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CHAPTER 5 and 6 OVERVIEW

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CHAPTER 5 and 6 OVERVIEW Powered By Docstoc
					ACADEMIC CHEMISTRY                                                      NAME:
PERIOD:                                                                 DATE:
                           TYPES OF BONDING in COMPOUNDS

         IONIC BONDS                                       COVALENT BONDS


                       What happens to
                       the electrons in     POLAR COVALENT                  PURE COVALENT
                         this bond?



                       What is the
                       electro-negativity
                       difference of the
                       atoms involved?




                       What type of
                       elements are
                       involved?



                       What is the
                       simplest unit of
                       this compound
                       called?




                       How is bond
                       strength
                       measured?
ACADEMIC CHEMISTRY                              NAME:
PERIOD:                                         DATE:


             IOINIC BONDS                   COVALENT BONDS




                              What are
                                some
                            properties of
                            compounds
                              with this
                               type of
                              bonding?
ACADEMIC CHEMISTRY                                                                                                NAME:
PERIOD:                                                                                                           DATE:



                                                           IONIC COMPOUNDS
Ion Formation
Why do ions form?

       Atoms become stable by gaining or losing VALENCE electrons to achieve a full outer ENERGY level.

           o Losing electrons produces a POSITIVE charge and the ion is called a CATION.

           o Gaining electrons produces a NEGATIVE charge and the ion is called an ANION.

       Normally, atoms will try to achieve a full octet and noble gas electron configuration to become stable.

Ex. Li loses 1 electron and forms Li+ which has the same electron configuration as Helium.

[He]1s1        [He]



Ex. S gains 2 electrons and forms S–2 which has the same configuration as Argon.

[Ne] 3s23p4  [Ne]3s23p6  [Ar]




       Exceptions to the octet rule:

   1.     Hydrogen and He only need 2 (not 8) to be stable, since they only include an S orbital.


   2.     TRANSITION METALS do not have to achieve noble gas configuration to become more stable.
ACADEMIC CHEMISTRY                                                                       NAME:
PERIOD:                                                                                  DATE:

Ex. Fe can have charges of +2, +3, or +6

Energy level diagram: (see board)


Lose the 2 electrons in the 4 s orbital: +2
Lose the 2 electrons in the 4s and 1 from 3d: +3
Lose the 6 electrons in the 3d: +6



      Ions have different properties than the parent atom.
          o Recall, the size of the ion changes.

                    Cations are smaller than the parent atom.

                    Anions are larger than the parent atom.


          o Ions are more stable than the parent atom.


Ex. Once sodium loses an electron, it is no longer reactive when it becomes +1.

We can eat Na+ in NaCl, but we could never eat Na (s) – it explodes in water!!!



                                                  POLYATIOMIC IONS
      A group of two or more atoms COVALENTLY bonded together with an OVERALL CHARGE.

See polyatomic ion sheet.
Memorize those with *. You will be quizzed on their names and formulas (with charges).
ACADEMIC CHEMISTRY                                                                                       NAME:
PERIOD:                                                                                                  DATE:
Ex. NH4+

Nitrogen and hydrogen share electrons.
(copy drawing from the board)

NH4+ - covalent inside


NH4Cl – ionic compound


*** THE BONDING WITHIN A POLYATOMIC ION IS COVALENT….but WHEN A POLYATOMIC ION IS IN A COMPOUND, IT WILL
BE AN IONIC COMPOUND

                                        IONIC FORMULAS AND NAMING IONIC COMPOUNDS

      Ionic compounds must have a net charge equal to zero.

      Subscripts are used to balance positive and negative charges. Use parentheses when a polyatomic ion has a subscript.

      A short cut method for balancing charges is the criss-cross method.

       1. Write the cation and its charge followed by the anion with its charge.

       Ex.

       2. Criss-cross the charges to form subscripts on the opposite ion.

       Ex.
ACADEMIC CHEMISTRY                                                                                            NAME:
PERIOD:                                                                                                       DATE:
Rules for Naming Ionic Compounds



      1. Name the cation first.

             a. Use the name of the element from the periodic table.

             Ex.

             b. If the cation is a transition metal, you must indicate the charge (oxidation state) on the ion using Roman numerals in

      parentheses following the cation name.

      2. Name the anion second.

             a. Use the name of the anion form the periodic table, drop the ending and add –ide.

