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Matter

VIEWS: 484 PAGES: 158

									Compounds and Their Bonds
           Chapter 4
A molecule is an aggregate of two or more atoms in a
definite arrangement held together by chemical bonds

  A diatomic molecule contains only two atoms
                H2, N2, O2, Br2, HCl, CO

A polyatomic molecule contains more than two atoms
                  O3, H2O, NH3, CH4




                   H2         H2O          NH3   CH4
Chemical Bonds



   Formed by an attraction between two or more
    (charged) atoms
       Ionic bonds
   Or by the interaction between valence
    electrons
       Covalent bonds
An ion is an atom, or group of atoms, that has a net
positive or negative charge.
cation – ion with a positive charge
     If a neutral atom loses one or more electrons
     it becomes a cation.

              11 protons                  11 protons
      Na      11 electrons         Na+    10 electrons


anion – ion with a negative charge
     If a neutral atom gains one or more electrons
     it becomes an anion.
              17 protons                  17 protons
       Cl     17 electrons         Cl-    18 electrons
A monatomic ion contains only one atom
           Na+, Cl-, Ca2+, O2-, Al3+, N3-




A polyatomic ion contains more than one atom
              OH-, CN-, NH4+, NO3-
  Formation of Magnesium Ion
Magnesium atom       Magnesium ion
 
 Mg         – 2e        Mg2+

2-8-2                         2-8 (=Ne)

     12 p+                    12 p+
     12 e-                    10 e-
      0                       2+
Ions from Nonmetal Ions
 In ionic compounds, nonmetals in 5A, 6A, and 7A
  gain electrons from metals

 Nonmetals add electrons to achieve the octet
  arrangement

 Nonmetal ionic charge:
  3-, 2-, or 1-
         Do You Understand Ions?


                                        27 3+
How many protons and electrons are in   13 Al   ?

         13 protons, 10 (13 – 3) electrons

                                        78 2-
How many protons and electrons are in   34 Se   ?

         34 protons, 36 (34 + 2) electrons
Octet Rule

 An octet in the outer shell makes atoms stable
 Electrons are lost, gained or shared to form an
  octet
 Unpaired valence electrons strongly influence
  bonding
    Formation of Ions from Metals
   Ionic compounds result when metals react with
    nonmetals
   Metals lose electrons to match the number of valence
    electrons of their nearest noble gas
   Positive ions form when the number of electrons are
    less than the number of protons
             Group 1A metals       ion 1+
             Group 2A metals       ion 2+
             Group 3A metals       ion 3+
Formation of Sodium Ion
Sodium atom            Sodium ion
 Na      – e       Na +

 2-8-1             2-8 ( = Ne)

  11 p+                   11 p+
  11 e-                  10 e-
   0                      1+
     Learning Check
       
A.    X would be the electron dot formula for
      1) Na             2) K             3) Al
        
B.     X    would be the electron dot formula
        

      1) B              2) N             3) P
  Learning Check

A. Why does Ca form a Ca2+ ion?




B. Why does O form O2- ion?
     Solution

A. Why does Ca form a Ca2+ ion?

       Loses 2 electrons to give octet
       2-8-8-2          2-8-8 (like Ar)

B.     Why does O form O2- ion?

       Gains 2 electrons to give octet
       2-6 + 2e-       2-8 (like Ne)
SKIP: Molecular vs. Empirical
  A molecular formula shows the exact number of
  atoms of each element in the smallest unit of a
  substance
  An empirical formula shows the simplest
  whole-number ratio of the atoms in a substance

           molecular          empirical
              H2O                H2O
            C6H12O6             CH2O

               O3                 O
              N2H4               NH2
ionic compounds consist of a combination of cations
and anions
• the formula is always the same as the empirical formula
• the sum of the charges on the cation(s) and anion(s) in each
       formula unit must equal zero
                The ionic compound NaCl
Formula of Ionic Compounds
   2 x +3 = +6           3 x -2 = -6

                 Al2O3
       Al3+               O2-

   1 x +2 = +2           2 x -1 = -2

                 CaBr2
       Ca2+               Br-

   1 x +2 = +2           1 x -2 = -2

              Na2CO3
       Na+                 CO32-
Chemical Nomenclature
   Ionic Compounds
     often a metal + nonmetal
     Monatomic anion (nonmetal): add “ide” to element name
     Use polyatomic anion name for polyatomic anions


