Molecular Models - Download as DOC by 4AtD2su

VIEWS: 10 PAGES: 6

									     CHEM 130                                              NAME: ____________________

                                        Molecular Models

                                                      Prelab

     In this lab, we will learn and practice predicting molecular structures from molecular formulas.

                                    The Periodic Table of the Elements

IA                                                                                                                   VIIIA
 1                                                                                                                     2
H     IIA                                                                            IIIA   IVA   VA    VIA VIIA     He
3      4                                                                               5     6     7     8   9        10
Li Be                                                                                B      C     N     O      F     Ne
11    12                                                                             13     14    15    16     17     18
Na Mg                                                                                Al     Si    P     S      Cl Ar
19    20     21      22    23     24     25    26     27     28     29     30        31     32    33    34     35     36
K Ca Sc Ti                 V     Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
37    38     39      40    41     42     43    44     45     46     47     48        49     50    51    52     53     54
Rb Sr        Y       Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te                                                     I    Xe
55    56     57-71   72    73     74     75    76     77     78     79     80        81     82    83    84     85     86
              La
Cs Ba       series   Hf Ta W Re Os Ir                        Pt Au Hg Tl Pb Bi Po At Rn
87    88    89-103
              Ac
Fr Ra       series
             57      58    59     60     61    62     63     64     65     66        67     68    69    70     71
             La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
             89      90    91     92     93    94     95     96     97     98        99     100   101   102    103
             Ac Th Pa             U Np Pu Am Cm Bk Cf Es Fm Md No Lr



     Part I: Valence Electrons
     To draw a Lewis Structure, we must know how many valence electrons each atom possesses. For main
     group atoms, we can get this information by looking at the periodic table.

     Rule 1: A main group atom has a number of valence electrons equal to the position of its main
             group column from the left of the periodic table (first main group column has one
             valence electron, second main group column has two valence electrons, etc.).

     Example: Lithium (Li) is in the first column, so it has one valence electron.

     The columns of the periodic table are often labeled at the top with a number. This number is called the
     group number, and is written in roman numerals.
Rule 2: A main group atom will have a number of valence electrons equal to its group number.

Determining the number of valence electrons for transition metal, lanthanide, and actinide atoms is much
more complicated because of the d and f orbitals. In this course, we will confine ourselves to main group
atoms, and we will ignore electrons in d and f orbital when counting valence electrons.

Exercise 1:       Complete the first column only of Table 1 by writing the number of valence
                  electrons in each of the atoms listed.

Part II: Electron Dot Pictures
We can write an abbreviated picture of a main group atom with its valence electrons by using the atomic
symbol surrounded by dots. An example for neon is shown below.


                                                    Ne
Neon has 8 valence electrons, so we surround the atomic symbol for neon (Ne) with 8 dots. Notice that the
dots are written in pairs at four positions around the symbol: right, left, bottom, and top.

Rule 3: When writing the electron dot picture of an atom, add one electron dot to all four
        positions before adding a second electron dot to any positions.

An example is shown below.

                                   Nitrogen (5 valence electrons)

        N                      N                      N                      N                   N
   1st electron           2nd electron            3rd electron           4th electron        5th electron

Atoms in groups I, II, and III will have empty positions; this will influence their bonding and chemistry.

Exercise 2:       Complete the second column only of Table 1 by writing the electron dot pictures
                  of the atoms listed.

Part III: Bonding Positions and Lone Pairs
Examining the electron dot pictures in Table 1, we can see that atoms can have pairs of electrons at some
positions and single electrons at other positions. The paired electrons are called lone pairs; in future
chemistry courses, especially organic chemistry, lone pairs have a strong effect on the chemistry of an atom
and its molecules. The single electrons are called unpaired electrons. Usually, an atom wants to have the
same number of bonds attached to it as it has unpaired electrons in its electron dot picture.


Exercise 3:       Complete columns 3, 4, and 5 of Table 1 by writing in the number of unpaired
                  electrons, lone pairs, and desired bonds, for each of the listed atoms.
Part IV: Lewis Structures
Step 1: Count the Electrons
1. Count the number of valence electrons in the molecule by adding up the number of valence
   electrons on all the atoms in the formula, and adjusting for the overall charge:

    Number of Valence Electrons = [sum of valence electrons on atoms] – [overall charge]

2. Calculate 6N + 2, where N = number of non-hydrogen atoms. 6N + 2 is the number of
   valence electrons required to have all single bonds.

3. Compare Number of Valence Electrons to 6N + 2 to determine if there are multiple bonds:

    If [6N + 2] – [Number of Valence Electrons] = 2, then 1 double bond;
    If [6N + 2] – [Number of Valence Electrons] = 4, then 2 double bonds or 1 triple bond.

    For each additional 2, there will be an additional multiple bond in the structure. (Note: in
    molecules that contain boron, beryllium, or aluminum, a 2 might indicate that these atoms
    have less than a full octet.)


