1. Crystal Chemistry Mineral – “…defined, but generally not fixed, chemical composition…” Mineral – “…defined, but generally not fixed, chemical composition…” Modern geology – Geochemistry or Geophysics Modern geology – Geochemistry or Geophysics Geophysics – application of physical principles to study of earth Geophysics – application of physical principles to study of earth Geochemistry – application of chemical principles to study of earth Geochemistry – application of chemical principles to study of earth high T or low T high T or low T

2. Coming up… a review of basic chemistry Coming up… a review of basic chemistry Elements Elements protons, neutrons, electrons protons, neutrons, electrons Bond types and controls on bonds Bond types and controls on bonds

3. Nuclear Chemistry Atomic number (Z) – number of protons Atomic number (Z) – number of protons Specific for particular elements (periodic table) Specific for particular elements (periodic table) Neutrons – about same mass as protons, different number of neutrons make isotopes Neutrons – about same mass as protons, different number of neutrons make isotopes

4. Atomic weight – sum of weight of neutrons and protons Atomic weight – sum of weight of neutrons and protons Isotopes - Superscript in front of element symbol, atomic weight exact Isotopes - Superscript in front of element symbol, atomic weight exact Elements – atomic weight is the average of abundance of isotopes Elements – atomic weight is the average of abundance of isotopes

5. Stable Isotopes Oxygen – Z = 8; three isotopes Oxygen – Z = 8; three isotopes Average bulk earth abundance: Average bulk earth abundance: 16 O – 99.757% 16 O – 99.757% 17 O – 0.038% 17 O – 0.038% 18 O – 0.205% 18 O – 0.205% Materials (minerals, water, air, shells, etc) have variable ratios of these isotopes Materials (minerals, water, air, shells, etc) have variable ratios of these isotopes  18 O = ratio of 18 O/ 16 O sample to 18 O/ 16 O standard  18 O = ratio of 18 O/ 16 O sample to 18 O/ 16 O standard

6. Radioactive isotope Potassium (z=19) Potassium (z=19) 40 K has 21 neutrons 40 K has 21 neutrons Natural abundance = 0.0117% Natural abundance = 0.0117% Radioactive Radioactive Decays to 40 Ar, basis of one type of age dating Decays to 40 Ar, basis of one type of age dating Half life = 1.248 x 10 9 a Half life = 1.248 x 10 9 a 39 K has 20 neutrons 39 K has 20 neutrons Natural abundance = 93.3% Natural abundance = 93.3% Stable (not radioactive) Stable (not radioactive) 41 K has 22 neutrons 41 K has 22 neutrons Natural abundance = 6.7% Natural abundance = 6.7%

7. Chemical Reactions Based on electron transfers, charge balance Based on electron transfers, charge balance If number of electrons = number of protons, no electrical charge If number of electrons = number of protons, no electrical charge Orbit nucleus in systematic way Orbit nucleus in systematic way Organized according to energy levels Organized according to energy levels Shells filled according to energy Shells filled according to energy

8. Electron Quantum number Quantum number - reflects energy of electron Quantum number - reflects energy of electron Unique for each electron Unique for each electron No two electrons in atom can have same quantum number No two electrons in atom can have same quantum number Controls how electrons fill shells Controls how electrons fill shells Controls their chemical reactivity Controls their chemical reactivity

9. Formation of ions Ions – excess or deficit of electrons relative to protons Ions – excess or deficit of electrons relative to protons Anions – net negative charge Anions – net negative charge Cations – net positive charge Cations – net positive charge Valence or Oxidation state is the value of the charge on the ion Valence or Oxidation state is the value of the charge on the ion

10. Configuration of valence electrons controls whether gain or lose electron Configuration of valence electrons controls whether gain or lose electron Metals – typically lose one or two valence electron: form cations Metals – typically lose one or two valence electron: form cations Non-metals – typically require a few electrons to fill valence shells: form anions Non-metals – typically require a few electrons to fill valence shells: form anions

11. Valence shells fill systematically – see table 3-3 for how shells filled Valence shells fill systematically – see table 3-3 for how shells filled Atomic number 1-20 and 31-38 – fill s & p subshells Atomic number 1-20 and 31-38 – fill s & p subshells Between atomic number 20 and 31 – shells fill from internal subshells – fill 3d shell (4s shell filled) Between atomic number 20 and 31 – shells fill from internal subshells – fill 3d shell (4s shell filled) Transition metals Transition metals Elements may have differing numbers of shells filled Elements may have differing numbers of shells filled E.g. Ferrous and Ferric iron E.g. Ferrous and Ferric iron

