EXPERIMENT 17 : Lewis Dot Structure / VSEPR Theory Materials: Molecular Model Kit INTRODUCTION Although it has recently become possible to image molecules and even atoms using a high-resolution microscope, most of our information about molecular structure comes from often this information enables us to piece together a 3-dimensional picture of the molecule. On paper, one of the best methods we have of representing this model is by drawing a Lewis Dot Structure of the molecule or ion. The ability to draw Lewis structures for covalently bonded compounds and polyatomic ions is essential for understanding of polarity, resonance structures, chemical reactivity, and isomerism. Molecular molecules are a useful tool to help you visualize structures, especially when the ion or molecule is not planar. It is not the intention here to teach you all aspects of drawing Lewis structures or constructing molecular models. The goal here is to supplement your textbook by guiding you through some examples and providing a few insights and tactics where difficulty is often encountered. Be sure to refer to your chemistry textbook for more complete instruction and details. You will use a molecular model kit to construct molecules as they are discussed in this exercise. For each model, you will draw a Lewis dot structure, including nonbonding electrons. The Lewis dot structure is a two-dimensional representation that shows the arrangement of atoms in a molecule. The Lewis dot structure includes both bonding and nonbonding electrons. When drawing covalent molecules, remember that the electrons are shared between two atoms, forming a covalent bond. In this experiment you will be asked to determine the polarity of certain bonds and, after considering the geometry of the molecule, decide whether it is polar or nonpolar. Lewis Diagram (Electron Dot) In most stable molecules or polyatomic ions, each atom tends to acquire a noble-gas structure by sharing electrons. This tendency is often referred to as the octet rule. One way to show the structure of an atom or a molecule is using dots to represent the outermost s-and p-electrons (the so-called valence electrons). For the group IA elements, the number of valence electrons is the same as the group number in the periodic table. Element Group Valence electrons Lewis dot ionic form 1
In drawing Lewis diagrams we usually do not attempt to show from which atom the valence electrons come we simply indicate a shared pair of electrons as either two dots or a straight line connecting the atoms. Unshared pairs, also called lone pairs, are indicated by dots drawn around the elemental symbols. For example: Atoms in polyatomic ions are held together by covalent bonds. In the ions we will consider in this assignment, all atoms have a noble-gas structure. Occasionally there are too few electrons available in a species to allow an octet to exist around each atom with only a single bond. In this instance, multiple (double or triple) bonds will form. For example: Carbon dioxide Hydrogen cyanide Acetylene Rules for drawing Lewis Dot Structures I. Determine the total number of valence electrons and valence electron pairs in all the atoms in the molecule. For polyatomic 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. II. Connect the central atom to the surrounding atoms with a single bond. Each bond represents two electrons. III. Place electrons about the outer atoms so each (except hydrogen with two electrons) has an octet. IV. Count the total pairs of electrons; if: (a) the number of pairs matches the total number of valence electron pairs calculated earlier; this is the correct Lewis dot structure. (b) the number of pairs is less than the total valence electron pairs calculated earlier; add a pairs to the central atom to match the total number of valence pairs. (c) the number of pairs is more than the total valence electron pairs calculated earlier, form one double bond (one extra pair) or one triple bond (two extra pairs) around central atom. In the case of a triple bond, you can place two double bonds instead. 2
