Molecular Geometry and VSEPR We gratefully acknowledge Portland Community College for the use of this experiment.

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1 Molecular and VSEPR We gratefully acknowledge Portland ommunity ollege for the use of this experiment. Objectives To construct molecular models for covalently bonded atoms in molecules and polyatomic ions To use valence bond theory to account for the bonding in covalently bonded systems To apply the valence shell electron pair repulsion theory and valence bond theory to the geometries and polarities of molecules Discussion An understanding of the structure of a m olecule is f undamental to an explanation of its ch emical and physical properties. For example, water is a liquid at room temperature, dissolves innumerable salts and sugars, is more dense than ice, and has a low vapor pressure, in part, because its molecules are bent rather than linear. Lewis Theory of Bonding Because structure is so important, chemists have developed a number of theories to explain and predict molecular geometries. In 1916, G. N. Lewis developed a theory that focused on the significance of valence electrons (electrons in the outer shell) in chemical reactions and in bonding. e proposed the "octet rule", in which atoms form bonds by losing, gaining, or sharing enough electrons to have the same number of valence electrons (eight) as the nearest noble gas in the Periodic Table. The bond formed is ionic or covalent depending on whether the electrons are transferred or shared between atoms. The octet rule is valid for nearly all compounds formed between atoms of the second period and for a large number of compounds formed between atoms of other representative elements. There are three types of exceptions to the octet rule: molecules having an atom with less than eight electrons, molecules having an atom with more than eight electrons and molecules having an odd number of electrons. Elements of the second and third family may also form stable molecules in which they have less than eight electrons surrounding the central element. For example the molecule BF 3. Elements beyond the third period may form stable molecules with more than eight electrons surrounding the central element. For example Pl 5. Radicals have an odd number of electrons therefore they have unpaired electrons, for example the molecule NO. More than 8 electrons Less than 8 electrons Odd number of Electrons- Radical Lewis emphasized the importance of valence electrons using a Lewis symbol to represent an atom of the element-the symbol for the element surrounded with its corresponding number of valence electrons. For main group elements the number of valence electrons is the same as their family number. For example, Na is the Lewis symbol for the sodium atom, and a is the Lewis symbol for calcium. The Lewis symbols for the atoms are used to account for the bonding in a compound; the resulting representation is called the Lewis formula (or Lewis structure) for the compound. The Lewis formula for water shows that by sharing valence electrons between the oxygen and hydrogen, all three atoms obtain the same number of valence electrons as the nearest noble gas; that is, the hydrogen atoms are isoelectronic with helium atoms and the oxygen atom is isoelectronic with the neon atom.

2 Often it is quite easy to construct an octet rule structure for a molecule. Given that an oxygen atom has six valence electrons (Group 6) and a hydrogen atom has one, it is clear that one O and two atoms have a total of eight valence electrons. Sometimes in these structures a pair of bonding electrons is substituted by a line, while non-bonding electrons or lone are represented with dots. Structures like that of 2 O, involving only single bonds and nonbonding electron, are common. Sometimes, however, there is a "shortage" of electrons; that is, it is not possible to construct an octet rule structure in which all the electron are either in single bonds or are nonbonding. 2 4 is a typical example of such a species. In such cases, octet rule structures can often be made in which two atoms are bonded by two, rather than one pair, of electrons. The two of electrons form a double bond. In the 2 4 molecule, shown above, the atoms each get four of their electrons from the double bond. The assumption that electrons behave this way is supported by the fact that the = double bond is both shorter and stronger than the - single bond in the 2 6 molecule. Double bonds, and triple bonds, occur in many molecules, usually between, O, N, and/or S atoms. For some molecules with a given molecular formula, it is possible to satisfy the octet rule with different atomic arrangements. A simple example would be, 2 6 O: O O The two molecules are called isomers of each other, and the phenomenon is called isomerism. Although the molecular formulas of both substances are the same, 2 6 O, their properties differ markedly because of their different atomic arrangements. Isomerism is very common, particularly in organic chemistry, and when double bonds are present, isomerism can occur in very small molecules as in, 2 2 l 2 : l l l l l l The first two isomers result from the fact that there is no rotation around a double bond, although such rotation can occur around single bonds. The third isomeric structure cannot be converted to either of the first two without breaking bonds. With certain molecules, given a fixed atomic geometry, it is possible to satisfy the octet rule with more than one bonding arrangement. The classic example is benzene, whose molecular formula is 6 6 :

