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Homework answers / question archive / Unit 2- Activity 9- Lab - Shapes of Molecules Instructions LAB Shapes of Molecules and Polyatomic ions Introduction: One of our objectives in chemistry is to explain the properties of macroscopic samples in terms of the nature and behavior of the molecules that make up the sample

Unit 2- Activity 9- Lab - Shapes of Molecules Instructions LAB Shapes of Molecules and Polyatomic ions Introduction: One of our objectives in chemistry is to explain the properties of macroscopic samples in terms of the nature and behavior of the molecules that make up the sample

Chemistry

Unit 2- Activity 9- Lab - Shapes of Molecules Instructions

LAB

Shapes of Molecules and Polyatomic ions

Introduction:

One of our objectives in chemistry is to explain the properties of macroscopic samples in terms of the nature and behavior of the molecules that make up the sample. As you might expect, there are wide variations in the properties of individual molecular substances. Properties such as melting points and boiling points depend on attractions between molecules. Such properties are influenced by the polarity of the individual molecules. We would expect the properties of substances composed of polar covalent molecules to differ from those composed of non-polar molecules. Polarity, in turn, depends on the shape of the molecule and the relative electronegativities of the atoms involved. Shapes of molecules, however, also influence the properties of non-polar substances. Thus, structural isomers whose chemical compositions are identical, differ in structure (shape) and in properties.

The shape of a molecule depends on the angle between atoms bonded to a central atom. This angle, known as the bond angle, is the angle formed by two imaginary lines which join, respectively, the nuclei of each of two outlying atoms to the nucleus of the central atom. The distance between the nuclei of two bonded atoms is known as the bond length. The shape of the molecule is outlined by imaginary lines which join the nucleus of each of the outlying atoms to the nuclei of adjacent atoms.

One simplified theory of molecular geometry, called the "valence-shell, electron-pair repulsion theory," assumes that the shapes of molecules (and other species) are related to the positions of the valence-electron clouds on the central atom of the molecule. Each negative charge cloud, which contains two electrons of opposing spins, tends to repel all other charge clouds in the vicinity. To achieve a condition of minimum potential energy, it is necessary that the charge clouds be located so that they will be as far apart as possible. In this position, the electrostatic repulsion between the clouds is at a minimum.

The spatial orientation of the charge clouds depends on the number that are present and on their size. The number is equal to the total number of electron-pairs (bonded and unbonded) in the valence level of the central atom as indicated by the Lewis structure of the species. A double or triple bond involving, respectively, two or three electron-pairs is counted as or considered equivalent to a single electron cloud for purposes of determining spatial orientation. The relative size depends on whether the electron-pair is a bonded or a lone (unbonded) pair. We would expect clouds associated with bonded electrons to be rather localized between nuclei and to take up less space than those associated with unbonded electrons.

In this activity, you will be given a set of Styrofoam models and asked to identify structural features of a chemical species associated with the model. You should be aware, of course, that the shapes and bond angles you identify are idealized. You are to use the symmetry of the molecular models as a guide to the polarity of the species. Assume that the "atoms" attached to the central "atom" are identical.

Models representing shapes of molecules

Purpose:

  1. To relate the shapes and bond angles of chemical species to the number of bonded and unbonded electron-pairs in the valence level of the central atom in the species.
  2. To relate the polarity of molecules to the shape and symmetry of the species.

Materials: Collect the following on a blue tray.

Metal frames of the fundamental shapes (one of each), 6 plastic connectors, 6 Styrofoam balls, 4 toothpicks

Procedure: (Allow 3 minutes per model)

  1. Prepare a model of shape "A" using the appropriate metal frames , plastic connectors, and Styrofoam balls. (repeat for models  B through N as you complete Table I)  For the each model identify:
    1. the number of electron-pairs in the valence level of the central atom (assume no multiple bonds),
    2. the number of bonded atoms (do not count the central atom),
    3. the number of lone-pairs (unbonded electron-pairs),
    4. the bond angles; where the bond angle is distorted by non bonded pairs simply use the designation < (less than).
  2. When identifying the shape of the species represented by each of the models in your set. This list may help.

Possible Shapes

1.

linear

7.         seesaw

2.

trigonal planar

8.         T-shaped

3.

angular (or bent or v-shaped)

9.         octahedral

4.

tetrahedral

10. square planar

5.

trigonal pyramidal

11. square-based pyramidal

6.

trigonal bipyramidal

12. pentagonal bipyramidal

  1. Assuming all the models represent neutral molecules, indicate whether the species is polar or non polar. (identify non polar species due to resonance)
  2. a) Write the general formula for a chemical species with each of the shapes listed in the table. Use the symbol X to represent the central atom and the symbol Y to represent the atoms (all identical to each other) bonded to the central atom.

b) Write a formula showing all atoms and the number of lone pairs. Represent the lone pairs of electrons with the symbol E. For example, the formula for a species composed of two atoms of Y bonded to an X atom which still has two unbonded pairs of electrons in its valence level would be XY2E2.

