H O H C H H N H H. Valence Shell Electron Pair Repulsion: Predicting Shape & Polarity

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1 Valence Shell Electron Pair Repulsion: Predicting Shape & Polarity BJECTIVES Students will develop the ability to: 1. Predict the arrangement that valence e pairs assume around an atom (e pair geometry) 2. Draw & name the e pair geometry and molecular shape of molecules & compound ions 3. Predict the polarity of molecules & compound ions INTRDUCTIN The main approach we will use to study/predict the shape of molecules and compound ions is Valence Shell Electron Pair Repulsion Theory (VSEPR). (See pp in Timberlake, 10 th ed.) The basis for this theory is that valence electron pairs, being negatively charged, tend to be arranged in space so that they are as far apart as possible. You might expect from your knowledge of the Lewis ctet Principle that atoms (other than, e, Li, Be, and B) would always have 8 valence electrons (4 valence e pairs) and therefore that structure prediction would be very simple. (See examples with C 4, N 3, 2, and F below.) In fact, the situation is complicated by formation of multiple bonds (double & triple) and the tendency of period 3 and higher elements to sometimes have more than 8 valence electrons. We are interested in valence electron pair orientation because knowing the shape of a molecule and the electronegativity of its atoms allows us to make predictions about its polarity. Knowledge of molecular structure and polarity in turn allows us to make predictions about many properties (boiling point, solubility, pharmacological etc.) of materials. See Appendix A, which illustrates some molecular structure drawing conventions. It is sometimes difficult for students to distinguish between the electron pair geometry of a molecule and the overall molecular shape of that molecule. The electron pair geometry refers to the orientation of valence e pairs around a specific (often central) atom. In the most direct sense, the e pair geometry of the central atom does not depend on the existence of the outer atoms. The molecular shape, however, refers to the overall space occupied by the entire molecule, and this is clearly dependent on the existence of any outer atoms as well as the central atom. A simple example which clearly makes this distinction concerns the case in which the central atom of a molecule has four valence e pairs. Look at the Lewis structures of the following four molecules: methane (C 4 ), ammonia (N 3 ), water ( 2 ), and hydrogen fluoride (F). C N In each case the central atom has 4 valence e pairs, and therefore the non- atom of each molecule has tetrahedral e pair geometry. owever, because the outer atom complement is different for each molecule, they all have different molecular shapes. C 4 is tetrahedral, N 3 is trigonal pyramidal, 2 is bent, and F must be linear, because it has only two atoms. See Appendix B for information on e pair geometry and molecular shape. If you want to see different ways to represent molecular structures with computers modeling, look at the last page of the lab. F 1

2 Geometry Review We need a little general background in geometry to predict e pair geometries. Recall that we commonly divide a circle into 360. The figure below shows how electron pairs would arrange themselves around a nucleus to achieve maximal separation. rbital Geometry Lesson 2 orbitals 3 orbitals 4 orbitals 6 orbitals 180 o 120 o o 90 o linear trigonal planar tetrahedral octahedral Note the angles between the electron pairs and the symmetry of each structure. ere, symmetry means that the structure looks the same from more than one viewing position. The vast majority of biochemicals contain C, N, and atoms with trigonal planar or tetrahedral e pair geometry. C, N, and are all in Period 2 (the second row) of the Periodic Table. A few relatively important compounds (e.g., C 2 ) do contain Period 2 atoms with linear e pair geometry. P (phosphorous) and S (sulfur) are the elements in Period 3 that are the most important in forming covalent bonds in biomolecules. While they behave somewhat differently than Period 2 elements, we can generally treat their e pair geometries much like those of C, N, and. P, S, and some of the biochemically important transition metals (e.g., Fe, Co, Ni, Cu, and Zn) often have bonding arrangements involving 4, 5, or 6 orbitals. We will spend only a little time with the 5 & 6 orbital cases. The most difficult transition for some students is realizing that when you go from 3 orbitals to 4, the geometric figure goes from 2 dimensions (planar) to 3 dimensions (tetrahedral). Molecular Type In today s exercise, you will need a shorthand method to describe specific electron pair arrangements (called molecular type) around the central atom of a molecule. The general form of molecular type is AB x E y, where A represents the central atom, B represents the atoms directly attached to the central atom (by bonding electron pairs), and E represents the non-bonding (unshared) valence electron pairs associated with the central atom. The x represents the number of atoms attached to and the y represents the number of non-bonding pairs of the central atom. The molecular type must be determined from the Lewis dot structure. From the Lewis dot structures on the previous page it can be seen that F is ABE 3, 2 is AB 2 E 2, N 3 is AB 3 E, and C 4 is AB 4. 2