12. Fig. 3-3 Noble Gases, He, Ne, Ar, Kr Lose electrons (cations) to become noble gas core Note the various oxidation states for the transition metals Ferric Fe (+3) Ferrous Fe (+2) Metallic Fe Gain electrons (anions)

13. Clearly – gain or loss of electrons important Clearly – gain or loss of electrons important “Quantified” as property called Electronegativity “Quantified” as property called Electronegativity

14. Electronegativity Defined by Linus Pauling Defined by Linus Pauling Propensity of element to gain or lose electron Propensity of element to gain or lose electron Based on arbitrary scale: Li = 1, C = 2.5, F = 4 Based on arbitrary scale: Li = 1, C = 2.5, F = 4 Low electronegativity - the more likely to lose electron form cations Low electronegativity - the more likely to lose electron form cations High electronegativity – likely to gain electron to form anions High electronegativity – likely to gain electron to form anions See Table 3-4 for values See Table 3-4 for values

15. Coming up 1)Abundance of elements on earth 2)Types of electron sharing bonds – ionic, covalent, metallic 3)How to estimate bond types from electronegativity

16. Earth abundances of elements What elements are most abundant? What elements are most abundant? These elements will make up common minerals These elements will make up common minerals What part of earth do they occur? What part of earth do they occur? Crust? Crust? Bulk earth? Bulk earth?

17. Crust - 8 common elements Crust - 8 common elements O 2-, Si 4+, Al 3+, Fe 2+,3+, Ca 2+, Na +, K + and Mg 2+ O 2-, Si 4+, Al 3+, Fe 2+,3+, Ca 2+, Na +, K + and Mg 2+ Most minerals are made of these elements Most minerals are made of these elements Determination of crustal abundance – simply collect large number of samples and measure Determination of crustal abundance – simply collect large number of samples and measure Bulk earth composition Bulk earth composition The same 8 elements are common in bulk, but different ratios The same 8 elements are common in bulk, but different ratios

18. Table 3.6 90% of atoms are Si, Al, & O Crust Core/ Mantle

19. Bulk Earth composition: Bulk Earth composition: Difficult to assess – impossible to directly sample mantle or core Difficult to assess – impossible to directly sample mantle or core Estimated by Estimated by Mass and density based on geophysical measurements Mass and density based on geophysical measurements Composition of mantle magmas and xenoliths Composition of mantle magmas and xenoliths Composition of meteorites Composition of meteorites

20. Chemical bonding Eight common elements (plus all others) bond to form minerals Eight common elements (plus all others) bond to form minerals Bonding controls spatial arrangement of atoms Bonding controls spatial arrangement of atoms Two categories Two categories Sharing of valence electrons: ionic, covalent and metallic Sharing of valence electrons: ionic, covalent and metallic No sharing: van der Waals and hydrogen No sharing: van der Waals and hydrogen These 5 types of bonds are “end members” These 5 types of bonds are “end members” Rarely just one type or the other Rarely just one type or the other However: We’ll consider most minerals to be ionically bonded However: We’ll consider most minerals to be ionically bonded

21. Ionic Bonding Transfer of electron(s) from one element to another Transfer of electron(s) from one element to another Results in filled valence shells of both Results in filled valence shells of both The electrostatic attraction keep atoms together The electrostatic attraction keep atoms together The distance between ions depends on attractive forces (Coulomb law) and repulsive forces (Born repulsion) The distance between ions depends on attractive forces (Coulomb law) and repulsive forces (Born repulsion)

22. Fig. 3-4 Attractive forces Repulsive forces Face centered cubic lattice arrangement of halite Bonding in Halite Equilibrium distance = 2.8 Å Fig. 2-10

23. Ionic bonding Ions bond so that positive = negative charges Ions bond so that positive = negative charges Minerals must be electrically neutral Minerals must be electrically neutral NaCl (Halite), Na(Mg,Fe,Li,Al) 3 Al 6 [Si 6 O 18 ](BO 3 ) 3 (O,OH,F) 4 (tourmaline – not all ionic bonds here) NaCl (Halite), Na(Mg,Fe,Li,Al) 3 Al 6 [Si 6 O 18 ](BO 3 ) 3 (O,OH,F) 4 (tourmaline – not all ionic bonds here) Characteristics: Characteristics: Ions act like spheres Ions act like spheres Alternating cations and anions Alternating cations and anions One of the strongest bonds One of the strongest bonds Brittle because like ions repel Brittle because like ions repel Cleavage is common Cleavage is common

