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Chemistry Lecture Notes Ionic and Molecular Compounds

Chemistry Lecture Notes Ionic and Molecular Compounds. New evidence for inorganic origin of oil

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Chemistry Lecture Notes Ionic and Molecular Compounds

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  1. Chemistry Lecture Notes Ionic and Molecular Compounds

  2. New evidence for inorganic origin of oil The origin of petroleum has long been an issue of debate. Although many people believe crude oil formed from plant and animal material in near-surface sedimentary rocks at high temperatures and pressures, a abiogenic origin has also been proposed. Proposed in the 1950s, the theory attributes oil formation to inorganic carbonate rocks at high temperatures and pressures found only at great depths. In the latest development in this debate, researchers at Gas Resources, Corp., Houston, and the Joint Institute of Earth Physics, Moscow, have predicted the thermodynamic conditions under which the hydrocarbons found in crude oil form and tested those conditions in the lab. They reacted iron oxide, marble (CaCO3), and water at condition reaching 1,500ºC and 50,000 atm. Hydrocarbons ranging from methane to decane were formed in proportions that mirror naturally occurring petroleum. With the exception of methane, hydrocarbons did not form at pressures below 30,000 atm, which corresponds to about 100 km below the Earth’s surface. Taken together, the theoretical and experimental results make the biogenic theory untenable, the researchers conclude.

  3. Lewis (electron-dot) symbols: show the valence electrons of atoms or ions octet rule: atoms tend to gain, lose, or share electrons until they reach the nearest noble gas configuration

  4. N a + C l N a + C l Ionic compounds _ + +  Na Cl Na+Cl Ionic “bond”: electrostatic attraction between positive ions (cations) and negative ions (anions). metal + nonmetal ionic compound (e.g., NaCl, MgF2, K2S, CaO, AlBr3) electrolytes! low IE low EA gives up e– high IE high EA takes e– transfer of electrons

  5. Metals tend to lose electrons to form cations: alkali metals: +1 alkaline earths: +2 group 3A: +3 transition metals: variable Nonmetals tend to gain electrons to form anions: halogens: -1 chalcogens: -2 group 5A -3 P3–

  6. We can usually predict the formulas of ionic compounds  empirical formula (formula unit) only Rules: 1. cation written first 2. sum of all charges = 0 3. empirical formula has smallest set of whole numbers Li, Br  Na, O  Mg, Cl  K, P  Al, F  Ca, O  Be, N  but Ca, C 

  7. Molecular compounds: covalent bonds +  C O2 CO2 Two or more atoms bound tightly together. nonmetal + nonmetal molecular compound (covalent bonds) e.g., N2O, CO2, H2S, PCl3, SO3 Can’t always predict formula: N2O, NO, NO2, N2O5, etc. nonelectrolytes!

  8. Covalent bond: attraction of two nuclei to one or more shared pairs of • electrons • nonmetal + nonmetal  covalent (molecular) compound equal attraction for electrons covalent bond; shared electrons each atom “sees” two electrons (He) each atom has an octet lone pair electrons bonding electrons duet octet

  9. Multiple bonds double bond triple bond octets

  10. Electronegativity and bond polarity • Electronegativity: ability of an atom in a compound to attract electrons to itself • ~ follows IE, EA • arbitrary scale (Pauling) most C increases least

  11. H—H C = 2.1 2.1 DC = 0  equal sharing of electrons Cl—Cl = nonpolar covalent bond C = 3.0 3.0 H—Cl DC = 0.9  unequal sharing of electrons C = 2.1 3.0 = polar covalent bond Na+Cl–DC = 2.1  transfer of electrons C = 0.9 3.0 = ionic bond generally: when DC < 1.9 covalent > 1.9 ionic d+d– (approximate!)

  12. Empirical, molecular, and structural formulas Empirical formula: smallest whole number ratio of atoms in a compound e.g. CH2O Molecular formula: actual number of atoms in a molecule of the compound CH2O formaldehyde C2H4O2 acetic acid all have same empirical formula C6H12O6 glucose

  13. Structural formula: shows how atoms are bonded together CH3CO2H HOCH2CHO Condensed formulas

  14. Drawing Lewis structures • 1. Sum all valence electrons. • add one electron for each negative charge • subtract one electron for each positive charge • 2. Draw single bonds from the central atom to the outer atoms. • 3. Place pairs of electrons on the outer atoms until they have octets. • exception: H (duet) • 4. Place any remaining electrons on the central atom. • 5. If the central atom has less than an octet, form multiple bonds with pairs from the outer atoms.

  15. NI3 SO42– CH2O PCl5 SeF6 C2H4 SO2 NO+ C2H5OH NO3– HCN CH3CO2H

  16. Formal charge Divide the electrons in each bond equally between the two atoms, then compare each atom to its normal valence. 7 e– 1 e– 4 e– (redo previous page examples)

  17. When more than one Lewis structure is possible, the better (more stable) • one will have: • 1. fewer (or smaller) formal charges. • 2. the negative charge on the more electronegative atom, and vice versa.

  18. Exceptions to the octet rule 1. odd number of electrons 2. less than an octet: Groups 1, 2, and 3 3. more than an octet: 3rd row and below can “expand” their octets better very bad

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