2.1 Crystal chemistry

2.1 Crystal chemistry - Crystal Chemistry Mineral...

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Unformatted text preview: Crystal Chemistry Mineral "...defined, but generally not fixed, composition..." Modern geology Geochemistry or Geophysics Geophysics application of physical principles to study of earth Geochemistry application of chemical principles to study of earth high T or low T Coming up... a review of basic chemistry Elements Bond types and controls on bonds protons, neutrons, electrons Nuclear Chemistry Atomic number (Z) number of protons Neutrons about weight of proton, different number of neutrons make isotopes Specific for particular elements (periodic table) Atomic weight sum of weight of neutrons and protons Isotopes Superscript in front of element symbol, atomic weight exact Elements atomic weight is the average of abundance of isotopes Stable Isotopes Oxygen Z = 8; three isotopes 16O 99.757% 17O 0.038% 18O 0.205% Materials (minerals, water, air, shells, etc) have variable ratios of these isotopes 18O = ratio of 18O/16Osample to 18O/16Ostandard Radioactive isotope Potassium (z=19) 40K has 21 neutrons 39K has 20 neutrons Natural abundance = 0.0117% Radioactive Decays to 40Ar, basis of one type of age dating Half life = 1.248 x 109 a Natural abundance = 93.3% Stable (not radioactive) 41K has 22 neutrons Natural abundance = 6.7% Chemical Reactions Based on electron transfers, charge balance If number of electrons = number of protons, no electrical charge Orbit nucleus in systematic way Organized according to energy levels Shells filled according to energy Electron Quantum number Quantum number e.g., n, l, ml, and ms Controls energy of electron Unique for each electron No two electrons in atom can have same quantum number Controls how electrons fill shells Controls their chemical reactivity Formation of ions Ions excess or deficit of electrons relative to protons Valence or Oxidation state is the value of the charge on the ion Anions net negative charge Cations net positive charge Configuration of valence electron controls whether gain or lose electron Metals typically lose one or two valence electron: form cations Nonmetals typically require a few electrons to fill valence shells: form anions Valence shells fill systematically Atomic number 120 and 3138 fill s & p subshells Between atomic number 20 and 31 shells fill from internal shells fill 3d shell (4s shell filled) Elements may have differing numbers of shells filled Transition metals E.g. Ferrous and Ferric iron Lose electrons (cations) to become noble gas core Note the various oxidation states for the transition metals Ferric Fe (+3) Ferrous Fe (+2) Gain electrons (anions) Metallic Fe Noble Gases, He, Ne, Ar, Kr Fig. 33 Clearly gain or loss of electrons important "Quantified" as property called Electronegativity Electronegativity Defined by Linus Pauling Propensity of element to gain or lose electron See Table 34 for values Based on arbitrary scale: Li = 1, C = 2.5, F = 4 Low electronegativity the more likely to lose electron form cations High electronegativity likely to gain electron to form anions 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 Earth abundances of elements What elements are most abundant? What part of earth do they occur? These elements will make up common minerals Crust? Bulk earth? Crust 8 common elements Determination of crustal abundance simply collect large number of samples and measure Bulk earth composition O, Si, Al, Fe, Ca, Na, K and Mg Most minerals are made of these elements The same 8 elements are common in bulk, but different ratios Bulk Earth composition: Estimated by Difficult to assess impossible to directly sample mantle or core Mass and density based on geophysical measurements Composition of mantle magmas and xenoliths Composition of meteorites Table 3.6 Crust Core/Mantle Mantle Chemical bonding Eight common elements (plus all others) bond to form minerals Two categories These are "end member" types Sharing of valence electrons: ionic, covalent and metallic No sharing: van der Waals and hydrogen Rarely just one type or the other Ionic Bonding Transfer of electron(s) from one element to another The distance between ions depends on attractive forces (Coulomb law) and repulsive forces (Born repulsion) Results in filled valence shells of both The electrostatic attraction keep atoms together Ionic bonding Ions bond so that positive = negative charges Characteristics: Minerals must be electrically neutral NaCl (Halite) and CaF2 (fluorite) Ions act like spheres Alternating cations and anions One of the strongest bonds Brittle because like ions repel Cleavage is common Attractive forces Bonding in Halite Equilibrium distance = 2.8 Repulsive forces Face centered cubic lattice arrangement of halite Fig. 34 Fig. 210 Covalent Bonds Electrons shared when orbitals of two different elements overlap Shared only by two atoms, differs