Chemistry Unit 5 by pptfiles

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									  Chemistry Unit 2

Chemical Bonding
         IONIC                   COVALENT
High electroneg.            Low electroneg
difference                  difference
Transfer of electrons       Sharing of electrons
Crystalline solids @ r.t.   Mostly liquids & gases
                            @ r.t.
Good conductor of           Poor conductor of
electricity when melted     electricity
or dissolved in water
High solubility in water    Low solubility in water
High melting & boiling      Lower melting & boiling
point                       point
Atomic Structure REVIEW
l   An atom of sodium has
    11 protons                       11
    l   Atomic # 11                  Na
    l   Found in nucleus
l   It has 11 electrons          Sodium
    l   Same as # of protons    22.98977
    l   Atoms are neutral
    l   Found outside nucleus         e-
l   It has 12 neutrons                e- e-
                                      e- e-
    l   At. mass – At. #
                                e-    11 p+   e-
    l   Only protons and        e-    12 n    e-
        neutrons have mass
    l   Found in nucleus              e- e-
Lewis Dot Structures REVIEW
l   Lewis dot structures – atomic models that
    show only the valence electrons for an atom.
l   Only the valence electrons are involved in
    bonding, therefore Lewis Structures are good
    models to use when showing how atoms
    bond to form compounds.
Lewis Dot Structures REVIEW
l   Sodium
         e- e-
         e- e-

    e-   11 p+   e-
    e-   12 n    e-

         e- e-
Lewis Dot Structures REVIEW
l   Dots (electrons) should be added to the
    element’s symbol in this order.
                       1 2
                   8         3
                       Ne    6
                       7 4
Lewis Dot Structures REVIEW
The “A” groups tell you the number of electrons in the valence shell.

  Li                 Be                  B                   C

   N                   O                 F                    Ne
Why Do Atoms Bond?
l   To become more stable
    l   like the noble gases.

l   Octet Rule – atoms tend to gain, lose or
    share electrons in order to acquire a full shell
    of valence electrons. (usually 8 – except H
    and He who only need 2!!)
What Does This Picture Tell
Three Main Types of Bonds
l   Ionic Bond – Atoms transfer electrons to fill
    their valence shells, oppositely charged
    ions are formed, opposites attract.
    l   Occurs between a metal and a nonmetal
l   Covalent Bond – Atoms share electrons to
    fill their valence shells.
    l   Occurs between nonmetals
l   Metallic Bonds – Atoms share a “sea of
    l   Occurs between atoms of a metal
          K           F
In an IONIC bond, electrons are lost
or gained, resulting in the formation
of IONS in ionic compounds.
K   F
K   F
K   F
K   F
K   F
K   F
K       F
      CATION                    ANION

               K         F
•An atom that loses electrons gets a
positive charge and is a cation
(Think of the t in cation as a plus sign)

•An atom that gains electrons gets a
negative charge and is an anion
Ionic Bonding
l   Ion – a charged particle
    l   A neutral atom becomes an ion when it loses or gains an

        l   If an atom loses an electron, it becomes a (+) ion
            called a cation.
        l   If an atom gains an electron, it becomes a (-) ion
            called an anion.
Anionson& Cations most of our atoms
l Based  the fact that
 want 8, determine the ions that each of
 the following atoms will form, then say
 whether it is a cation or an anion

   Na          S          O          Mg
   Li          Ne         Si         F
   C           He         N          Cl

 Have me check your answers when you’re
 finished. If I’m busy, move on and I will
 check later.
Ionic Bonding
l   Example

              Na         Cl
         To become      To become
         more stable,   more stable,
         sodium         chlorine
         must loose     must gain
         one electron   one electron
Ionic Bonding
l   Example

              Na          Cl
      Sodium loses      Chlorine gains
      an electron       an electron and
      and becomes       becomes a Cl-1
      an Na+1 ion.      ion.
Opposites attract, and an ionic compound is formed…

Try Another Example

  Aluminum will         Bromine will
  become more           become more stable
  stable if it gets     if it receives one
  rid of three          electron.
Are both atoms more stable as a result of this
transfer? No, Al must donate two more… where?
Aluminum & Bromine


