# 5. Structure, Geometry, and Polarity of Molecules

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1 5. Structure, Geometry, and Polarity of Molecules What you will accomplish in this experiment This experiment will give you an opportunity to draw Lewis structures of covalent compounds, then use those structures to visualize the three-dimensional shapes of the molecules. You ll also learn how to determine the polarity of the covalent bonds within the molecules. You need to build these skills so that you ll ultimately be able to make predictions about the behavior of molecules, specifically their ability to interact with neighboring molecules via intermolecular attractive forces. These interactions will dictate the physical properties of these covalent compounds (their physical state and solubility). In summary, this experiment will help you to practice the following skills: Draw Lewis structures of simple molecules (containing one to three central s). Use the number of electron groups around each central in the structure to determine the threedimensional geometry (spatial arrangement) of the bonding s and lone pairs. Assess the polarity of the covalent bonds in the structure by comparing the electronegativities of the two s sharing electrons. Concepts you need to know to be prepared You ve learned that compounds are a chemical combination of elements, meaning that they re created when two or more elements chemically react with one another. The driving force for that reaction is that the elements are trying to achieve the stability of a noble gas electron configuration (for most elements, this means an s 2 p 6 octet of electrons. In Lab 4, you saw that when the reacting elements are a metal and a nonmetal (elements with an electronegativity difference greater than 2.0), the metal achieves its octet when it loses one or more electrons (to become a positively-charged ion, a cation ), and the nonmetal gains those electrons (to become a negatively-charged ion, an anion ). The chemical formula for the ionic compound that s formed describes the fixed proportion in which vast numbers of the oppositely-charged ions must combine in order to appropriately balance out the positive and negative charges. C. Graham Brittain Page 1 of 13 10/5/2010

2 Covalent Bonds and Lewis Structures When the reacting elements are nonmetals with a small electronegativity difference (less than 2.0), the s achieve their octet NOT by gaining or losing electrons, but by sharing electrons by forming covalent bonds. For most nonmetal s, this means sharing enough electrons to achieve the octet (the s 2 p 6 electron configuration of all noble gases except helium). Lewis structures are illustrations of the covalent bond connections between s. When drawn correctly, these structures will help you to envision the three-dimensional geometry of each bonding. The polarity of each covalent bond, considered with the geometries of the bonding s, will enable you to make predictions about the behavior of the molecule, specifically its ability to interact with other molecules via intermolecular attractive forces. Drawing Lewis Structures The s that need the most electrons to achieve a noble gas electron configuration (and thus form the most covalent bonds) should be used as central s in a Lewis structure. The s that need very few electrons to achieve a noble gas electron configuration (and thus form very few bonds) should be used as peripheral s in the structure. Examples of central s: Carbon has 4 valence electrons. It needs 4 more electrons to achieve an octet. It gets those 4 electrons by sharing: by forming 4 covalent bonds. CARBON S octet: 4 BONDS and NO LONE PAIR. Nitrogen has 5 valence electrons. It needs 3 more electrons to achieve an octet. It gets those 3 electrons by sharing: by forming 3 covalent bonds. NITROGEN S octet: 3 BONDS and 1 LONE PAIR. Oxygen has 6 valence electrons. It needs 2 more electrons to achieve an octet. It gets those 2 electrons by sharing: by forming 2 covalent bonds. OXYGEN S octet: 2 BONDS and 2 LONE PAIR. Examples of peripheral s: Fluorine has 7 valence electrons. It needs 1 more electron to achieve an octet. It gets that 1 electron by sharing: by forming 1 covalent bond. FLUORINE S octet (and that of the other halogens chlorine, bromine, and iodine): 1 BOND and 3 LONE PAIR. ydrogen has 1 valence electron. It needs 1 more electron to achieve a duet (like elium). It gets that 1 electron by sharing: by forming 1 covalent bond. YDROGEN S duet: 1 BOND. You should indicate the covalent bonds between the s with lines, and the lone pairs of electrons (unshared electrons) with dots. And you ll need to arrange the bonds and lone pairs so that each achieves its noble gas electron configuration (an octet, or a duet for hydrogen s). A single bond is one pair of electrons shared between two s, and is represented in the Lewis structure by a single line between the bonding s. A double bond is two pairs of electrons shared between two s, and is represented by drawing two lines between the bonding s. Double bonds are most common between s of carbon, nitrogen, oxygen, and sulfur. (In organic and biological molecules, there are many examples of carbon-carbon, carbon-nitrogen, and carbon-oxygen double bonds.) A triple bond is three pairs of electrons shared between two s, and is represented by drawing three lines between the bonding s. Triple bonds are most common between carbon and nitrogen s. Be sure to draw NO MORE and NO FEWER electrons than the TOTAL number of valence electrons that the s bring with them to the molecule. C. Graham Brittain Page 2 of 13 10/5/2010

