Effect of unshared pairs on molecular geometry

Transcription

1 Chapter 7 covalent bonding Introduction Lewis dot structures are misleading, for example, could easily represent that the electrons are in a fixed position between the 2 nuclei. The more correct designation is shown in figure 7.1. At any given point the two electrons may be located at any of the various points around the 2 nuclei. owever, they are more likely to be found between them than at the far ends. If we look at the graph on page 164 of the interactions. 1. at large distances they do not interact with each other, but as they get closer together they experience an attraction that leads to an energy minimum. The distance is 0.074nm which has an attractive energy of 436 kj. The system is in its most stable state and is referred to as the 2 molecule. When they are forced closer together, the repulsion forces are too great and the energy rises. 2. When two hydrogen atoms come together to form a molecule, the electrons are spread over the entire volume of the molecule. Quantum mechanics tells us that increasing the volume available to an electron decreases its kinetic energy. This is described by saying that the 2 1s orbitals overlap to form a new bonding orbital. Main points of the chapter 1. distribution of outer level (valence) electrons. Lewis structures 2. molecular geometries VSEPR model can be used to predict angles between covalent bonds formed by a central atom 3. polarity of covalent bonds and the molecules they form, positive and negative poles 4. distribution of valence electrons among atomic orbitals, using the valence bond approach 7.1 Lewis structures and the octet rule G. N. Lewis in 1916 proposed that nonmetal atoms, by sharing electrons to form an electron pair bond, can acquire a stable noble gas structure Example F, Fluorine has 7 normal electrons from its configuration 1s 2 2s 2 2p 5 and shares one with ydrogen Lewis structure rules 1. a pair of electrons shared between 2 atoms is a covalent bond, ordinarily shown as a straight line between bonded atoms 2. unshared pairs of electrons, owned entirely by one atom, is shown as a pair of dots on that atom (lone pair) Examples of single bonds hydroxide ion, water, ammonia, ammonium ion double bonds ethylene C 2 4 triple bonds acetylene C octet rule nonmetals, except for hydrogen, achieve a noble-gas configuration by sharing an octet of electrons, hydrogen shares 2 Writing Lewis structures 1. Count the number of valence electrons 2. Draw a skeleton structure for the species, joining atoms by single bonds, central atom is the one written first in the formula, put it in the center. Terminal atoms are most often hydrogen, oxygen, or a halogen; bond these to the central atom. 3. Determine the number of valence electrons still available for distribution

2 4. Determine the number of valence electrons required to fill out an octet for each atom (except ) in the skeleton Resonance forms Take for example, SO 2 the double bond can be on either oxygen, but what actually occurs is an intermediate between the two. Other examples include NO 3, and benzene Resonance species rules 1. resonance forms do not imply different kinds of molecules with electrons shifting eternally between them. There is only one type of SO 2 molecule structure; its structure is intermediate between those of the two resonance forms drawn for the sulfur dioxide 2. resonance can be anticipated when it is possible to write 2 or more Lewis structures that are about equally plausible. 3. resonance forms differ only in the distribution of electrons, not in the arrangement of atoms Formal charge Often it is possible to write 2 different Lewis structures for a molecule differing in the arrangement of atoms. There are several ways to choose 2 possible structures differing in their arrangement of atoms. One is known as formal charge, which is the difference between the number of valence electrons in the free atom and the number assigned to that atom in the Lewis structure. Assigned electrons include: all the unshared electrons owned by that atom. one half of the bonding electrons shared by that atom. C f (formal charge) = X (Y + Z/2) X = number of valence electrons in the free atom Y = number of unshared electrons owned by the atom in the Lewis structure Z = number of bonding electrons shared by the atom in the Lewis structure Example:.... C O C O We will calculate the formal charges on the carbons and oxygens For C: X = 4, Y = 0, Z = 8 For C: X = 4, Y = 2, Z = 6 C f = 0 C f = -1 For O: X = 6, Y = 4, Z = 4 For O: X = 6, Y = 2, Z = 6 C f = 0 C f = +1 Ordinarily, the more likely Lewis structure is the one in which: 1. the formal charges are as close to zero as possible 2. any negative formal charge is located on the most strongly electronegative atom. Compare 1 to 2 Structure 1 C f are both 0 on structure 2 C has a -1, which is the less electronegative atom of the 2 Exceptions to the octet rule: electron deficient molecules

