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CHAPTER 6 Representing Molecules

CHAPTER 6 Representing Molecules (Note: Sections 6.1 to 6.3 are covered on Exam2; Sections 6.4 to 6.6 are covered on Exam 3). Octet Rule - Covalent Bonding

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CHAPTER 6 Representing Molecules

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  1. CHAPTER 6 Representing Molecules (Note: Sections 6.1 to 6.3 are covered on Exam2; Sections 6.4 to 6.6 are covered on Exam 3)

  2. Octet Rule - Covalent Bonding Octet rule - Main group elements tend to gain or lose electrons so that they end up with a noble gas electron configuration. This will either completely fill or completely empty the valence shell of electrons. For most atoms, filling the valence shell means having a total of eight electrons - hence, the name octet rule. Hydrogen (H) will tend to add one electron to obtain the same electron configuration as He. For covalently bonded molecules the octet rule means that atoms will share one or more pairs of electrons. Atoms then achieve a filled valence shell by using the shared electrons. There are exceptions to the octet rule in covalent bonding, but for many molecules, particularly those in the first two rows of the periodic table, the rule is usually followed.

  3. Covalent Bonding and Lewis Structures Covalent bonding is bonding of atoms by the sharing of one or more pairs of electrons. The representation of the arrangement of electrons in a covalently bond-ed molecule is called a Lewis structure. In a Lewis structure: 1) Bonding electron pairs are indicated by lines. 2) Nonbonding electrons, called lone pair electrons, are indicated by dots.

  4. Multiple Bonds It is possible for atoms to share more than one pair of electrons. Consider the diatomic molecules F2, O2, and N2. F 1s2 2s2 2p5 O 1s2 2s2 2p4 N 1s2 2s2 2p3 Bond order = number of pairs of shared electrons.

  5. Polyatomic Molecules and Ions Lewis structure for polyatomic molecules indicate which atoms are bonded together, the bond orders for these bonds, and the number of lone pairs electrons. Lewis structures do not directly indicate molecular geometry (the arrangement of atoms in three dimensions). As previously noted, we often write molecular formulas in such a way that they provide information about the arrangement of atoms in a polyatomic molecule.

  6. There are often several ways in which a Lewis structure may be drawn for a particular molecule. Example: CH3COOH (acetic acid) Although these Lewis structures look different, they are all the same, and communicate the same information about the bonding in acetic acid.

  7. Average Bond Length The average bond length is the average length of a particular type of bond (single, double, triple) between two atoms. In general, the average bond length for a covalent bond between two atoms decreases as we go from a single to a double to a triple bond. bond N-N N=N NN length (nm) 0.147 0.124 0.110 Bond strength is also related to the bond order. A double bond between two atoms is stronger than a single bond, and a triple bond is stronger than a double bond.

  8. Average Bond Lengths (Table)

  9. Lewis Structures For Ions Lewis structures for ions are drawn the same way as for molecules, except that the ion is placed within brackets, with the charge of the ion shown outside the brackets. NO2+ NH4+

  10. Comparison of Ionic and Covalent Compounds IonicCovalent Bonding by transfer of electrons Bonding by sharing of electrons Bonding is isotropic Bonding is directional Do not exist as molecules Exists as molecules Solids at room temperature May be solid, liquid, or gas at room temperature High melting point Low melting point High boiling point Low boiling point Strong electrolytes Usually nonelectrolytes

  11. Bond Polarity A covalent bond represents the sharing of one or more pairs of electrons. However, the electron pairs are not necessarily equally shared between the two bonded atoms. We call a bond where there is an unequal sharing of electrons a polar covalent bond. There are three general cases: 1) Equal sharing - Atoms bonded by electrons that are equally shared. The bond is nonpolar. 2) Unequal sharing - Atoms bonded by electrons that are unequally shared. The bond is polar. 3) Ionic bonding - Ions are formed from the transfer of one or more electrons to form cations and anions. There is nosharing of electrons in ionic bonding.

