Zumdahls Chapter 8 Chemical Bonding Stealing and Sharing

Zumdahl’s Chapter 8 Chemical Bonding: Stealing and Sharing of electrons

Chapter Contents • • • Types of Bonds Electronegativity Polarity & Dipoles Ions Binary Ionic Compounds • Polar Covalency • • The Covalent Bond Energies Local & Global Model G. N. Lewis Structures Octet or not Octet Resonance fudge Valence Shell Electron Pair Repulsion Theory

Types of Chemical Bonds • Ionic Bonds • Large differences in electronegativity hold atoms together by Coulombic potentials. • Radically polar poorly-shared binding electrons • Covalent Bonds • Near equally-shared, weakly polar binding pair • Dative Bonds (donor-receptor model) • Metal Bonding (macroscopic wavefunction)

Electronegativity, • Measure of ionization potential and electron affinity. The power to hold and attract. • Empirically, deviation of bond energies, DHX, from the geometrical average of DHH and DXX. • = | A – B| determines ionicity of AB • < ~1 “covalent” while > ~2 “ionic” • “Polar covalent” for between 1 and 2. • Not yet ionic, but still significantly polar.

Pauling’s Definition • DAB = [ DAA DBB ]½ + 96. 5( )² H 2. 1 Pauling Electronegativities He Li 1. 0 Be 1. 5 B 2. 0 C 2. 5 N 3. 0 O 3. 5 F 4. 0 Ne Na 0. 9 Mg 1. 2 Al 1. 5 Si 1. 8 P 2. 1 S 2. 5 Cl 3. 0 Ar soon K 0. 8 Ca 1. 0 Ga 1. 6 Ge 1. 8 As 2. 0 Se 2. 4 Br 2. 8 Kr 3. 0 Rb 0. 8 Sr 1. 0 In 1. 7 Sn 1. 8 Sb 1. 9 Te 2. 1 I 2. 5 Xe 2. 7 Cs 0. 7 Ba 0. 9 Tl 1. 8 Pb 1. 9 Bi 1. 9 Po 2. 0 At 2. 2 Rn

Dipole Moment, , & Polarity • 0 implies charge separation in bonds. • Charge separation means bond polarity: • Product of Q times RAB, bond distance. • Molecular dipole is vector sum of bond dipoles. • Symmetries cause vector cancellations! • Bond polarity is necessary but not sufficient cause of molecular polarity.

Ions • Electron configuration follows removal of electrons with highest n from parent atom. • Effective charge governs size of ion: • Cations are much smaller than parent atom. • Anions are much larger but • High charge anions are unstable except in compounds because 2 nd extra electron is repelled by 1 st anion’s negative charge.

Recipe for a Crystal • Gasify both the electropositive and the • • electronegative elements. ( Hphase ) Dissociate both to atoms. ( DXX ) Make the cation and anion. ( IP & EA ) Condense the crystal from the infinitely remote ions. ( Ulattice ) Add + and – components Hformation

Madlung Constants • Lattice energies can be calculated by crystal geometries. Nearest neighbors attract, next nearest neighbors repel, and so on forever. • Madlung constants (k = 1. 5– 4. 2) sum those all up as a scale factor for binary ionic potentials. • While O 2– can’t exist by itself (O– has a positive electron it is stabilized in Mg 2+O 2–, say, by its very negative lattice energy, 4 that of Na. Cl due to Q 1 Q 2, with which it shares a crystal structure (same k). affinity),

Polar Covalency • Most compounds have heteronuclear bonds, so 0 and the bond is polar. • But few will have so high (>~2) that their bonds are truly ionic. • Slight inequality in electron sharing makes them “polar covalent. ” • % ionic = 100% [ measured / ionic ] • Not to be confused with ionic % in .

Covalent Bonds • Electron pairs shared (more or less equally) between bonding partners. • Build-up of e– density between nuclei. • Counteracts nuclear (proton) repulsion • Increases attractions (to neighboring protons) • Stretches ; helps minimize e– kinetic energy • Reduces e– density near nucleus, reducing the e– –e– repulsions there.

Two H atoms NOT Bonding Notice the extent of overlap.

H 2 with Bonding turned ON Deeper Overlap

“Bond Energies” • While it takes a definite energy to break a particular bond in a unique molecule, “bond energies” refers to some weighted average of a particular bond in all the molecules in which it is found. • As such, they are correct for no molecule. • Still energies can be estimated from – D, for bonds changed during a reaction.

Bond Enthalpy • H ~ – D since H measures differences in the formation of compounds from elements while D shows the differences in the destruction of compounds. Opposites. • So H ~ D broken – D formed • If D formed > D broken , products are more stable than reactants, and we expect exothermic reaction, heat evolved as potential lowers.

Covalency is Local • In the Local Electron model, electrons are • Shared in pairs between adjacent atoms, • unshared entirely as lone pairs, • or hidden deep and uninvolved as core. • This ignores the reality of global electrons, free to roam over more than one atom pair. • G. N. Lewis had a patch for this difficulty with the LE model, but we need a better model.