            BaCl2               barium chloride
            K2O                 potassium oxide
            Mg(OH)2             magnesium hydroxide

            KNO3                potassium nitrate
    Transition metal ionic compounds
        indicate charge on metal with Roman numerals




FeCl2      2 Cl- -2 so Fe is +2          iron(II) chloride

FeCl3      3 Cl- -3 so Fe is +3          iron(III) chloride

Cr2S3      3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide
Ionic Compounds
   Electrons go from metals to nonmetals
        Attraction between + ions and - ions


electron transfer
    metal               nonmetal       ion+     ion–


          Electrons lost = Electrons gained
Writing a Formula
Write the formula for the ionic compound that will
form between Ba2+ and Cl.
Solution:
1. Balance charge with + and – ions
2. Write the positive ion of metal first, and the
  negative ion         Ba2+           Cl
                                      Cl
3. Write the number of ions needed as
  subscripts                  BaCl2
             Learning Check
Write the correct formula for the compounds
containing the following ions:
A. Na+, S2-
   1) NaS         2) Na2S         3) NaS2
B. Al3+, Cl-
   1) AlCl3       2) AlCl         3) Al3Cl
C. Mg2+, N3-
   1) MgN         2) Mg2N3        3) Mg3N2
 Naming Binary Ionic Compounds
 Contain 2 different elements
 Name the metal first, then the nonmetal as -ide.
 Use name of a metal with a fixed charge
           Groups 1A, 2A, 3A
           and Ag, Zn, and Cd
  Examples:
         NaCl            sodium chloride
         ZnI2            zinc iodide
         Al2O3           aluminum oxide
Learning Check
Complete the names of the following binary
compounds:
Na3N           sodium       ________________
KBr            potassium ________________

Al2O3          aluminum   ________________

MgS            _________________________
Solution
Complete the names of the following binary
compounds:
Na3N           sodium nitride
KBr            potassium bromide

Al2O3          aluminum oxide

MgS            magnesium sulfide
   Learning Check
A. The formula for the ionic compound of
   Na+ and O2- is
     1) NaO              2) Na2O           3) NaO2

B. The formula of a compound of aluminum and chlorine
  is
     1) Al3Cl           2) AlCl2        3) AlCl3

C. The formula of Fe3+ and O2- is
     1) Fe3O2           2) FeO3            3) Fe2O3
       Transition Metals
Many form 2 or more positive ions
  1+        2+          1+ or 2+          2+ or 3+
  Ag+       Cd2+        Cu+, Cu2+         Fe2+, Fe3+
  silver    cadmium     copper(I) ion     iron(II) ion
  ion       ion         copper (II) ion   iron(III) ion

            Zn2+
            zinc ion
Names of Variable Ions
Use a roman number after the name of a metal
that forms two or more ions
   Transition metals and
   the metals in groups 4A and 5A
FeCl3        (Fe3+)    iron (III) chloride
CuCl         (Cu+ )    copper (I) chloride
SnF4         (Sn4+)    tin (IV) fluoride
PbCl2        (Pb2+)    lead (II) chloride
Fe2S3        (Fe3+)    iron (III) sulfide
      The Ionic Bond


    Li + F         Li+ F -
1s22s1 1s22s22p5    1s2 1s22s22p6
                   [He]    [Ne]
           Li       Li+ + e-
    e- +   F        F -

   Li+ +   F -      Li+ F -
   Molecular compounds (covalent)
    •   nonmetal + nonmetal or
        nonmetals + metalloids
         •   common names
              • H2O, NH3, CH4
    •   if more than one compound can be
        formed from the same elements, use
        prefixes to indicate number of each
        kind of atom
    •   last element ends in ide
                     Molecular Compounds

        HI            hydrogen iodide

        NF3           nitrogen trifluoride

        SO2           sulfur dioxide

        N2Cl4         dinitrogen tetrachloride

        NO2           nitrogen dioxide           TOXIC!