             Ex.    Oxygen becomes oxide.

                    Phosphorus becomes phosphide.

                    Nitrogen becomes nitride.

                    Sulfur becomes sulfide.

      3. If a polyatomic ion is present, use the name given on the polyatomic ion sheet.

             Ex.



Write the names of the following compounds.

      BeI2
ACADEMIC CHEMISTRY                                                     NAME:
PERIOD:                                                                DATE:

       NiS
       NiCl3
       SnO
       (NH4)Cr2O7




Writing Formulas From the Name



1. Write the cation first with its charge.

       a. The charge on transition metals is given in the name.

       b. If a polyatomic ion exists, use your polyatomic ion sheet.

2. Write the anion next with its charge.

       a. Find the charge using your periodic table.

       b. If a polyatomic ion exists, use your polyatomic ion sheet.

3. Criss-cross the charges to find subscripts.

4. Rewrite the fomula without charges.



Ex. Write the formulas for the following compounds.

       1. cesium nitride
       2. strontium sulfate
ACADEMIC CHEMISTRY                                                                             NAME:
PERIOD:                                                                                        DATE:

      3. copper (II) oxide
      4. ammonium fluoride
      5. barium carbonate



                                                     COVALENT COMPOUNDS


 Formed when two atoms share electrons
         o Usually involve 2 or more non-metal atoms
         o Similar electronegativity values (electronegativity difference is less than 1.7).


                  Can be pure covalent – equal sharing of e (<.3)
                  Polar covalent – uneven sharing of electrons (.3-1.7)


         o Called molecules


         o There are more covalent compounds than ionic compounds.


                  Binary covalent – include 2 elements only


                  Organic molecules (which include carbon) are covalent
                         Glucose       C6H12O6
                         Sucrose       C12H22O11
                         Hydrocarbon         C6H6 etc.
ACADEMIC CHEMISTRY                                                                                        NAME:
PERIOD:                                                                                                   DATE:

 The shared valence electrons are found in a molecular orbital, formed by overlapping atomic orbitals of the atoms involved in
  bonding.


 Recall the properties of covalent molecules listed previously in chapter 5 notes (flow chart).


                                                  NAMING COVALENT COMPOUNDS


 Covalent molecules are NOT named like ionic compounds!!!


 Steps to naming binary covalent molecules: (only include 2 elements)
   1. the first element (least electronegative) is named first
   2. the second element has the ending –ide
   3. Prefixes are used before the element name to indicate the number of each atom in the molecule (the subscript)




             # of atoms                          Prefix
                  1                     mono
                  2                     di
                  3                     tri
                  4                     tetra
ACADEMIC CHEMISTRY                                                                                  NAME:
PERIOD:                                                                                             DATE:
            5                          penta
                      6                hexa
                      7                hepta
                      8                octa
                      9                nona
                  10                   deca


Do NOT use a prefix for the first atom when there is only one present.


When the element name begins with a vowel, the (a) and (o) are dropped from the end of the prefix




Examples:
CCl4
carbon tetrachloride


CO
carbon monoxide
(eliminate the o from the prefix)
P2Cl5
diphosphorus pentachloride (more than one of the first element)
N2O3
dinitrogen trioxide
ACADEMIC CHEMISTRY                                                  NAME:
PERIOD:                                                             DATE:

Si3N4
trisilicon tetranitride


H2O
dihydrogen monoxide




Steps for writing covalent formulas from the name.


    1. the elements appear in the same order as in the name
    2. the prefix indicates the subscript in the chemical formula




Examples:


Boron trifluoride           BF3


Dinitrogen monoxide         N2O


Dinitrogen tetroxide        N2O4
ACADEMIC CHEMISTRY                                                                                      NAME:
PERIOD:                                                                                                 DATE:
                                           ENERGY and STABILITY of COVALENT BONDS


 Most atoms have relatively LOW STABILITY and HIGH POTENTIAL ENERGY.


 When a compound forms, the atoms become MORE STABLE and the potential energy is at a MINIMUM.




Example:

The potential energy curve for hydrogen.



 When the nuclei are farthest apart, the potential energy is ZERO.


 As they get closer, the energy DECREASES.


 When the potential energy is LOWEST (at -436 kJ//mole), the atoms bond.


