; notes ch 14 polar mc and complex ions
Documents
Resources
Learning Center
Upload
Plans & pricing Sign in
Sign Out
Your Federal Quarterly Tax Payments are due April 15th Get Help Now >>

notes ch 14 polar mc and complex ions

VIEWS: 7 PAGES: 6

  • pg 1
									                                     Chapter 14
                                   Polar Molecules

Introduction
         We know that ions are held together in the crystal
         lattice by the strong attractive forces of the
         oppositely charged ions. But how do neutral
         molecules of solids and liquid “stick together” if
         there are not opposite charges? These
         INTERMOLECULAR FORCES are a function
         of internal molecular structure and vary widely
         in their strength.

I. Polarity (the existence of poles)

   •   Polar molecules are generally the result of polar bonds
   •   not all polar bonds produce polar molecules
   •   poles within a molecule are indicated with a  (+ or -)
   •   polar molecules (dipoles) have a dipole moment
         def: numerical measure of the strength of dipole
         formula: µ = Qd
                     µ = dipole moment (coulomb•meters)
                     Q = size of charge in coulombs
                     d = distance between charges in meters

           higher dipole moment = higher melting/boiling pts.
                                  when similar mass




Mary Payton, Carroll High School          1
   • Practice:
       Determine which of the following molecules are
          dipoles. Indicate the + and - when applicable.

       NH3                         HF        CO2




       CCl4                        CH3Cl       H2O




   HBr                             SO2         BF3




Mary Payton, Carroll High School         2
II.        Weak Forces (Intermolecular Forces)
           aka: Van der Waals forces

       • forces BETWEEN molecules
         (as opposed to WITHIN molecule (bond))

       • these forces responsible for MANY of the different
           physical properties of substances.
           ex. boiling/melting point, vapor pressure, etc.

       • effective only over short distances (1/d6)

       • TYPES OF WEAK FORCES
         1) Dipole-Dipole forces

               • attraction between two polar molecules
               • special case in molecules where H bonded to
                 F, O, or N ===> hydrogen bonding
                    1) F,O,N very small & high EN
                            => suck e- from H
                    2) H now acts like naked proton
                         (very attractive)
                    3) + H attracts to lone pairs on other
                    F,O,N atoms rather strongly
                    4) Hydrogen bond energy approx. 5 x
                       regular dipole-dipole energies

           2) Dipole-Induced dipole forces

                   • polar molecules influence nonpolar
                   molecules to form an induced dipole

Mary Payton, Carroll High School   3
                   • repulsion between - end of dipole induces
                     non-polar mc’s e- to run to other side
                     creating an induced dipole

           3) London dispersion forces

                   • forces between 2 non-polar molecules
                   • largest contributor to van der Waals forces
                   • e- are moving could end up on one side of
                     mc creating a temporary (instantaneous)
                     dipole.
                   • temporary dipole can induce another temp.
                     dipole etc. holding substance together

       • Summary Table p. 354


III. Coordination Chemistry

       • polar molecules surround ions in soln. (complex
                ions)
       • polar molecules or anion (ligands) attach to metal
          by coordinate covalent bond
           • ligand donates a lone pair of e-
                 one pair = unidentate
                 two pair = di or bidentate
                 many pair = polydentate
           • ex. NH3, H2O, halides, hydroxides,EDTA

       • # of ligands attached = coordination number
           • if 6 then octahedral ex: [PtCl6]2-

Mary Payton, Carroll High School   4
               • if 4 then tetrahedral [CoI4]2- or square planar
                   [NiBr4]2-
               • if 2 then linear [Ag(CN)2]-




       • NAMING COMPLEX IONS

           1) For all complex ions:
                a) name the ligand first
                b) use prefix for # ligands
                c) use -o if ligand is anion
                d) alphabetical order of ligands (ignore prefix)
                e) Roman numeral for metal’s oxidation state
                f) word “ion”

                  Common Ligands
    Neutral Molecules                Anions
                            -      -
CO           Carbonyl    Br , Cl …       Bromo,
                                         Chloro
NH3          Ammine      CN-             Cyano
                              -
NO           Nitrosyl    OH              Hydroxo
                           2-
H2O          Aqua        O               Oxo
                          2-
                         S               Thio
                                -
                         SCN             Thiocyanato
                         S2O32-          Thosulfato
                                2-
                         C2O4            (oxalato)
                                         organic in ()



           2) Negative complex ions ONLY
Mary Payton, Carroll High School   5
                   a) use Latin root for metal + ate ending

           • Practice problems p. 361


       • WRITING COMPLEX ION FORMULAS

           1) write the central ion first
           2) then negative ligands (alphabetical order)
           3) then neurtal molecules (alphabetical order)
           4) polyatomics always in parentheses (even if 1)

           • practice problems p. 362


       • COORDINATION COMPOUNDS
         • name positive ion, then negative ion


       • BONDING IN COMPLEXES
         • most complex ions from transition metals
         • small ions w/ high oxidation # = high charge
                density ==> favorable for complexes
         • involve d orbitals
         • when in complex ions, previously degenerate d
                orbitals split into two groups (2 orbitals high
                energy/ 3 orbitals low energy)
         • e- transitions between split orbitals = bright color
              EX: [Cu(H2O)6]2+ = bright blue

IV. Chromatography
      SELF STUDY SECTION -- READ ON OWN

Mary Payton, Carroll High School    6

								
To top