Step 2: Draw a Skeleton for the Molecule
A skeleton has all the atoms connected together by single bonds, but lacks lone pairs and multiple bonds. It
can be difficult to guess the “correct” skeleton for a molecule, and often there are several possible
skeletons, all of which are correct. Try to get clues about the skeleton from the order that the atoms are
written in the molecular formula. The following set of guidelines may also be useful, but there are
exceptions to all of these guidelines so use them with care.

1. Look for symmetry in the skeleton.

2. Generally, the less electronegative atoms are closer to the center of the molecule (except
   hydrogen).

3. When there are several oxygens in the formula along with a different atom, put the different
   atom in the center and surround it with oxygens.

4. Avoid oxygen to oxygen bonds if possible. Obviously, this won’t be possible for O2 or O3.

5. Try to give each atom its desired number of bonds. This was determined for each atom in
   Part III.

6. Add hydrogens last. Each hydrogen should have only one bond.


Step 3: Complete the Skeleton by Adding Multiple Bonds and Lone Pairs

7. Add the number of multiple bonds calculated in Part I to the skeleton. Consider the desired
   number of bonds for each atom when deciding where to place multiple bonds.
8. Determine the remaining number of electrons:

    Remaining Electrons = [Number of Valence Electrons] – 2 x [number of bonds]

9. Complete the octets of the atoms in the structure by distributing the Remaining Electrons as
   lone pairs, starting with the most electronegative atoms first.




Part V: Molecular Structures
We will use the Lewis Structures to build models and look at the shapes of atoms in molecules. Most of
the molecules that we will deal with in this class follow the octet rule. Because their electrons are generally
paired, we will be building molecules with 4 pairs of electrons around each atom. To represent these
atoms, we will use balls with 4 equally spaced holes into which we will insert our "electrons". Noteworthy
exceptions are hydrogen and helium. Because they desire only 2 valence electrons, and these electrons will
be paired, these atoms will be represented by single-holed balls.

The pairs of electrons that form bonds between atoms will be represented by a short stick. A single stick
represents one bond, which is made of two shared electrons. To form a double bond, it will be necessary to
make two connections between two atoms. You will find, as you build your models that you cannot use the
sticks to make multiple bonds because they are inflexible. To form a double bond, use two pieces of
flexible spring to make the connections. Each spring represents one bond (two electrons), so two springs
are two bonds (a double bond, 4 electrons.) A triple bond is formed by three springs, and contains six
electrons.

A lone pair is a bond that only has an atom on one end. If we wish to represent lone pairs, use bonds but
only connect an atom on one end. This will be useful for Part VI below.




Part VI: VSEPR (Valence Shell Electron Pair Repulsion)
VSEPR theory is used to predict the shapes around atoms in a molecule. VSEPR assumes that pairs of
electrons repel each other. Because the bonds around atoms are pairs of electrons, they will repel each
other to be as far apart as possible, which will force the atom to have a particular shape. We can see how
many pairs of electrons are around a given atom in a molecule by looking at the Lewis Structure.

There are two different descriptions of the structure around an atom: the geometry and the shape. The
geometry refers to the arrangement of all of the pairs of electrons around an atom, including lone pairs.
The shape refers only to the arrangement of bonds around an atom, ignoring lone pairs. Sometimes the two
will be the same; often they will be different.
When a molecule has multiple bonds in its Lewis Structure, we count the multiple bond as a single pair of
electrons for VSEPR.

The geometries and shapes around atoms for a variety of different possible arrangements are given in the
following table:
 Total
           Bonding                   Electron         Electron      Bond     Molecular      Molecular
Electron             Lewis Diagram                                                                       Example
           Groups                    Geometry      Geometry Model   Angle     Shape        Shape Model
Groups




   2         2                         Linear                       180o       Linear                     BeF2




                                      Trigonal                                Trigonal
   3         3                                                      120o                                  BF3
                                       Planar                                  Planar




   4         4                       Tetrahedral                    109.5o   Tetrahedral                  CH4




                                                                              Trigonal
   4         3                       Tetrahedral                    109.5o                                NH3
                                                                              Pyramid




   4         2                       Tetrahedral                    109.5o      Bent                      H2O
               Table 1: Atomic Electron Dot Pictures and Bonding Properties

                            Number of                  Number of                Number of
                                        Electron Dot               Number of
             Atom            Valence                   Unpaired                  Desired
                                           Picture                 Lone Pairs
                            Electrons                  Electrons                 Bonds

Beryllium (Be)



Boron (B)



Carbon (C)



Nitrogen (N)



Oxygen (O)



Fluorine (F)



Aluminum (Al)



Silicon (Si)



Phosphorus (P)



Sulfur (S)



Chlorine (Cl)



Selenium (Se)

								
To top