3 Example Lewis Structure: Formaldehyde, C 2 O The TOTAL number of valence electrons the s bring with them to the molecule: 1 C with 4 valence electrons 4 2 s, each with 1 valence electron 2(1) 1 O with 6 valence electrons 6 Total available electrons 12 Consider the bonding patterns of our component s: Carbon is always a central because it forms 4 bonds. Oxygen tends to form 2 bonds and 2 lone pairs. ydrogen s form just 1 bond. Since carbon forms the most bonds, we ll place it at the center and attach the O and s to it. Since the C needs four bonds, and the O s needs two bonds, we ll place a double bond between them. 2 lone pairs of electrons are added to the oxygen to complete its octet. The total number of valence electrons the s brought into the molecule is 12. This Lewis structure accounts for ALL 12 of those electrons: 2 electrons in each of the 4 bonds (8 bonding electrons), plus 2 lone pairs (4 unshared/non-bonding electrons). O C Molecular Geometry The repulsive forces between bonding and non-bonding electrons determine the three-dimensional geometry (spatial arrangement) of the groups of electrons around a central. Because the negative charges repel one another, the electron groups arrange themselves so that they are as far apart from one other as possible. This idea is known as the Valence Shell Electron Pair Repulsion (VSEPR) Theory. Thus the geometry of the electrons around each central in a Lewis structure can be determined by: 1. Simply counting the groups of electrons that surround that, and then 2. Considering how these groups would arrange themselves so as to be as far apart as possible. It s essential to remember that a group of electrons is one of the following: A single bond, a double bond, a triple bond, or a lone pair of electrons For example, the Lewis structure of formaldehyde indicates 3 groups of electrons around the central carbon : 2 single bonds (between the carbon and each of the two hydrogen s), and 1 double bond (between the carbon and oxygen s). As shown in the list of potential Electron Group Geometries, in order for three groups of electrons to be as far apart as possible, they would all have to lie in the same plane and be directed toward the corners of a triangle. This electron group geometry is called trigonal planar. The bond angles in this geometry would be 120 o. TWO groups of electrons around a central : LINEAR TREE groups of electrons around a central : TRIGONAL PLANAR FOUR groups of electrons around a central : TETRAEDRAL C The table on the next page indicates the overall shape (or Molecular Geometry ) of a central depending on whether the electron groups around it are covalent bonds to other s or simply lone pairs of electrons. O C. Graham Brittain Page 3 of 13 10/5/2010

4 IF the electron groups are bonds, then s are present to mark the corners of each of the spatial arrangements: the two ends of the linear geometry, the three corners of the trigonal planar geometry, or the four corners of the tetrahedral geometry. But IF an electron group is a lone pair, then there is no visible to mark that corner of the geometry. Thus the Electron Group Geometry is determined by the number of groups of electrons and the arrangement of those groups around the central. But the Molecular Geometry depends on whether or not there is an to visibly mark each corner of the Electron Group Geometry. Remember that you can only see s; you cannot see lone pairs of electrons. # of Groups of Electrons Electron Group Geometry Number of Lone Pairs Molecular Geometry Approximate Bond Angles Example Compound 2 Linear 0 Linear 180 o carbon dioxide, CO 2 3 Trigonal Planar 0 Trigonal Planar 120 o formaldehyde, C 2 O 4 Tetrahedral 0 Tetrahedral o methane, C 4 4 Tetrahedral 1 Trigonal Pyramid 107 o ammonia, N 3 4 Tetrahedral 2 Angular (Bent) o water, 2 O C. Graham Brittain Page 4 of 13 10/5/2010