  12. Bond Polarity - Examples H - H H+ - F- [ Na+] [F-] A covalent bond formed by electrons that are not equally shared is called a polar bond. In such a bond, one atom achieves a partial positive charge, and the other atom a partial negative charge.

  13. Electronegativity Electronegativity (EN) is a number (between 0 - 4) that is assigned to an element that represents the tendency of atoms of that element to attract electrons. The larger the value for electronegativity the greater the tendency for atoms of that element to attract electrons. Note that a value for electronegativity is not assigned to most noble gases. The general trends for electronegativity are as follows: 1) Within a group of elements electronegativity increases from bottom to top. 2) Within a period electronegativity increases from left to right. For a polar covalent bond between two atoms, the atom with the larger value for electronegativity will have a partial negative charge (-) and the atom with the smaller value for electronegativity will have a partial positive charge (+).

  14. Electronegativity Chart

  15. Use of Electronegativity The difference in electronegativity between two atoms can be used to predict the type of bonding that exists between the atoms. EN < 0.5 Bond is nonpolar or only slightly polar. EN 0.5 - 2.0 Bond is polar covalent, more electronegative atom has a partial negative charge EN > 2.0 Bond is ionic; more electronegative atom forms an anion Above guidelines are approximate. Note that electronegativity is also related to metallic character. A small value for electronegativity indicates a metallic element, and a large value indicates a nonmetallic element.

  16. EN = 2.1 EN = 0.9 EN = 0.0 ionic polar covalent nonpolar covalent

  17. Dipole Moment () Dipole moment () is a measure of how polar a bond or a molecule is. By definition, dipole moment is given by the expression  = Qr Q = size of separated +/- charge r = distance of separation of charge As q increases and r increases  also increases. Dipole moment is measured in units of Debye 1 Debye = 3.336 x 10-30 Cm |e-| = 1.602 x 10-19 C Molecule EN  (Debye) F2 0.0 0.00 HF 1.9 1.82 LiF 3.0 6.33

  18. Percent Ionic Character The percent ionic character is a measure of the approximate amount of ionic character in a bond between two atoms. It is defined as % ionic character = (observed) 100 % (assuming discrete charges) where (observed) = experimentally observed dipole moment (assuming discrete charges) = dipole moment calculated using the bond distance and assuming complete electron transfer

  19. Percent Ionic Character - Example The experimental values for dipole moment and bond distance for the molecule HCl are  = 1.08 D r = 0.127 nm What is the percent ionic character of the bond?

  20. The experimental values for dipole moment and bond distance for the molecule HCl are  = 1.08 D r = 0.127 nm What is the percent ionic character of the bond? (assuming discrete charges) = dipole moment calculated using the bond distance and assuming complete electron transfer = (1)(1.602 x 10-19 C)(0.127 x 10-9 m) 1 D 3.336 x 10-30 Cm = 6.10 D % ionic character = 1.08 D 100% = 18% 6.10 D

  21. Electronegativity and Percent Ionic Character

  22. Guidelines For Drawing Lewis Structures The following general guidelines are useful in drawing Lewis structures for molecules and ions. 1) The central atom is usually the least electronegative atom (excluding hydrogen, which will never be the central atom). 2) For molecules or ions that obey the octet rule number of bonds = (# e- needed for octet rule) - (# valence e-) 2 3) There is usually one electron from each atom making a covalent bond. 4) Common number of covalent bonds formed H - 1 bond (no exceptions) O - 2 bonds F - 1 bond (no exceptions) N - 3 bonds Cl, Br, I - 1 bond C - 4 bonds (almost no exceptions)