Gilbert Newton Lewis • Developed a method to manipulate valence electrons to satisfy local atomic “needs. ” • Devised Lewis Structures and rules for the placement of electron pairs about them to • Put at least one pair between all bonded atoms. • Complete “rare gas electronic configurations” about all the molecule’s atoms. • Arrange multiple bonds (multi-electron pairs) between atoms to minimize “formal charges. ”

Lewis Structure Rules • • Sum all valence and ionic electrons = N. Pick an atomic skeletal structure. Place two of N between all bonded atoms. Distribute remainder as lone pairs to achieve an octet (only duet for hydrogen). • Minimize formal charge by stealing lone pairs to make additional (multiple) bond pairs. If FC is dumb, pick another skeleton.

Formal Charges • Shared electrons count toward BOTH atoms’ octets. • But shared electrons are divided equally between their bonded atoms for FC. • Lone pairs count fully in FC of an atom. • FC = # valence electrons – sum of above. • Best value is FC=0 for all atoms. • But FC = ion’s charge, so some won’t be 0. • Negative FC goes to highest electronegativity.

Egregious Example, NOCl • Nitrosyl chloride, “NOCl, ” is a horrific nonaqueous solvent for some food processing applications; O=N–Cl seems more likely. • Try to find a good NOCl Lewis Structure: • # = 5+6+7 = 18 valence electrons • N–O –Cl uses 4, so use 7 lone pairs (14 e–) • : N: –: O: –: Cl: : fails to deliver a N octet • : N: =O: –: Cl: : gives everyone an octet; OK?

NOCl Bombs in Formal Charge! • Since N, O, and Cl are expecting 5, 6, and 7 valence electrons, respectively, • : N: =O: –: Cl: : then shows “formal charges” of • – 1 +1 0 • putting the positive charge on the most electronegative atom? !? • It’s more likely we had the skeleton wrong, so let’s try ONCl.

ONCl, a Better Lewisite • • • O–N–Cl still needs 7 lone pairs : : O: –N: –: Cl: : still needs N octet help : : O=N: –: Cl: : satisfies octets & gives FC of 0 0 0 Perfect! So Lewis Structures permit us to correct an incorrect molecular formula into one truly reflecting the geometry, Cl. NO.

Octet Trumps Formal Charge • If you can’t get both, sacrifice good formal charges to securing an octet. For example, • CO has 4+6=10 valence electrons • C–O needs 4 more lone pairs • : C: –: O: gives nobody an octet • : C O: isoelectronic with : N N: , but FC are • – 1 +1 not what one would expect at all! • but that inobvious polarity is correct.

Breaking the Octet • “Rare Gas Configuration” is 8 valence electrons (save for He) ending in ns 2 np 6. • Atoms beyond row 2 have d orbitals which empower them to adopt more than 8. • Often “central, ” such atoms can surround themselves with typically 12 but sometimes 14 or more valence electrons in their molecules.

Meta Arsenic Acid • (HO)As. O 2 has As as its central atom. • # = 5+3(6)+1 = 24 valence electrons • : : O = As = O: : distributes all 24 • | and puts octets on all oxygens • : : O – H (and the duet on hydrogen) • with zero formal charges everywhere, • but requires 10 not 8 electrons on As. • HNO 3 exists too but has a N octet and FCs!

The Impoverished nd 2 Row • But perchloric acid, (HO)Cl. O 3 • : : O – H • | • : : O = Cl = O: : • || • : O: • has no fluorine (octet) analogue. • No 2 d orbitals & F can’t take +3 formal charge!

Newest Rare Gas Compound • HAr. F (sounds like a hairball), recently synthesized, leaves only He and Ne as truly noble gases. • But neither has d orbitals. • It’s Ar’s d orbitals that permit H–Ar: : : –F: : : with its 10 electrons around Ar. • Don’t hold your breath for Ne compounds.

Resonance • Often several different legitimate Lewis Structures are possible. NO 3– ion is all of • : O: : _ • | and | • : : O=N–O: : : O–N=O: : • with FCs of +1 (N), – 1 (O–), and 0 (O=) • representing delocalized electrons, greater stability, and uniform NO bond orders of 1. 33!

Valence Shell Electron Pair Repulsion Theory (VSEPR) • Lewis Structures show e– pairs on & between atoms. • VSEPR uses those to predict shapes as a consequence of e––e– repulsions such that • Lone pairs repel more strongly than bond pairs. • All pairs seek geometries that distance them. • Pairs in multiple bonds are treated as if they are a single bond pair for directional purposes. • Nuclei follow blindly where bonding e–s point.

# of directional electron pairs: • 1: linear A-B molecule • 4: tetrahedral • 109. 47° • 2: collinear • 180° • 3: trigonal • 120° • 5: trigonal bipyramid • 120° and 90° • 6: octahedral • 90°

Molecular Shapes • Determined only by the direction of the bonding pairs since only they terminate in atoms. • Lone pairs still dictate where bonding pairs go, but lone pair directions aren’t involved in describing molecular shapes. • A Xn Em used to count bonding directions (n) and lone pairs (m). A is central atom.

Examples of 5 Directional Pairs • AX 5 • Trigonal bipyramid • PCl 5 • AX 4 E • See Saw • SF 4 • AX 3 E 2 • T-shaped • Br. F 3 • AX 2 E 3 • Linear • I 3–

Central Atoms • … aren’t terminal atoms in the molecule. • Molecules may have several central atoms. • Geometry is determined at each by VSEPR. Acetic acid Mint Histamine Nitroglycerin