5 The molecule used as the final example on the preceding table is water, 2 O. The Lewis structure for water indicates that the oxygen is surrounded by four groups of electrons: two bonds and two lone pairs. These four groups will be arranged in the tetrahedral Electron Group Geometry. ydrogen s mark the positions at two of the corners of this tetrahedron, but the lone pairs of electrons at the other two positions make those corners invisible. Thus the shape ( Molecular Geometry ) of the water molecule is actually angular or bent. O Polarity of Molecules When the two s sharing electrons in a covalent bond have significantly different electronegativity values (a difference greater than 0.4, but less than 2.0), the electrons in the bond are NOT shared EQUALLY between the two s. The element that has the higher electronegativity has greater electron pulling power; that is, the is able to attract the electrons in the bond toward itself. The result is that the covalent bond is polarized. Or to put it more simply, we say the bond is polar. Since the bonding electrons spend more of their time in the region around the element with the higher electronegativity, that takes on a partial negative charge (represented by the symbol - ). As the bonding electrons spend very little time in the vicinity of the less electronegative element, that takes on a partial positive charge (represented by the symbol + ). So, similar to a battery, a polar covalent bond has TWO poles: one partial positive + ), and the other partial negative ( - ). Thus we can say that it s a dipole, and we can represent the polarized bonding electrons in an alternative way: by using a dipole moment vector. That is, we can draw a vector (simply an arrow) pointing in the direction that the electrons are being pulled: toward the with the higher electronegativity. To emphasize the fact that the at the other end of the bond has a partial positive charge, we make the tail of the arrow into a plus sign. Thus the very polar hydrogen fluoride bond (electronegativities: = 2.1, F = 4.0) could be represented either by labeling the two s with the partial charge symbols or by drawing a dipole moment vector: F F The geometry of a molecule is tremendously important when evaluating the polarity of each covalent bond around a central (in order to determine the type of intermolecular attractive forces that a molecule is capable of using). Quite often the three-dimensional geometry of the central s bonding and nonbonding electron groups can enhance the overall polarity of a cluster of s. For example, consider the water molecule, 2 O. C. Graham Brittain Page 5 of 13 10/5/2010

6 The hydrogen-oxygen single bonds in the angular 2 O molecule are polar (electronegativities: = 2.1, O = 3.5). Thus the electrons in each bond are drawn toward the oxygen. This can be illustrated using either the partial charge symbols or dipole moment vectors. O O But now consider the tetrahedral geometry of the four groups of electrons around the oxygen (2 bonds, 2 long pairs). This geometry greatly enhances the overall polarity of the angular (bent) water molecule. The dipole moment vectors are extremely useful in illustrating this additive effect of the two polar bonds. The partial negative charge of both dipoles is focused at the oxygen. Additionally, the oxygen s two lone pairs of electrons contribute significantly to its overall partial negative charge. The result is that the water molecule is extremely polar, with the partial negative charge centered at the oxygen, and a partial positive charge at each of the two hydrogen s. Occasionally, for very small molecules, the symmetry of the bonding electrons around the central can cause two or more dipole moment vectors to completely cancel one another, so that the molecule as a whole has no NET dipole. For example, consider the carbon dioxide (CO 2 ) molecule. The carbon-oxygen double bonds in the linear CO 2 molecule are polar (electronegativities: C = 2.5, O = 3.5). The electrons in each of the double bonds are drawn toward the oxygens, so both oxygen s have a partial negative charge. This can be illustrated using either the partial charge symbols or dipole moment vectors. The vectors are especially helpful in illustrating that, because of the linear geometry of the molecule, the bond dipoles are completely symmetric to one another. They are equal and opposite, so they cancel (the CENTERS of partial-negative and partial-positive charge are at the same place in the molecule). Thus the carbon dioxide molecule as a whole is nonpolar, even though it contains polar bonds. O C O O C O C. Graham Brittain Page 6 of 13 10/5/2010

7 Procedure that you will follow You and your lab partner will be supplied with a molecular modeling kit. The colors of the plastic balls in the kit correspond to s of different elements as follows: carbon s (black) hydrogen s (white) oxygen s (red) nitrogen s (blue) chlorine s (green) bromine s (orange) For each chemical formula specified on the report sheet, you will: 1) Draw the Lewis structure of the molecule. (Note that for one of the chemical formulas, TWO different Lewis structures are possible. These different structures that have the same chemical formula are said to be structural isomers of one another.) 2) Determine the Electron Group and Molecular Geometry of each central. 3) Verify that geometry by constructing a model of the molecule using the s (colored balls) and covalent bonds (plastic connectors) in the kit. You ll then: 4) Assess the polarity of the covalent bonds by comparing the electronegativities of the bonding s. Electronegativity values of common Representative Elements are provided in the table below: C. Graham Brittain Page 7 of 13 10/5/2010

