Lecture 22: Lewis Dot Structures

Lecture 22: Lewis Dot Structures • Reading: Zumdahl 13.9-13.12 • Outline • Lewis Dot Structure Basics • Resonance • Those annoying exceptions

Partial Ionic Compounds (cont.) • We can define the ionic character of bonds as follows: (dipole moment X-Y)experimental % Ionic Character = x 100% (dipole moment X+Y-)calculated

Partial Ionic Compounds (cont.) Covalent Polar Covalent Increased Ionic Character Ionic

Covalent Compounds (cont.) • The same concept can be envisioned for other covalent compounds: Think of the covalent bond as the electron density existing between the C and H atoms.

Covalent Compounds (cont.) • We can quantify the degree of stabilization by seeing how much energy it takes to separate a covalent compound into its atomic constituents.

Covalent Compounds (cont.) • Since we broke 4 C-H bonds with 1652 kJ in, the bond energy for a C-H bond is: • We can continue this process for a variety of compounds to develop a table of bond strengths.

Covalent Compounds (cont.) • Example: It takes 1578 kJ/mol to decompose CH3Cl into its atomic constituents. What is the energy of the C-Cl bond? CH3Cl: 3 C-H bonds and 1 C-C bond. 3 (C-H bond energy) + C-Cl bond energy = 1578 kJ/mol 413 kJ/mol 1239 kJ/mol + C-Cl bond energy = 1578 kJ/mol C-Cl bond energy = 339 kJ/mol

Covalent Compounds (cont.) • We can use these bond energies to determine DHrxn: DH = sum of energy required to break bonds (positive….heat into system) plus the sum of energy released when the new bonds are formed (negative….heat out from system).

Covalent Compounds (cont.) • Example: Calculate DH for the following reaction using the bond enthalpy method. CH4(g) + 2O2 (g) CO2 (g) + 2H2O (g) Go to Table 13.6: 4 x 4 x O-H 467 C=O 745 C-H 413 O=O 495 2 x 2 x

Covalent Compounds (cont.) CH4(g) + 2O2 (g) CO2 (g) + 2H2O (g) = 4D(C-H) + 2D(O=O)- 4D(O-H) - 2D(C=O) = 4(413) + 2(495) - 4(467) - 2(745) = -716 kJ/mol • Exothermic, as expected.

Covalent Compounds (cont.) CH4(g) + 2O2 (g) CO2 (g) + 2H2O (g) • As a check: 0 = DH°f(CO2(g)) + 2DH°f(H2O(g)) - DH°f(CH4(g)) - 2 DH°f(O2(g)) = -393.5 kJ/mol + 2(-242 kJ/mol) - - (-75 kJ/mol) = -802.5 kJ/mol

Localized Bond Models • Consider our energy diagram for H2 bonding:

Localized Model Limitations • It is important to keep in mind that the models we are discussing are just that…..models. • We are operating under the assumption that when forming bonds, atoms “share” electrons using atomic orbitals. • Electrons involved in bonding: “bonding pairs”. Electrons not involved in bonding: “lone pairs”.

Lewis Dot Structures (cont.) • Developed by G. N. Lewis to serve as a way to describe bonding in polyatomic systems. • Central idea: the most stable arrangement of electrons is one in which all atoms have a “noble” gas configuration. • Example: NaCl versus Na+Cl- Na: [Ne]3s1 Cl: [Ne]3s23p5 Na+: [Ne] Cl-: [Ne]3s23p6 = [Ar]

LDS Mechanics • Atoms are represented by atomic symbols surrounded by valence electrons. Lone Pair (6 x) • Electron pairs between atoms indicate bond formation. Bonding Pair

LDS Mechanics (cont.) • Three steps for “basic” Lewis structures: Sum the valence electrons for all atoms to determine total number of electrons. Use pairs of electrons to form a bond between each pair of atoms (bonding pairs). Arrange remaining electrons around atoms (lone pairs) to satisfy the “octet rule” (“duet” rule for hydrogen).

LDS Mechanics (cont.) • An example: Cl2O 20 e- 16 e- left

LDS Mechanics (cont.) • An example: CH4 8 e- 0 e- left Done!

LDS Mechanics (cont.) • An example: CO2 16 e- 12 e- left Octet Violation 0 e- left CO double bond

LDS Mechanics (cont.) + • An example: NO+ + 10 e- 8 e- left +

Resonance Structures • We have assumed up to this point that there is one correct Lewis structure. • There are systems for which more than one Lewis structure is possible: • Different atomic linkages: Structural Isomers • Same atomic linkages, different bonding: Resonance

Resonance Structures (cont.) • The classic example: O3. Both structures are correct!

Resonance Structures (cont.) • In this example, O3 has two resonance structures: • Conceptually, we think of the bonding being an average of these two structures. • Electrons are delocalized between the oxygens such that on average the bond strength is equivalent to 1.5 O-O bonds.

Structural Isomers • What if different sets of atomic linkages can be used to construct correct LDSs: • Both are correct, but which is “more” correct?

Formal Charge • Formal Charge: Compare the nuclear charge (+Z) to the number of electrons (dividing bonding electron pairs by 2). Difference is known as the “formal charge”. #e- 7 6 7 7 6 7 Z+ 7 6 7 7 7 6 Formal C. 0 0 0 0 +1 -1 • Structure with less F. C. is more correct.

Formal Charge • Example: CO2 e- 6 4 6 6 4 6 7 4 5 Z+ 6 4 6 6 6 4 6 6 4 FC 0 0 0 0 +2 -2 -1 +2 -1 More Correct

Beyond the Octet Rule • There are numerous exceptions to the octet rule. • We’ll deal with three classes of violation here: • Sub-octet systems • Valence shell expansion • Odd-electron systems

Beyond the Octet Rule (cont.) • Some atoms (Be and B in particular) undergo bonding, but will form stable molecules that do not fulfill the octet rule. • Experiments demonstrate that the B-F bond strength is consistent with single bonds only.

Beyond the Octet Rule (cont.) • For third-row elements (“Period 3”), the energetic proximity of the d orbitals allows for the participation of these orbitals in bonding. • When this occurs, more than 8 electrons can surround a third-row element. • Example: ClF3 (a 28 e- system) F obey octet rule Cl has 10e-

Beyond the Octet Rule (cont.) • Finally, one can encounter odd electron systems where full pairs will not exist. • Example: Chlorine Dioxide. Unpaired electron

Summary • Remember the following: • C, N, O, and F almost always obey the octet rule. • B and Be are often sub-octet • Second row (Period 2) elements never exceed the octet rule • Third Row elements and beyond can use valence shell expansion to exceed the octet rule. • In the end, you have to practice…..a lot!

Lecture 22: Lewis Dot Structures