 The distance between the two nuclei at this lowest energy is called the BOND LENGTH and is 75 pm for the H2 molecule.


 When the REPULSION of the two atoms perfectly balance the ATTRACTIVE FORCES between the two nuclei, a COVALENT BOND
  forms.



 Since the potential energy DECREASED, energy has been RELEASED. For H 2, the energy released is -436 kj/mol.
ACADEMIC CHEMISTRY                                                                                    NAME:
PERIOD:                                                                                               DATE:




 The energy released when forming the bond is the exact amount of energy that would be needed to break the bond. Energy needed
  = +436 kj/mol



 The energy required to break a bond is known as the BOND ENERGY



                                                        BOND ENERGY
                                                            and
                                                        BOND LENGTH

                        Bond energy           Bond length (pm)       Electronegativity Difference
                          (kj/ mol)
      HF                    570                      92                          1.8
      CF                    552                     138
      OO                    498                     121
      HH                    436                      75
      HCl                   432                     127                          1.0
      CCl                   397                     177
      HBr                   366                     141                          0.8
      HI                    299                     161                          0.5


 BOND ENERGY (STRENGTH) is inversely related to the BOND LENGTH. The higher the bond energy, the SHORTER the bond.
ACADEMIC CHEMISTRY                                                                                     NAME:
PERIOD:                                                                                                 DATE:
 Bond length is actually an AVERAGE DISTANCE between the two nuclei since the distance is constantly changing due to the bond
  being able to vibrate and bend.



BOND STRENGTH and POLARITY


 The HIGHER the electronegativity difference, the STRONGER the bond.

          o Ex. HF has a much stronger bond than HI.

          o HF is nearly IONIC and HI is almost PURE COVALENT.




                                                    LEWIS DOT STRUCTURES

 Named after Gilbert Newton Lewis, who in 1920 came up with a way to represent valence electrons in an atom using dots.

SINGLE ATOM LEWIS STRUCTURES

 Only show the valence electrons around the element symbol.
 A maximum of 2 electrons per side of the element symbol.
 To find valence electrons, use the group number from the periodic table.

Ex.
ACADEMIC CHEMISTRY                                                                                           NAME:
PERIOD:                                                                                                      DATE:
LEWIS DOT STRUCTURES for COVALENTLY BONDED MOLECULES

 Recall the octet rule – most atoms need 8 valance electrons to be satisfied.
 Exceptions to the octet rule:
          o H and He
          o B
          o Elements beyond period 3 can EXCEED the octet rule. This can happen since these atoms have d-orbitals available to
              store the extra electrons.

                Ex.


Steps to completing a Lewis dot structure:
   1. Use a PENCIL. You may have to move electrons around and this becomes frustrating and messy if you use a pen.

      2. Count the total number of valence electrons available.
           a. For ions, a negative charge increases the number of electrons and a positive charge decreases the number of electrons.

Ex.


      3. Try to find a central atom and place all other atoms around it. Use symmetry whenever possible. Other times, the order in which
         the formula is written will allow you to determine which atoms should be connected.




      4. Use dots to represent the electrons shared between atoms. 2 electrons represent a BONDING PAIR or a single bond. Use the
         available valence electrons to make bonds between the atoms. These electrons would actually exist in molecular orbitals.
ACADEMIC CHEMISTRY                                                                                           NAME:
PERIOD:                                                                                                       DATE:
  5. Fill in the leftover electrons as LONE PAIRS around the atoms to achieve an octet (remember the exceptions). Electrons must
     always be paired. Lone pair electrons do not participate in the bonds, they are held in the atomic orbitals of the atoms involved.




   6. Count the electrons and make sure all atoms have been satisfied. You may need to make some lone pair electrons into bonding
      electrons (in double or triple bonds) to satisfy some of the atoms involved.




   7. Any extra electrons should be placed as lone pairs on the central atom. (This will happen if the central atom is one that can
      exceed the octet rule.)




   8. Polyatomic ions should have their structure placed in a bracket with the overall charge placed outside of the bracket.




RESONANCE

      when 2 or more equivalent Lewis structures exist for a molecule or compound

      all structures should be represented with a double arrow between
ACADEMIC CHEMISTRY                                                                                                         NAME:
PERIOD:                                                                                                                    DATE:
     the actual bond strength is an average of all the bonds.