8 5. Report Sheet: Structure, Geometry, and Polarity of Molecules Student Lab Partner Date Lab Performed Section # Lab Instructor Date Report Received I. ydrogen Cyanide, CN ow many valence electrons does this have? And how many electrons does this need to get an octet (duet for )? So how many covalent bonds will this form? And how many lone pairs will this have? What is the electronegativity value of this? a) TOTAL number of valence electrons in the CN molecule = C N b) Draw a Lewis structure for the CN molecule, including all covalent bonds and all lone pairs of electrons. Your structure should have no more and no fewer electrons than the TOTAL number above. c) ow many groups of electrons are around your central? (Remember that a group of electrons can be a: single bond, double bond, triple bond, or lone pair.) d) What is the electron group geometry of the central? After noting whether the electron groups are covalent bonds or lone pairs: e) What is the molecular geometry of the central? Now use your kit to construct a molecular model of the CN molecule to verify the molecular geometry. Compare the electronegativities of the two s participating in each covalent bond. f) Are any of the bonds POLAR? If YES, clearly indicate the partial negative and partial positive s in your Lewis structure above. g) Do the bond polarities reinforce one another (to make a net dipole), or cancel one another (to make no net dipole)? Explain. C. Graham Brittain Page 8 of 13 10/5/2010

9 II. Difluoromethane, C 2 F 2 ow many valence electrons does this have? Each C Each F And how many electrons does this need to get an octet (duet for )? So how many covalent bonds will this form? And how many lone pairs will this have? What is the electronegativity value of this? a) TOTAL number of valence electrons in the C 2 F 2 molecule = b) Draw a Lewis structure for the C 2 F 2 molecule, including all covalent bonds and all lone pairs of electrons. Your structure should have no more and no fewer electrons than the TOTAL number above. c) ow many groups of electrons are around your central? (Remember that a group of electrons can be a: single bond, double bond, triple bond, or lone pair.) d) What is the electron group geometry of the central? After noting whether the electron groups are covalent bonds or lone pairs: e) What is the molecular geometry of the central? Now use your kit to construct a molecular model of the C 2 F 2 molecule to verify the molecular geometry. Compare the electronegativities of the two s participating in each covalent bond. g) Are any of the bonds POLAR? If YES, clearly indicate the partial negative and partial positive s in your Lewis structure above. h) Do the bond polarities reinforce one another (to make a net dipole), or cancel one another (to make no net dipole)? Explain. C. Graham Brittain Page 9 of 13 10/5/2010

10 II. Dichloroethylene, C 2 2 Cl 2 ow many valence electrons does this have? Each C Each Cl And how many electrons does this need to get an octet (duet for )? So how many covalent bonds will this form? And how many lone pairs will this have? What is the electronegativity value of this? a) TOTAL number of valence electrons in the C 2 2 Cl 2 molecule = b) Draw a Lewis structure for the C 2 2 Cl 2 molecule, including all covalent bonds and all lone pairs. Recall that the video your lab instructor played for you showed two DIFFERENT Lewis structures for this chemical formula. The Lewis structure you draw ERE should be DIFFERENT from those shown in the video. (And remember that your structure should have no more and no fewer electrons than the TOTAL number above.) In order to answer the questions below, number the two carbon s in your structure (1 and 2). c) ow many groups of electrons are around carbon #1? d) What is the electron group geometry of carbon #1? e) What is the molecular geometry of carbon #1? f) ow many groups of electrons are around carbon #2? g) What is the electron group geometry of carbon #2? h) What is the molecular geometry of carbon #2? Now use your kit to construct a molecular model of the C 2 2 Cl 2 molecule to verify the molecular geometries. Compare the electronegativities of the two s participating in each covalent bond. i) Are any of the bonds POLAR? If YES, clearly indicate the partial negative and partial positive s in your Lewis structure above. j) Do the bond polarities reinforce one another (to make a net dipole), or cancel one another (to make no net dipole)? Explain. C. Graham Brittain Page 10 of 13 10/5/2010