  1. Obtain the name and formula of the "unknowns" from your instructor and use the styrofoam spheres and sticks in your set to construct a model of it. Record the information requested in Table II. Unlike the species represented by the models in your set, your unknown may involve a double bond. It is suggested that you first draw a Lewis (electron-dot) formula of your unknown. This should reveal the number of lone electron-pairs and thus, help you identify the shape of the species. Reminder: Sometimes the central atom does not follow the octet rule.

Observations:

  1. Draw the following table for model codes A to N.

 

Table I  Structural Characteristics of Chemical Species

Model

code

number of

atoms

bonded to

Central

Atom

number of

lone pairs

in valence

level of

bonded

central

atom

total  #

e pairs

in valence

level of

bonded

central

atom

bond

angle

shape

electric

nature

polar or

non-

polar

molecular

formula

formula

including

lone pairs

A

B

 

  1. Draw the following table for the unknowns

4-         -

a) NH3        b) ICl5 c) SiO4              d) NO3 e)TeO3

 

"Unknowns"

Table II

formula

valence e

pairs in

central

atom(s)

atoms

bonded to

central

atom(s)

lone pairs

on central

atom

Lewis

structure

number

and kind

of multiple

bonds

shape

NH

3

 

Additional Information:

In this activity you used the number of atoms bonded to a given atom and the number of lonepair valence electrons as a guide to the shape of a molecule. You found that the shapes of species in which there are lone pairs of valence electrons are related to the shapes of species in which all pairs of valence electrons are involved in bond formation. For example, the shapes of all the species in which the central atom has five pairs of valence electrons can, in theory, be related to the trigonal bipyramidal structure. Removal of one atom from a vertex of the equilateral triangle in the central plane of a bipyramid leaves a lone pair of electrons and yields a seesaw or teetertotter structure. It should be noted that the angles between the atoms in the central plane of the bipyramid are identical but differ from the angles between the atoms at the apices of the pyramid and the atoms in the central plane. It can be shown that minimum repulsion exists when lone pairs of electrons in the valence level of the central atom occupy positions in the central plane of the bipyramid.

For the purpose of this experiment we assumed that some of the models represented double bonded structures. It is possible to predict the presence of a double bond in the Lewis structure of the species by comparing the total number of valence electrons in the central atom with the sum obtained by adding the number of lone pairs to the number of atoms bonded to the central atom. For example, if there are four pairs of valence electrons, three atoms bonded to a central atom and no lone pairs, then the Lewis structure for the species must contain a double bond, be trigonal planar and non polar.

Another assumption we made in this activity was that identical atoms were bonded to a central atom. With this assumption, four atoms bonded to a central atom having no lone pairs gives rise to a symmetrical tetrahedral species which is non-polar. If the attached atoms were not identical, then the symmetry would be destroyed, the molecule would be polar, and its shape would resemble a distorted tetrahedron. You should also keep in mind that the geometry of polyatomic ions can be described in the same terms as those for neutral molecules. Polarity, however, is a term reserved for describing the electric nature of neutral molecules, not ions. Although partial charges are associated with molecular dipoles, ions carry a unit charge or a multiple of the unit

2-

charge. For example, sulfate ions  (SO4 )  although tetrahedral shaped with no lone electronpairs on the sulfur atom, carry a net charge of -2.

As we noted earlier, the degree of molecular polarity has an important bearing on the behavior of substances. Quantitatively, the degree of molecular polarity may be expressed in terms of a dipole moment. Dipole moments can be experimentally determined by means of electric measurements and then used to calculate the percent ionic character in a bond and to help determine the geometry of a molecule. For example, knowing that carbon dioxide has a zero dipole moment enables us to predict that the molecule is linear.

Discussion Questions:

  1. Which of these species are polar?  Draw a Lewis structure for each molecule. Use a table with the headings Formula, Lewis Structure, Polarity
    1. CH4       b) CH3Cl        c)         CCl4     d) C2H4             e) SO3

f)   SCl2     g) PH3 h)         BCl3     i)          SO2      j)          TeCl4

  1. Which of these involve double bonds in their Lewis structures?  Assume that all atoms follow the octet rule.  Draw Lewis structures for each species.

Use a table with the headings Formula, Lewis Structure, Multiple Bonds

2-         2-           2-         2-

    1. CO3      b) SO4 c)         SO3      d) CO2 e) SiO3              f)         SO3
  1. What is the shape of the following species?  Draw Lewis structures for each species.  Use a table with the headings Formula, Lewis Structure, Shape

3-         -

    1. CS2       b) BO3                c)        TiCl4   d) TeO2             e) IO3

2-           -

f)   SCl2     g) SiO3              h)         SF4       i)          AsF5     j)      AuCl4

-

k) XeF2      l)          MoF4   m) ICl4               n) IF5

  1. Measurements have shown that hydrogen peroxide (H2O2,) is a polar molecule.  Sketch a

Lewis structure for the hydrogen peroxide molecule.

Is H2O2  a symmetrical molecule? Explain.  Why is it a polar molecule?

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