3 Normally odd electron species, these are sometimes called free radicals, these are impossible to write Lewis structures in which each atom obeys the octet rule. Examples are NO and NO 2 in both cases the unpaired electron goes on the nitrogen. Elementary oxygen, has experimental evidence which shows that there are 2 unpaired electrons and a double bond. There are other species where the central atom violates the octet rule, BeF 2, BF 3 and boric acid, 3 BO 3 which is an insecticide and fungicide. Exceptions: expanded octets The largest class of molecules that violate the octet rule are those that have a central atom with more than 8 electrons. Examples include PCl 5 with 10 electrons, and SF 6 with 12 electrons. In molecules of this type the terminal atoms are mostly halogens and in a few examples oxygen is the terminal atom. The central atom is a nonmetal in the third, fourth, or fifth period of the periodic table. Most frequently: P, As, Sb, S, Se, Te, Cl, Br, I, Kr, and Xe. All of these atoms have d orbitals available for bonding (3d, 4d, 5d). Because there is no 2d sublevel, C, N, and O never form expanded octets. When you look at some examples it is obvious that there is an expanded octet. owever, ClF 3, and XeF 4 look straightforward, but you can not draw a Lewis structure without having left over electrons. When this occurs they are to be placed around the central atom as unshared pairs. Molecular Geometry: The geometry of a diatomic molecule is very easy to describe in that 2 points define a straight line. Cl, 2 etc. owever, when molecules have more than 2 the geometry is not as obvious. ere bond angles, must be considered. For a molecule with the atoms YX 2 there are 2 possibilities a. linear, with a bond angle of 180 o X Y X Y b. bent, bond angles less than 180 o X X The major way to describe molecular geometries is based on electron pair repulsion. This is the basic principle of the valence shell electron-pair repulsion (VSEPR) model. According to VSEPR the valence electron pairs surrounding an atom repel one another. Consequently, the orbitals containing those electron pairs are oriented to be as far apart as possible. Ideal Geometries with 2 6 electron pairs on the central atom We will use the species AX 2 AX 6. There are no unshared pairs around atom A. Species Orientation Predicted angles Examples AX 2 Linear 180 o BeF 2 AX 3 Triangular Planar 120 o BF 3 AX 4 Tetrahedron o C 4 AX 5 Triangular bipyramid 90 o, 120 o, 180 o PF 5 AX 6 Octahedron 90 o, 180 o SF 6 Effect of unshared pairs on molecular geometry

4 1. electron pair geometry is approximately the same as that observed when single bonds are involved, ordinarily a little smaller bond angles 2. molecular geometry is quite different when one or more unshared pairs are present. When describing molecular geometry, we refer only to the positions of the bonded atoms. These can be determined experimentally; positions of unshared pairs cannot be established by experiment. The locations of unshared pairs are not specified in describing molecular geometry. Examples N 3 bond angles are 107 o due to the lone electron pair (triangular pyramid), and water with 2 lone electron pairs is bent at 105 o. The reason for this is an unshared pair is attracted by one nucleus, that of the atom to which it belongs. In contrast, a bonding pair is attracted to 2 nuclei. No. of terms Atoms (X) Species Type Ideal Bond Molecular And unshared pairs (E) Angles Geometry Examples 2 AX o Linear BeF 2, CO 2 3 AX o Triangular planar BF 3, SO 3 AX 2 E <120 o Bent GeF 2, SO 2 4 AX o Tetrahedron C 4 AX 3 E <109.5 o Triangular Pyramid N 3 AX 2 E 2 <109.5 o Bent 2 O 1. In molecules of the type AX 4 E 2, the 2 lone pairs occupy opposite rather than adjacent vertices of the octahedron 2. In molecules AX 4 E, AX 3 E 2, and AX 2 E 3 the lone pairs occupy successive positions in the equilateral triangle at the center of the triangular bipyramid Multiple bonds 1. The VSEPR model is readily extended to multiple bond species. Insofar as molecular geometry is concerned, a multiple bond behaves like a single bond. 2. This is because the shared electrons must be between the 2 atoms. 3. The geometry of multiple bond species depends on: a. the number of terminal atoms, X, bonded to the central atom, irrespective of whether the bonds are single, double or triple. b. the number of unshared pairs, E, around the central atom. 5 electron paired species Species Type Description Bond Angles Examples AX 5 Triangular bipyramidal 90 o, 120 o, 180 o PF 5 AX 4 E See-saw 90 o, 120 o, 180 o SF 4 AX 3 E 2 T-shaped 90 o, 180 o ClF 3 AX 2 E 3 Linear 180 o XeF 2 6 electron paired species Species Type Description Bond Angles Examples AX 6 Octahedral 90 o, 180 o SF 6 AX 5 E Square pyramidal 90 o, 180 o ClF 5 AX 4 E 2 Square planar 90 o, 180 o XeF 4

