Lecture 3 Crystal Chemistry Part 2: Bonding and Ionic Radii

1. Lecture 3Crystal Chemistry Part 2: Bonding and Ionic Radii Salt, Calcite and Graphite models

2. Chemical Bonding in Minerals Bonding forces related to electrically charged particles – negative attracts positive Bond strength controls most physical and chemical properties of minerals In general, the stronger the bond, the harder the crystal, higher the melting point, and the lower the coefficient of thermal expansion

3. Chemical Bonding in Minerals Five general types bonding types: Ionic Covalent Metallic van der Waals Hydrogen Commonly different bond types occur in the same mineral

4. Chemical Bonds Electrical in nature- responsible for most mineral properties 1) Ionic Na: low 1st Ionization Potential 1s2 2s2 2p6 3s1 Na  e-+ Na+(Sodium ion has a Neon configuration) 1s2 2s2 2p6 Cl: high e-neg takes the e- Cl-(Cl- ion has Argon configuration) Now they have opposite charges & attract = bond Bonding is strong (e.g. Salt has high melting point) But easily disrupted by polarized solvents (e.g. water) Poor electrical conductors; electron strongly held by anion Strength  (1/bond length) & valence Also non-directional so symmetrical packing is possible (Isometric crystal system is common in Alkali Metal – Halogen Salts). If electronegativity of anion and cation differs by 2.0 or more will be mostly ionic , say about 70%.

5. Halite (NaCl)- An Example of Ionic Bonding Na+ fits into interstices Na+ lost an electron shell, smaller; Cl- gained an electron, repels nucleus, larger

6. Ionic Bonding Example: NaCl Na (1s22s22p63s1) –> Na+(1s22s22p6) + e- Cl (1s22s22p63s23p5) + e- –> Cl- (1s22s22p63s23p6)

7. Problem 1 • Write down the electron configuration for neutral Chlorine Cl and for Chloride Ion Cl- using the info from lecture 2.

8. Chemical Bonds 2) Covalent Consider 2 close Cl atoms, each = 1s2 2s2 2 p6 3s2 3p5 If draw closer until overlap an outer orbital, can share whereby 2 e- "fill" the remaining 3p shell of each Cl Low energy condition causes electrons to stay overlapped; results in a strong bond  Cl2 This is the covalent or shared electron bond Usually stronger than Ionic bond

9. Covalent bonding – sharing of valence electrons Cl:1s2 2s2 2p6 3s2 3p5 so 7 electrons in outer shell “The sharing of an electron pair … constitutes a single bond” S&P p54.

10. Chemical Bonds 3) Metallic Bonding Metals have few, loosely held valence electrons If closely pack them can get up to 12 nearest neighbors This causes a high density of valence e- around any given atom & also a high density of neighbor atoms around the loose valence e- These become a sea of mobile electrons Metals are excellent conductors

11. Chemical Bonds 4) Van der Waals Bonds Weakest bond – due localized excess charge Usually between neutral molecules (even large ones like graphite sheets) Weakness of the bond is apparent in graphite cleavage Caused by momentary correlations in the charge polarity of adjacent atoms

12. More Detail • Now let’s look at the bond types in more detail

13. Ionic Bonds Dominate Most Mineral Geometry • Most minerals have a strong ionic component. • Mostly covalent Ion complexes SiO4-4, CO3 --, etc. are ionically bonded to metal ions to achieve neutrality. Calcite CaCO3

14. Ionic Bond Properties • Results in minerals displaying moderate degrees of hardness and specific gravity, moderately high melting points, high degrees of symmetry • Poor conductors • Strength of ionic bonds are related to: 1) the spacing between ions 2) the charge of the ions Stronger bond has a higher melting point

15. Compound Bond Strength = Melting Point vs. interionic distance, ionic charge 9 17 35 53 Sodium Na+ with various anions Small inter-ionic distance = higher melting point 12 20 38 56 +2 cations Small inter-ionic distance = higher melting point A (ångström) = 10 -10m 3 11 19 37 +1 cations Li F is an exception