Example 1 - Methane CH 4 1. This compound is covalent. Determine the total number of valence electrons available. One carbon has 4 valence electrons, four hydrogen, each with one valence electron, totals 4. This means there are 8 valence electrons, making 4 pairs, available. 2. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by outer atoms. Hydrogen is never the central atom. Determine a provisional electron distribution by arranging the electron pairs (E.P.) until all available pairs have been distributed. The formal charge (F) on the central atom is zero. Example 2 - Ammonia NH 3 1. This compound is covalent. Determine the total number of valence electrons available. One nitrogen has 5 valence electrons. Three hydrogen, each with one valence electron, totals 3. This means there are 8 valence electrons, making 4 pairs, available. 2. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by outer atoms. Hydrogen is never the central atom. It does not matter which of the three sides you use to put hydrogens on. Determine a provisional electron distribution by arranging the electron pairs (E.P.) until all available pairs have been distributed. The formal charge (F) on the central atom is zero. 3
Example 3 - Water H 2 O 1. This compound is covalent. Determine the total number of valence electrons available. One oxygen has 6 valence electrons. Two hydrogen, each with one valence electron, totals 2. This means there are 8 valence electrons, making 4 pairs, available. 2. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by outer atoms. Hydrogen is never the central atom. It does matter which of the two sides you use to put hydrogens on. Use sides that are next to each other. DO NOT put the hydrogens 180 degrees apart. Determine a provisional electron distribution by arranging the electron pairs (E.P.) until all available pairs have been distributed. The formal charge (F) on the central atom is zero. Example 4 Carbonate ion, CO 3 2- Total valence pair electrons = 1 (C) + 3 (O) + 2 e - = 1 (4 e - ) + 3 ( 6e - ) + 2 e - = 24 e - = 12 electron pairs Drawing: (13 E.P.) (12 E.P.) (12 E.P.) (12 E.P.) The carbonate ion, CO 3 2- has three resonance structures. Delocalization of pi electrons forms resonance structures. 4
Exceptions to Octet Rule and Lewis Dot Structure There are compounds that cannot be represented by these rules (octet rules) for Lewis dot structures. The central atom may have either less than eight electrons ( BF 3 ) or more than eight electrons ( PCl 5, SF 6, XeF 4, etc). For most of these compounds, the central atom and each outer atom are bonded by single bonds consisting of one pair of electron. If there are any extra electrons on the central atom, they are grouped as shared pairs on the central atom. For example: BF 3 = 1(3e - ) + 3(7 e - ) = 24 e - = 12 p e - PCl 5 = 1 (5 e - ) + 5 ( 7 e - ) = 40 e - = 20 p e - ( 12 p e - ) ( 20 p e - ) Valence Shell Electron Pair Repulsion(VSEPR) Theory (Electron Pair and Molecular Geometry) VSEPR stands for Valence Shell Electron Pair Repulsion. The whole concept revolves around the idea that the electrons in a molecule repel each other and will try and get as far away from each other as possible. VSEPR explains a lot about molecular geometry and structure. The electrons (both in pairs and singles as you will see) are "attached" to a central atom in the molecule and can "pivot" freely on the atom's surface to move away from the other electrons. Electrons will come in several flavors: a) bonding pairs - this set of two electrons is involved in a bond, so we will write the two dots BETWEEN two atoms. This applies to single, double, and triple bonds. b) nonbonding pairs - this should be rather obvious. c) single electrons - in almost every case, this single electron will be nonbonding. Almost 100% of the examples will involve pairs, but there are a significant number of examples that involve a lone electron. 5
VSEPR uses a set of letters to represent general formulas of compounds. These are: a) A - this is the central atom of the molecule (or portion of a large molecule being focused on). b) X - this letter represents the atoms attached to the central atom. No distinction is made between atoms of different elements. For example, AX 4 can refer to CH 4 or to CCl 4. c) Eand e - this stands for nonbonding electron pairs. Each area where electrons exist is called an "electron domain" or simply "domain." It does not matter how many electrons are present, from one to six, it is still just one domain. Now a domain with six electrons in it (a triple bond) is bigger (and more repulsive) than a lone-electron domain. However, it is still just one domain. The more electrons in a domain, the more repulsive it is