        N2O           dinitrogen monoxide               Laughing Gas
Nitric oxide (NO) should not be confused with nitrous
   oxide (N2O) or with nitrogen dioxide (NO2)
Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that
particpate in chemical bonding.
(n is the principal quantum number)

             Group              e- configuration   # of valence e-
               1A                      ns1              1
               2A                      ns2              2
               3A                     ns2np1            3
               4A                     ns2np2            4
               5A                     ns2np3            5
               6A                     ns2np4            6
               7A                     ns2np5            7
     Valence Electrons
   Valence electrons are the electrons in the highest
    (outer) electron level
   Have the most contact with other atoms
   Outer shells of noble gases contain 8 valence
    electrons (except He = 2)
       Example:   Ne    2, 8
                   Ar    2, 8, 8
            SKIP: Electrostatic (Lattice) Energy
Lattice energy (E) is the energy required to completely separate
one mole of a solid ionic compound into gaseous ions.

                           Q+ is the charge on the cation
            Q+Q-
        E=k                Q- is the charge on the anion
             r
                           r is the distance between the ions


                                   cmpd     lattice energy
                                   MgF2          2957 Q= +2,-1
 Lattice energy (E) increases      MgO           3938 Q= +2,-2
    as Q increases and/or
        as r decreases.            LiF          1036
                                                       r F < r Cl
                                   LiCl         853
SKIP: Born-Haber Cycle for Determining Lattice Energy
         o
       DHoverall = DHo + DHo + DHo + DHo + DHo
                     1     2     3     4     5
SKIP:
   Covalent Bonds
 Formed between two nonmetals in 4A, 5A, 6A, and 7A
 Nonmetals have high electronegativity values
 Electrons are shared
     single bond shares one pair of electrons
     double bond shares two pairs of electrons
     triple bond shares three pairs of electrons
A covalent bond is a chemical bond in which two or more
electrons are shared by two atoms.

             Why should two atoms share electrons?

                      F    +       F               F F
                    7e-           7e-             8e- 8e-

                          Lewis structure of F2

         single covalent bond       lone pairs        F      F          lone pairs


                                                 single covalent bond
lone pairs      F F             lone pairs
            Covalent Bonds
Two nonmetal atoms form a covalent bond
because they have less energy after they bonded
H +   H        H : H = HH = H2




                    hydrogen molecule
Covalent Bonds in NH3
                 Bonding pairs
         H
         

     H   :   N   :   H
         
Lone pair of electrons
Lewis structure of water                    single covalent bonds

H   +   O +     H           H O H          or       H    O   H
                            2e- -2e-
                              8e

Double bond – two atoms share two pairs of electrons

            O C O           or         O        C        O
            8e- 8e- 8e-
           double bonds                double bonds

Triple bond – two atoms share three pairs of electrons

            N N             or              N       N
          triple-8e-
             8e bond
                                           triple bond
        SKIP: Lengths of Covalent Bonds
            Bond Lengths
Triple bond < Double Bond < Single Bond

                                                  Bond
                                          Bond   Length
                                          Type    (pm)

                                          C-C     154

                                          CC     133

                                          CC     120

                                          C-N     143

                                          CN     138

                                          CN     116
Naming Binary Covalent Compounds

Two nonmetals
 Name each element
 End the last element in -ide
 Add prefixes to show more than 1 atom
     Prefixes
            mon        1         penta    5
            di         2         hexa     6
            tri        3
            tetra      4
Learning Check
 Fill in the blanks to complete the following names
 of covalent compounds.
    CO          carbon ______oxide
    CO2         carbon _______________
    PCl3        phosphorus _______chloride
    CCl4        carbon ________chloride
    N2O         _____nitrogen _____oxide
Solution
    CO     carbon monoxide
    CO2    carbon dioxide
    PCl3   phosphorus trichloride
    CCl4   carbon tetrachloride
    N2O    dinitrogen monoxide
               Polyatomic Ions
A group of atoms with an overall charge.
  NH4+     ammonium OH-      hydroxide

  NO3-     nitrate          NO2-   nitrite

  CO32-    ______________ (carbonate)

  HCO3-    hydrogen carbonate (bicarbonate)
 More Polyatomic Ions
Sulfur
SO42-       sulfate            SO32- sulfite
HSO4-       hydrogen sulfate
HSO3-       hydrogen sulfite
Phosphate
PO43-       phosphate          PO33- phosphite
HPO42-      ________________
H2PO4-      dihydrogen phosphate
    Naming Ternary Compounds
   Contain at least 3 elements
   Name the nonmetals as a polyatomic ion
   Examples:
    NaNO3          Sodium nitrate
    K2SO4         Potassium sulfate
    Al(HCO3)3     Aluminum bicarbonate
                  or
                  Aluminum hydrogen carbonate
     Learning Check
Match each set with the correct name:
A.   Na2CO3              1) magnesium sulfite
     MgSO3               2) magnesium sulfate
     MgSO4               3) sodium carbonate
B.     Ca(HCO3)2       1) calcium carbonate
       CaCO3           2) calcium phosphate
       Ca3(PO4)2       3) calcium bicarbonate
     Solution