3 These two structures are called resonance structures, and molecules such as benzene, which have two or more resonance structures, are said to exhibit resonance. The actual bonding in such molecules is thought to be an average of the bonding present in the resonance structures. The stability of molecules exhibiting resonance is found to be higher than that anticipated for any single resonant structure. Although a Lewis formula accounts for the bonding based on the valence electrons on each atom, it does not explain how the valence electrons are shared, nor does it predict any three-dimensional structure for a molecule. The valence bond (VB) theory and the valence shell electron pair repulsion (VSEPR) theory provide some insight into the nature of the bonding between atoms and the three-dimensional structure of the molecule. Valence Bond (VB) Theory of Bonding The basic postulate of VB theory is that a covalent bond forms when a pair of electrons is shared by overlapping atomic orbitals between bonding atoms. When these overlapping orbitals point directly at one another, creating a cylindrical symmetry of electron density along the internuclear axis, the bond is a sigma ( ) bond. For example, a bond forms between a hydrogen atom and a fluorine atom making the F molecule when the 1s atomic orbital of the hydrogen atom (having a single electron) overlaps with a 2p atomic orbital of the fluorine atom (also having a single electron). The result is a pair of shared valence electrons along the internuclear axis between the two atoms. 1s 2p bond F F In 1931, Linus Pauling proposed that the orientation of orbitals involved in a bonding determines the three-dimensional structure of a polyatomic molecule or ion. A detailed look at the bonding and structure of methane, 4, supports this. The valence shell electron configuration for carbon is 2s 2 2p 2. Only two "free" 2p electrons are available for bonding on the isolated carbon atom. For methane, however, all four - bonds are equivalent, all -- bond angles are and the geometric structure is tetrahedral. To account for the properties of 4, VB theory states that the bonded carbon atom must not maintain the use of the same 2s and 2p atomic orbitals as on the "isolated" atom. But rather, it must adjust its valence shell orbitals to allow each of its four valence electrons to bond; in so doing, an independent region of space for each electron is available for bonding. 2s 2p As we began with four valence shell atomic orbitals, one s orbital and three p orbitals, we name the four new orbitals in the bonded atom as a combination of these four, we say that each of the four new orbitals is an sp 3 orbital. The word hybrid is often tagged to the name of these new orbitals; the four sp 3 hybrid orbitals are a hybrid of the one s and the three p orbitals from the isolated carbon atom; each has a single electron available for sharing in the formation of a covalent bond. 2p hybridization sp 3 2s ybridization of the valence shell orbitals on the carbon atom.

4 Each sp 3 hybrid orbital bond overlaps with a 1s atomic orbital on the hydrogen atom to form four equivalent to bonds. The four - bonds are oriented in space as a tetrahedron to lessen electrostatic interaction between the hydrogen nuclei. s orbitals of hydrogen sp 3 hybrid orbitals of carbon Four sp 3 hybrid orbitals on a carbon in methane produce a tetrahedral structure. For some molecules one or more of nonbonding electrons in the valence shell of the central atom of the molecule occupy hybrid orbitals. Because these hybrid orbitals are the same as the hybrid orbitals forming the bonds, they also contribute to the geometry of the molecule. Therefore, the orientation of all hybrid valence shell orbitals (those forming the bonds and those containing of nonbonding electrons) determines the geometry of the molecule. ybridization modes for valence shell orbitals in bonding atoms and their corresponding geometries are summarized in Table 1. ybridization igh Electron Density Areas Around entral Atom Bonding Electron Pairs Lone Pairs Molecular Bond Angle Example sp Linear 180 BeF 2 sp Trigonal Planar 120 BF 3 sp Bent / Angular <120 GeF 2 sp Tetrahedral sp Trigonal Pyramidal <109.5 N 3 sp Bent / Angular < O sp 3 d Trigonal Bipyramidal 90, 120 Pl 5 sp 3 d SeeSaw 180, SF 4 120, 90 sp 3 d T-Shape 180, Pl 3 90 sp 3 d Linear 180 XeF 2 sp 3 d Octahedral 90 SF 6 sp 3 d Square Pyramidal 90, 180 IF 5 sp 3 d Square Planar 90, 180 XeF 4 Valence Shell Electron Pair Repulsion Theory of Bonding