and it will push other domains farther away than if all domains were equal in strength. Keep in mind that the domains are all attached to the central atom and will pivot so as to maximize the distance between domains. An important point to mention in this introduction is that an element's electronegativity will play an important role is determining its role in the molecule. For example, the least electronegative element will be the central atom in the molecule. The more electronegative the element, the more attractive it is to its bonding electrons This will play a very important role, especially in five domains. The most important domain numbers at the introductory level are 3, 4, 5, and 6. Domains of 1 and 2 exist, but are simple to figure out. 6
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 are counted for each single atom. In the following examples, CO 3 2- has three regions, CO 2 has two regions, and PCl 5 has five regions. The following electronic geometries are expected for the following numbers of regions of high electron density. Some examples follow. Regions 2 3 4 5 6 Drawing X X X X X = X = X = X Geometry linear trigonal tetrahedral trigonal octahedral planar pyramidal Example CO 2 BF 3 CCl 4 PCl 5 SF 6 Molecular Polarity (Dipole Moment) If its molecular geometry is completely symmetrical, a molecule is nonpolar. If the molecular geometry is unsymmetrical, the molecule is polar because of the lone pair of electrons on the central atom. Polar bonds (due to differences in the electronegativities) may reinforce or oppose the effect of the lone pairs of electrons. Some examples are as follow: Symmetrical(polar) Unsymmetrical (non-polar) 7
Molecular Hybridization Once the Lewis structure is drawn, it is possible to apply either thevsepr theory or a hybridization to predict the shape of the species. Both systems when appropriately used will predict (with very few exceptions the same shape. The table below summarizes the theories. Electron Pair and Molecular Geometry Table of Three to Six Electron Domains Electron Domains Arrangement of Electron Domains General Molecular Formula Molecular Shape Examples Hybrid Orbitals (central atom) Polar? 2 Linear AX 2 Linear CO 2 sp No 3 4 5 Trigonal Planar (3 electron domains) Tetrahedral (4 electron domains) Trigonal bipyramidal (5 electron domains) AX 3 Trigonal planar BF 3, AlCl 3 sp 2 No AX 2 E Angular SnCl 2 Yes AX 4 Tetrahedral CH 4, SiCl 4 sp 3 No AX 3 E Trigonal pyramidal NH 3, PCl 3 Yes AX 2 E 2 Angular (Bent) H 2 O, SCl 2 Yes AX 5 Trigonal bipyramidal PCl 5, AsF 5 sp 3 d No AX 4 E Seesaw SF 4 Yes AX 3 E 2 T-shaped ClF 3 Yes AX 2 E 3 Linear XeF 2 No 6 Octahedral (6 electron domains) AX 6 Octahedral SF 6 sp 3 d 2 No AX 5 E Square pyramidal BrF 5 Yes AX 4 E 2 Square planar XeF 4 No 8
EXPERIMENT 17: Lewis Dot Structure / VSEPR Theory REPORT FORM Name Instructor Date Partner s Name: Formula Total Valence electrons Bonding pair electrons Lone pair electrons Lewis Dot Structure Electron Geometry and Shape Bond Angle Hybridization NH 3 8 3 1 H.. N H H Tetrahedral Trigonal Pyramidal 107⁰ sp 3 107⁰ CF 4 FNO 2 9
HCN IBr 2 OH H 3 O + ClO 3 10
POCl 3 ClO 2 CCl 4 NCl 3 11
SiF 4 CO 3 2- NH 4 + I 3-12
EXPERIMENT 17 : Lewis Dot Structure / VSEPR Theory Name: Pre- laboratory Questions and Exercises Due before lab begins. Answer in space provided. 1. Write the Lewis dot structure for the following atoms or ions. Al Cl K + O 2- Si 2. What is the electronic geometry about a central atom which has the following number of regions of electron density? a) Three regions of electron density: b) Five regions of electron density: 3. Draw Lewis structures for CO 2 and SO 2. Compare the geometries, angles, and polarities. 4. Define resonance and give an example. 5. Are BF 3 and NF 3 symmetrical or unsymmetrical molecules? Explain your reasoning. 13
EXPERIMENT 17 : Lewis Dot Structure / VSEPR Theory Name: Post- laboratory Questions and Exercises Due after completing lab. Answer in space provided. 1. Describe polar and non-polar bonding. Give an example of each. 2. Draw Lewis structure for SO 3 and SO 3 2. 3. Are CH 3 + and CH 3 polar or non-polar molecules? Explain your reasoning. 4. Define dipole moment and give an example. 5. Draw the Lewis structures for ethanol, (C 2 H 5 OH), and dimethylether (CH 3 OCH 3 ). Determine the electron pair geometry and molecular geometry around the oxygen atom in each. 14