A.     Na2CO3      3) sodium carbonate
       MgSO3       1) magnesium sulfite
       MgSO4       2) magnesium sulfate

B.     Ca(HCO3)2   3) calcium bicarbonate
       CaCO3       1) calcium carbonate
       Ca3(PO4)2   2) calcium phosphate
  Learning Check
A. aluminum nitrate
   1) AlNO3    2) Al(NO)3            3) Al(NO3)3
B. copper(II) nitrate
    1) CuNO3        2) Cu(NO3)2 3) Cu2(NO3)
C. Iron (III) hydroxide
    1) FeOH         2) Fe3OH         3) Fe(OH)3
D. Tin(IV) hydroxide
    1) Sn(OH)4      2) Sn(OH)2       3) Sn4(OH)
           Electronegativity
 The attraction of an atom for electrons is
  called its electronegativity.
 Fluorine has the greatest electronegativity.
 The metals have low electronegativities.
           Bond Polarity: Nonpolar
Nonpolar covalent bond
 Electrons are shared equally between atoms

  with the same electronegativity values.
 Electronegativity Difference = 0

 Examples:


  N2      Br2
Polar covalent bond or polar bond is a covalent
bond with greater electron density around one of the
two atoms




                                        electron rich
                        electron poor
                                           region
e- poor   e- rich          region

    H     F                  H               F
    d+    d-
SKIP:
Electronegativity is the ability of an atom to attract
toward itself the electrons in a chemical bond.
       Electron Affinity - measurable, Cl is highest

                     X (g) + e-   X-(g)


         Electronegativity - relative, F is highest
           Bond Polarity: Polar
Polar covalent bond
 Electrons are shared (unequally) between

  different nonmetal atoms. Examples:
            O-Cl       O-S        N-Cl
             Bond Polarity: Ionic
Ionic bond
 Electrons are transferred between metal and

  nonmetal atoms. Example:
     NaCl            KF
 Classification of bonds by difference in electronegativity

                Difference            Bond Type
                    0             Non-Polar Covalent
                   2                     Ionic
               0 < and <2            Polar Covalent


           Increasing difference in electronegativity


Covalent                Polar Covalent                Ionic

share e-              partial transfer of e-      transfer e-
    Classify the following bonds as ionic, polar covalent,
    or covalent: The bond in CsCl; the bond in H2S; and
    the NN bond in H2NNH2.


Cs – 0.7      Cl – 3.0      3.0 – 0.7 = 2.3    Ionic

H – 2.1       S – 2.5       2.5 – 2.1 = 0.4    Polar Covalent

N – 3.0       N – 3.0       3.0 – 3.0 = 0      Covalent
Learning Check
Indicate whether a bond between the following would
be 1) Ionic     2) covalent

____     A.   sodium and oxygen
____     B.   nitrogen and oxygen
____     C.   phosphorus and chlorine
____     D.   calcium and sulfur
____     E.   chlorine and bromine
SKIP:
          SKIP: Writing Lewis Structures

1. Draw skeletal structure of compound showing
   what atoms are bonded to each other. Put least
   electronegative element in the center.
2. Count total number of valence e-. Add 1 for
   each negative charge. Subtract 1 for each
   positive charge.
3. Complete an octet for all atoms except
   hydrogen
4. If structure contains too many electrons, form
   double and triple bonds on central atom as
   needed.
SKIP:Write the Lewis structure of nitrogen trifluoride (NF3)
Step 1 – N is less electronegative than F, put N in center
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
            5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete
         octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons




F      N      F

       F
SKIP:Write the Lewis structure of the carbonate ion (CO32-)
Step 1 – C is less electronegative than O, put C in center
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
         -2 charge – 2e-
           4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete
         octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-

                        2 single bonds (2x2) = 4
                               1 double bond = 4
O      C      O            8 lone pairs (8x2) = 16
                                        Total = 24
       O
SKIP:
A resonance structure is one of two or more Lewis structures
for a single molecule that cannot be represented accurately by
only one Lewis structure.
                   +      -           -    +
             O   O     O                O    O   O


        What are the resonance structures of the
        carbonate (CO32-) ion?