Now, each atom has a full valence shell… all are
more stable.
 Aluminum and Bromine
donated 3 e
-, so it                  Br
becomes       Al

                     Br        Each bromine
              Br               accepted 1 e-, so they
                               each become Br-1

      The compound that forms is AlBr3
Let’s Wrap it Up
l   Ionic bonds are held together by electrostatic
    forces. (Opposite charges attract.)
l   The result of an ionic bond is called an ionic
l   Ionic bonds form between a metal and a
    nonmetal atom due to large differences in
    electronegativity. (1.7 or greater)
    l   Electronegativity is the ability of an atom to pull shared
        electrons to itself. (Think of it like the strength of an
        atom in a tug-of-war for bonded electrons.)
l   The nonmetal’s EN is so much greater than the
    metal’s, that it removes the electrons, forming
    oppositely charged ions.
For Example: Na and O

    EN of Na = 0.9   EN of O = 3.5
Why does Sodium and Oxygen
form an ionic bond?

             3.5 EN of O
            - 0.9 EN of Na
             2.4 Difference in EN
l   Difference in electronegativity is 2.4(>1.7)
l   An ionic bond will form.
l   Oxygen has a greater electronegativity,
    and is able to yank electrons away from
How will two chlorine atoms react?

       Cl          Cl
      Cl            Cl
Each chlorine atom wants to gain one
electron to achieve an octet
       Cl            Cl
Neither atom will give up an electron

What’s the solution – what can they
do to achieve an octet?
Cl   Cl
Cl Cl
Cl Cl
Cl Cl
Cl Cl
        Cl Cl
     Cl Cl
The octet is achieved by
each atom sharing the
electron pair in the middle
Cl Cl
  Cl Cl
This is the bonding pair
    Cl Cl
It is a single bonding pair
    Cl Cl
It is called a SINGLE   BOND
      Cl Cl
Single bonds are abbreviated
         with a dash
Covalent Bonding

             O              O

         Each atom of Oxygen needs
         two more electrons to become
         more stable, so they will share
         two pairs of electrons.
  A diatomic molecule of oxygen is formed.

Try another example
                             Each atom of
Oxygen                H      hydrogen needs
needs two                    one more electron
electrons to   O             to become more
become more                  stable.
stable.               H
         All atoms become more stable
         (have full valence shells). A
         molecule of water is made.

Let’s Wrap it Up… Again!
l   Covalent bonds are held together by a
    mutual attraction for the shared electrons
l   Covalent bonding occurs when a sharing
    of electrons results in an overlap of
    valence orbitals. Each electron is attracted
    to the positive charge of the opposite
l   The result of a covalent bond is called a
 • Properties of metals
         – malleable                       -ductile
         – conduct heat                    -conduct electricity

     hy?                                      e   -                             e-

W             e-
                                   2         e-
                                                              2+           2+
         e-             e-
                                                                  -         e-
         2+             2+                   2+                       2+
                   e-                                     -                e-
e-                                                    e
• Why?
  – The valence electrons of metals are held
  – In metallic bonding metals don’t lose electrons.
  – Metal atoms release valence electrons in a sea
    of electrons shared by all metal atoms.
  – Valence electrons are free to move.
To determine if a bond will be
ionic or covalent, you can use
electronegativity (an atom’s desire
to get electrons)
Lets face it…some atoms are just
more attractive than others.
These atoms get more electrons
and are called more
electronegative . The charge of
an electron is negative, so if
there are more electrons near
you, you are more negative.
l   Every element has an electronegativity

l   The higher the #, the more the element wants
    the electron(s) and the harder it will pull the
    electron(s) towards itself

