# EXPERIMENT 9 Dot Structures and Geometries of Molecules

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1 EXPERIMENT 9 Dot Structures and Geometries of Molecules INTRODUCTION Lewis dot structures are our first tier in drawing molecules and representing bonds between the atoms. The method was first published by Gilbert N. Lewis in 1916, shortly after the announcement of the Bohr model of the atom in which the electrons move in orbits around the nucleus in a manner like the planets orbit the sun. Using this concept, Lewis made the distinction between inner electrons, occupying inner orbits closer to the nucleus, and the outermost, or valence electrons. In Lewis concept of bonding, valence electrons of different atoms are shared between both atoms when the atoms are at an optimal distance from each other. This sharing of valence electrons constitutes a covalent bond. Lewis recognized that the noble gases are especially stable and that other atoms can achieve a more stable state by attaining the same number of valence electrons as the noble gases when the react with other elements to form compounds. Since all of the noble gases except helium have 8 valence electrons (helium has 2 valence electrons), this rule of eight became known as the octet rule. When atoms after the element helium bond together, they most commonly will form a sufficient number of bonds to acquire 8 valence electrons. Lewis dot structures include both bonding and nonbonding electrons. The bonding electrons are the ones shared by both atoms. A single bond consists of two shared electrons. Lewis represented two shared electrons (a bonding pair) with a single line between the atoms. EXAMPLE 1. Consider the HCl molecule, hydrogen chloride. A neutral hydrogen atom, by itself, has one valence electron, while similarly, a neutral, single chlorine atom has 7 valence electrons. The total number of valence electrons in this molecule is therefore = 8 valence electrons. We can represent the formation of a hydrogen chloride molecule from a hydrogen atom and a chlorine atom as follows: H + Cl H Cl or H Cl Hydrogen is in Group 1A and therefore has one valence electron, while chlorine is in Group 7A and therefore has 7 valence electrons. In the HCl molecule, hydrogen is now seen to have two electrons associated with it (the same number of valence electrons as helium, the nearest noble gas to hydrogen). Counting the dots around the chlorine atom, we see that it is surrounded by 8 dots in the HCl molecule, the same number of valence electrons as the noble gas argon (the noble gas closest to chlorine in the periodic table). The shared electrons are counted twice; once for the hydrogen and once as part of the chlorine s electrons. The HCl molecule thus has one bonding pair of electrons, making a single bond between the hydrogen and chlorine atoms, and in addition, the chlorine atom has three nonbonding pairs, or lone pairs, associated with it. In molecular compounds, hydrogen atoms will have only one bond. Other atoms will have enough bonds and lone pairs to give a total of 8 electrons around them, satisfying the octet rule, although exceptions to the octet rule do occur. CHEM 1405 Experiment 9 1

2 EXAMPLE 2. Next, let us examine the O 2 molecule. The total number of valence electrons is 12, 6 from each oxygen atom. If we first try putting a single bond between the oxygen atoms, and then add enough lone pairs to obtain an octet of electrons around each oxygen atom, we will have 14 electrons in our structure, two too many: O O Therefore, let us try using a double bond between the oxygen atoms: O=O Then, we will add enough lone pairs to obtain an octet of electrons around each oxygen atom as before: O=O The result is an octet around each oxygen atom, and we have the correct total number of valence electrons in our structure, 12. Therefore, this is the correct dot structure for the O 2 molecule. EXAMPLE 3. The N 2 molecule has a total of 10 valence electrons, 5 valence electrons contributed from each nitrogen atom. When we draw the dot structure of this molecule, we will discover that we cannot satisfy the octet rule on each nitrogen atom and have the correct total of 10 valence electrons using a single or double bond between the nitrogen atoms. In this molecule, a triple bond is required: N N N N Molecules Containing More than Two Atoms When drawing dot structures of molecules, a certain amount of trial and error is common. However, for simple molecules consisting of three or more atoms, it is common for one atom to occupy a central position with the other atoms bonded to this central atom. The following steps generally work well for arriving at a correct dot structure for molecules that have an even number of valence electrons: 1) Add up the total number of valence electrons from all the atoms in the molecule. For ions, add or subtract electrons to arrive at the appropriate charge. For example, if the charge is 1, you need to add one electron, and if the charge is +1, you need to subtract one electron. 2) Place the unique atom (in which there is only one) in the center and connect the other atoms to this central atom using single bonds (single lines). 3) Place enough lone pairs on each atom to obtain an octet of electrons on each atom (except hydrogen). 4) Count the valence electrons in your structure: dots plus two electrons per line. If the total agrees with the total from Step 1, you have a valid dot structure. 5) If the total is more than the correct number from Step 1, try placing one double bond in your structure, and then repeat Steps 2 through 4. 6) If one double bond does not work, try using two double bonds, or one triple bond, and repeat Steps 2 through 4. Eventually, you will arrive at a correct dot structure. Exceptions to the Octet Rule 1) If the molecule an odd number of valence electrons, such as in the molecule NO 2, you will not be able to obtain an octet on each atom. 2) The molecule or ion may not have enough total valence electrons to attain an octet, such as in the methyl cation CH 3 + ( = 6 valence electrons total). 3) Atoms in Row 3 or below in the periodic table can have expanded octets of more than 8 electrons, in order to obtain lower formal charges on the atoms. CHEM 1405 Experiment 9 2

