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Ch. 6 - Chemical Bonds I. Why Atoms Combine (p.298-302) Chemical Formula Chemical Bond Stability A. Chemical Formula • Shows: 1) elements in the compound 2) ratio of their atoms H2O 1 oxygen atom 2 hydrogen atoms B. Chemical Bond • Strong attractive force between atoms or ions in a molecule or compound. • Formed by: – transferring e- (losing or gaining) – sharing e- C. Stability • Octet Rule – most atoms form bonds in order to have 8 valence e– full outer energy level – like the Noble Gases! Ne Stability is the driving force behind bond formation! C. Stability • Transferring e- Sharing e- Ch. 6 - Chemical Bonds II. Kinds of Chemical Bonds (p.304-308) Ionic Bond Covalent Bond Comparison Chart A. Ionic Bond • Attraction between 2 oppositely charged ions – Ions - charged atoms – formed by transferring efrom a metal to a nonmetal A. Ionic Bond – ions form a 3-D crystal lattice NaCl B. Covalent Bond • Attraction between neutral atoms – formed by sharing e- between two nonmetals B. Covalent Bond – covalent bonds result in discrete molecules Cl2 NH3 H2O B. Covalent Bond Nonpolar Covalent Bond • e- are shared equally • usually identical atoms B. Covalent Bond Polar Covalent Bond • e- are shared unequally between 2 different atoms • results in partial opposite charges + B. Covalent Bond Nonpolar Polar Ionic View Bonding Animations. C. Comparison Chart IONIC transferred from metal to nonmetal COVALENT shared between nonmetals Melting Point high low Soluble in Water yes usually not Electrons Conduct Electricity Other Properties yes no (solution or liquid) crystal lattice of ions, molecules, odorous liquids & gases crystalline solids Ch. 6 - Chemical Bonds III. Naming Molecular Compounds Molecular Names Molecular Formulas A. Molecular Names Write the names of both elements. Change the final ending to -ide. Add prefixes to indicate subscripts. Only use mono- prefix with oxide. A. Molecular Names PREFIX monoditritetrapentahexa- SUBSCRIPT 1 2 3 4 5 6 A. Molecular Names CCl4 • carbon tetrachloride N2O • dinitrogen monoxide SF6 • sulfur hexafluoride B. Molecular Formulas Write the more metallic element first. Add subscripts according to prefixes. B. Molecular Formulas phosphorus trichloride • PCl3 dinitrogen pentoxide • N2O5 dihydrogen monoxide • H2O B. Molecular Formulas The Seven Diatomic Elements Br2 I2 N2 Cl2 H2 O2 F2 Ch. 6 - Chemical Bonds IV. Naming Ionic Compounds (p. 314-320) Oxidation Number Ionic Names Ionic Formulas A. Oxidation Number • The charge on an ion. • Indicates the # of e- gained/lost to become stable. 1+ 2+ 3+ 4+ 3- 2- 1- 0 B. Ionic Names Write the names of both elements, cation first. Change the anion’s ending to -ide. Write the names of polyatomic ions. For ions with variable oxidation #’s, write the ox. # in parentheses using Roman numerals. Overall charge = 0. B. Ionic Names NaBr • sodium bromide Na2CO3 • sodium carbonate FeCl3 • iron(III) chloride C. Ionic Formulas Write each ion. Put the cation first. Overall charge must equal zero. • If charges cancel, just write the symbols. • If not, crisscross the charges to find subscripts. Use parentheses when more than one polyatomic ion is needed. Roman numerals indicate the oxidation #. C. Ionic Formulas potassium chloride • K+ Cl- KCl magnesium nitrate • Mg2+ NO3- Mg(NO3)2 copper(II) chloride • Cu2+ Cl- CuCl2 C. Ionic Formulas calcium oxide • Ca2+ O2- CaO aluminum chlorate • Al3+ ClO3- Al(ClO3)3 iron(III) oxide • Fe3+ O2- Fe2O3 Ch. 6 - Chemical Bonds V. Naming Acids Definition Acid Names Acid Formulas A. Definition Acid • Compound that forms H+ in water. • Formula usually begins with ‘H’. Examples: • HCl, HNO3, H2SO4 B. Acid Names Anion Ending Acid Name -ide hydro-(stem)-ic acid -ate (stem)-ic acid -ite (stem)-ous acid B. Acid Names HBr • -ide hydrobromic acid carbonic acid sulfurous acid H2CO3 • -ate H2SO3 • -ite C. Acid Formulas hydrofluoric acid • -ide HF H2SO4 HNO2 sulfuric acid • -ate nitrous acid • -ite