Health Science Chemistry I CHEM-1180 Experiment No. 15 Molecular Models (Revised 05/22/2015)

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1 (Revised 05/22/2015) Introduction In the early 1900s, the chemist G. N. Lewis proposed that bonds between atoms consist of two electrons apiece and that most atoms are able to accommodate eight electrons in their valence shells (the octet rule). These simple assumptions proved very successful at predicting the shapes, polarity, stability and other properties of most molecular and ionic species. Remember, however, that many real molecules and ions do not obey the octet rule. A few compounds, such as BeCl2 and BCl3, classed as electron deficient species, have only four or six electrons around their central atoms. Many heavier atoms in species such as SCl6 and PF5 and I 3 accommodate five or six pairs of electrons, rather that four. Most structures in this experiment will obey the octet rule, however. This molecular model experiment will give you insight into the geometries of molecules and polyatomic ions. Two-dimensional Lewis structures on paper often do not adequately represent the three dimensional structure of many molecules. Three-dimensional molecular models will enable you to see the actual shapes of molecules predicted by the Valence Shell Electron Pair Repulsion (VSEPR) theory. Molecular models are especially useful in visualizing tetrahedral structures and those of molecules with non-bonded electron pairs such as ammonia and sulfur dioxide. Seeing the true geometry of molecules allows more insight into a molecule s polarity and its other properties. Your instructor will show examples such the apparent planar structure of the CH4 molecule drawn on paper and its true tetrahedral shape represented by a molecular model. Typical molecular model kits consist of a set of wooden or plastic spheres or polygons with holes for accepting pegs that represent bonds or electron pairs. Representations of second and third period atoms have four tetrahedrally oriented holes to represent the four available valance orbitals. Representations of the hydrogen atom have one hole since hydrogen forms only one bond. Some kits have pieces with five or six holes to represent atoms such as sulfur, arsenic and iodine that at times have expanded octets, that is, more than eight electrons in their valence shells. It is important to understand the limitations of ball-and-stick models. They only depict bond angles and molecular symmetries. Interatomic distances are greatly exaggerated by the sticks that hold the atoms together. Multiple bonds are also poorly depicted. Two identical springs or curved pegs representing a double bond ignore the fact that a double bond consists of one sigma and one pi bond having completely different shapes and orientations. Nevertheless, Lewis structures and their ball and stick models make surprisingly accurate predictions about the stability and polarity of most molecules. Drawing Lewis Structures To write the Lewis structure of any molecule or ion, first add up the number of valence electrons of all atoms in the formula. Correct the electron total for the ionic charge when necessary. Draw a skeleton structure for the species, joining the atoms with single bonds. The more symmetrical and compact a skeleton structure is, the more likely it is to represent the actual structure. Be careful here to use the number of bonds commonly formed by each kind of atom, i.e., one for hydrogen, four for carbon, two for oxygen (usually), three for nitrogen (usually,), one for fluorine, one for the other halogens (except when Cl, Br and I atoms are bonded to O), etc. For polyatomic oxoanions such as nitrate, phosphate,

2 chlorate and sulfate, attach the oxygen atoms only to the central atom, not to each other. (Oxygen rarely forms chains with itself.) Do not attempt to draw cyclic structures for inorganic species. Ring structures usually occur only in organic (carbon-containing) compounds. Refer to your text for help in answering the prestudy questions and in filling out the practice sheet on pages 4 and 5 before you come to lab. Depending on how much molecular structure theory was discussed in lecture, your instructor will give you additional examples of Lewis structures at the start of the lab period to enable you to construct models and learn from them. After you draw a skeleton structure, count the electrons utilized by the single bonds and distribute remaining electrons as unshared pairs and multiple bonds as needed to give each atom an octet. Example: Draw the Lewis structure of HCN, the hydrogen cyanide molecule, Hydrogen cyanide has = 10 valence electrons For the skeleton structure, draw the atoms in the order H, C, N with the C atom in the middle since C is required to form the most bonds. H C N Six electrons remain after drawing the two single bonds. No more electrons can fit on hydrogen. To satisfy the need for four bonds on the C atom, it is necessary to put two more bonds (electron pairs) between the C and N atoms, forming a triple bond. H C N To satisfy the octet rule for nitrogen, put the two remaining electrons on the N atom as a lone pair. H C N: This is the completed Lewis structure for HCN showing that it has a linear structure. As an exercise, verify that C and N each has an octet of electrons and that H has its maximum of two electrons. If necessary, review the rules for drawing Lewis structures in your textbook. Bond Polarity and Molecular Polarity A bond between two dissimilar atoms is, in principle, polar because any two different atoms are likely to have different electronegativity values. In a bond between two dissimilar atoms, the two electrons of the bond are attracted more strongly to the more electronegative atom, giving that atom a small partial negative charge, which is marked in a Lewis structure with a lower-case Greek delta followed by a minus sign, δ-. The other atom then assumes a small partial positive charge, indicated by the δ+. This slight charge separation gives rise to a charge dipole. Whether an entire molecule is polar, however, depends on its symmetry. If two or more polar bonds are oriented in opposing directions, their dipole contributions will cancel and the molecule will be nonpolar. For instance, the highly symmetrical tetrahedral CCl4 molecule is nonpolar, even though its C-Cl bonds are polar. The linear molecules BeF2 and CO2 are nonpolar while the bent molecules OF2 and SO2 are polar. Your instructor will illustrate some of these examples before the lab starts. During the lab, examine the symmetry of every model carefully before deciding on the allover polarity of the structure. 2