11 IV. Methyl amine, C 5 N ow many valence electrons does this have? And how many electrons does this need to get an octet (duet for )? So how many covalent bonds will this form? And how many lone pairs will this have? What is the electronegativity value of this? Each C N a) TOTAL number of valence electrons in the C 5 N molecule = b) Draw a Lewis structure for the C 5 N molecule, including all covalent bonds and all lone pairs of electrons. Your structure should have no more and no fewer electrons than the TOTAL number above. c) ow many groups of electrons are around the CARBON? d) What is the electron group geometry of the CARBON? e) What is the molecular geometry of the CARBON? f) ow many groups of electrons are around the NITROGEN? g) What is the electron group geometry of the NITROGEN? h) What is the molecular geometry of the NITROGEN? Now use your kit to construct a molecular model of the C 5 N molecule to verify the molecular geometries. Compare the electronegativities of the two s participating in each covalent bond. i) Are any of the bonds POLAR? If YES, clearly indicate the partial negative and partial positive s in your Lewis structure above. j) Do the bond polarities reinforce one another (to make a net dipole), or cancel one another (to make no net dipole)? Explain. C. Graham Brittain Page 11 of 13 10/5/2010

12 V & VI. Ethanol and Dimethyl Ether, C 2 6 O ow many valence electrons does this have? Each Each C O And how many electrons does this need to get an octet (duet for )? So how many covalent bonds will this form? And how many lone pairs will this have? What is the electronegativity value of this? a) TOTAL number of valence electrons in the C 2 6 O molecule = b) Draw a Lewis structure for the C 2 6 O molecule, including all covalent bonds and all lone pairs of electrons. Your structure should have no more and no fewer electrons than the TOTAL number above. In order to answer the questions below, number the two carbon s in your structure (1 and 2). c) ow many groups of electrons are around CARBON #1? d) What is the electron group geometry of CARBON #1? e) What is the molecular geometry of CARBON #1? f) ow many groups of electrons are around CARBON #2? g) What is the electron group geometry of CARBON #2? h) What is the molecular geometry of CARBON #2? i) ow many groups of electrons are around the OXYGEN? j) What is the electron group geometry of the OXYGEN? k) What is the molecular geometry of the OXYGEN? Now use your kit to construct a molecular model of the C 2 6 O molecule to verify the molecular geometries. Compare the electronegativities of the two s participating in each covalent bond. l) Are any of the bonds POLAR? If YES, clearly indicate the partial negative and partial positive s in your Lewis structure above. m) Do the bond polarities reinforce one another (to make a net dipole), or cancel one another (to make no net dipole)? Explain. C. Graham Brittain Page 12 of 13 10/5/2010

13 V & VI. Ethanol and Dimethyl Ether, C 2 6 O ow many valence electrons does this have? Each Each C O And how many electrons does this need to get an octet (duet for )? So how many covalent bonds will this form? And how many lone pairs will this have? What is the electronegativity value of this? a) TOTAL number of valence electrons in the C 2 6 O molecule = b) Draw a Lewis structure for the C 2 6 O molecule, including all covalent bonds and all lone pairs of electrons. Your structure should have no more and no fewer electrons than the TOTAL number above. In order to answer the questions below, number the two carbon s in your structure (1 and 2). c) ow many groups of electrons are around CARBON #1? d) What is the electron group geometry of CARBON #1? e) What is the molecular geometry of CARBON #1? f) ow many groups of electrons are around CARBON #2? g) What is the electron group geometry of CARBON #2? h) What is the molecular geometry of CARBON #2? i) ow many groups of electrons are around the OXYGEN? j) What is the electron group geometry of the OXYGEN? k) What is the molecular geometry of the OXYGEN? Now use your kit to construct a molecular model of the C 2 6 O molecule to verify the molecular geometries. Compare the electronegativities of the two s participating in each covalent bond. l) Are any of the bonds POLAR? If YES, clearly indicate the partial negative and partial positive s in your Lewis structure above. m) Do the bond polarities reinforce one another (to make a net dipole), or cancel one another (to make no net dipole)? Explain. C. Graham Brittain Page 13 of 13 10/5/2010

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