-                 -                        -   -
    O    C    O           O    C       O           O   C       O

        O                      O                       O
                                   -                       -
         SKIP: Exceptions to the Octet Rule

The Incomplete Octet

                         Be – 2e-
           BeH2        2H – 2x1e-       H      Be      H
                              4e-




          B – 3e-                           3 single bonds (3x2) = 6
       3F – 3x7e-      F    B       F
BF3                                            9 lone pairs (9x2) = 18
             24e-                                           Total = 24
                            F
           SKIP: Exceptions to the Octet Rule

Odd-Electron Molecules

              N – 5e-
  NO          O – 6e-        N       O
                 11e-

The Expanded Octet (central atom with principal quantum number n > 2)


                                 F
                             F       F
              S – 6e-                        6 single bonds (6x2) = 12
 SF6        6F – 42e-            S           18 lone pairs (18x2) = 36
                 48e-                                        Total = 48
                             F       F
                                 F
Valence shell electron pair repulsion (VSEPR)
model:
Predict the geometry of the molecule from the electrostatic
repulsions between the electron (bonding and nonbonding) pairs.

          # of atoms       # lone
          bonded to       pairs on     Arrangement of    Molecular
 Class   central atom   central atom    electron pairs   Geometry

 AB2          2              0             linear            linear

                                                         B            B
       Cl    Be    Cl


  0 lone pairs on central atom
2 atoms bonded to central atom
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of    Molecular
Class   central atom   central atom    electron pairs   Geometry

AB2          2              0             linear         linear
                                         trigonal       trigonal
AB3          3              0
                                          planar         planar
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of     Molecular
Class   central atom   central atom    electron pairs    Geometry

AB2          2              0             linear          linear
                                         trigonal        trigonal
AB3          3              0
                                          planar          planar
AB4          4              0         tetrahedral       tetrahedral
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of     Molecular
Class   central atom   central atom    electron pairs    Geometry

AB2          2              0             linear           linear
                                         trigonal         trigonal
AB3          3              0
                                          planar           planar
AB4          4              0         tetrahedral       tetrahedral
                                        trigonal          trigonal
AB5          5              0
                                      bipyramidal       bipyramidal
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of     Molecular
Class   central atom   central atom    electron pairs    Geometry

AB2          2              0             linear           linear
                                         trigonal         trigonal
AB3          3              0
                                          planar           planar
AB4          4              0         tetrahedral       tetrahedral
                                        trigonal          trigonal
AB5          5              0
                                      bipyramidal       bipyramidal
AB6          6              0         octahedral        octahedral
bonding-pair vs. bonding     lone-pair vs. bonding
     pair repulsion      <       pair repulsion    < lone-pair vs. lone pair
                                                           repulsion
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of    Molecular
Class   central atom   central atom    electron pairs   Geometry

                                         trigonal       trigonal
AB3          3              0
                                          planar         planar
                                         trigonal
AB2E         2              1                             bent
                                          planar
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of     Molecular
Class   central atom   central atom    electron pairs    Geometry

AB4          4              0         tetrahedral       tetrahedral
                                                         trigonal
AB3E         3              1         tetrahedral
                                                        pyramidal
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of     Molecular
Class   central atom   central atom    electron pairs    Geometry


AB4          4              0         tetrahedral       tetrahedral
                                                         trigonal
AB3E         3              1         tetrahedral
                                                        pyramidal

AB2E2        2              2         tetrahedral          bent
                                                              O
                                                          H       H
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of     Molecular
Class   central atom   central atom    electron pairs    Geometry
                                        trigonal          trigonal
AB5          5              0
                                      bipyramidal       bipyramidal
                                        trigonal          distorted
AB4E         4              1
                                      bipyramidal       tetrahedron
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of        Molecular
Class   central atom   central atom    electron pairs       Geometry
                                        trigonal          trigonal
AB5          5              0
                                      bipyramidal       bipyramidal
                                        trigonal          distorted
AB4E         4              1
                                      bipyramidal       tetrahedron
                                        trigonal
AB3E2        3              2                            T-shaped
                                      bipyramidal
                                                              F

