Bonding in Minerals

1. Bonding in Minerals The Glue That Holds Minerals Together GLY 4200 –Fall, 2012

2. Types of Bonds • Intramolecular • Ionic • Covalent • Metallic • Intermolecular • Hydrogen • Van der Waals

3. Definition of Bonding • A chemical bond is an attraction between atoms brought about by: • A sharing of electrons between two atoms or, • A complete transfer of electrons • When a chemical bond is formed, energy is released • Breaking chemical bonds requires energy

4. Substances Formed by Bonding • When two or more atoms of the same element bond together, a molecule is formed – example, hydrogen H2 • When 2 or more atoms of different elements combine together chemically, a compound is formed – example, water H2O • Most minerals are compounds

5. Ionic Bonding • Ionic bonding is the result of electrostatic attraction between two oppositely charged ions • Positive ions are formed from metals (usually) and negative ions are usually formed from non-metals

6. Halite • Halite, NaCl, is a classical example of an ionically bonded substance • The sodium donates an electron to chlorine to complete the eight-electron subshell on chlorine

7. Physical Properties of Ionically Bonded Crystals • Ionic bonding is non-directional • Ionically bonded minerals may yield ions to solution • Moderate hardness • Fairly high to very high melting points & boiling points • Poor thermal & electrical conductors except near the melting points

8. Polarization • Polarity is the distortion of the electron cloud of one atom by another. • A standard example is often hydrogen chloride (HCl)

9. Does Size Affect Polarizing Power? • Yes, and so does electronegativity • The greater the electronegativity, the greater the polarizing power • So for hydrogen halogen compounds: • Bond polarity has a huge hand in determining chemistry

10. Relative Size of Ions • The size mismatch of the anions and cations is of importance also • If two ions are similar in size, then they exist quite happily • If there is a size mismatch, then is it quite likely that covalent bonding will occur

11. Size Mismatch • NaCl melts at 801°C, strong attraction between particles in solid lattice structure (Ionic bonding likely) • AlCl3 sublimes (goes from solid to gas not via the liquid phase) at 180°C, so there are no strong attractions present (Covalent bonding likely)

12. Polarizing Cations • If the cation is small and highly charged, it has a large polarizing power • If the anion is large and has a relatively low charge, then it is said to have a large polarizability • In the first case, the anion is being polarized by the cation • There will be a significant degree of covalent character to the bond

13. Non-Existent Compounds • There are some ionic compounds that do not exist at all • Aluminum carbonate is an example • The aluminum 3+ cation is so small and highly polarizing that is completely distorts the large CO32- ion into self-decomposition • Instead of Al2(CO32-)3, carbon dioxide is driven off, leaving aluminum oxide

14. Ionic Bond Nomenclature • Compounds ending in –ide are simple binary compounds containing 2 elements - even if there is no metal • e.g H2S – hydrogen sulfide • Ending in –ate means oxygen is present • e.g. CaS  = calcium sulfide • CaSO4  = calcium sulfate

15. Ionic Bond Nomenclature II • Ending in –ite less oxygen present than in –ate compounds • e.g. NaS = sodium sulfide • NaSO4 = sodium sulfate • NaSO3 = sodium sulfite

16. Covalent Bonding • Covalent bonds involve a complete sharing of electrons and occur most commonly between atoms that have partially filled outer shells or energy levels • Thus, if the atoms are similar in electronegativity then the electrons will be shared

17. Carbon • Carbon forms covalent bonds • The electrons are in hybrid orbitals formed by the atoms involved as in this example: ethane • Diamond is strong because it involves a vast network of covalent bonds between the carbon atoms in the diamond C2H6

18. Physical Properties of Covalently Bonded Crystals • Covalent bonds are directional and molecules are often formed. • Covalently bonded crystals do not yield ions to solutions, as ionically bond crystals sometimes do • Covalent crystals have very high melting points & boiling points

19. Octet Rule • The idea that the noble-gas configuration is a particularly favorable one which can be achieved through formation of electron-pair bonds with other atoms is known as the octet rule • Present-day shared electron-pair theory is based on the premise that the s2p6 octet in the outermost shells of the noble gas elements above helium represents a particularly favorable configuration

20. Basis of Octet Rule • By allowing each nucleus to claim half-ownership of a shared electron, more electrons are effectively “seeing” more nuclei, leading to increased electrostatic attractions and a lowering of the potential energy

21. Fluorine • Noble gas configuration (in this case, that of neon, s2p6) is achieved when two fluorine atoms (s2p5) are able to share an electron pair,which becomes the covalent bond • Only the outer (valence shell) electrons are involved

22. Covalent Bonds Between Different Elements • Hydrogen chloride (aka hydrochloric acid) • The hydrogen has a helium structure, and the chlorine an argon structure

23. Octet Limitations – Light Elements • For the lightest atoms the octet rule must be modified, since the noble-gas configuration will be that of helium, which is simply s2 rather than s2p6 • Thus we write LiH as Li:H, where the electrons represented by the two dots come from the s orbitals of the separate atoms

24. Octet Limitations – Heavy Elements • The octet rule applies quite well to the first full row of the periodic table (Li through F), but beyond this it is generally applicable only to the non-transition elements, and even in many of these it cannot explain many of the bonding patterns that are observed • The principal difficulty is that a central atom that is bonded to more than four peripheral atoms must have more than eight electrons around it if each bond is assumed to consist of an electron pair • In these cases, we hedge the rule a bit, and euphemistically refer to the larger number of electrons as an “expanded octet”

