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Covalent bonds

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									         HISTORY OF THE ATOM

460 BC   Democritus develops the idea of atoms



              He pounded up materials in his pestle and

              mortar until he had reduced them to smaller

              and smaller particles which he called



                      ATOMA
                  (Greek for indivisible)
       HISTORY OF THE ATOM

1808   John Dalton



            suggested that all matter was made up of

            tiny spheres that were able to bounce around

            with perfect elasticity and called them



                     ATOMS
       HISTORY OF THE ATOM

1898   Joseph John Thompson



           found that atoms could sometimes eject a far

           smaller negative particle which he called an




                 ELECTRON
         HISTORY OF THE ATOM
1904
Thompson develops the idea that an atom was made up of

electrons scattered unevenly within an elastic sphere surrounded

by a soup of positive charge to balance the electron's charge

                             like plums surrounded by pudding.



                                 PLUM PUDDING
                                       MODEL
       HISTORY OF THE ATOM

1910   Ernest Rutherford


            Oversaw Geiger and Marsden carrying out his
            famous experiment.

            They fired Helium nuclei at a piece of gold
            foil which was only a few atoms thick.

            They found that although most of them
            passed through. About 1 in 10,000 hit
      HISTORY OF THE ATOM

                                               gold foil
                   helium nuclei




                                             helium nuclei


They found that while most of the helium nuclei passed
through the foil, a small number were deflected and, to their
surprise, some helium nuclei bounced straight back.
         HISTORY OF THE ATOM


Rutherford’s new evidence allowed him to propose a more
detailed model with a central nucleus.


He suggested that the positive charge was all in a central
nucleus. With this holding the electrons in place by electrical
attraction



However, this was not the end of the story.
       HISTORY OF THE ATOM

1913   Niels Bohr


              studied under Rutherford at the Victoria
              University in Manchester.

              Bohr refined Rutherford's idea by adding
              that the electrons were in orbits. Rather
              like planets orbiting the sun. With each
              orbit only able to contain a set number of
              electrons.
Bohr’s Atom

                        electrons in orbits




              nucleus
           HELIUM ATOM
                                Shell
proton




                +
                    N
                         -
                    +
           -    N




electron                     neutron
             ATOMIC STRUCTURE

    Electrons are arranged in Energy Levels or

    Shells around the nucleus of an atom.


•     first shell         a maximum of 2 electrons

•     outer shell         a maximum of 8 electrons


Valence electrons          are the electrons in the last


                       shell or energy level of an atom.
Periodic Table Group                      Valence Electrons
Group 1 (I) (alkali metals)               1
Group 2 (II) (alkaline earth metals)      2

Groups 3-12 (transition metals)           #*

Group 13 (III) (boron group)              3
Group 14 (IV) (carbon group)              4
Group 15 (V) (nitrogen group)             5
Group 16 (VI) (chalcogens)                6
Group 17 (VII) (halogens)                 7
Group 18 (VIII) (noble gases)             8**


                         #* The general method for counting valence electrons is
                         generally not useful for transition metals.

                         ** Except for helium, which has only two valence electrons.
   Atoms with a complete shell of valence electrons
    tend to be chemically inert.

   Atoms with one or two valence electrons more than
    a closed shell are highly reactive because the extra
    electrons are easily removed to form positive ions.

      Atoms with one or two valence electrons less
    than a closed shell are also highly reactive because
    of a tendency either to gain the missing electrons
    and form negative ions, or to share electrons and
    form covalent bonds.
Octet Rule = atoms tend to gain, lose or share electrons so
as to have 8 electrons
                    H would like to Lose 1 electron
                    N would like to Gain 3 electrons
                    O would like to Gain 2 electrons
                IONIC BONDING




Electron from Na is transferred to Cl, this causes a charge
imbalance in each atom. The Na becomes (Na+) and the
Cl becomes (Cl-), charged particles or ions.
Covalent bonding-bond formed by the
sharing of electrons. Between nonmetallic
elements of similar electronegativity.
       Two Types of Covalent Bonds

Non-polar covalent bond – equal sharing of
electrons


Polar covalent bond – unequal sharing of
electrons
What does the word polar mean?
Symbols of atoms with dots to represent the valence-shell electrons

  Bonds in all the polyatomic ions and diatomic molecules are all
                           covalent bonds


Cl2
H2



                                       Ammonium NH4+
   Electron Configuration Model
Oxygen Atom          Oxygen Atom




         Oxygen Molecule (O2)
Polar Covalent Bonds: Unevenly
 matched, but willing to share.
- water is a polar molecule because oxygen is more
electronegative than hydrogen, and therefore electrons are
pulled closer to oxygen.
                                   Lewis Dots
Molecular
Formula              Structural
                     Formula




            Ball-and-Stick
                                  Space-Filling
   (Greek isos = "equal", méros = "part") are
    compounds with the same molecular formula
    but different structural formula

A simple example of isomerism is given by
propanol: it has the formula C3H8O (or C3H7OH) and
occurs as two isomers: propan-1-ol (n-propyl alcohol;
I) and propan-2-ol (isopropyl alcohol; II) (methyl-
ethyl-ether; III)
The difference in electronegativity of the
elements can be used to determine the type of
bond that is present.
1) What is electronegativity?
It is a measure of how strongly an atom
attracts the bonding electrons in a chemical
bond. The higher the electronegativity, the
stronger an atom's attraction for bonding
electrons.
Intermolecular Forces-attraction between
molecules
Pauling Scale of Electronegativities
Electronegativity   Type of      Example
difference          Bond
(approx.)

0.0-0.4             Non-polar    H-H (0.0)
                    covalent
0.4-1.0             Moderately   H-Cl (0.9)
                    polar
                    covalent

1.0-2.0             Very polar   H-F (1.9)
                    covalent
≥ 2.0               Ionic        Na+ Cl-(2.1)

								
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