                                                        F     Cl

                                                              F
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of     Molecular
Class   central atom   central atom    electron pairs    Geometry
                                        trigonal          trigonal
AB5          5              0
                                      bipyramidal       bipyramidal
                                        trigonal          distorted
AB4E         4              1
                                      bipyramidal       tetrahedron
                                        trigonal
AB3E2        3              2                            T-shaped
                                      bipyramidal
                                        trigonal
AB2E3        2              3                             linear
                                      bipyramidal
                                                            I

                                                            I

                                                            I
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of     Molecular
Class   central atom   central atom    electron pairs    Geometry

AB6          6              0         octahedral        octahedral

                                      octahedral         square
AB5E         5              1
                                                        pyramidal
                                                            F
                                                         F    F
                                                             Br
                                                         F        F
                           VSEPR
         # of atoms       # lone
         bonded to       pairs on     Arrangement of     Molecular
Class   central atom   central atom    electron pairs    Geometry

AB6          6              0         octahedral        octahedral

                                      octahedral         square
AB5E         5              1
                                                        pyramidal
                                                         square
AB4E2        4              2         octahedral
                                                         planar
                                                         F        F
                                                             Xe
                                                         F        F
Predicting Molecular Geometry
  1. Draw Lewis structure for molecule.
  2. Count number of lone pairs on the central atom and
     number of atoms bonded to the central atom.
  3. Use VSEPR to predict the geometry of the molecule.

     What are the molecular geometries of SO2 and SF4?

          O    S     O                F
                                                   AB4E
              AB2E              F    S    F
                                                  distorted
              bent                              tetrahedron
                                      F
        How does Lewis theory explain the bonds in H2 and F2?

       Sharing of two electrons between the two atoms.

      Bond Dissociation Energy    Bond Length      Overlap Of

H2          436.4 kJ/mole             74 pm          2 1s

F2          150.6 kJ/mole            142 pm          2 2p


Valence bond theory – bonds are formed by sharing
of e- from overlapping atomic orbitals.
     Valence bond theory (hybridization) is one of two
     attempts to explain HOW covalent bonds form!
Change in electron
density as two hydrogen
atoms approach each
other.
Valence Bond Theory and NH3

N – 1s22s22p3

 3 H – 1s1

  If the bonds form from overlap of 3 2p orbitals on nitrogen
  with the 1s orbital on each hydrogen atom, what would
  the molecular geometry of NH3 be?
                                                If use the
                                              3 2p orbitals
                                               predict 900

                                             Actual H-N-H
                                             bond angle is
                                                107.30
Hybridization – mixing of two or more
atomic orbitals to form a new set of
hybrid orbitals.
1. Mix at least 2 nonequivalent atomic orbitals (e.g. s
   and p). Hybrid orbitals have very different shape
   from original atomic orbitals.
2. Number of hybrid orbitals is equal to number of
   pure atomic orbitals used in the hybridization
   process.
3. Covalent bonds are formed by:
   a. Overlap of hybrid orbitals with atomic orbitals
   b. Overlap of hybrid orbitals with other hybrid
      orbitals
Predict correct
bond angle
Formation of sp Hybrid Orbitals
Formation of sp2 Hybrid Orbitals
   How do I predict the hybridization of the central atom?

         Count the number of lone pairs AND the number
         of atoms bonded to the central atom

  # of Lone Pairs
         +
# of Bonded Atoms        Hybridization       Examples
        2                      sp               BeCl2

        3                     sp2                BF3

        4                     sp3         CH4, NH3, H2O

        5                     sp3d              PCl5

        6                    sp3d2               SF6
Pi bond (p) – electron density above and below plane of nuclei
Sigma bond (s) – electron density between the 2 atoms
              of the bonding atoms
Sigma (s) and Pi Bonds (p)

Single bond                 1 sigma bond

Double bond           1 sigma bond and 1 pi bond

Triple bond           1 sigma bond and 2 pi bonds

  How many s and p bonds are in the acetic acid
  (vinegar) molecule CH3COOH?
              O




       H
                            s bonds = 6 + 1 = 7
  H    C      C   O    H
                            p bonds = 1
       H
                        Experiments show O2 is paramagnetic
            O
       O
   No unpaired e-
Should be diamagnetic