25. Metallic Bonding • A metallic bond occurs whenpositive metal ions like Cu+2 or Fe+3 are surrounded by a "sea of electrons" or freely-moving valence electrons. • The valence electrons are not bound to any particular cation, but are free to move throughout the metallic crystal

26. Sea of Electrons • In the picture, the red circles are metal cations packed in a crystal lattice • The black dots represent the "sea" of freely moving valence electrons 

27. Minerals with Metallic Bonding • Only native metals display metallic bonding • Alkaline metals are for too reactive to be found uncombined in nature • Only a few minerals, such as gold, silver, copper and the platinum group are metallically bound

28. Conductivity Properties • Metals are good conductors of electricity • Electric current is a movement of free electrons • Substances with partial metallic bonding may be semiconductors • Metals are good conductors of heat • Heat is transferred by the increased speed of electrons

29. Flexibility Properties of Metallic Bonding • The model of metallic bonding explains the flexibility properties of metals • Metals are ductile - They can be drawn into wires because electrons are mobile. • Metals are malleable - They can be hammered into sheets due to mobility of electrons • Metals are tenacious – they do not break easily

30. Electronic Forces in Metals • Strong attraction between positive nuclei and the electrons • The positive ions repel as do the negative electrons • The electrons move constantly, but some electrons will always be between the layers creating an attraction and keeping them attracted to one another

31. Explanation of Metallic Properties • An impact will allow a shearing effect as there is a degree of repulsion between layers • The sea of electrons allows movement of ions, therefore pure metals are not brittle

32. Other Physical Properties • Low hardness • Low melting point & boiling point

33. Optical Properties • Metallically bonded minerals are opaque – • This is often true at very small thicknesses, such as the 30 micron thickness of a thin section • Metallically bonded substance usually show metallic luster • Weathering may make this luster dull Thin section of llanite, a hypabyssal rhyolite porphyry dike – opaque mineral grains are magnetite

34. Intermolecular Bonds • Bonds which hold molecules together are called intermolecular bonds • In minerals, the concept of a “molecule” is often inapplicable, but the term is still used

35. Hydrogen Bonding • In some substances, hydrogen is bonded to elements which are quite electronegative, and which possess “lone pairs” of electrons • Examples include water and ammonia • Hydrogen bonding leads to the many anomalous properties of water and ammonia

36. Hydrogen Bond Image • The δ+ hydrogen is so strongly attracted to the lone pair that it is almost as if you were beginning to form a co-ordinate bond • It doesn't go that far, but the attraction is significantly stronger than an ordinary dipole-dipole interaction

37. Relative Bond Strength • Hydrogen bonds have about a tenth of the strength of an average covalent bond, and are being constantly broken and reformed in liquid water • If you liken the covalent bond between the oxygen and hydrogen to a stable marriage, the hydrogen bond has "just good friends" status • On the same scale, van der Waals attractions represent mere passing acquaintances!

38. Relative Boiling Points

39. Relative Boiling Points • The boiling point of the hydride of the first element in each group is abnormally high • In the cases of NH3, H2O and HF there must be some additional intermolecular forces of attraction, requiring significantly more heat energy to break • These relatively powerful intermolecular forces are described as hydrogen bonds

40. Water • Each water molecule can potentially form four hydrogen bonds with surrounding water molecules • There are exactly the right numbers of δ+ hydrogens and lone pairs so that every one of them can be involved in hydrogen bonding

41. Ammonia and Hydrogen Fluoride • In the case of ammonia, the amount of hydrogen bonding is limited by the fact that each nitrogen only has one lone pair • In a group of ammonia molecules, there aren't enough lone pairs to go around to satisfy all the hydrogens • In hydrogen fluoride, the problem is a shortage of hydrogens • In water, there are exactly the right number of each • Water could be considered as the "perfect" hydrogen bonded system

42. Hydrogen Bonding in Biology • Hydrogen bonding also holds the DNA double helix together • During sexual reproduction, the hydrogen bonds break, allowing each parent to pass on a strand of DNA • The strands recombine to form a new double helix, a combination of genetic material from each parent

43. Residual Bonding Forces • All molecules experience intermolecular attractions, although in some cases those attractions are very weak • Even in a gas like hydrogen, H2, if you slow the molecules down by cooling the gas, the attractions are large enough for the molecules to stick together eventually to form a liquid and then a solid

44. Hydrogen and Helium • In hydrogen's case the attractions are so weak that the molecules have to be cooled to 21 K (-252°C) before the attractions are enough to condense the hydrogen as a liquid • Helium's intermolecular attractions are even weaker - the molecules won't stick together to form a liquid until the temperature drops to 4 K (-269°C)

45. Van der Waals Bonding • There are two types of Van der Waals forces • Dispersion forces are also known as "London forces" (named after Fritz London who first suggested how they might arise) • Dipole-dipole interactions

46. Electrical Attractions • Attractions are electrical in nature • In a symmetrical molecule like hydrogen, however, there doesn't seem to be any electrical distortion to produce positive or negative parts • But that's only true on average

47. Distortion of Electron Cloud • The lozenge-shaped diagram represents a small symmetrical molecule - H2, perhaps, or Br2 • The even shading shows that on average there is no electrical distortion

48. Mobile Electrons • But the electrons are mobile, and at any one instant they might find themselves towards one end of the molecule, making that end δ- • The other end will be temporarily short of electrons and so becomes δ +

49. Temporary Fluctuating Dipoles • An instant later the electrons may well have moved up to the other end, reversing the polarity of the molecule

50. Momentary Dipoles • This constant "sloshing around" of the electrons in the molecule causes rapidly fluctuating dipoles even in the most symmetrical molecule • It even happens in monatomic molecules - molecules of noble gases, like helium, which consist of a single atom