from metallic Electrons move around nucleus of both atoms Examples Diamond and Graphite Diamond Stable Ne configuration by either gain or loss of 4 electrons Ionic bonding not possible because all electrons exactly the same electronegativity One carbon won't "steal" electron from another Instead share electrons very strong bonding Covalent bonding in diamond 4 orbitals shown as bonds, call bonds bonds distorted Each bold line represents another similar bond Fig. 35 Graphite Additional Sharing electrons, bonds. Essentially metallic bonds Similar bonds, but only in layers Fig. 36 Metallic bonds A type of covalent bond Electrons shared without systematic change in orbitals Formed with low electronegativity weakly held valence electrons Free to move throughout crystal structure Relation between valence dependent bonds Most bonds not purely ionic, covalent or metallic Amount of bond type depends on electronegativity Greater difference in electronegativity between ions means more ionic characteristic Only 1 anion (of 8 comment elements) Oxygen Electronegativity of O = 3.5 Electronegativity of other common elements range from 0.8 (K) to 1.8 (Si) Difference in electronegativity Possible to quantify % ionic bonding: OK = 3.5 0.8 = 2.7, more ionic characteristics OSi = 3.5 1.8 = 1.7, less ionic characteristics Oelement bonding of 8 common elements ranges from 50% ionic (SiO) to 80% ionic (K O) OK ~80 % ionic OSi ~50 % ionic % ionic character = 1 e 2 0.25(Xa Xc) Note negative sign, typo in book X = electronegativity of a, anion and c, cation Fig. 310 Native elements Examples: S, Fe, Au... No differences in electronegativity Bonding intermediate between covalent and metallic Low electronegativity values (Cu, Ag, Au) favor metallic bonding High electronegativity values (nonmetals, C, S) favor covalent bonding Range of possible mixtures of bond types Lim tio ns d ite ari a va tio ria nti nu ou s v 100% covalent, metallic or ionic 50 % covalent & 50% metallic Part covalent, part metallic, and part ionic ns Co Continuous variations Fig. 39 Physical Properties caused by Valence bonds Electrical conductance Solubility Ionic and covalent have little conductance Metallic highly conductive Ionic highly soluble (think halite) Ionic highly brittle cleavage common Brittleness Halite perfect {001} cubic cleavage Hardness Malleable Covalent tightest bonding, so hardest. Think diamond Metallic easily worked Nonvalence bonds Result of asymmetric charge distribution Create electrostatic forces Two types Van der Waals and Hydrogen Hydrogen Bonding Ice example H2O is polar molecule The asymmetric charges allow solidifying liquid when T < 0 C @ 1 atm P O is more electronegative than H O = 3.5, H = 2.1 O "claims" more of the electron Net negative charge on O side of molecule Asymmetrical charge polar Ice viewed down c axis Hydrogen bond Hexagonal symmetry Fig. 311 & 182 Van der Walls Carbon example Graphite carbon bonded in sheets Physical properties Bonding within sheets is covalent bonds Over time electrons evenly distributed At given time, excess electrons on one side of sheet Creates weak electrostatic attraction Typically soft Graphite good lubricant Covalent bonds within the sheets Van der Waal forces between the sheets, Caused by bonds on top of sheets Other examples: talc serpentine/smectite Fig. 312 Atoms and ion size Assume that atoms are spheres Atoms pack together in regular arrangement Clear simplification electron distributions are not spherical Assumption works well for arrangement in solids If we assume the ions are spheres Effective radius a measure of size of the atoms Can assume an effective radius Measure distance between adjacent atoms in the solid Measured with Xray diffraction, d spacing Very important one control of how atoms pack together Bond length sum of effective radius of two adjacent atoms Metallic bonds: all same effective radius Ionic bonds: effective radius different between two atoms distance between nuclei Not distance between nuclei Metallic bonding Covalent and ionic bonding Fig. 313 Primary variables to control ionic radius: Oxidation state i.e. charge on ion Coordination number i.e. number of ions surrounding central ions Oxidation state Inversely related Lower oxidation state means larger effective radius Cations smaller than anions Positive charge holds electron closer to nucleus Decreasing radius Increasing charge (higher oxidation state) Fig. 315 Coordination Positive correlation Think of solids as large anions surrounding small spaces filled by cations Size of space determined by effective radius of anions Cation effective radius changes to fill space Increasing radius Increasing coordination number Fig. 316 ...
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This note was uploaded on 07/06/2011 for the course GLY 5245 taught by Professor Staff during the Spring '11 term at University of Florida.

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