5 VSEPR theory proposes that the geometry of a molecule is determined by the repulsive interaction of electron in the valence shell of its central atom. The orientation is such that the distance between the electron is maximized so that electron pair-electron pair interactions are minimized. onstruction of the Lewis formula of a molecule provides the first link in predicting the geometry of the molecule. Methane, 4, has four bonding electron in the valence shell of its carbon atom (the central atom in the molecule). Repulsive interactions between these four electron are minimized when the electron are positioned at the vertices of a tetrahedron. One can generalize that all molecules having four electron in the valence shell of their central atom have a tetrahedral arrangement of these electron. The nitrogen atom in ammonia, N 3, and the oxygen atom in water, 2 O, also have four electron in their valence shell! N O The arrangement of the bonding and nonbonding electron around the central atom gives rise to the corresponding electronic geometry, i.e., tetrahedral. The bonding electrons give rise to the molecular geometry and actual shape of the molecule, trigonal pyramidal for N 3 and bent for 2 O. Various molecular shapes can be determined from Lewis formulas and s which use the following notations: A Xm En Refers to the central atom Refers to "m" number of bonding of electrons on A Refers to "n" number of nonbonding of electrons on A If a molecule has the formula AXmEn, it means there are m + n electron in the valence shell of A, the central atom of the molecule; m are bonding and n are nonbonding electron. For N 3 it would AX 3 E and for 2 O it would be : AX 2 E 2 It should be noted here that electrons on the central atom contributing to a multiple bond do not affect the geometry of a molecule. For example in SO 2, the sulfur atom is sp 2 hybridized, the is AX2E, and the molecule is V-shaped. Further applications of these theories are presented in more advanced chemistry courses. Table 1 summarizes the geometries of molecules or ions and the corresponding bond angle(s) predicted by the VB and VSEPR theories. Polarity of Molecules or Ions Once the three-dimensional shape of a molecule is determined, its polarity can be qualitatively understood. A molecule is polar if an unsymmetrical distribution of electrons exists in the molecule resulting in a partial separation of charges. An unsymmetrical distribution of charge occurs when bonded atoms have different electronegativities; the atoms having a higher electronegativity (electronegativity increases as you move up and to the right in the periodic table) more strongly attract bonding electrons, acquiring a greater electron density and a partial negative charge, -, relative to another portion of the molecule, +. Thus all heteronuclear diatomic molecules are polar. The greater the electronegativity differences of the atoms, the greater the distortion of the electron density, thus, the more polar is the molecule. The direction and magnitude of the polarity are represented by a vector, drawn in the direction of the greater electron density; the + center of the molecule is indicated by a plus sign on the tail of the vector. For example, the F molecule is more polar than the l molecule because the fluorine atom is more electronegative than the chlorine; both are more electronegative than the hydrogen atom. Therefore, the + is on the hydrogen; the vector is drawn in the direction of the fluorine and chlorine. -F -l If more than one polar bond exists in a molecule, the entire molecule may be polar or nonpolar, depending on the geometry of the molecule since the polarity from one bond may be cancelled by the polarity of another. onsider BF 3 and OF 2. The BF 3 molecule has a of AX3, where each B-F bond is polar. Table 1 shows that the BF 3 molecule has a trigonal planar geometry (Figure 4). The

6 three more electronegative fluorine atoms attract the bonding electron from the boron atom with equal magnitude, that is all dipoles are cancelled. Polarity of bonds in BF 3 A geometric sum of the magnitude (all the same) and direction of the three vectors equals zero. Because in the BF 3 molecule there is no resultant vector that can be drawn to indicate a resultant + or - electron density in the molecule, it is said to be a nonpolar molecule, even though each B-F bond is polar. The OF 2 molecule has the AX 2 E 2 and each O-F bond is polar. Table 1 shows that OF 2 has a V-shaped geometry. The geometric sum of the magnitude (all the same) and direction of the two vectors produces a resultant vector that is not equal to zero. A resultant vector can be drawn indicating a + and - electron density in the molecule. In such cases the molecule is polar; therefore, OF 2 is a polar molecule. Experimentally, polar molecules are attracted to an electric field; nonpolar molecules are not. In an experiment using an electrostatically charged rod, a polar liquid, such as water, flowing near an electrically charged rod, is deflected slightly whereas a nonpolar liquid, such as benzene, is unaffected by the electric charge. Although the conclusions we have drawn regarding molecular geometry and polarity can be obtained from Lewis structures, it is much easier to draw such conclusions from models of molecules and ions. The rules we have cited for octet rule structures transfer readily to models. In many ways the models are easier to construct than are the drawings of Lewis structures on paper. In addition, the models are threedimensional and hence much more representative of the actual species. Using the models, it is relatively easy to see both geometry and polarity, as well as to deduce Lewis structures. In this experiment you will assemble models for a sizeable number of common chemical species and interpret them in the ways we have discussed.

7 ALL INFORMATION FOR TESE MOLESULES IS TO BE PALED DIRETLY INTO YOUR LAB NOTEBOOKS Using an appropriate set of molecular models, construct the following molecules/polyatomic ions, and write their Lewis formulas. Determine the number of bonding orbitals (or electron involved in bonding), nonbonding orbitals (or number of nonbonding electron ), and degree of hybridization of the valence shell orbitals on the central atom. Also determine the, the geometry of the molecule/polyatomic ion, and its polarity. Molecule or Polyatomic Lewis Structure Data Sketch of the 3-D Ion 1. F 3 l Bonding orbitals or 4 0 ybridization sp3 AX4 tetrahedral nonpolar no none 2. N3 Bonding orbitals or ybridization 3. 2 O Bonding orbitals or ybridization - 4. NO 3 Bonding orbitals or ybridization Bonding orbitals or ybridization

8 Molecule or Polyatomic Lewis Structure Data Sketch of the 3-D Ion Ol Bonding orbitals or ybridization 3-7. PO 4 Bonding orbitals or ybridization - 8. BF 4 Bonding orbitals or ybridization Ol Bonding orbitals or ybridization 10. P1 2 F2 Bonding orbitals or ybridization 11. SF6 Bonding orbitals or ybridization

9 Molecule or Polyatomic Lewis Structure Data Sketch of the 3-D Ion 12. Pl 2 F 3 Bonding orbitals or ybridization 13. BrF 3 Bonding orbitals or ybridization 14. SF 4 Bonding orbitals or ybridization IF 4 Bonding orbitals or ybridization SO 3 Bonding orbitals or ybridization

10 Molecule or Polyatomic Lewis Structure Data Sketch of the 3-D Ion 17. XeF 2 Bonding orbitals or ybridization 18. SF 2 Bonding orbitals or ybridization