 Molecular orbital theory – bonds are formed from
 interaction of atomic orbitals to form molecular
 orbitals.
3 models for molecular bonding!
   Three different theories used to explain how molecules bond:
        VSEPR Theory (with Lewis Dot Structures) – predicts WHAT shape they take
        Valence Bond Theory (with hybridization) – explains HOW they bond (1st theory)
        Molecular Orbital Theory – explains HOW they bond (2nd theory)
   A good theory should predict physical and chemical properties of the molecule
    such as shape, bond energy, bond length, and bond angles.
   One model does not describe all the properties of molecular bonds.
   Each model desribes a set of properties better than the others.
   The final test for any theory is experimental data.
   The Molecular Orbital Theory does a good job of predicting electronic spectra and
    paramagnetism, when VSEPR and the V-B Theories don't.
   The MO theory does not need resonance structures to describe molecules, as well
    as being able to predict bond length and energy.
   The major draw back is that we are limited to talking about diatomic molecules
    (molecules that have only two atoms bonded together), or the theory gets very
    complex.
   The MO theory treats molecular bonds as a sharing of electrons between nuclei.
    Unlike the V-B theory, which treats the electrons as localized balloons of electron
    density, the MO theory says that the electrons are delocalized. That means that
    they are spread out over the entire molecule.
   Now, when two atoms come together, their two atomic orbitals react to form two
    possible molecular orbitals. One of the molecular orbitals is lower in energy. It is
    called the bonding orbital and stabilizes the molecule. The other orbital is called an
    anti-bonding orbital. It is higher in energy than the original atomic orbitals and
    destabilizes the molecule.
                      MO Theory’s Rules:
   The MO Theory has five basic rules:
    1.   The number of molecular orbitals = the number of atomic
         orbitals combined
    2.   Of the two MOs, one is a bonding orbital (lower energy) and
         one is an anti-bonding orbital (higher energy)
    3.   Electrons enter the lowest orbital available
    4.   The maximum number of electrons in an orbital is 2 (Pauli
         Exclusion Principle)
    5.   Electrons spread out before pairing up (Hund's Rule)
Energy levels of bonding and antibonding molecular
              orbitals in hydrogen (H2).

A bonding molecular orbital has lower energy and greater
stability than the atomic orbitals from which it was formed.

An antibonding molecular orbital has higher energy and
lower stability than the atomic orbitals from which it was
formed.
Molecular Orbital (MO) Configurations

1. The number of molecular orbitals (MOs) formed is always
   equal to the number of atomic orbitals combined.
2. The more stable the bonding MO, the less stable the
   corresponding antibonding MO.
3. The filling of MOs proceeds from low to high energies.
4. Each MO can accommodate up to two electrons.
5. Use Hund’s rule when adding electrons to MOs of the
   same energy.
6. The number of electrons in the MOs is equal to the sum of
   all the electrons on the bonding atoms.
                       Number of          Number of
  bond order =
               1
               2   (   electrons in
                       bonding
                       MOs
                                      -   electrons in
                                          antibonding
                                          MOs
                                                             )



bond
         ½             1              ½                  0
order
Delocalized molecular orbitals are not confined between
two adjacent bonding atoms, but actually extend over three
or more atoms – Benzene!
Electron density above and below the plane of the
benzene molecule.
Chemistry In Action: Buckyball Anyone?
An acid can be defined as a substance that yields
hydrogen ions (H+) when dissolved in water.
  HCl
    •Pure substance, hydrogen chloride
    •Dissolved in water (H+ Cl-), hydrochloric acid

An oxoacid is an acid that contains hydrogen,
oxygen, and another element.

  HNO3          nitric acid
  H2CO3         carbonic acid
  H2SO4         sulfuric acid
                                         HNO3
A base can be defined as a substance that yields
hydroxide ions (OH-) when dissolved in water.


            NaOH          sodium hydroxide
            KOH           potassium hydroxide
            Ba(OH)2       barium hydroxide
Chemistry In Action:
  Sodium Chloride




      Mining Salt      Solar Evaporation for Salt
  Two possible skeletal structures of formaldehyde (CH2O)

