CHM 101 Lewis Structure Lab Report

Lewis Structure and Molecular GeometryObjectives
1. Write Lewis electron dot formulas for molecular compounds.
2. Draw and/or name the VSEPR molecular geometry representation of a given molecular
compounds, including its angles and polarity.
3. Predict whether nonionic molecules will be polar or nonpolar. Identify the type of polar bond by
using the electronegativity difference.
Lewis Structures
Lewis electron dot structures are a type of symbol used to represent only the valence
electrons of an element or ion in a 2 dimensional (2-D) plane (a flat surface). The symbol if the
element represents the core electrons, those electrons not involved in bonding. Valence electrons
are symbolized by dots – • – one for each valence electron; these dots may be singular or in pairs.
Molecular Compounds and Polyatomic Ions
These Lewis structures always involve covalent bonding – the sharing of a pair of
valence electrons between two atoms. The goal in this type of bonding is to achieve eight
valence electrons around each atom – except for hydrogen which only needs two. This is an
application of the Octet Rule. The elements involved include the nonmetals and metalloids and,
in some polyatomic ions, a few metals like manganese and chromium.
Every valence electron is placed either as one-half of a bond pair or as lone pairs around
the element’s symbol. Bonding between atoms can be symbolized by one or more pairs of dots
or by bars – one bar per pair. E.G., H∶Cl
or H— Cl
Each element can form certain numbers and types of bonds. Carbon has the distinction of
forming 4 different bond types: 4 single bonds; 2 single bonds plus 1 double bond; 2 double
bonds; or 1 triple bond and 1 single bond. This is because carbon has 4 valence electrons.
However, the other members of Group 4A, Si and Ge, cannot form triple bonds. Only 1 single
bond occurs with any halogen (F, Cl, Br or I) and with hydrogen. The number and type of
possible bonds nonmetals and metalloids can form are given in the following table.
LERNER
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Electronegativity
Electronegativity values are based on the 1st Ionization Potential of an element and gives
a relative measure of how strongly an atom wants to retain its bonding electrons, and, therefore,
how attractive it becomes to other atoms with available bonding electrons. Fluorine (9F ) with a
value of 4.0 is the most electronegative element. Put a fluorine atom with any other element that
has available bonding electrons and it will attract those electrons to complete its outer shell octet.
It’s practical opposite, or least electronegative element, is cesium (55Cs).
Why does this matter? For some molecular compounds and polyatomic ions, it sets up
dipoles – temporary small electrical charges symbolized by δ+ (for the electron poor element)
and δ- (for the electron rich element) – at either end of a bond. This will contribute to the overall
molecular geometry, and, therefore, reactivity for a compound or polyatomic ion.
There are 2 subtypes of covalent bonds: polar covalent and nonpolar covalent. In
addition, there is also the ionic bond. What type of bond two elements may form can be
predicted from their electronegativity difference value, ΔEneg, according to the following rule:
LERNER
Bond Type
ΔEneg value
Ionic
ΔEneg ≥ 1.9
Polar covalent
0.5< ΔEneg < 1.9 Nonpolar covalent 0.5 ≤ ΔEneg Page 2 E.G. For Na @ 0.9 and Cl @ 3.0, ΔEneg = 2.1 and the bond is ionic. For C @ 2.5 and H @ 2.1, ΔEneg = 0.4 and the bond is nonpolar covalent. For H @ 2.1 and O @ 3.5, ΔEneg = 1.4 and the bond is polar covalent. Polar bonds are depicted as a signed arrow, either < + or + >, with the point
showing the more electronegative element. E.G. For H and O, this would be H+
>O.
Molecular Geometry/Shape
An electron pair always produces an electromagnetic field that can repel another electron
pair. This produces the 3-dimensional (3-D) – or real – structure of the molecule or polyatomic
ion. The Valence Shell Electron Pair Repulsion, VSEPR, Theory is used to model the 3-D
geometry of these structures. Using VSEPR, you get 3 basic geometric patterns as given by the
following table:
Total #
# of
# of
Geometry Name
of
Approximate
Bonding
Lone Pairs
Shape
Examples
Groups of
e-
Bond Angle
Directions
(# of E)
(VSEPR class)
2
180o
0
linear
BeH2,
(AB2)
CO2
(# of B)
2
trigonal planar
3
120o
3
0
(AB3)
BF3, NO3
2
1
bent
SO2
–
(AB2E)
4
109.5o
4
0
tetrahedral
CH4
(AB4)
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107.3o
3
1
trigonal
pyramidal
NH3
(AB3E)
Derivative of
tetrahedral
104.5o
2
2
bent
H2O
(AB2E2)
Derivative of
tetrahedral
***Pay attention to the wedge (solid triangle) and dashed appearance in how we draw the shape.
Table 2: Drawing out the structures.