     Examples: SO2 NO2

  (see back of notes)

MOLECULARGEOMETRY: ARRANGEMENT and SHAPE

     The three dimensional shape of molecules can be predicted using the VALENCE SHELL ELECTRON PAIR REPULSION THEORY.

         o   maximizes space between electron pairs to predict shape.

         o   VSEPR

     Arrangement of electrons – based on the number of regions of electron density around the central atom.

     Shape of the molecule – can be predicted by counting the number of bonding pair electrons and lone pair electrons.

     Treat double and triple bonds as single electron pairs when determining VSEPR shape.

      once the shape is determined, certain bond ANGLES can be predicted. Bond angles are altered by the present of LONE PAIR ELECTRONS.

      When present, the bond angles are compressed due to the REPULSION of the.

                 o      lone pairs require more space since they are only controlled by one nucleus.

     The arrangement of electrons and the shape of the molecule with be the same when there are NO LONE PAIR ELECTRONS.
ACADEMIC CHEMISTRY                                                   NAME:
PERIOD:                                                              DATE:
VSEPR and MOLECULAR GEOMETRY



BONDING PAIRS   LONE PAIR      TOTAL NUMBER   ARRANGEMENT   SHAPE   SKETCH and
of ELECTRONS    ELECTRONS      OF ELECTRON                          BOND ANGLES
                               PAIRS
1               0              1




2               0              2




3               0              3




2               1              3
ACADEMIC CHEMISTRY       NAME:
PERIOD:                  DATE:
4              0     4




3             1      4




2             2      4




5             0      5
ACADEMIC CHEMISTRY       NAME:
PERIOD:                  DATE:
4              1     5




3             2      5




2             3      5
ACADEMIC CHEMISTRY                                                NAME:
PERIOD:                                                            DATE:
BONDING PAIRS LONE PAIR    TOTAL NUMBER   ARRANGEMENT   SHAPE   SKETCH
of ELECTRONS   ELECTRONS   OF ELECTRON                          and
                           PAIRS                                BOND ANGLES
6             0            6




5             1            6




4             2            6
ACADEMIC CHEMISTRY                                                             NAME:
PERIOD:                                                                        DATE:
3              3                          6




2                    4                    6




HYBRIDIZATION
In order for bonds to have equivalent energy, mixing of orbitals must occur.

Ex. Methane (CH4) predicts 4 equivalent bonds. Which orbitals of C are used?



mix the orbitals used into a new “hybrid orbital” with a new name.
ACADEMIC CHEMISTRY                                                                                    NAME:
PERIOD:                                                                                               DATE:



Total number of e pairs around a     Arrangement                       Hybridization
         central atom
               2                         Linear                             sp
               3                    Trigonal planar                         sp2
               4                     Tetrahedral                            sp3
               5                        Trigonal                           sp3d
                                     bipyramidal
               6                      octahedral                           sp3d2



                                               POLARITY and DIPOLE MOMENT

POLAR BONDS:

 Result from high ELECTRONEGATIVITY DIFFERENCES between atoms of a molecule.


 Causes PARTIAL POSITIVE and PARTIAL NEGATIVE ends of the molecule (called DIPOLES).

 SHAPE AFFECTS POLARITY:
 DIPOLE MOMENT – overall direction of electron “pull” within a molecule. Show using a molecular model.
 Sometimes dipoles will CANCEL each other, and the result will be a molecule with NO NET DIPOLE due to the shape.

Molecule                Sketch                   Direction of dipole or no net dipole
HBr


BeCl2
ACADEMIC CHEMISTRY                                                                                        NAME:
PERIOD:                                                                                                   DATE:
BBr3


SeCl2


CO2



H2O




POLARITY AFFECTS PROPERTIES:


 For example:
         o CO2 is non-polar, so the ATTRACTIVE FORCE between CO2 atoms is VERY WEAK.
         o Only London Disperson Forces exist. Weakest intermolecular force (IMF).
         o This results in a lower MELTING POINT and BOILING POINT.


 H2O is POLAR, the molecules interact with each other and attractive forces are greater.


 This results in a higher MELTING POINT and BOILING POINT.


 A very strong INTERMOLECULAR force (called a dipole-dipole interaction) exists between water molecules due to its polarity.

       This force is called HYDROGEN BONDING

                 It is not the bond within water….it is a bond between water molecules.

				
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