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Basic Atomic Theory the Periodic Table

Basic Atomic Theory the Periodic Table. Order in the Court!. Atomic Theory. Atomic Theory. Atoms: smallest particles of a substance with the same chemical/physical properties as that substance. Protons —> stable, (+) charge. Electrons—> mobile, (-) charge. Neutrons—> stable, no charge.

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Basic Atomic Theory the Periodic Table

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  1. Basic Atomic Theory the Periodic Table Order in the Court!

  2. Atomic Theory

  3. Atomic Theory • Atoms: smallest particles of a substance with the same chemical/physical properties as that substance. • Protons —> stable, (+) charge. • Electrons—> mobile, (-) charge. • Neutrons—> stable, no charge. • Nucleus of atom contains neutrons and protons. • Electrons orbit the nucleus. • In an electrically neutral atom, the number of electrons = the number of protons.

  4. Atomic Theory • Atomic number – • The number of protons in the nucleus determines the atomic number • Change the number of protons and you change the element (also atomic number, atomic mass) • Neutral atoms—> protons = electrons

  5. Atomic Theory • Ions—> # protons not equal to the # electrons • Cation —> protons > electrons; (+) Charge; Ca+2 • Anion —> protons < electrons; (-) Charge; S-2 • Complex ion —> cation or anion group (more than 1 element); SiO4-4, CO3-2 • See anionic complexes, (more later) • Valence —> ionic charge (preferred ion configuration) • # of protons (minus) # of electrons

  6. Atomic Theory • Isotope – • A form of an element with a different number of neutrons • Some are stable (don’t spontaneously degenerate) • Some are unstable or radioactive (spontaneously degenerate by nuclear reactions to form different elements)

  7. The Periodic Table, Valence, Atomic/Ionic Radius, Bonding, and Coordination

  8. The Periodic Table of Elements(Order in the Court!) • Elemental symbols (1-2 letters); Ca, Si, etc • Atomic # • Number of protons (or electrons) in an uncharged atom • Horizontal rows are periods • Vertical columns are groups

  9. The Periodic Table of Elements(Order in the Court!) • Chemical variation in periods (rows) • Increasing # of electrons in outer electron shell to capacity (8) • Stable (chemically non-reactive) noble gases electron configuration (NGEC, rule of eight or octet rule) • Electrons fill lowest energy (inner most) electron shells first (except for transition metals; A# 21-31) • Outer most electrons are mobile and result in chemical reactivity; the valence electrons

  10. The Periodic Table of Elements(Order in the Court!) • Chemical properties of groups (columns) • Similar physical and chemical properties due to similar outer electron shell configuration • Similar tendency to gain (group 5A-7A, and become (-) anions) or loose (group 1A-4A, and become (+) cations) electrons to form ions with NGEC

  11. The Periodic Table of Elements(Order in the Court!) • Chemical properties of groups • Group 1A—> alkali metals (electropositive) • Valence of +1 • Group 2A—> alkaline earth metals (electropositive) • Valence of +2 • Groups 1B-7B and 8—> transition metals; • More complex chemical properties • Several valences, i.e. Fe+2 (ferrous iron) and Fe+3 (ferric iron) • Group 3A—> metalloids (electropositive) • Valence of +3 • Group 4A—> silica-carbon group (important mineral and organic compounds; electropositive) • Valence of +4

  12. The Periodic Table of Elements(Order in the Court!) • Chemical properties of groups (cont) • Group 5A—> nitrogen-phosphorous group (electropositive) • Valence of +5 • Group 6A—> oxygen group (non-metals, electronegative) • Valence of -2 • Group 7A—> halogen group (non-metals, electronegative) • Valence of –1 • Group 8ANoble (inert, non-reactive ) Gases

  13. Valence • Characteristic loss or gain of electrons from a neutral atom with an electric charge on resulting ion Dark shade are the most common elements in the crust. Make them your friends! Lighter shade are other important elements in earth materials.

