we tend to use a more explicit approach Fe 2 O 3 iron(III) oxide FeOiron(II) oxide Ionic compounds

25. More than two elements with polyatomic anions such as hydroxideOH - carbonate(CO 3 ) 2- nitrate(NO 3 ) - phosphate(PO 3 ) 3- The compound names follow pretty logically: KOHpotassium hydroxide CaCO 3 calcium carbonate Polyatomic anions can have more than one configuration, e.g. nitrate(NO 3 ) - nitrite(NO 2 ) -

26. CO 2 carbon dioxide(Greek "di" for 2) COcarbon monoxide(Greek "mono" for 1) CCl 4 carbon tetrachloride(Greek "tetra" for 4) Empirical Formulas Always used in ionic compounds NaClCaF 2 CH 2 O Covalent compounds

27. Used for largely covalent compounds While CH 2 O accurately describes the ratio of elements in glucose, it fails to characterize the whole molecule Molecular Formulas

28. Structural formulas show the geometry and bonding

29. You are already familiar with atomic mass… Recall that the atomic mass of an element is the mass of the individual isotopes proportionally distributed. Molecular mass is obtained by summing the atomic mass for the number of atoms of each element. CaF 2 at. masstot. mass 1 atom Ca40.1 au40.1 au 2 atoms F19.0 au38.0 au m.m.78.1 au

30. Avogadro's number Mg + ½ O 2  MgO A mole is the number of atoms or molecules (6.02 x 10 23 ) needed to make a mass (g) equivalent to the atomic or molecular mass.

31. Image from Gill, 1996

32. Generally, directionality of bonds may be important in covalent species - bonds are generally weaker* Directionality less important for ionic bonds - radius ratios determine structures. * exception diamond

33. As atoms approach, there is a coulombic attraction E p (potential energy) = -e 2 /r But as ions approach, their electron distributions overlap, so there is a repulsion. E p =-e 2 /r + (be 2 )/r 12

34. Balance between the overall attraction of atoms and the inter- electron repulsions Attraction F = k c q 1 q 2 r -2 Repulsion F is elastic rebound Images by E.B. Watson

35. Image from Gill, 1996

36. NaCl vs. NaF F:1s 2 2s 2 2p 5 Cl:1s 2 2s 2 2p 6 3s 2 3p 5 Both are strongly electronegative So what differences might you expect?

37. As it tuns out, the ions are not hard spheres, but more squishy! There is some variation allowed in the range of “acceptable” radius ratios (cation/anion) in this configuration. Image from Blackburn and Dennen, 1994

38. Ionic bonding is a reduction of excess energy because of (+) or (-) charges. Ionically-bonded compounds are charge balanced controls stoichiometry, or ratio of atoms in the compound Na + + Cl - = NaCl Ca ++ + 2F - = CaF 2

39. The Ionic Bond Model Changing coordination Bonding

40. The atoms are not stationary. TVibration 25 o C 0.025 nm 600 o C0.060 nm Image from Blackburn and Dennen, 1994

41. Coordination

43. Certain anion/cation ratios produce a predictable polyhedra in IONIC bonding.

44. No Yes Yes, but... Preferred

45. Symmetrically-packed structures Atoms are spheres Bonds must be either non-directional or highly symmetical Cubic packing Hexagonal packing

46. Molecular structures Composed of atoms characterized by strongly directional bonding with low symmetry. Result is clusters, chains, and layers of atoms that are strongly bonded internally, but weakly bonded to one another. Examples: ice, organic molecules

48. Cubic Closest Packing Colors here denote different layers of atoms

49. Hexagonal Closest Packing

52. Symmetrically-packed structures can be monatomic or multiatomic. Symmetrical packing applies to the larger anions. Monoatomic - Native Elements (Au, Ag, S) Multiatomic - Most oxides and many silicates

53. Quartz: Two oxygens for every silicon. Charge of Si = +4, O = -2 We then say the stoichiometery is 1:2 Important: Si, Ti, Al, Fe, Mg, Ca, Na, K, P, O, S Today’s $10K question: how do we reconcile the overall stoichiometry (i.e. the chemical formula SiO 2 ) with the coordination of SiO 4 ?

54. The tetrahedra link together such that the oxygens are shared This illustrates the general principles of the architecture of crystals

55. Endmember types Polyhedra-frame structures Bonding dominantly ionic Anions group around cations to make coordination polyhedra Silicates (SiO 2 minerals) built of silicate tetrahedra which share corners with one another - this makes a framework. It imparts certain properties to the minerals: strong, hard, etc.

57. Image from Klein and Hurlbut, 1985 Halite (NaCl) Two ways to model NaCl: left, atoms of Na and Cl shown, right, atoms of Cl shown with coordination polyhedra (octahedra)

59. How are these related? QUARTZQUARTZ Six-fold symmetrical arrangement of tetrahedra all joined by common oxygen atoms.

60. Rule 1. Interatomic Distance. A coordination polyhedron of anions is formed about each cat-ion. The cation-anion distance being determined by the radius sum and the coordination number of the cation by the radius ratio. Rule 2. Electrostaic Valency Principle. In a stable coordination structure, the total strength of the valence bonds that reach an anion from all neigh-boring cations is equal to the charge of the anion. Rule 3. Sharing of Polyhedral Elements I. The existence of edges. and particularly of faces, common to two coordination polyhedra decreases the stability of ionic structures. Rule 4. Sharing of Polyhedral Elements II. In a crystal containing different cations, those with large valence and small coordination tend not to share polyhedral elements with each other. Rule 5. Principle of Parity. The number of essentially different kinds of constituents in a crystal structure tends to be small. Pauling's Rules