Covalent Bonding Chapter 7 Basics of structure • Introduction – Lewis dot structures are misleading, for example H – H, 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. However, they are more likely to be found between them than at the far ends. Comparing H-H interactions • 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 H2 molecule. When they are forced closer together, the repulsion forces are too great and the energy rises. 1 • 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 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 H – F, Fluorine has 7 normal electrons from its configuration 1s 22s 22p 5 and shares one with Hydrogen 2 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) • 3. octet rule – nonmetals, except for hydrogen, achieve a noble-gas configuration by sharing an octet of electrons, hydrogen shares 2 Lewis Rules examples • Examples of single bonds hydroxide ion, water, ammonia, ammonium ion • double bonds ethylene C2H4 • triple bonds acetylene C2H2 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 • 4. Determine the number of valence electrons required to fill out an octet for each atom (except H) in the skeleton 3 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 SO2 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. 4 Formal Charge continued • Assigned electrons include: • all the unshared electrons owned by that atom. • one half of the bonding electrons shared by that atom. •Cf (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 Examples H | .. .. H C O H H C O H | ⋅⋅ | | H H H Calculating their formal charge • 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 •Cf = 0 Cf = -1 • For O: X = 6, Y = 4, Z = 4 For O: X = 6, Y = 2, Z = 6 •Cf = 0 Cf = +1 5 Comparing their F. Charges • 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 Cf are both 0, on structure 2 C has a -1, which is the less electronegative atom of the 2 As always exceptions • Exceptions to the octet rule: electron deficient molecules • 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, H 3BO 3 which is an insecticide and fungicide. More exceptions • 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 PCl5 with 10 electrons, and SF6 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. 6 Most common expanded octets • 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. However, 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 Geometries • The geometry of a diatomic molecule is very easy to describe in that 2 points define a straight line. HCl, H2 etc. However, when molecules have more than 2 the geometry is not as obvious. Here bond angles, must be considered. For a molecule with the atoms • YX 2 there are 2 possibilities 2 types explained a. linear, with a bond angle of 180 o X—Y—X Y b. bent, bond angles less than 180 o X X 7 VSEPR theory • 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. Types of molecules • 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 o • AX 2 Linear 180 BeF 2 o • AX 3 Triangular Planar 120 BF 3 o • AX 4 Tetrahedron 109.5 CH4 o o o • AX 5 Triangular bipyramid 90 , 120 , 180 PF 5 o o • AX 6 Octahedron 90 , 180 SF 6 Effect of unshared pairs on molecular geometry • 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. 8 Continued • 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. o • Examples NH 3 bond angles are 107 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. Molecular Geometries • No. of terms Atoms (X) Species Ideal Bond Molecular • And unshared pairs (E) Type Angles Geometry Examples o • 2 AX 2 180 Linear BeF 2, CO 2 o • 3 AX 3 120 Triangular planar BF 3, SO 3 o • AX 2E <<<120 Bent GeF 2, SO 2 o • 4 AX 4 109.5 Tetrahedron CH 4 o • AX 3E <<<109.5 Triangular Pyramid NH 3 < o • AX 2E2 <<109.5 Bent H2O How this works out • 1. In molecules of the type AX 4E2, the 2 lone pairs occupy opposite rather than adjacent vertices of the octahedron • 2.