3 The concept of polarity only applies to neutral molecules. Ions, even if they seem to contain polar bonds and are unsymmetrical are not considered polar because the net charge of the ion overshadows any effect due to small charge differences within the ion. Resonance If a Lewis structure containing a double bond can be written with the double bond in more than one position without changing the positions of the atoms, then the molecule is said to exhibit resonance. This is indicated by showing the possible structures separated by a double arrow. Resonance does not imply that a double bond moves between its various drawn positions but rather that the bonds are intermediate between single and double bonds. This is borne out by X-ray crystallographic data that show, for instance, that the two sulfur-oxygen bonds in SO2 are equivalent. There is no evidence of a long single S O bond and a short double S=O bond in the molecule. Resonance is an important concept because its existence means that a structure has delocalized electrons rather than electrons confined to discrete bonds within the molecule. Such a structure is more stable than a similar structure without resonance. Experimental Fill out the practice sheets on pages 4 and 5 in pencil before you come to the laboratory so you will have enough time to construct and appreciate your molecular models. To begin, write the total number of valence electrons in the upper right corner of each box. Then draw the Lewis structure of each species in pencil. State the shape of each molecule or ion. For every structure containing one or more multiple bonds, state whether or not any resonance forms exist. Indicate the polarity of all heteroatomic bonds in neutral molecules using the δ+ and δ- notation. (Note: Carbon and hydrogen atoms have similar electronegativity values so you may assume that all C-H bonds are nonpolar.) Based on bond polarities and molecular symmetry, decide whether each neutral molecule is polar or not. Remember that the concept of bond and molecular polarity does not apply to ions. At the beginning of the laboratory period, your instructor will check your pencil entries and tell you if any need correction. During the laboratory period, make a model of each species and show it to your instructor for discussion and approval. To save time, note that some molecular models can be used to represent more than one Lewis species on pages 4 and 5. Report Obtain an assignment sheet from your instructor with four unknown formulas and enter the formulas on the report page 6. For each unknown, write the total number of valence electrons, draw the Lewis structure, and state the shape of the molecule or ion. For the uncharged molecular species, indicate the polarity of each bond and the allover polarity of the molecule. (Do not address the issue of polarity for any ionic species.) If any multiple bonds are present, tell whether resonance forms exist. Do not discuss resonance if a structure has only single bonds. Hand the report page to your instructor for grading before you leave the lab. Your report will be given back to you with your score out of 20 points. Hand in the complete laboratory report with abstract and cover page at the start of the next laboratory period. 3

4 Practice Sheet 1 Complete this two-page practice form in pencil before you come to the laboratory to give you time to construct molecular models and learn from them. Write the total number of valence electrons in the upper right of each box. Draw the Lewis structure of each species. If multiple bonds are present, tell whether resonance forms exist. Do not mention resonance if there are no double bonds. Indicate the polarity of all heteroatomic bonds in molecules and state whether the neutral molecules are polar or not. Do not address bond polarity in polyatomic ions. State the shape of each molecule or ion. H2O Cl2O H2O2 C2H6 C2H4 C2H2 SO2 SO3 SO3 2- SO4 2- PO4 3- NO3-4

5 Practice Sheet 2 POCl3 PCl3 CHCl3 BCl4 - CN - SeO2 CO2 CO3 2- N2H2 ClO - ClO4 - NNO 5

6 Report On this page, write the information requested for each of the four unknown species assigned to you. To complete your laboratory report, attach a cover page, abstract page and the assignment sheet containing your four unknown species to this page. Write the number of valence electrons in each species. Draw the correct Lewis structure of the species. State the shape of each species. In the case of a neutral molecule, indicate the polarity of each bond. For each neutral molecule, indicate whether the species is polar or nonpolar. If any structure has a double bond, indicate whether or not it has resonance forms. (5 points) (5 points) (5 points) (5 points) 6

7 General Chemistry I CHEM-1030 Laboratory Prestudy Write the total number of valence electrons and draw the Lewis structure of each atom, molecule and ion. (One-half point each) NH3 CH4 Br - NH4 + Cl2 CH3Cl HI NO2 SF2 HCN O2 CO2 O 2- SO2 OH - BF3 H2O PCl3 H3O + CCl4 7

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