                                             H
      H     C       O   H                              C   O
                                             H

An atom’s formal charge is the difference between the
number of valence electrons in an isolated atom and the
number of electrons assigned to that atom in a Lewis
structure.
formal charge       total number
                                   total number        1       total number
on an atom in
a Lewis
structure
                =
                    of valence
                    electrons in -
                    the free atom
                                   of nonbonding
                                   electrons
                                                   -   2   (   of bonding
                                                               electrons  )
 The sum of the formal charges of the atoms in a molecule
 or ion must equal the charge on the molecule or ion.
       -1   +1                      C – 4 e-    2 single bonds (2x2) = 4
 H     C        O       H           O – 6 e-           1 double bond = 4
                                 2H – 2x1 e-       2 lone pairs (2x2) = 4
                                      12 e-                     Total = 12


formal charge           total number
                                       total number        1       total number
on an atom in
a Lewis
structure
                    =
                        of valence
                        electrons in -
                        the free atom
                                       of nonbonding
                                       electrons
                                                       -   2   (   of bonding
                                                                   electrons  )
     formal charge
         on C
                   = 4 -2 -½ x 6 = -1

     formal charge
         on O
                   = 6 -2 -½ x 6 = +1
    H    0      0               C – 4 e-    2 single bonds (2x2) = 4
         C      O               O – 6 e-           1 double bond = 4
    H                        2H – 2x1 e-       2 lone pairs (2x2) = 4
                                  12 e-                     Total = 12


formal charge       total number
                                   total number        1       total number
on an atom in
a Lewis
structure
                =
                    of valence
                    electrons in -
                    the free atom
                                   of nonbonding
                                   electrons
                                                   -   2   (   of bonding
                                                               electrons  )
     formal charge
         on C
                   = 4 - 0 -½ x 8 = 0

     formal charge
         on O
                   = 6 -4 -½ x 4 = 0
           Formal Charge and Lewis Structures
1. For neutral molecules, a Lewis structure in which there
   are no formal charges is preferable to one in which
   formal charges are present.
2. Lewis structures with large formal charges are less
   plausible than those with small formal charges.
3. Among Lewis structures having similar distributions of
   formal charges, the most plausible structure is the one in
   which negative formal charges are placed on the more
   electronegative atoms.
     Which is the most likely Lewis structure for CH2O?

             -1   +1                   H    0    0
       H     C    O    H                    C    O
                                       H
The enthalpy change required to break a particular bond in
one mole of gaseous molecules is the bond energy.

                                   Bond Energy
    H2 (g)         H (g) + H (g)   DH0 = 436.4 kJ
    Cl2 (g)        Cl (g) + Cl (g) DH0 = 242.7 kJ
   HCl (g)         H (g) + Cl (g) DH0 = 431.9 kJ
    O2 (g)         O (g) + O (g) DH0 = 498.7 kJ     O   O
    N2 (g)         N (g) + N (g) DH0 = 941.4 kJ     N   N


                           Bond Energies
              Single bond < Double bond < Triple bond
Average bond energy in polyatomic molecules

     H2O (g)      H (g) + OH (g) DH0 = 502 kJ

      OH (g)      H (g) + O (g)   DH0 = 427 kJ
                               502 + 427
      Average OH bond energy =           = 464 kJ
                                   2
Bond Energies (BE) and Enthalpy changes in reactions
         Imagine reaction proceeding by breaking all bonds in the
         reactants and then using the gaseous atoms to form all the
         bonds in the products.
          DH0 = total energy input – total energy released
              = SBE(reactants) – SBE(products)
H2 (g) + Cl2 (g)   2HCl (g)   2H2 (g) + O2 (g)   2H2O (g)
     Use bond energies to calculate the enthalpy change for:
                   H2 (g) + F2 (g)    2HF (g)

       DH0 = SBE(reactants) – SBE(products)


   Type of       Number of       Bond energy       Energy
bonds broken    bonds broken       (kJ/mol)      change (kJ)
 H     H             1             436.4             436.4
 F     F             1             156.9             156.9
   Type of       Number of       Bond energy       Energy
bonds formed    bonds formed       (kJ/mol)      change (kJ)
 H     F             2             568.2          1136.4

       DH0 = 436.4 + 156.9 – 2 x 568.2 = -543.1 kJ

                                                  Released!
Dipole Moments and Polar Molecules


                                        electron rich
                        electron poor
                                           region
                           region

                             H               F


                             d+              d


m=Qxr
Q is the charge
r is the distance between charges
1 D = 3.36 x 10-30 C m
  Which of the following molecules have a dipole moment?
  H2O, CO2, SO2, and CH4


       O                                S

 dipole moment                   dipole moment
 polar molecule                  polar molecule

                                         H

                                   H     C    H
   O    C   O

no dipole moment                         H
nonpolar molecule                no dipole moment
                                 nonpolar molecule
Does CH2Cl2 have
a dipole moment?
Chemistry In Action: Microwave Ovens

								
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