3 PRCEDURE Pay very careful attention to the number of holes and bond angles in the atoms you select. If you pick one atom incorrectly, the whole structure may be wrong. Get one sandwich bag containing: black atoms: one with 2 holes, two w/ 3 holes, 3 w/ 4 holes green atoms: nine (1 hole) R white atoms: nine (1 hole) short bonds: twelve gray and three white Electron Pair Geometries. Start by locating one each of the black atoms that have 2, 3, and 4 holes. Select 9 short grey bonds. Insert one bond into each hole on each of the three atoms. Use the models that you have constructed to fill in the blanks in the table on p. 5. For the purposes of this lab, a double bond counts as only one e pair, and likewise for a triple bond. We will explain this apparent contradiction later in the semester. Molecular Shapes. For each molecular structure that you create below, sketch the molecule and name the molecular shape and molecular type in the table on p. 6. Also, for each structure that you create, state whether the molecule would be polar or non-polar if all of the outer (i.e., non-central) atoms were identical. If you are unclear about bond polarity, refer to your electronegativity table and ask your instructor. 1) Linear e pair geometry attach atoms white (or green) with a single hole to each bond on your linear model from the previous section. This should require 2 white (or green) atoms. 2) Trigonal Planar a) attach atoms with a single hole to two of the bonds on your trigonal planar model from the previous section. This should require 2 atoms. Sketch and name this structure in the table on p. 5 in the upper of the two rows marked Trigonal Planar. In your drawing, place a pair of dots at the end of the bond with no atom attached to emphasize that this represents a non-bonding electron pair. b) Now place one more atom on the structure and use this model to fill in the lower of the two rows marked Trigonal Planar. 3) Tetrahedral a) attach atoms with a single hole to two of the bonds on your tetrahedral model from the previous section. This should require 2 atoms. Sketch and name this structure in the table on p. 5 in the upper row marked Tetrahedral. Remember the dots for the non-bonding e pairs. b) Now place one more atom on the structure and use this model to fill in the middle row marked Tetrahedral. c) Finally, place an atom on the fourth bond and fill in the final row in the table on p. 5. Unknown Structures Your instructor will assign each individual (not each group) three different molecules/compound ions. Fill in the table on p. 7* and construct models of your three molecules. If your structure has double or triple bonds, use one of the white bonds to represent the double or triple bond. ave these checked by the instructor before you leave the lab today. While orbital geometry tells you qualitatively whether or not a compound is polar, you need the Electronegativity Table when making decisions about how polar molecule is. *Note: Remember, when counting the number of electron pairs around the central atom, count the bonding and the non-bonding pairs. For this purpose, double and triple bonds, however, count as only one pair! rev. Aug

4 Pre lab Questions 1. What are valence electrons? 2. ow many valence electrons do the following atoms have? a) Li e) Si b) e f) N c) Be g) Ne d) C h) Na Which of the above atoms are considered stable? 3. What is the Lewis ctet Rule? 4. In Lewis Dot Structures, valence electrons are represented by dots arranged around the elemental symbol. Write Lewis Dot structures for the following atoms. a) phosphorous, P d) chlorine b) carbon e) oxygen c) potassium, K f) argon 4

5 Sec: Name: Table 1. Electron Pair Geometry # of e -- pairs around central atom sketch of your model bond angles name of this geometric shape Instructor s initials 5

6 Table 2. Molecular Shape e -- pair geometry sketch + name e pair geom (include non-bonding e -- pairs) sketch + name of molec shape (don t include non-bonding e - pairs) molecular type polar or not Linear Trigonal Planar Trigonal Planar Tetrahedral Tetrahedral Tetrahedral 6

7 Table 3. Unknowns Molecular Formula #1 #2 #3 Number of valence e Lewis Dot Structure Molecular Type # of valence e -- pairs around central atom * e -- pair geometry: name and sketch Molecular shape: name and sketch Polar or non-polar * Double and single bonds represent just one valence e - pair in this case. 7

8 Post lab Questions 1. Determine the total number of valence electrons for the following molecules/compound ions. 3- a) C 3 C d) P 4 b) 3 e) N 3 c) CBr 3 2. What possible molecular shapes can result from trigonal planar electron pair geometry? 3. Can a molecule have tetrahedral electron configuration and trigonal planar molecular shape? Explain. 4. If you had molecules with the following molecular types, which would be polar and which would be non-polar? a) AB 3 E b) AB 2 c) AB 2 E 2 d) Explain your answer choice 8

9 Appendix A # valence e - pairs VSEPR Geometry Table drawing name bond angle(s) linear trigonal planar tetrahedral octahedral 180 o 120 o o 90 o geometric figure geometric figure name line equilateral triangle tetrahedron octahedron Notes: 1. represents the central atom of a structure. 2. Is an e - pair extending toward you. Is an e - pair extending away from you. 3. The lines in front & top views represent valence electron pairs of the central atom. 4. The geometric figures indicate the shapes that would be outlined by outer atoms. 9

10 Appendix B Electron Pair Geometry & Molecular Shape formula C 4 N 3 2 F ball & stick molecular tetrahedral trigonal bent NA shape pyramidal name picture N F Space filling Notes: 1. All molecules shown above have tetrahedral electron pair geometry. owever, because they have different outer atom complements, they have different molecular shapes. 2. The molecules are not drawn exactly to scale, but are shown with appropriate bond angles. 10

11 line cylinder ball & stick space filling 11