AP Notes Chapter 8 Bonding and Molecular Structure: Fundamental Concepts Valence e- and Bonding Covalent Ionic Resonance & Exceptions to Octet Rule Bond Energy & Length Structure, Shape & Polarity of Compounds

What is a Bond? A force that holds atoms together. Why? We will look at it in terms of energy. Bond energy the energy required to break a bond. Why are compounds formed? Because it gives the system the lowest energy.

Covalent compounds? The electrons in each atom are attracted to the nucleus of the other. The electrons repel each other, The nuclei repel each other. The reach a distance with the lowest possible energy. The distance between is the bond length.

Thus Hydrogen is Diatomic! Bond Formation

Covalent Character e-

Why Isn’t Helium Diatomic? . . E He + He He2 Inter-nuclear Distance . .

2p ____ ____ ___ ___ ____ ____ 2p 2s ____ ____ 2s F + F F2 2p ____ ____ ___ ___ ____ ____ 2p 2s ____ ____ 2s F F

Ionic Bonding An atom with a low ionization energy reacts with an atom with high electron affinity. The electron moves. Opposite charges hold the atoms together.

Li + Cl 1s22s1 [Ne] 3s23p5 2s ___ 3p _____ _____ ___ 1s _____ 3s _____ [Ne]

Li + Cl 2s ___ 3P _____ _____ _____ 1s _____ 3s _____ [Ne]

LiCl 2s ___ 3P _____ _____ _____ 1s _____ 3s _____ [Ne]

Electronegativity Describes the relative ability of an atom within a molecule to attract a shared pair of electrons to itself.

Pauling electronegativity values, which are unit-less, are the norm.

Electronegativity Range from 0.7 to 4.0 Figure 9.9 – Kotz & Treichel

Bond: A - B DEN = | ENA - ENB |

Bond Character “Ionic Bond” - Principally Ionic Character “Covalent Bond” - Principally Covalent Character

Determining Principal Character of Bond covalent ionic EN ~0 ~4 1.7

F - F EN = 0 Non-polar

N - O EN = |3.0 - 3.5| = 0.5 O N Slightly polar

Ionic Bond with some covalent character Ca - O  EN = |1.0 - 3.5| = 2.5 Ca O Ionic Bond with some covalent character

Electronegativity The ability of an electron to attract shared electrons to itself. Pauling method Imaginary molecule HX Expected H-X energy = H-H energy + X-X energy 2 D = (H-X) actual - (H-X)expected

Electronegativity D is known for almost every element Gives us relative electronegativities of all elements. Tends to increase left to right. decreases as you go down a group. Noble gases aren’t discussed. Difference in electronegativity between atoms tells us how polar.

Electronegativity difference Bond Type Zero Covalent Covalent Character decreases Ionic Character increases Polar Covalent Intermediate Ionic Large

Dipole Moments A molecule with a center of negative charge and a center of positive charge is dipolar (two poles), or has a dipole moment. Center of charge doesn’t have to be on an atom. Will line up in the presence of an electric field.

How It is drawn H - F d+ d-

Which Molecules Have Them? Any two atom molecule with a polar bond. With three or more atoms there are two considerations. There must be a polar bond. Geometry can’t cancel it out.

Ionic Radii -- Cations

Ionic Radii -- Anions

Vector Sum of Bond Polarities Molecular Polarity Vector Sum of Bond Polarities

Covalent BOND w/much ionic character, BUT NON-POLAR molecule MgBr2 Mg - Br EN = |1.2 - 2.8| = 1.6 Mg Br Br Covalent BOND w/much ionic character, BUT NON-POLAR molecule

Lewis Structures

The most important requirement for the formation of a stable compound is that the atoms achieve noble gas e- configuration

Valence Shell Electron Pair Repulsion Model (VSEPR) The structure around a given atom is determined principally by minimizing electron-pair repulsions

VSEPR Electron pairs Bond Angles Underlying Shape 2 180° Linear 3 120° Trigonal Planar 4 109.5° Tetrahedral 5 90° & 120° Trigonal Bipyramidal 6 90° Octagonal

LEWIS STRUCTURES : draw skeleton of species : count e- in species : subtract 2 e- for each bond in skeleton : distribute remaining e-

Distinguish Between ELECTRONIC Geometry & MOLECULAR

CH4 Bond angle = 109.50 Electronic geometry: tetrahedral Molecular geometry: tetrahedral

H3O+ Bond angle ~ 1070 Electronic geometry: tetrahedral Molecular geometry: trigonal pyramidal

H2O Bond angle ~ 104.50 Electronic geometry: tetrahedral Molecular geometry: bent

NH2- Bond angle ~ 104.50 Electronic geometry: tetrahedral Molecular geometry: bent

“Octet Rule” holds for connecting atoms, but may not for the central atom.