24. Covalent Bonds Electrons shared when orbitals of two different elements overlap Electrons shared when orbitals of two different elements overlap Shared by only two atoms Shared by only two atoms Differs from metallic (later – all atoms share electrons) Differs from metallic (later – all atoms share electrons) Electrons move around nucleus of both atoms Electrons move around nucleus of both atoms

25. Examples – Diamond and Graphite Examples – Diamond and Graphite Diamond Diamond Stable Ne configuration by either gain or loss of 4 electrons Stable Ne configuration by either gain or loss of 4 electrons Ionic bonding not possible because all electrons exactly the same electronegativity Ionic bonding not possible because all electrons exactly the same electronegativity One carbon won’t “steal” electron from another One carbon won’t “steal” electron from another Instead share electrons – very strong bonding Instead share electrons – very strong bonding

26. Fig. 3-5 Covalent bonding in diamond 4 orbitals shown as bonds, call  bonds  bonds distorted  bonds distorted Each bold line represents another similar bond

27. Graphite Fig. 3-6 Similar  bonds, but only in layers Additional Sharing electrons,  bonds.

28. Metallic bonds A type of covalent bond A type of covalent bond Electrons shared without systematic change in orbitals Electrons shared without systematic change in orbitals Free to move throughout crystal structure Free to move throughout crystal structure Formed with low electronegativity – weakly held valence electrons Formed with low electronegativity – weakly held valence electrons

29. Relation between valence- dependent bonds Most bonds not purely ionic, covalent or metallic Most bonds not purely ionic, covalent or metallic Amount of bond type depends on electronegativity (tendency to give up electrons) Amount of bond type depends on electronegativity (tendency to give up electrons) Greater difference in electronegativity between ions means more ionic characteristic Greater difference in electronegativity between ions means more ionic characteristic

30. Only 1 anion (of 8 common elements) Only 1 anion (of 8 common elements) Oxygen Oxygen Electronegativity of O = 3.5 Electronegativity of O = 3.5 Electronegativity of other common elements range from 0.8 (K) to 1.8 (Si) Electronegativity of other common elements range from 0.8 (K) to 1.8 (Si)

31. Qualitative difference in electronegativity: Qualitative difference in electronegativity: O-K = 3.5 – 0.8 = 2.7, more ionic characteristics O-K = 3.5 – 0.8 = 2.7, more ionic characteristics O-Si = 3.5 – 1.8 = 1.7, less ionic characteristics O-Si = 3.5 – 1.8 = 1.7, less ionic characteristics Possible to quantify % ionic bonding: Possible to quantify % ionic bonding: O-element bonding of 8 common elements ranges from 50% ionic (Si-O) to 80% ionic (K-O) O-element bonding of 8 common elements ranges from 50% ionic (Si-O) to 80% ionic (K-O)

32. Eq. 3.4 Fig. 3-10 % ionic character = 1 – e -0.25(X a – X c ) 2 Note negative sign, typo in 1 st edition O-Si ~50 % ionic O-K ~80 % ionic X = electronegativity of a, anion and c, cation

33. Native elements Examples: S, Fe, Au… Examples: S, Fe, Au… No differences in electronegativity No differences in electronegativity Bonding intermediate between covalent and metallic Bonding intermediate between covalent and metallic Low electronegativity values (Cu, Ag, Au) favor metallic bonding Low electronegativity values (Cu, Ag, Au) favor metallic bonding High electronegativity values (non-metals, C, S) favor covalent bonding High electronegativity values (non-metals, C, S) favor covalent bonding

34. Fig. 3-9 Continuous variations Limited variations 100% covalent, metallic or ionic 50 % covalent & 50% metallic Part covalent, part metallic, and part ionic Range of possible mixtures of electron-sharing valence bond types Percentages Not Allowed

35. Physical Properties caused by Valence bonds Electrical conductance Electrical conductance Ionic and covalent have little conductance Ionic and covalent have little conductance Metallic highly conductive Metallic highly conductive Solubility Solubility Ionic highly soluble (think halite) Ionic highly soluble (think halite) Brittleness Brittleness Ionic highly brittle – cleavage common Ionic highly brittle – cleavage common Halite – perfect {001} cubic cleavage Halite – perfect {001} cubic cleavage