3 EXAMPLE 4. The SO 2 molecule. First we will add up the total number of valence electrons from one S and 2 O atoms: 6 + (2 X 6) = 18 valence electrons. Next, put the unique atom (the sulfur) in the middle and attach the two oxygen atoms to it: O S O Now fill in enough pairs of electrons to get an octet of electrons around each atom: O S O And then count the total number of valence electrons in the structure: 2 bonding pairs + 8 nonbonding pairs = 10 electron pairs total, which is 20 valence electrons in the structure, two too many. The octet rule is satisfied, but we have too many electrons. Try using one double bond: O=S O Now fill in enough electron pairs to get an octet around each atom: O=S O Finally, check the total: 3 bonding pairs + 6 nonbonding pairs = 9 electron pairs total = 18 valence electrons, the correct number. Since the octet rule is satisfied on all atoms and the structure has the correct total number of valence electrons, this is a valid dot structure of the SO 2 molecule. The SO 2 molecule also has an additional resonance structure which makes the two S O bonds equivalent. In resonance structures, the electrons are arranged differently but the atoms stay in the same positions. Resonance structures are representations of the same molecule. The true structure of the molecule is a blending of the individual resonance structures. USING DOT STRUCTURES TO DETERMINE MOLECULAR GEOMETRIES The geometry of a molecule is determined by the number of atoms, the number of bonds, and the type of bonds in the molecule. When there is a central atom, the geometry of the molecule as a whole is described by the geometry associated with the central atom. First we should note that the geometry of a molecule can be described in two ways. The first is the overall geometry, often called the electron pair geometry or electron domain geometry. In determining the kind of overall geometry a molecule has, we must consider the location of all atoms and lone pairs attached to the central atom. The second way to describe the geometry of a molecule is called simply the molecular geometry or molecular shape. To determine the molecular geometry, we must first obtain the overall geometry. The molecular geometry is a subcategory of the overall geometry and is obtained by looking at the locations of the atoms only that are attached to the central atom. Pretend that any lone pairs on the central atom are invisible. To determine the overall geometry of a molecule, and subsequently its molecular geometry, we first must draw a valid dot structure of the molecule. From the dot structure, we can see how many atoms and lone pairs are attached to the central atom. The overall geometry is determined by a very simple principle: the things (atoms and lone pairs) attached to the central atom will naturally spread themselves apart to attain the least amount of crowding. CHEM 1405 Experiment 9 3

4 This principle of minimizing the crowding of atoms and lone pairs attached to an atom is called the Valence Shell Electron Pair Repulsion theory, or VSEPR theory in short. It is that simple! The atoms and lone pairs attached to the central atom will spread out in such an arrangement that crowding is as low as possible. This arrangement is described by the overall geometry associated with the central atom, and by extension, the molecule as a whole. The VSEPR theory states that regions of high electron density will arrange themselves as far apart as possible around the central atom. The regions of high electron density around the central atom result from each atom and lone pair attached to the central atom. In the following examples, CO 3 2 has three such regions, CO 2 has two regions, and PCl 5 has five regions. The following overall geometries are expected for the numbers of regions of high electron density: Number of Regions Possible X X X X X Bonding = X = Patterns X = X Overall Linear Trigonal Tetrahedral Trigonal Octahedral Geometry Planar Bipyramidal Example CO 2 BF 3 CCl 4 PCl 5 SF 6 The next table lists the bond angles that occur in molecules with various overall geometries. These are ideal angles, which would be the case if all repulsions between atoms and lone pairs are the same: Overall Geometry Bond Angles Present Linear 180 Trigonal Planar 120 Tetrahedral Trigonal Bipyramidal 90, 120 and 180 Octahedral 90 and 180 Geometries of molecules verified by experiment are in good agreement with the geometries predicted by VSEPR theory. However, small distortions from the ideal geometries are observed, evidently due to the fact that repulsions between lone pairs are greater than repulsions between atoms that are attached to the central atom. For example, from the dot structure of the water molecule, H 2 O, we see that the central O atom has 4 electron regions, two atoms (the H atoms) and two lone pairs. From the above table, the overall geometry that spreads apart 4 things attached to the central atom as much as possible is the tetrahedral geometry, in which the bond angles are all In the water molecule, however, the actual H O H bond angle is 104.5, a little less than the ideal angle of 109.5, because the lone pairs spread out from each other to a greater extent, crowding the hydrogen atoms a little closer to each other. Study the following tables to visualize the different overall geometries (called electron pair geometries in the tables) and molecular geometries. CHEM 1405 Experiment 9 4

