Basic Atomic Theory: The Periodic Table

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

2. Basic Chemistry Review • With your neighbor, define the following terms: • Element • Atom (define, also sketch and name internal “parts”) • Atomic number, mass • Isotope • Ion (types) • Ionic radius • Valence • Bonding (include types, if you know them)

3. Atomic Theory

4. Atomic Theory • Element: basic, indivisible substance of matter • Atom: smallest particle of an element with the same chemical/physical properties as that element • Protons —> stable, (+) charge • Neutrons—> stable, no charge • Electrons—> mobile, (-) 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

5. Atomic Theory • Atomic number • = Number of protons in the nucleus • Change the number of protons and you change the element • Atomic mass • = Number of protons + number of neutrons • Varies with isotopes of the same element (more later) • Atomic weight (gram atomic weight) • = Weight of particles (minus some energy loss) • Standardized to 12C = 12.0000 • A mole of an element always has 6.023 x 1023 atoms

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

7. Atomic Theory Bohr models: • Bohr model of an atom: • Electrons orbit the nucleus in discrete energy levels (1-7) • Each level contains sublevels that can hold different number of electrons (s, p, d, f, g) • Electrons fill orbitals from lowest to highest energy • Valence electrons in outermost shell affect chemical properties • Modern models (Schrödinger) are more complex!

8. 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 • In minerals, anion complexes (negative charge) are important! • Valence —> ionic charge (preferred ion configuration) • # of protons (minus) # of electrons

9. Atomic Theory • Why make an ion? • To achieve stability (filled outer shell) • By gaining (- charge) or losing (+ charge) electrons • How do you make an ion? • Oxidation: • Lose 1+ electrons • Fe+2 Fe +3 • Reduction: • Gain 1+ electrons • O2 2O-2

10. The Periodic Table: Order in the Court! • Elemental symbols (1-2 letters) • Horizontal rows are periods • Vertical columns are groups

11. The Periodic Table • Chemical variation in periods (rows) • Ordered by increasing atomic weight • Row # (period) indicates shells occupied by electrons • Across the row, # of electrons in outer electron shell increases to capacity (8, rule of eight or octet rule) • Stable (chemically non-reactive) noble gases electron configuration is reached when outer shell filled • If outer shell is partially filled (<8), elements give up, receive, or share electrons from other elements to achieve stability • Significance: outermost electrons are mobile and result in chemical reactivity; the valence electrons

12. The Periodic Table • Chemical properties of groups (columns) • Similar physical and chemical properties due to similar valence configuration • Similar tendency to gain electrons and become anions (group 5A-7A) • Similar tendency to lose electrons and become cations (group 1A-4A)

13. The Periodic Table • 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 (filling of inner shells) • 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

14. The Periodic Table • 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

15. Bonding • 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 • 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

16. Bonding • Electrically neutral chemical compounds (such as minerals) are stoichiometric • Have equal positive and negative charge  balanced (no net charge) • Stoichiometry determines the relative proportions of elements in stable mineral compounds determined by valence of the elements involved Fe(+3)2O(-2)3 Iron oxide (rust) hematite

17. Bonding • “Goals” of bonding • Attain stoichiometry (electrical neutrality) • Attain full outer electron shell • Particular mineral species form in order to minimize internal (molecular scale chemical) energy in accordance with external conditions of: • T (temperature) • P (pressure) • Chemical environment (available elements)

18. Bond Types • Ionic • Covalent • Metallic • Hydrogen • Van der Waal’s • Bond types are not mutually exclusive within any mineral compound

19. Bond Types – Ionic: • Dominant bond type in 90% of all minerals • Consideration of ionic bonding explains most mineral properties • Electron exchange between cations and anions • Each ion is surrounded by oppositely charged ions to satisfy the octet rule • Structure of ionic compounds determined by ionic radius and valence of constituents • Size of elements and how they are packed together results in crystal structure • Stoichiometry required balanced charges

20. Bond Types - Covalent • Only a few minerals • Results from shared electrons to satisfy octet rule • Example: diamond

21. Bond Types - Metallic • 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

22. 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

23. Origin of Earth Elements • ~3000 named minerals; only about a dozen are common in the earth’s crust – why? • In the universe • Original H & He: the big bang • Nuclear fusion to form heavier elements in stars • 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 last, to form the atmosphere

24. 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)

25. Abundant 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

26. Abundance Red are the most common elements in the crust. Make them your friends! Green are other important elements in earth materials.

27. Collaborative activity In groups, answer the following questions. Show your work and turn in. • Calculate: • The mass (in g) of an Al atom • The mass (in g) of a molecule of NaCl • The number of K atoms in a sample weighing 5 x 10-8 g 2. Predict the simplest chemical formula for: a. Calcium fluoride b. Sodium phosphate (hint: “phosphate” is an anion complex)