  23. Example: Draw the Lewis structure for the following molecules a) NOF b) CH2O c) CH3CHO

  24. Example: Draw the Lewis structure for the following molecules a) NOF b) CH2O c) CH3CHO NOF N 5 valence e- central atom = N (least electronegative) O 6 valence e- F 7 valence e- 18 valence e- total (8 + 8 + 8) = 24 e- needed for octet rule number of covalent bonds = [ 24 - 18 ]/2 = 6/2 = 3

  25. Examples of Covalent Bonding NOF central atom = N # bonds = ( 24 - 18 ) = 3 2 CH2O central atom = C # bonds = ( 20 - 12) = 4 2 CH3CHO # bonds = ( 32 - 18) = 7 2

  26. Organic Molecules For organic molecules, we can often use the way in which the formula for the molecule is written as a guide to its Lewis structure. Example: What is the Lewis structure for diethyl ether, whose chemical formula is CH3CH2OCH2CH3?

  27. For organic molecules, we can often use the way in which the formula for the molecule is written as a guide to its Lewis structure. Example: What is the Lewis structure for diethyl ether, whose chemical formula is CH3CH2OCH2CH3?

  28. End of Material For Second Hour Exam

  29. Coordinate Covalent Bond In most covalent bonds each atom contributes the same number of electrons to the bond. In a coordinate covalent bond, (sometimes called a dative bond) both electrons come from the same atom. One example of a coordinate covalent bond is in the ammonium ion (NH4+). Coordinate covalent bonds can form when there is an atom that has one or more available lone pairs of electrons. Note that in the ammonium ion nitrogen is making more bonds than usual, and that all four of the N-H bonds are identical.

  30. Formal Charge Formal charge (FC) is a number that represents in an approximate way the electron density around a particular atom in a molecule or ion. Note that it does not represent the real charge on the atom - it is a bookkeeping device to keep track of electron density in the molecule or ion. It is similar to, but not the same, as oxidation number. Formal charge is assigned as follows: FC = (# valence e- in atom) - [ (# nonbonding e-) + 1/2 (# bonding e-) ] The sum of the formal charges must add up to the charge of the molecule or ion. Example: FC(N) = 5 - [ 0 + 1/2 (8) ] = +1 FC(H) = 1 - [ 0 + 1/2 (2) ] = 0

  31. Use of Formal Charge Formal charge can be used to determine which structure is most important in representing a molecule or ion for cases where the resonance structures are not equivalent to one another. The best structure is the one which: 1) Makes the formal charges of all atoms as close to zero as possible. 2) If there are nonzero formal charges, places negative formal charges on the more electronegative atoms and positive formal charges on the less electronegative atoms. Example: What is the best structure for the molecule N2O?

  32. Example: Which of the following structures for N2O is the best structure? # bonds = (24 - 16) = 4 2

  33. Example: Which of the following structures for N2O is the best structure? # bonds = (24 - 16) = 4 2 0 +1 -1 -1 +1 0 -2 +1 +1 The structure on the right can be ruled out because it has a formal charge that is more different from zero than the other two structures. Of the two remaining structures, the one on the left places the negative formal charge on oxygen, the more electronegative atom, and so is better than the one in the middle.

  34. Resonance Structures It is not always possible to represent bonding in a molecule or ion with a single Lewis structure. Example: What is the Lewis structure for O3?

  35. Example: What is the Lewis structure for O3? Experimentally, the two bonds in O3 are the same, and inter-mediate between a single and a double bond. That suggests the real structure is some combination of the above two Lewis structures. Resonance structure - Two or more valid Lewis structures for a molecule or ion. The actual structure is an “average” of the resonance structures. We use an arrow () to indicate the Lewis structures that contribute to the representation of the molecule or ion.

  36. Resonance structures differ only in the arrangement of electrons, and not in the arrangement of the atoms making up the molecule or ion. Example: NO3- (nitrate ion) # bonds = ( 32 - 24) = 4 N = central atom 2 In this case there are three equivalent Lewis structures that differ only in the location of the N=O double bond. The N-O bond is equal to 1 1/3 of a covalent bond. Note that in the above structures N has a formal charge of +1, both single bonded oxygens have a formal charge of -1, and the double bonded oxygen has a formal charge of 0.

