Los Angeles Mission College Molecular Geometry Experiment 3

Experiment 3: Molecular GeometryObjective
The goal of this experiment is to use Lewis dot structures to determine the geometric shapes of a variety of
molecules and polyatomic ions. You will then use the molecular model kits to build and visualize the geometry
of each.
Background
Lewis electron dot structures are a useful device for keeping track of valence electrons for the representative
elements. In this notation, the nucleus and core electrons are represented by the atomic symbol and the valence
electrons are represented by dots around the symbol. Lewis structures can be drawn for individual atoms,
monatomic ions, molecules, and polyatomic ions.
Molecules and Polyatomic Ions: Lewis structures for molecules and polyatomic ions emphasize the principle that
atoms in covalently bonded groups achieve the noble gas configuration, ns2np6. Since all noble gases except
helium have eight valence electrons, this is often called the octet rule. Although many molecules and ions have
structures which support the octet rule, it is only a guideline. There are many exceptions. One major exception
is the hydrogen atom which can covalently bond with only one atom and share a total of two electrons to form
a noble gas configuration like helium. All of the examples in this experiment follow the octet rule except
hydrogen.
A Lewis structure for covalently bonded atoms is a two-dimensional model in which one pair of shared electrons
between two atoms is a single covalent bond represented by a short line; unshared or lone pairs of electrons
are shown as dots. Sometimes two pairs of electrons are shared between two atoms forming a double bond and
are represented by two short lines. It is even possible for two atoms to share three pairs of electrons forming a
triple bond, represented by three short lines.
For a polyatomic ion, the rules are the same except that the group of atoms is enclosed in brackets and the
overall charge of the ion is shown. For example:
The rules for writing Lewis structures for molecules and polyatomic ions will be provided in the procedure section
so you can use your Lewis structures to build three-dimensional models.
Chem 102 / Cuyamaca College
Chemistry 102
23 | P a g e
21
Molecular Model Building
The three-dimensional structure of a molecule is difficult to visualize from a two-dimensional Lewis structure.
Therefore, in this experiment, a ball-and-stick model kit (molecular “tinker toys”) is used to build models so the
common geometric patterns into which atoms are arranged can be seen. Each model that is constructed must
be checked by the instructor and described by its geometry and its bond angles on the report form.
Molecular Geometry
Atoms in a molecule or polyatomic ion are arranged into geometric patterns that allow their electron pairs to
get as far away from each other as possible (which minimize the repulsive forces between them). The theory
underlying this molecular model is known as the valence shell electron pair repulsion VSEPR) theory. All of the
geometric structures in this experiment fall into the following patterns:
1. Tetrahedral: four pairs of shared electrons (no pairs of lone (unshared) electrons) around a central atom.
2. Trigonal pyramidal: three pairs of shared electrons and one pair of unshared electrons around a central
atom.
3. Trigonal planar: three groups of shared electrons around a central atom. Two of these groups are single
bonds and one group is a double bond made up of two pairs of shared electrons. There are no unshared
electrons around the central atoms.
4. Bent: two groups of shared electrons (in single or double bonds) and one or two pairs of unshared
electrons around a central atom.
5. Linear: two groups of shared electrons, usually double bonds with two shared electron pairs between two
atoms, and no unshared electrons around a central atom. When there are only two atoms in a molecule or
ion, and there is no central atom (HBr, for example), the geometry is also linear. These patterns are
described more extensively in Section E, which follows.
Note: There are other electron arrangements and molecular geometries. Since they do not follow the
octet rule, they are not included in this experiment.
Bond Angles
Bond angles always refer to the angle formed between two end atoms with respect to the central atom.
Central atom
Bond angle
The size of the angle depends mainly on the repulsive forces of the electrons around the central atom. The
molecular model kits are designed so that these angles can be determined when sticks representing electron
pairs are inserted into pre-drilled holes.
Chem 102College
/ Cuyamaca College
22  Cuyamaca
24 | P a g e
1. Bond angles for atoms bonded to a central atom without unshared electrons
on the central atom.
a.
For four pairs of shared electrons around a central atom (tetrahedral
geometry) the angle between the bonds is approximately 109.5°.
b.
For three atoms bonded to a central atom, (trigonal planar) the angle is
120°. The shared electron pairs can be arranged in single or double bonds.
c. For two atoms bonded to a central atom (linear) the angle is 180°. The
shared electrons are usually arranged in double bonds.
d. Linear diatomic molecules or ions with no central atom do not have a
bond angle.
2. Bond angles for atoms bonded to a central atom with unshared electrons on the central atom.
When some of the valence electrons around a central atom are unshared, the VSEPR theory can be used to
predict changes in spatial arrangements. An unshared pair of electrons on the central atom has a strong influence
on the shape of the molecule. It reduces the angle of bonding pairs by squeezing them toward each other.
For example:
Bond Polarity
‘
Chem 102 / Cuyamaca College
Chemistry 102
••
••
Electrons shared by two atoms are influenced by the positive attractive forces of both atomic nuclei. For like
atoms, these forces are equal. For example, in diatomic molecules such as H2, or Cl2 the bonded
atoms have exactly the same electronegativity (affinity for the bonding electrons). Electronegativity H ─ H
••
values for most of the elements have been assigned as seen in table below. You can estimate ••
Cl ─ Cl
••
••
the polarity of a bond by calculating the difference between the electronegativity values for the
bonded atoms. The greater the difference the more polar the bond. If the electronegativity
difference between the atoms is:
• Between 0.0 – 0.3 the bond is considered nonpolar
• Between 0.4 – 1.7 the bond is considered polar
• Between 1.8 + the bond is considered ionic
25 | P a g e
23
In general, electronegativity increases as we move across a period and up a group on the periodic table. Identical
atoms with identical attractions for their shared electron pairs form non-polar covalent bonds. Unlike atoms
exert unequal attractions for their shared electrons and form polar covalent bonds.