5 Tables are copyright Prentice-Hall, Inc. CHEM 1405 Experiment 9 5

6 CHEM 1405 Experiment 9 6

7 POLARITY IN MOLECULES If a molecule (or polyatomic ion) is completely symmetrical about the central atom, the molecule will be nonpolar. If the geometry about the central atom is not completely symmetrical, the molecule will be polar. In the latter case, this asymmetry about the central atom results from the effect of lone pairs of electrons on the central atom. Some examples are as follows: O = C = O N H H H Symmetrical about the central atom (non-polar) Unsymmetrical about the central atom (polar) Polarity results from an unsymmetrical distribution of electron density in the molecule or ion. The reason this can occur is because different atoms have different attractions for electrons they are sharing in a covalent bond. If in atom has a high attraction for electrons it is sharing in a covalent bond, the atom is said to have a high electronegativity. The bond between atoms with different electronegativities is called a polar covalent bond. In a polar covalent bond, the electrons making the bond are not shared equally. Rather, the electrons are pulled toward the atom with the higher electronegativity. If the electronegativitites of the atoms are the same or nearly the same, the bond is classified as a nonpolar covalent bond, and the electrons are shared equally, or nearly so. Taken together, the polarities in the different bonds in a molecule may cancel each other, as they do in symmetrical molecules such as CO 2, to give a molecule that is nonpolar overall. On the other hand, if these individual bond dipoles of the polar bonds in the molecule do not cancel, the molecule will have a net or overall polarity, as is the case with NH 3 and H 2 O. To better understand this principle, examine the molecular geometries illustrated in the above tables and see if you can tell whether the molecule is polar or nonpolar. EXPERIMENTAL PROCEDURE 1. Draw the Lewis dot structure for each molecule or ion listed on the report form. 2. Based on your dot structure, and using the tables above as a guide, determine the overall geometry and the molecular geometry of each molecule or ion. 3. Using a model kit, build a model of the molecule. Use the model to verify the geometry you predicted. If the model and your prediction do not agree, find out why. DO NOT proceed to the next compound until you thoroughly understand each structure. 4. Based on its molecular geometry, identify whether the molecule or ion is polar or nonpolar. CHEM 1405 Experiment 9 7

8 EXPERIMENT 9 Dot Structures and Geometries of Molecules REPORT FORM Name Instructor Date Molecular Number of Lewis Dot Overall Molecular Polar/ Formula Valence e Structure Geometry Geometry Nonpolar _ H 2 O 8 Tetrahedral Bent Polar CCl 4 NF 3 CS 2 OF 2 CHEM 1405 Experiment 9 8

9 EXPERIMENT 9 Name Molecular Number of Lewis Dot Overall Molecular Polar/ Formula Valence e Structure Geometry Geometry Nonpolar _ HCN NH 4 + SO 3 (resonance structures) NO 3 (resonance structures) AlH 3 (less than an octet) PF 5 (expanded octet) CHEM 1405 Experiment 9 9

10 EXPERIMENT 9 Name Pre-Laboratory Questions and Exercises Due before lab begins. Answer in the space provided. 1. Write the Lewis dot structure for the following atoms or ions: Al Cl K + O 2 Si 2. What is the overall geometry about a central atom which has the following number of regions of electron density? a) Three regions of electron density: b) Four regions of electron density: 3. Draw Lewis structures for BH 3 and NH 3. Compare the geometries, angles, and polarities. Give the overall geometry and the molecular geometry of BF 3 and NF 3 according tovsepr theory, and identify if the molecule is polar or nonpolar. Compound: BH 3 NH 3 Dot Structure: Overall Geometry: Molecular Geometry: Polar / Nonpolar: CHEM 1405 Experiment 9 10

11 EXPERIMENT 9 Name Post-Laboratory Questions and Exercises Due after completing the lab. Answer in the space provided. 1. a) What is a polar covalent bond? b) What is a nonpolar covalent bond? c) What is an ionic bond? 2. Draw the Lewis dot structure for ethanol, CH 3 CH 2 OH, and dimethyl ether, CH 3 OCH 3. Determine the overall geometry and molecular geometry around the oxygen atom in each, and identify the molecules as being polar or nonpolar. Compound: CH 3 CH 2 OH CH 3 OCH 3 Dot Structure: Overall Geometry: (around the O atom) Molecular Geometry: (around the O atom) Polar / Nonpolar: CHEM 1405 Experiment 9 11

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