BaI2 Bond angle =1800 Electronic geometry: linear Molecular geometry: linear

BF3 Bond angle =1200 Electronic geometry: trigonal planar Molecular geometry: trigonal planar

PF5 Bond angle = 1200 & 900 Electronic geometry: trigonal bipyramidal Molecular geometry: trigonal bipyramidal

SF4 Bond angle = 1200 & 900 Electronic geometry: trigonal bipyramidal Molecular geometry: see-saw

ICl3 Bond angle <= 900 Electronic geometry: trigonal bipyramidal Molecular geometry: T-shape

I3- Bond angle = 1800 Electronic geometry: trigonal bipyramid Molecular geometry: linear

PCl6- Bond angle = 900 Electronic geometry: octahedral Molecular geometry: octahedral

BrF5 Bond angle ~ 900 Electronic geometry: octahedral Molecular geometry: square pyramidal

ICl4- Bond angle = 900 Electronic geometry: octahedral Molecular geometry: square planar

Actual shape Non-BondingPairs ElectronPairs BondingPairs Shape 2 2 linear 3 3 trigonal planar 3 2 1 bent 4 4 tetrahedral 4 3 1 trigonal pyramidal 4 2 2 bent

Actual Shape Non-BondingPairs ElectronPairs BondingPairs Shape 5 5 trigonal bipyrimidal 5 4 1 See-saw 5 3 2 T-shaped 5 2 3 linear

Actual Shape Non-BondingPairs ElectronPairs BondingPairs Shape 6 6 Octahedral 6 5 1 Square Pyramidal 6 4 2 Square Planar 6 3 3 T-shaped 6 2 1 linear

What happens when there are not enough electrons to “satisfy” the central atom?

EXAMPLES Ethene Acetic Acid Oxygen Nitrogen

RESONANCE & FORMAL CHARGE

Resonance Sometimes there is more than one valid structure for an molecule or ion. NO3- Use double arrows to indicate it is the “average” of the structures. It doesn’t switch between them. NO2- Localized electron model is based on pairs of electrons, doesn’t deal with odd numbers.

EXAMPLES Nitrate ion Ozone

FORMAL CHARGE the charge assigned to an atom in a molecule or polyatomic ion FC atom = Family# - [LPE + ½(BE)] Sum FC’s atoms = ion charge Closer sum FC’s is to zero more stable

Formal Charge For molecules and polyatomic ions that exceed the octet there are several different structures. Use charges on atoms to help decide which. Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.

Formal Charge The difference between the number of valence electrons on the free atom and that assigned in the molecule. We count half the electrons in each bond as “belonging” to the atom. SO4-2 Molecules try to achieve as low a formal charge as possible. Negative formal charges should be on electronegative elements.

Assignment of e- 1. Lone pairs belong entirely to atom in question 2. Shared e- are divided equally between the two sharing atoms

The sum of the formal charges of all atoms in a species must equal the overall charge on the species.

A useful equation (happy-have) / 2 = bonds POCl3 P is central atom SO4-2 S is central atom SO3-2 S is central atom PO4-2 P is central atom SCl2 S is central atom

Exceptions to the octet BH3 Be and B often do not achieve octet Have less than and octet, for electron deficient molecules. SF6 Third row and larger elements can exceed the octet Use 3d orbitals? I3-

Exceptions to the octet When we must exceed the octet, extra electrons go on central atom. ClF3 XeO3 ICl4- BeCl2

If nonequivalent Lewis structures exist, the one(s) that best describe the bonding in the species has...

FAVORED LEWIS STRUCTURES 1. formal charges closest to zero 2. negative formal charge is on the most electronegative atom

EXAMPLES Carbon dioxide Thiocyanate ion Sulfate ion

BOND ENERGY & LENGTH

Bond Energies E = (Bonds Broken) – (Bonds Made)

Bonds form between atoms because bonded atoms exhibit a lower energy. Thus, energy is required to break bonds and energy is released when bonds are formed.

Bond Order = # bonds to a specific set of elements C-C the BO=1 C=C the BO=2 C C the BO=3 Fractions are possible

COVALENT BONDS Bond Dissociation Energy Table 9.9 (text)

Bond Energy (kJ/mol) H-F 565 H-Cl 432 H-Br 366 H-I 299

Bond Energy (kJ/mol) Cl-Cl 242 Br-Br 193 I-I 151

Bond Energy (kJ/mol)

Bond Energy (kJ/mol)

Use bond energies to predict Hc for acetylene (C2H2).

Energy Internuclear Distance

Energy Internuclear Distance

Energy Internuclear Distance

Energy Internuclear Distance

Energy Bond Length Internuclear Distance

Energy Bond Energy Internuclear Distance

Bond: Energy Length (kJ/mol) (pm)

Bond: Energy Length (kJ/mol) (pm)

Binary Ionic Compounds metal(s) + non-metal (g) ---> salt(s)

M+(g) + NM-(g) --> M-NM Lattice Energy Energy change occurring when separated gaseous ions are packed together to form an ionic solid M+(g) + NM-(g) --> M-NM

What is the lattice energy of NaCl(s)? Na+(g) + Cl-(g) ---> NaCl(s)

Lattice Energies LiCl 834 : BeCl2 3004 NaCl 769 : MgCl2 2326 KCl 701 : CaCl2 2223 Li2O 2799 : BeO 4293 Na2O 2481 : MgO 3795 K2O 2238 : CaO 3414

LE = Lattice Energy Where: k = proportionality constant dependent on structure of solid and on electron configuration of the ions Where: Q1 & Q2 = charges on the ions Where: r = the shortest distance between the centers of the cation and anion