Electronegativity is used to determine the direction of bond polarity which can be indicated in the Lewis
structure by showing an arrow ( ) pointed towards the more electronegative atom. For example, nitrogen
and hydrogen have electronegativity values of 3.0 and 2.1, respectively. The N ─ H bond is thus represented
as N H with the arrow directed toward the more electronegative nitrogen atom. Then, the Lewis structure
can be redrawn with arrows point towards the most electronegative atom. Example: NH3
EN of N-H = 3.0 – 2.1 = 0.9
bond is polar
Nitrogen is more electronegative thus arrow points towards N atom
Chem 102College
/ Cuyamaca College
24  Cuyamaca
26 | P a g e
Molecular Dipoles
When there are several polar covalent bonds within a molecule as in NH3, the polar effect of these bonds around a
central atom can be cancelled if they are arranged symmetrically as shown in CCl4 below. On the other hand, if
the arrangement of the polar bonds is asymmetrical, as in the bent water molecule, H2O, the resulting molecule
has a definite positive end and oppositely charged negative end, and the molecule is called a dipole. The symmetry,
or lack of symmetry of molecules and polyatomic ions, can generally be seen in the three-dimensional model.
Molecular Geometry Summary
e- Groups*
Type of e- Groups
Bonding
Lone Pair
Groups
Type of Geometry
Electron
Molecular
Geometry
Geometry
Bond Angle
Examples
2
2
0
Linear
Linear
180°
CO2
3
3
0
Trigonal
Planar
Trigonal
Planar
120°
CH2O
3
2
1
Trigonal
Planar
Bent
< 120° SO2 4 4 0 Tetrahedral Tetrahedral 109.5° CH4 4 3 1 Tetrahedral Trigonal Pyramidal < 109.5° NH3 4 2 2 Tetrahedral Bent < 109.5° H2O * e- Groups: Types of electrons (shared or non-shared) in central atom: 1 – Lone Pair, 2 – Single bond, 3 – Double Bond, 4 – Triple Bond. Chem 102 / Cuyamaca College Chemistry 102 27 | P a g e 25 Chemistry 102 Name______________________________________ Experimental Procedure Section______________________________________ For each of the following molecules or polyatomic ions, fill out columns A through G using the instructions provided in the procedure section. These instructions are summarized briefly below. A. Calculate the total # of valence electrons in each formula. B. Draw a Lewis structure for the molecule or ion which satisfies the rules provided in the procedure. C. From the Lewis structure determine the # of electron groups and type of groups when looking at the central atom/s D. From the Electron groups, determine the Molecular Geometry E. Determine the bond angle between the central atom and the atoms bonded to it. If there are only two atoms write “no central atom” in the space provided. F. Use the electronegativity table to determine the electronegativity of the bonded atoms. If the bonds are polar, indicate this with a modified arrow ( ) pointing to the more electronegative element. If there are two or more different atoms bonded to the central atom, include each bond. G. Use your model and your knowledge of the bond polarity to determine if the molecule as a whole is polar or nonpolar. Explain your reasoning. A Molecule or Polyatomic Ion # of Valence Electrons B Lewis Structure and 3-D models C D E # of Electron Groups & Type of Electron Molecular Bond Angles Groups at Central Geometry atom/s F Calculate Bond Polarity Is bond Polar? G Is Molecule Polar? Explain! Use Arrows to show direction of dipole. (Caution: Just because bond is polar, does not mean molecule is automatically polar) Ammonia NH3 Chloromethane CH3Cl Water H2O 27 Chem 102 / Cuyamaca College 29 | P a g e 28  Cuyamaca College Name______________________________________ Section______________________________________ A Molecule or Polyatomic Ion # of Valence Electrons B Lewis Structure and 3-D models C D E # of Electron Groups & Type of Electron Molecular Bond Angles Groups at Central Geometry atom/s F Calculate Bond Polarity. Is bond Polar? G Is Molecule Polar? Explain! Use Arrows to show direction of dipole. (Caution: Just because bond is polar, does not mean molecule is automatically polar) Sulfur Dioxide SO2 * Nitrogen Gas N2 Methanol CH3OH Carbon Dioxide CO2 Ethane C2 H 6 Ethene C2 H 4 * More than one possible Lewis Structure can be drawn. Chem 102 / Cuyamaca College 30 | P a g e Chemistry 102 Name______________________________________ Section______________________________________ A Molecule or Polyatomic Ion # of Valence Electrons B Lewis Structure and 3-D models C D E # of Electron Groups & Type of Electron Molecular Bond Angles Groups at Central Geometry atom/s F Calculate Bond Polarity. Is bond Polar? G Is Molecule Polar? Explain! Use Arrows to show direction of dipole. (Caution: Just because bond is polar, does not mean molecule is automatically polar) Azanone HNO Formaldehyde CH2O Acetonitrile HCN Carbonate ion CO32- This is an ion not a molecule, N/A Chlorotrifluoro methane CF3Cl Nitrogen trifluoride NF3 * More than one possible Lewis Structure can be drawn. 29 Chem 102 / Cuyamaca College 31 | P a g e