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Pictorially, the corresponding shapes, angles, and some examples angles are:
Linear (< = 180°) Trigonal planar (< = 120°) Tetrahedral (< = 109.5°) LERNER Page 5 Trigonal pyramidal (< = 107.3°) Bent or V-shaped (< = ~104.5°) There are more shapes than these 5 patterns. Here is sulfur hexafluoride which assumes an octahedral shape. LERNER Page 6 In addition, adding or removing atoms and/or forming double or triple bonds can also affect the 3-D geometry. E.g., shown below are the geometries for ethane, a single-bonded hydrocarbon, and ethene, a double-bonded hydrocarbon; note the differences when 2 H atoms are replaced by a second bond between the carbons. Prelaboratory Assignment A. In your own words, define the following terms: 1. Valence electron 2. Lone pair electron 3. Polarity 4. Lewis structure 5. VSEPR B. Complete the table. Determine the: NH3 H2CCl2 Central atom Total # valence electrons # of bonds formed # of lone pairs LERNER Page 7 Procedures (Your instructor will demonstrate these procedures; use the appropriate line in the Data Table to record this.) Drawing a Lewis Structure 1. Determine the central atom or atoms. 2. Find the total number of valence electrons. 3. Place the other atoms around the central atom(s). Remember – H and halogens (F, Cl, Br, I) are allowed only one bond so they have to be on the outside. 4. Join pairs of atoms with a bar – this is equal to 2 electrons. 5. Make sure each atom – except for H – has 8 electrons around its symbol. If the bonds are not sufficient, add lone pairs to complete each atoms octet OR double or triple a bond. 6, Count the bond electrons and lone pairs; these must add to the total number of valence electrons found in step 2. Making a Ball-and-Stick Model 1. Using the kit provided, choose the appropriate color and numbers of balls for your elements and the appropriate bars for the bonds. 2. Construct your model. 3. When done with all the models, call your instructor over to verify that they are correct. Suggestion: Do one and confirm its correctness before doing another. Drawing the Molecular Geometry 1. From your Lewis structure, count the number of bonds (type does not matter) and the number of lone pairs. 2. Using the table on page 5, determine the 3-D shape. 3. Write the shape’s name and its angle in the table. 4. Sketch the shape as follows: a. Linear and Trigonal planar are shown in a plane (2-D). b. For other 3-D types, 2 bonds are in a plane and these use solid bars. For the out-of-plane bonds, use a dark wedge ( ► ) from the central atom to the atom in “front” of the plane and a dashed line ( ) from the central atom to the atom in “back” of the plane. Use a “balloon” ( ) to show lone pairs if present. LERNER Page 8 DATA TABLE Formula (# of val . e-s) Lewis Structure (with non-zero formal charge) Electron Geometry/ Molecular Geometry Bond Angle Molecular Geometry/3-D Shape (with molecular polarity) Instr. Demo CS2 H2S PH3 SO3 CH3Cl C2H6 LERNER Page 9 Formula (# of val . e-s) Lewis Structure (with non-zero formal charge) Electron Geometry/ Molecular Geometry Bond Angle Molecular Geometry/3-D Shape (with molecular polarity) CH4 OF2 C2H2Cl2 BF3 Polarity of Compounds 1. For the compounds in the Data Table, determine if the molecule is polar or nonpolar. 2. Polar molecules have either polar covalent bonds and asymmetry (lack of a “mirror image” bondwise) or at least one lone pair. If its polar, put dipoles in place on the shape drawing. 3. Complete the table by placing a YES in the appropriate column next to the formula. Molecular Formula CS2 H2S PH3 SO3 CH3Cl LERNER POLAR? NONPOLAR? Molecular Formula POLAR? NONPOLAR? C2H6 C2H4 OF2 C2H2Cl2 BF3 Page 10 PostLaboratory Assignment NAME_____________________________________________ SECTION #______________________ 1. There are 3 possible Lewis structures for C2H2Cl2. You drew one for your data table. Now draw the other two. 2. Draw one of the Lewis structures for the nitrate ion, NO3-1, and give its molecular shape and bond angle. You may also sketch the geometry. 3. Complete the table. Molecular Formula Polarity Lewis Structure Molecular Shape (NAME) Sketch of shape and angle CO2 HCN CH3F AsCl3 LERNER Page 11 4. For each of the following pairs, give the ΔEneg value and identify the bond type. Elements ΔEneg value Bond type Fe - O Cs - F Cl - Cl Ca - Cl As - O 5. Draw the Lewis structure for each: a. Calcium ion b. Phosphorous atom c. Selenide ion d. Magnesium sulfide, MgS e. Barium phosphide, Ba3P2 LERNER a. b. c. d. e. Page 12