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Atoms • Element = substance that can not be chemically broken down into smaller parts; has unique chemical properties • Atom = smallest subdivision of an element that still retains all of the element’s chemical properties • Protons = positively charged particles; clustered in the nucleus of the atom • Neutrons = particles with a neutral charge; clustered in the nucleus of the atom • Electrons = negatively charged particles; located in a cloud around the nucleus • Atomic number = Z = number of protons in the atom; each element has a unique number • Isotope = same number of protons, but a different number of neutrons • Atomic weight = neutrons + protons + electrons Crystal Chemistry Chapter 3 E. Goeke, Fall 2006 • For a neutral charge, the # of electrons = # of protons • Ion = atom with either more or less electrons than protons – Anion = extra e; - charge – Cation = missing e; + charge • The classic method of imaging the relationship between electrons & the nucleus is the Bohr Model – Each shell has a specific # of electrons – Electrons in outer shells have more energy; the excess energy is released as photons when the electrons fall into inner shells Electrons http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm E. Goeke, Fall 2006 Quantum Mechanics • Electrons not in a specific location, but in a zone of probability • The electron location is described using 4 factors: – n = principal quantum number (similar to the shell in the Bohr model) – l = angular momentum quantum number = designates which kind of subshell shape; btw 0 and n-1; 1 = s, 2 = p, 3 = d, 4 = f, and so on – ml = magnetic quantum number = distinguishes between orbitals of the same l value with different orientations; btw -l and +l – ms = spin quantum number = distinguishes between electrons spinning to the right and those to the left; +1/2 or -1/2 • Pauli exclusion principle = no two electrons can have the same four quantum numbers; each suborbital can have 2 electrons, but they must spin in opposite directions E. Goeke, Fall 2006 s orbital; increasing width & energy with increasing n; maximum of 2 electrons per n p orbitals; three orientations; only present when n = 2+; max of 6e per n d orbitals; five orientations; only present when n = 3+; max of 10e per n http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm Mineralogy 12:041, Fall 2006, E. Goeke E. Goeke, Fall 2006 E. Goeke, Fall 2006 Orbital Energy Levels • Lower energy shells & subshells are filled first • There is a large energy gap between the 1s and 2s orbitals, but a smaller energy difference between 6s and 7s • Note that 4s has a lower energy level than 3d--it will fill first • Order shells & subshells fill: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm E. Goeke, Fall 2006 1 s Periodic Table p Ions K d L M N O P Q f http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm Noble gases • Atoms tend to lose or gain electrons to fill their outer electron shells – Valence state = the charge on an ion; elements can have more than 1 common valence state (e.g. Fe 2+ or Fe 3+ ) – Elements in the same column will tend to have similar valence states – Elements that lose electrons have a positive charge and can be called “metals” – Elements that gain electrons have a negative charge and are referred to as “non-metals” • The oxidization state of the environment plays a large role in determining the valence state of a given ion; the more oxygen present = more electrons removed from metals E. Goeke, Fall 2006 E. Goeke, Fall 2006 Electronegativity Bonding • Electronegativity = ability of an atom within a crystal structure to attract electrons into its outer shell – Low EN = electron donors – High EN = electron acceptors – Noble gases have an EN = 0 http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm • There are four kinds of bonds that we will consider in mineralogy: – Ionic – Covalent – Metallic – Van der Waals and hydrogen bonding • The type of bond will make a significant contribution to the physical properties of a given mineral • Though presented as four separate types, most bonds actually transitional between two end members • EN plays a role in determining what kind of bond will form • Charge balance must be maintained E. Goeke, Fall 2006 E. Goeke, Fall 2006 Ionic Bonds Covalent Bonds • Formed from two oppositely charged ions that have either given up or gained an electron to fill their outer shells • Non-directional • Dissolve easily in polar solvents (e.g. water) • Commonly form high symmetry xtals that have moderate hardness and density as well as high melting temperatures • Usually poor conductors of heat & electricity http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm Mineralogy 12:041, Fall 2006, E. Goeke E. Goeke, Fall 2006 • • • • • • Two atoms share 1+ electrons in order to be more stable Very strong directional bonds Relatively insoluble in polar solvents High melting temperatures & high hardness Generally form low symmetry crystals Usually poor conductors of heat & energy http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm E. Goeke, Fall 2006 2 Covalent or Ionic? • Elements on the right hand side of the periodic table tend to be covalent, while elements on the left hand side are more likely to be ionic • Most bonds are a mix between covalent or ionic determined by the differences in electronegativity between the two atoms • Bonding between two atoms of the same element will produce a pure covalent bond • Ionic bonded crystals tend to be closely packed anions & cations http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm E. Goeke, Fall 2006 Metallic Bonds • Hydrogen bonds = because O has high electronegativity, it has a greater pull on the shared e than the H, which causes an electrical polarity of H2O molecules – OH- layers result in perfect cleavage along 1 plane • Van der Waals bonds = polarized atoms or molecules result in concentrations of positive & negative charges that causes a slight attraction – Weak bonds that are easy to break http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm E. Goeke, Fall 2006 Atomic & Ionic Radii • Size of an atom/ion will depend on the size of the nucleus plus the number of electrons – Those with greater numbers of electrons ~ have larger radii – Ions have different radii than atoms of the same element – As the charge becomes more positive -> fewer electrons -> smaller radii – Larger atomic number -> larger radii – Radii also depend on type of bonds and how many other ions are linked to the given ion • Ions with similar radii are more likely to substitute for each other • Coordination number = CN = # of anions packed around a cation E. Goeke, Fall 2006 Mineralogy 12:041, Fall 2006, E. Goeke E. Goeke, Fall 2006 http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm Van der Waals & Hydrogen bonds • Positively charged atomic nuclei share electrons amongst themselves • Conduct electricity & heat easily • Low to moderate hardness • Normally malleable & ductile • Soluble only in acids • Crystals usually have high symmetry http://www.tulane.edu/~sanelson/eens211/crystal_chemistry.htm • Some elements produce hybrid s and p orbitals into sp3 orbitals during covalent bonds – Though Si-O bonds are about 50% covalent, the bonds produce hybrid orbitals – The orientation of the hybrid orbitals causes the typical silica tetrahedron arrangement E. Goeke, Fall 2006 Ion R(Å) C.N. = 6 R(Å) C.N. = 8 La+ 0.74 0.92 Na+ 1.02 1.18 K+ 1.38 1.51 Rb+ 1.52 1.61 Cs + 1.67 1.74 Fe2+ 0.75 1.06 Fe3+ 0.69 0.92 Nb4+ 0.82 0.93 Nb5+ 0.78 0.88 U6+ 0.87 1.00 E. Goeke, Fall 2006 3