l   Think of it as the stronger player in a game of
    tug-a-war or someone trying to steal the blanket
Electronegativity & the
Periodic Table
l   The further apart the atoms are on the
    l   the greater the electronegativity difference
    l   the greater the chance for unequal sharing
 Electronegativity &
 Type of Bond
All bonds have a certain degree of
  sharing of electrons - even ionic!
l Ionic - so unequally shared that it
  is considered transferred
l Polar Covalent - unequal sharing
l Nonpolar Covalent - equal sharing
  Sharing is
sooooo unequal
that one atom
   takes the
  How rude.
   Sharing is a
  little unequal
 that one atom
pulls harder on
the electron(s)
Sharing is a
you, think of the north and south
poles. They are at opposite ends of
the Earth. So a POLAR covalent
molecule has opposite ends. One
end has more of a negative charge
because it is pulling the electrons
more towards itself, and one end
has more of a positive charge.
Using Electronegativity Difference
     to Figure Out Bond Type
l   You will be given a compound
l   Look up the electronegativity values of
    each element
l   Subtract the values
l   Take the absolute value of your answer
l   Use your answer to figure out the bond
                     Bond Type
        COVALENT                    IONIC

 nonpolar      polar
0      0.5                1.7                        3.3
                                  1.0 - 4.0 = -3.0
       0.9 - 3.0 = -2.1
                                Absolute Value = 3.0
    Absolute Value = 2.1
                 Bond Type
        COVALENT           IONIC

 nonpolar    polar
0      0.5           1.7           4.0

         CS2                 Cl2
    2.5 - 2.5 = 0
                 Bond Type
        COVALENT           IONIC

 nonpolar    polar
0      0.5           1.7           4.0

         H 2O                SO2

l   H & Cl
l   H&O
l   Na & F
l   Li & Cl
l   Na & Cl
l   C&O
For Example: N and O

 EN of N = 3.0   EN of O = 3.5
Why does Nitrogen and
Oxygen form a Covalent
Bond?     3.5 EN of Oxygen
               -    3.0 EN of Nitrogen
                   0.5 = difference in EN

Difference in EN is less than 1.7, therefore a covalent
bond will form.
Difference in EN is greater than 0, therefore the
covalent bond will be polar. (Unequal sharing of e-)
One Final Example

  If Chlorine bonds with Chlorine (a diatomic
  molecule Cl2), the difference in EN would be
  “0”, thus a nonpolar covalent bond will
  form. (Equal sharing of e-)
Lewis Dot & Ionic Compounds
 l   Draw the Lewis dot diagram for each
     element and show how the electrons move
     using arrows
 l   Show the ions that result
 l   The first one is done for you as an example

                                +     -
 l   Na      +     Cl   --> Na + Cl
 l   Li      +     Cl   -->
 l   Mg      +     O    -->
 l   K       +     F    -->
 l   K       +     O    -->

l   Lewis structures may also be used to show
    how atoms position themselves around one
    another and how the electrons are shared

l   For example instead of writing H2O for
    water, a Lewis structure can be drawn…
What do the dots represent?

What do you think the dashes represent?

Who is sharing electrons with who?
How to Draw Lewis Structures For Molecules
   1.   Find out the # of electrons needed
                  Every atom wants 8 (octet) except H only wants 2
   2.   Find out the # of valence electrons available
   3.   Subtract (electrons needed - electrons available)
   4.   Divide by 2 to determine the # of bonds you’ll have
   5.   Arrange the symbols as symmetrically as possible.
       l   The single atom goes in the middle with all other atoms
           attached to it.
   7. Connect the atoms with single bonds (----)
   8. Put in dots so that all atoms have eight electrons (except H
        which only needs 2)
   *Bonds count as two electrons!*
   9. Use double or triple bonds if possible
That’s a lot of steps…so to summarize…

   l   electrons Needed - electrons Available = electrons Shared

   l   Divide by two to get the number of bonds
       because each bond is made up of 2 electrons
How to Draw Lewis Structures
1. Do the math
                      ll o ctet
  2         +8            = 10 electrons needed

  1         +7            = 8 electrons available

                              2 e- shared
      len ons

                              2e- /2 =1 bond
How to Draw Lewis Structures

     2 e-         8 e-
How to Draw Lewis Structures

    A single bond      unshared
      (a pair of        pairs of
      electrons        electrons
Try Drawing the Structure for
H2 O

(2 x 2) + 8   = 12 electrons needed

(1 x 2) + 6   = 8 electrons available
                  4 e- shared
                  4e- /2 = 2 bonds
 pairs of   A single bond
electrons     (a pair of
Try Drawing the Structure for

