Chapter 9 Chemical Bonding I: Lewis Theory

Determining the Number of Valence Electrons in an Atom the column number on the Periodic Table will tell you how many valence electrons a main group atom has Transition Elements all have 2 valence electrons; Why? 1A 2A 3A 4A 5A 6A 7A 8A Li Be B C N O F Ne 1 e-1 2 e-1 3 e-1 4 e-1 5 e-1 6 e-1 7 e-1 8 e-1

Lewis Symbols of Atoms aka electron dot symbols use symbol of element to represent nucleus and inner electrons use dots around the symbol to represent valence electrons pair first two electrons for the s orbital put one electron on each open side for p electrons then pair rest of the p electrons

Lewis Symbols of Ions Li• Li+1 Cations have Lewis symbols without valence electrons Lost in the cation formation Anions have Lewis symbols with 8 valence electrons Electrons gained in the formation of the anion Li• Li+1

Stable Electron Arrangements And Ion Charge Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas The noble gas electron configuration must be very stable

Rules when atoms bond, they tend to gain, lose, or share electrons to result in 8 valence electrons ns2np6 noble gas configuration Duet Rule: sharing of 2 electrons E.g H2 H : H Octet Rule: sharing of 8 electrons Carbon, oxygen, nitrogen and fluorine always obey this rule in a stable molecule E.g F2, O2

Exceptions many exceptions H, Li, Be, B attain an electron configuration like He He = 2 valence electrons Li loses its one valence electron H shares or gains one electron though it commonly loses its one electron to become H+ Be loses 2 electrons to become Be2+ though it commonly shares its two electrons in covalent bonds, resulting in 4 valence electrons B loses 3 electrons to become B3+ though it commonly shares its three electrons in covalent bonds, resulting in 6 valence electrons expanded octets for elements in Period 3 or below using empty valence d orbitals

Lewis Theory the basis of Lewis Theory is that there are certain electron arrangements in the atom that are more stable octet rule bonding occurs so atoms attain a more stable electron configuration more stable = lower potential energy no attempt to quantify the energy as the calculation is extremely complex Bonding pair: two of which are shared with other atoms Lone pair or nonbonding pair: those that are not used for bonding

Electron-Dot Structures Chapter 7: Covalent Bonds and Molecular Structure 4/26/2017 Electron-Dot Structures H • O •• H O •• Think of this section as an introduction. It is much easier to write electron-dot structures using the rules listed in the next section. Copyright © 2008 Pearson Prentice Hall, Inc.

Electron-Dot Structures Chapter 7: Covalent Bonds and Molecular Structure 4/26/2017 Electron-Dot Structures We have single, double, and triple bonds. Copyright © 2008 Pearson Prentice Hall, Inc.

Covalent Bonding Predictions from Lewis Theory Lewis theory allows us to predict the formulas of molecules Lewis theory predicts that some combinations should be stable, while others should not because the stable combinations result in “octets” Lewis theory predicts in covalent bonding that the attractions between atoms are directional the shared electrons are most stable between the bonding atoms resulting in molecules rather than an array

Electronegativity measure of the pull an atom has on bonding electrons increases across period (left to right) and decreases down group (top to bottom) fluorine is the most electronegative element francium is the least electronegative element the larger the difference in electronegativity, the more polar the bond negative end toward more electronegative atom

Bond Polarity covalent bonding between unlike atoms results in unequal sharing of the electrons one atom pulls the electrons in the bond closer to its side one end of the bond has larger electron density than the other the result is a polar covalent bond bond polarity the end with the larger electron density gets a partial negative charge the end that is electron deficient gets a partial positive charge

Polar Covalent Bonds: Electronegativity Chapter 7: Covalent Bonds and Molecular Structure 4/26/2017 Polar Covalent Bonds: Electronegativity NaCl Cl2 HCl Copyright © 2008 Pearson Prentice Hall, Inc.

Electronegativity and Bond Polarity If difference in electronegativity between bonded atoms is 0, the bond is pure covalent equal sharing If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent If difference in electronegativity between bonded atoms 0.5 to 1.9, the bond is polar covalent If difference in electronegativity between bonded atoms larger than or equal to 2.0, the bond is ionic 0.4 2.0 4.0 4% 51% Percent Ionic Character Electronegativity Difference “100%”

Bond Dipole Moments 1 D = 3.34 x 10-30 C•m the dipole moment is a quantitative way of describing the polarity of a bond a dipole is a material with positively and negatively charged ends measured dipole moment, m, is a measure of bond polarity it is directly proportional to the size of the partial charges and directly proportional to the distance between them m = (q)(r) r = radius q = 1.6 x 10-19 C 1 D = 3.34 x 10-30 C•m

Polarity and Dipole Moment a vector quantity from the center of the positive charge to the center of negative charge Represents with an arrow

Example Determine whether bond formed between the following pair is ionic, covalent, or polar covalent N and O Sr and F N and Cl E.g Draw the dipole moment for HF HCl OF

Lewis Structures of Molecules shows pattern of valence electron distribution in the molecule useful for understanding the bonding in many compounds allows us to predict shapes of molecules allows us to predict properties of molecules and how they will interact together

Rules for writing Dots Lewis structures Write the correct skeletal structure for molecule Least electronegative atom will be in the center Hydrogen will always be the terminal Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule If polyatomic ions, charges must be considered when calculating the total valence electrons Distribute the electrons among the atoms, giving octets (or duet for hydrogen) to as many atoms as possible If any atoms lack an octet, form double or triple bonds as necessary to give them octets.

Lewis Structures B C N O F use common bonding patterns C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs often Lewis structures with line bonds have the lone pairs left off their presence is assumed from common bonding patterns structures which result in bonding patterns different from common have formal charges B C N O F

FC = valence e- - nonbonding e- - ½ bonding e- Formal Charge during bonding, atoms may wind up with more or less electrons in order to fulfill octets - this results in atoms having a formal charge FC = valence e- - nonbonding e- - ½ bonding e- sum of all the formal charges in a molecule = 0 in an ion, total equals the charge

Examples Draw a Lewis formula then assign formal charge for the following molecules and/or ions HBr CO2 NH4+ SO32-

Resonance when there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures the actual molecule is a combination of the resonance forms – a resonance hybrid it does not resonate between the two forms, though we often draw it that way look for multiple bonds or lone pairs

Rules of Resonance Structures Resonance structures must have the same connectivity only electron positions can change Resonance structures must have the same number of electrons Second row elements have a maximum of 8 electrons bonding and nonbonding third row can have expanded octet Formal charges must total same Better structures have fewer formal charges Better structures have smaller formal charges Better structures have − formal charge on more electronegative atom

Drawing resonance Any compound for which more than one Lewis structure may be written is accurately described by no single structure. The actual structure is a resonance hydrid of them all (NOT “flipping back and forth” between resonance forms). The various structures are called contributing structure or resonance forms

Drawing Resonance Structures draw first Lewis structure that maximizes octets assign formal charges move electron pairs from atoms with (-) formal charge toward atoms with (+) formal charge if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet. -1 -1 +1 -1 +1

Exceptions to the Octet Rule expanded octets elements with empty d orbitals can have more than 8 electrons odd number electron species e.g., NO will have 1 unpaired electron free-radical very reactive incomplete octets B, Al

Drawing Resonance Structures -1 draw first Lewis structure that maximizes octets assign formal charges move electron pairs from atoms with (-) formal charge toward atoms with (+) formal charge if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet. +2 -1

Examples Identify Structures with Better or Equal Resonance Forms and Draw Them O3 NO2- PO43-