16. Interionic Distance vs. Hardness 4 12 20 38 56 22 21 12 11 Closer Interionic Distance = Increased Bond Strength (Hardness)

17. Covalent Bonding • formed by sharing of outer shell electrons • strongest of all chemical bonds • most covalent minerals are insoluble in acids • high melting points, • hard, nonconductive • have low symmetry due to multi-directional bonding. • common among elements with high numbers of vacancies in the outer shell (e.g. C, Si, Al, S) Diamond

18. Tendencies for Ionic vs. Covalent Pairing Ionic Pairs Covalent Pairs Si-O, C-O, S-O, N-O, P-O

19. Difference in electronegativity of the elements involved tells us if one member is more attractive to electrons i.e. forms ionic bonds. F to Na 4.1 – 1 = 3.1, very different, so Na-F bond very ionic in character. Si-O difference 3.5-1.8 = 1.7 ~ 50% covalent Covalent-Ionic continuum Covalent Ionic

20. Metallic Bonding • Atomic nuclei and inner filled electron shells in a “sea” of electrons made up of unbound valence electrons. • Typical of elements with low ionization potential. Valence electrons easily stripped. • Yields minerals with minerals that are soft, ductile/malleable, highly conductive (due to easily mobile electrons). • Non-directional bonding produces high symmetry

21. Van der Waals (Residual) Bonding • created by weak bonding of oppositely depolarized electron clouds • commonly occurs around covalently bonded elements • produces solids that are soft, very poor conductors, have low melting points, with low symmetry crystals and strong cleavage.

22. Hydrogen Bonding example ICE H+ • Electrostatic bonding between an H+ ion with an anion or anionic complex or with a polarized molecules • Weaker than ionic or covalent; stronger than Van der Waals Close packing of polarized molecules Anions polarized H2O molecule Ice One Hydrogen bond shown as red line above

23. Summary of Bonding Characteristics

24. Crystal Chemistry Crystals can be classified into 4 types: 1. Molecular Crystals Neutral molecules held together by weak van der Waals bonds Rare as minerals Mostly organic Weak and readily decompose, melt, cleave, etc. Example: graphite

25. Crystal Chemistry 2. Covalent Crystals Atoms of similar high e-neg and toward right side of Periodic Table Also uncommon as minerals (but less so than molecular) Network of strong covalent bonds with no weak links Directional bonds  low symmetry and density Example: diamond

26. hard-sphere model Crystal Chemistry The diamond structure All carbon atoms in IV coordination FCC unit cell polyhedral model blue C only ball-and-stick model

27. Crystal Chemistry 3. Metallic Crystals Atoms of similar e-negand toward left side of Periodic Table Metallic bonds are directionless bonds  high symmetry and density Pure metals have same sized atoms Closest packing 12 nearest mutually-touching neighbors Cubic Closest Packing (CCP) abcabcabc stacking = FCC cell (face-centered cubic AKA cubic close packed) Hexagonal Closest Packing (HCP) ababab = hexagonal cell Also BCC in metals, but this is not Closest Packing More on coordination and closest packing next time

28. Crystal Chemistry 4. Ionic Crystals Most minerals First approximation: Closest-packed array of oxygen atoms Cations fit into interstices between oxygens, balance the negative charges. Negative charges mostly due to oxide ions O- Different types of interstitial sites available Cations occupy only certain sites where can fit Only enough cations to attain electrical neutrality

29. Multiple Bonding in Minerals • Graphite – covalently bonded sheets of C loosely bound by Van der Waals bonds. • Mica – strongly bonded silica tetrahedra sheets (mixed covalent and ionic) bound by weak ionic and hydrogen bonds • Calcite: Cleavage planes commonly correlate to planes of weak ionic bonding versus strong covalent bonds • in CO3--