8 + 2 + (8 x 3) = 34 electrons needed

4 + 1 + (7 x 3) = 26 electrons available
                  8 e- shared
                  8e- /2 = 4 bonds
Try Drawing the Structure for
Try Drawing the Structure for

8+   (2 x 3)   = 14 electrons needed

5+   (1 x 3)   = 8 electrons available
                 6 e- shared
                 6e- /2 = 3 bonds
Try Drawing the Structure for
How to Draw Lewis Structures

 8 + (2 x 2) + 8 = 20 electrons needed

 4 + (1 x 2) + 6 = 12 electrons available
                   8 e- shared
                   8e- /2 = 4 bonds
How to Draw Lewis Structures

(8 x 2) + (2 x 2) = 20 electrons needed

(4 x 2) + (1 x 2) = 10 electrons available
                    10 e - shared
                    10e- /2 = 5 bonds
5 bonds

          A triple bond: three pairs
          of electrons are shared by
          two atoms
 How to Draw Lewis Structures
Polyatomic Ions

   8   + (8 x 3)            = 32 needed

   6   + (6 x 3) = 24 + 2 = 26 available
                              6 e-
                              6e- /2 = 3 bonds
Polyatomic Ions

3 bonds
Polyatomic Ions

   8    + (2 x 4)       = 16 needed

   5    + (1 x 4) = 9 - 1 = 8 available
                           8 e-
                           8e- /2 = 4 bonds
Polyatomic Ions

4 bonds
Exceptions to the Octet Rule
A Central Atom With Less Than 8

  Beryllium only
    forms two single

  Boron only forms
    three single
Exceptions to the Octet Rule - A Central Atom
With More Than 8 (An expanded octet)
Sulfur can form six

Phosphorus can
  form five bonds.
    Additional Items to Remember
l   Polyatomic ions
    l   positive: subtract e- from available
    l   negative: add e- to the available
    l   add brackets & charge to finished structures
l   Too many bonds? Double and triple up.
l   Exceptions:
    l   B & Be - less than 8
    l   S & P - more than 8
Try to draw the Lewis structures for these
     ICl                 N2
     HBr                 F2
     CH2Cl2               CH4
     O2                   H2
Molecular, Empirical and
Structural Formulas

l   Molecular formula: Tells how many atoms are in a
    single molecule of a compound.
         ex. C6H12O6,(glucose), C3H6O3 (lactic acid)
l   Empirical formula: gives the simplest whole-number
    ratio of atoms of elements in a molecule. (Covalent
         ex. CH2O (for glucose and lactic acid)
l   Structural formula: specifies which atoms are bonded
    to each other in a molecule.
Molecular Geometry
l   Linear molecules: atoms are connected in a
    straight line.
    l   All molecules with only 2 atoms are linear.
    l   Many molecules with 3 atoms are also linear.
    l   The central atom will not have unbonded electron
        pairs. (All bonding domains.)
    l   Ex. O2, HCl, CO2
                                         O     C O
Molecular Geometry
l   Bent: connected atoms have a bent shape
    due to unbonded pairs of electrons.
    l   Unbonded electron pairs still exert a repulsion
        force (they still take up space) and contribute to
        the shape of the molecule.
    l   Ex. H2O


                                     H            H
Molecular Geometry
l   Tetrahedral: one atom bonded to four other
    l   The angle between any two bonds is 109.5o.
    l   Ex. CH4 (methane)

                                      C H
Molecular Polarity
l   Molecules can be polar or non-polar, even if
    they contain polar covalent bonds.
l   Polar molecules are sometimes called
l   If a molecule contains polar bonds that
    directly oppose each other in position, the
    molecule is non-polar. (The dipoles cancel.)
l   Perfectly symmetrical molecules tend to be
Polar vs. Nonpolar Molecules
l   Draw Lewis Structures for each of the
    following molecules to determine whether
    they are polar or nonpolar.
l   HCl
l   CCl4
l   NH3
l   BF3
l   CH3Cl
Polar vs. Nonpolar Molecules

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