5 The VSEPR model applies equally well to molecules with no single central atom. For example acetylene is linear having a triple bond between the carbons. In ethylene there is a double bond between the 2 carbon atoms. The molecule has the geometry to be expected if each carbon atom had only 3 pairs of electrons around it. Describe bond angles of ethylene. 7.3 Polarity of molecules 1. Covalent bonds and molecules held together by such bonds may be a. polar as a result of unsymmetrical distribution of electrons, the bond has a positive and negative pole and is therefore a DIPOLE. 1. Occurs in bonds between unlike atoms, due to electronegativity b. nonpolar a symmetrical distribution of electrons leads to a bond or molecule without a positive or negative pole 1. occurs in molecules created by the same atom, 2 2. Polar molecules orient themselves in an electrical field, which is a measure of their dipole moment. Figure 7.11 to show how to draw polarity of molecules 3. There are 2 major criteria for determining the polarity of a molecule: bond polarity and molecular geometry. If the polar A-X bonds in a molecule AX m E n are arranged symmetrically around the central atom A, the molecule is nonpolar. a. molecules of the type AX 2 (linear), AX 3 (triangular planar), and AX 4 (tetrahedral) are nonpolar. Ex: CO 2, BF 3, CCl 4 b. molecules of the type AX 2 E (bent), AX 2 E 2 (bent), and AX 3 E (triangular pyramid) are polar. Ex: SO 2, 2 O, N Atomic orbitals; hybridization 1. The valence bond model, a covalent bond consists of a pair of electrons of opposed spin within an orbital. ( ) a. developed by Linus Pauling, Nobel prize winner in chemistry in 1954 for this model and also Nobel peace prize for his efforts to stop nuclear testing in According to this theory a full orbital with 2 opposing electrons should not form any bonds on its atom, but experiments show that this is wrong (page 184). 3. To explain these discrepancies, simple valence bond theory must be modified, hybrid orbitals. ybrid orbitals: sp, sp 2, sp 3, sp 3 d, sp 3 d 2 Ex: Be should not form any bonds because its 2s orbital is full. What occurs is the 2s orbital is hybridized with the 2p orbitals to form 2 new sp hybrid orbitals. One s atomic orbital + one p atomic orbital two sp hybrid orbitals In the BeF 2 molecule, there are 2 electron pair bonds, located in the 2 sp hybrid orbitals. In each orbital, one electron is a valence electron contributed by Be; the other from the F. Borons 3 bonds and carbons 4: One s atomic orbital + 2 p atomic orbitals 3 sp 2 hybrid orbitals One s atomic orbital + 3 p atomic orbitals 4 sp 3 hybrid orbitals Notice a pattern to the naming? In general unshared as well as shared electron pairs can be located in hybrid orbitals.

6 1. The number of hybrid orbitals formed is always equal to the number of atomic orbitals mixed. 2. the geometries, as found mathematically by quantum mechanics, are exactly as predicted by VSEPR theory. In each case, the hybrid orbitals are directed to be as far apart as possible. ybrid orbitals and their geometries # Electron Atomic ybrid orientation examples pairs orbitals orbitals 2 s, p sp linear BeF 2, CO 2 3 s, 2p sp 2 triangular planar BF 3, SO 3 4 s, 3p sp 3 tetrahedron C 4, N 3, 2 O 5 s, 3p, d sp 3 d triangular Bipyramid PCl 5, SF 4, ClF 3 6 s, 3p, 2d sp 3 d 2 octahedron SF 6, ClF 5, XeF 4 The extra electron pairs in a multiple bond (one pair in a double bond, two pairs in a triple bond) are not located in hybrid orbitals. Consider ethylene and acetylene hybridized electron pairs ethylene to all and one CC bond hybridized electron pairs on acetylene to all and one CC bond In both cases, only one of the electron pairs in the multiple bond occupies a hybrid orbital. Sigma and Pi bonds The extra electron pairs in a multiple bond are not hybridized and have no effect on molecular geometry. What happened to those electrons? Consider the spatial distribution of the orbitals filled by bonding electrons in molecules. We can distinguish between 2 types of bonding orbitals. The most common is the sigma bonding orbital, it is a single lobe, where the highest electron density is between the 2 atoms. A sigma (σ) bond consists of an electron pair occupying a sigma bonding orbital. Pi (Π) bonds, consist of 2 lobes, one above The extra electron pairs in a multiple bond are not hybridized and have no effect on molecular geometry. What happened to those electrons? the bond axis and the other below it. Along the bond axis itself, the electron density is zero. The electron pair of a pi bond occupies a pi bonding orbital. Draw a sigma and pi bond In general, to find the number of sigma and pi bonds, remember that: 1. all single bonds are sigma bonds

7 2. one of the electron pairs in a multiple bond is a sigma bond; the others are pi bonds.