  37. Resonance Structures for Benzene (C6H6) A molecule of great importance in organic chemistry is the benzene molecule, C6H6. Experimentally, the molecule is found to have a ring structure. Benzene can be represented by two resonance structures. Many organic molecules contain one or more of the above benzene rings. Such molecules are classified as aromatic molecules.

  38. Formal Charge and Resonance Structures We can sometimes use formal charge to identify the most important resonance structure for a molecule or ion. Consider our previous case of the N2O molecule. 0 +1 -1 -1 +1 0 -2 +1 +1 The average structure will have a larger contribution from the resonance structure on the left, and a smaller contribution from the resonance structure on the right, based on the formal charges observed.

  39. Exceptions to the Octet Rule While the octet rule works well for many substances, there are several important exceptions to the general rule. 1) Compounds of beryllium (Be) and boron (B). Beryllium (which often forms covalent compounds) usually makes two covalent bonds, boron usually makes three covalent bonds. Be 1s2 2s2 B 1s2 2s2 2p1

  40. 2) Molecules or ions with an odd number of electrons. In this case, it is impossible to pair up the electrons so that every atom satisfies the octet rule. In this case, the less electronegative atom generally is the one that will be one electron short of an octet. Examples: NO and ClO.

  41. Examples: NO and ClO. NO ClO (16-11)/2 = 2.5 (16-13)/2 = 1.5 3.0 3.5 3.0 3.5 Compounds with an odd number of electrons are usually very reactive, as they would like to acquire an additional electron to satisfy the octet rule.

  42. 3) Elements below the second period of the periodic table. Since these elements have d orbitals in addition to their s and p orbitals, there is room in the valence shell for more than 8 electrons. The elements can form compounds with an expanded octet (more than 8 valence electrons). N 1s2 2s2 2p3 P 1s2 2s2 2p6 3s2 3p3 3d0 NF3 PF3 PF5 NF5 - does not occur We usually consider Lewis structures that contain atoms with more than an octet of electrons only if no reasonable structure obeying the octet rule can be found. We also use formal charge to guide us (discussed later).

  43. Example: What are the Lewis structures for SF2, SF4, and SF6?

  44. Example: What are the Lewis structures for SF2, SF4, and SF6?

  45. Formal Charge and the Octet Rule We sometimes see cases where a resonance structure that is better based on formal charge is worse based on satisfying the octet rule. In these cases, how do we decide which resonance structure is more important?

  46. All the atoms in the structure at left satisfy the octet rule. Sulfur violates the octet rule in the structure at right, but the formal charges are all equal to zero. So both are likely to be important.

  47. Bonding in Metals Metals have a general tendency to give up electrons to form cations. This indicates that the valence electrons in a metallic element are only loosely attracted to the nucleus of the metal. A simple model for bonding in metals is to treat the valence electrons as forming a “sea” of electrons which can easily move about in the metal. Because of this, this suggests that it should be easy to move electrons through the volume of the metal. In fact, metals are good conductors of electricity because these loosely bound electrons can easily move through the metal. The above also accounts for the malleability and ductility of metals. Because there are usually no strong localized bonds in metals, it becomes easy to alter their shape (hammer them into thin sheets or draw them into wires).

  48. End of Chapter 6 “The underlying physical laws necessary for the mathematical theory of a large part of physics and the whole of chemistry are thus completely known, and the difficulty is only that the application of these laws leads to equations much too complicated to be soluble.” - P. A. M. Dirac “The great importance in Lewis’ theory is that it provided chemists with a valuable way of visualizing the electronic structures of atoms and molecules, and for practical purposes his ideas are still used today.” - Keith J. Laidler “My name is Bond - Covalent Bond.” - anonymous

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