  14. Oxidation and Reduction • Oxidation: • Process of loosing an electron • Fe+2 Fe +3 • Reduction: • Process of gaining an electron(s) • O2 2O-2

  15. Bonding in Minerals • Most atoms are unstable because they have unfilled outer electron shells (rule of eight or octet rule) • Exception: Noble (inert) gases • Most elements ionize; gain or loose electrons and become charged

  16. Bonding • Electrostatic force of attraction that holds cations, anions, and/or complex ions together in chemical compounds and especially, in the rigid geometric structures of minerals • Opposite electric charges attract • Like electric charges repel

  17. Bonding • Electrically neutral chemical compounds (such as minerals) are stoichiometric (equal positive and negative chargeuncharged) • Stoichiometry determines the relative proportions of elements in stable mineral compounds determined by valence of the elements involved, e.g. Fe(+3)2O(-2)3 Iron oxide (rust) hematite

  18. Bonding • Ions combine to attain stoichiometry (electrical neutrality) and noble gas (full outer electron shell) electron configuration • Particular mineral species form in order to minimize internal (molecular scale chemical) energy in accordance with external conditions of: • oT (temperature), • P (pressure), and • Chemical environment (composition of available constituents or raw materials)

  19. Bond Types • Ionic bond • Covalent bond • Metallic bond • Hydrogen bond • Van der Waal’s bond

  20. Bond Types • Bond types are not mutually exclusive within any mineral compound

  21. Bond Types • Ionic bond • Fundamental (dominant) bond type in 90% of all minerals • Consideration of ionic bonding explains most mineral properties • Electron exchange between cations and anions; ions are surrounded by oppositely charged ions to satisfy the octet rule

  22. Bond Types • Ionic bond • Structure of ionic compounds determined by ionic radius and valence of constituents • Size and packing of constituent elements results in regular rigid structure • Electrical neutrality requires balanced ionic charges; Stoichiometry

  23. Bond Types • Covalent bond • Only a few minerals • Results from shared electrons to satisfy octet rule • See diamond

  24. Bond Types • Metallic bonding • A special case of covalent bonding) • Outer electrons free to move and are shared over a wider range in a crystal lattice (long range covalent bonding) • Common in the native elements especially metallic minerals • High electrical conductivity • High thermal conductivity • Lustrous • Ductile

  25. Bond Types • Hydrogen bond • Weak bonds due to asymmetry within a crystal structure (or molecule) i.e. Hydrogen bonding of polar water molecules • Van der Waal’s bond • Weak forces due to electron mobility and temporary polarization of charge • Determines cleavage directions in soft minerals

  26. The Origin of Earth Materials, Bonding, the Coordination Principle, and an Introduction to the Unit Cell Relationship between Minerals and Chemistry: Crystal Chemistry

  27. Origin of Earth Elements • In the universe • Original H & He: the big bang • Nuclear fusion to form heavier elements in stars

  28. Origin of Earth Elements • In our solar system • Remnant heavier elements concentrated in the terrestrial (Mercury, Mars, Earth, and Venus) planets through • Sequential planetesimal amalgamation • Fe-Ni rich core first • Si rich mantle and crust • Volatile elements (easily vaporized) last, to form the atmosphere

  29. Origin of Earth Elements • During Earth formation • Early molten period • Density differentiation, cooling, and partial solidification • Goldschmidt’s classification • Predictability of elemental distribution • Siderophile (elements associated with iron) • Chalcophile (elements associated with sulfur) • Lithophile (elements associated with silica) • Atmophile (elements that form a gas)

  30. Abundance of Elements in the Crust and Mantle; Lithophile Elements • O, Si, Al, Fe, Ca, Na, K, Mg • Mantle: • Si, O, Fe, and Mg • Crust: • Si, Al, Ca, Na, K • Segregation and concentration of elements through various Earth processes • Partial melting- magma formation • Surface weathering

  31. Cations—> generally smaller AR Anions—> generally larger AR Variable dependant on atomic number and interaction with other ions Ionic Radius: • Atomic radii (AR)~ 1 D (angstrom, 10-10 meters)

  32. The Coordination PrincipleGeometry of Atomic Building Blocks • In an ionically bonded substance (all minerals for our purposes) cations are surrounded by anions (or anionic complexes) • In stable mineral crystals • The number and arrangement of anions surrounding a cation forms a Coordination polyhedron

  33. The Coordination PrincipleGeometry of Atomic Building Blocks • Coordination polyhedron • The size and shape of the coordination polyhedron is determined by the relative size (ionic radius) of the cation and anion (anion complex) involved • Radius ratio (cation radius/anion radius) • This shape is described by the number of anions surrounding a cation called the Coordination number (C.N.)

  34. Coordination Polyhedron, Radius Ratio, and Coordination Number

  35. The Coordination Principle • Oxygen (O-2) is the most common anion in coordination polyhedron • Silica tetrahedra (SiO4)-4 with coordination number of 4, radius ratio of 0.30. • It is very important

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