36. Hardness Hardness Covalent – strongest bonding, so hardest. Think diamond Covalent – strongest bonding, so hardest. Think diamond Malleable Malleable Metallic easily worked Metallic easily worked

37. Non-valence bonds Result of asymmetric charge distribution Result of asymmetric charge distribution Create electrostatic forces Create electrostatic forces Two types Two types Van der Waals and Hydrogen Van der Waals and Hydrogen

38. Hydrogen Bonding Ice example Ice example H 2 O is polar molecule H 2 O is polar molecule 2 H atoms at angle to O atom (not straight line) 2 H atoms at angle to O atom (not straight line) O is more electronegative than H O is more electronegative than H O = 3.5, H = 2.1 O = 3.5, H = 2.1 O “claims” more of the electron O “claims” more of the electron Net negative charge on O side of molecule Net negative charge on O side of molecule The asymmetric charges allow solidifying liquid when T < 0º C @ 1 atm P The asymmetric charges allow solidifying liquid when T < 0º C @ 1 atm P

39. Fig. 3-11 & 18-2 Asymmetrical charge - polar Hydrogen bond Hexagonal symmetry Ice – viewed down c axis

40. Van der Walls Carbon example Carbon example Graphite – carbon bonded in sheets Graphite – carbon bonded in sheets Bonding within sheets is covalent –  bonds Bonding within sheets is covalent –  bonds Over time electrons evenly distributed Over time electrons evenly distributed At given time, excess electrons on one side of sheet At given time, excess electrons on one side of sheet Creates weak electrostatic attraction Creates weak electrostatic attraction Physical properties Physical properties Typically soft Typically soft Graphite good lubricant Graphite good lubricant

41. Fig. 3-12 Covalent bonds within the sheets Van der Waal forces between the sheets, Caused by  bonds on top of sheets Other examples: talc serpentine/smectite

42. Atoms and ion size Assume that atoms are spheres Assume that atoms are spheres Clear simplification – electron distributions are not spherical Clear simplification – electron distributions are not spherical Assumption works well for arrangement in solids Assumption works well for arrangement in solids Atoms pack together in regular arrangement Atoms pack together in regular arrangement

43. If we assume the ions are spheres If we assume the ions are spheres Can assume an effective radius Can assume an effective radius Measure distance between adjacent atoms in the solid Measure distance between adjacent atoms in the solid Measured with X-ray diffraction, d spacing Measured with X-ray diffraction, d spacing Effective radius a measure of size of the atoms Effective radius a measure of size of the atoms

44. Very important – one control of how atoms pack together Very important – one control of how atoms pack together Bond length – sum of effective radius of two adjacent atoms Bond length – sum of effective radius of two adjacent atoms Metallic bonds: all same effective radius Metallic bonds: all same effective radius ½ distance between nuclei ½ distance between nuclei Ionic bonds: effective radius different between two atoms Ionic bonds: effective radius different between two atoms Not ½ distance between nuclei Not ½ distance between nuclei

45. Fig. 3-13 Metallic bonding Bond length = d spacing Ionic radius = ½*d spacing Covalent and ionic bonding Bond length = d spacing d spacing = R a + R c

46. Clearly – what types of ions present control ionic radius Clearly – what types of ions present control ionic radius Primary variables controlling ionic radius: Primary variables controlling ionic radius: Oxidation state – i.e. charge on ion Oxidation state – i.e. charge on ion Coordination number – i.e. number of ions surrounding central ions Coordination number – i.e. number of ions surrounding central ions

47. Oxidation state Inversely related Inversely related More oxidized (less negative, more positive) means smaller effective radius More oxidized (less negative, more positive) means smaller effective radius e.g., Fe 3+ or Fe 2+ e.g., Fe 3+ or Fe 2+ Cations smaller than anions, O 2- very large Cations smaller than anions, O 2- very large Positive charge holds electron closer to nucleus Positive charge holds electron closer to nucleus

48. Fig. 3-15 Charge (higher oxidation state) Ionic radius (Å) Coordination numbers

49. Coordination Positive correlation – high coordination number, smaller ions Positive correlation – high coordination number, smaller ions Think of solids as large anions surrounding small spaces filled by cations Think of solids as large anions surrounding small spaces filled by cations Size of space determined by effective radius of anions Size of space determined by effective radius of anions Cation effective radius changes to fill space Cation effective radius changes to fill space

50. Fig. 3-16 Coordination number Ionic radius Note – increase in ionic radius independent of charge on ion