Covalent Bonding and the Octet Rule
 Bonds form to lower atoms energy.
 For representative elements, an octet of electrons around the atom is preferred.
 Consider F2, HBr.. and let’s use electron-dot notation. Only use valence electrons.
 There are exceptions to this rule. o Hydrogen (One electron) o Boron Group
 Several elements starting in period 3 and lower can expand their octet to include up to 12 e- around a central atom.
Covalent Bonding
 Based on valence electron count.
 When representing bonding, lines are often used to represent a bond.
 Each line represents 2e-
Single Bonds
 Represented by one line.
 Consider CI2, H2O, NH3
Lone Pairs
 Consider the diatomic molecule Cl2 o How many e- are covalently bonded? (2) o How many e- are left unbonded? (12)
 The unbonded e- pairs are known as “lone pairs,” “unshared pairs, or “non-bonded pairs.”
 Sometimes these e- can engage in bonding themselves.
 Coordinate covalent bonds (dative bonding)
 Another way to say this: one atom donates both e- to a bonding situation. CO, NH4+
 When dealing with charges on polyatomic ions and the covalent bonds that make them up, simply add/take away the number of e- that corresponds to the charge. o Ex) OH-

Covalent Bonding- A Simple System for Lewis Dot Structures
 No-no’s with LDS o Terminal halogens NEVER get multiply bonded. o Terminal atoms NEVER get bonded to each other. o When you draw LDS, do NOT cross lines… all multiple bonds must be parallel/
 Yes-yes’s with LDS o All atoms should generally have 8 e- in the final drawing
(barring the exceptions.)
Covalent Bonding- Double and Triple Bonds
 Depending on the valence e- count, atoms can sometimes bond more than once.
 Double bonds- 2 pairs (4 e-) o O2
 Triple bonds- 3 pairs (6 e-)
Simple Guidelines for Writing out Lewis Dot Structure
1.
Count the number of total valence e-.
2.
Draw the central atom (Usually elements from groups IV, V).
3.
Put the other elements around the central atom.
4.
Draw single lines connecting the central atom to the elements.
5.
Subtract the number of e- used thus far from your total valence e- count.
6.
Pepper the remaining e- around whatever atoms need lone pairs to get an octet.
 You might find that in certain cases, you might need to make multiple bonds.
VSEPR
 Valence Shell Electron Pair Repulsion (VSEPR) determines the shape of the molecule.
 Atoms arrange themselves in molecules to minimize e- e- repulsion.

o LP-LP interactions- the worst (can cause significant distortion of molecular shape.) o LP-BP interactions- not so bad. o BP-BP interactions- barely felt.
 Get the lone pairs as far apart as possible from one another!
VSEPR- How to Draw in 3D
 There are three types of lines used. o Regular lines –In Plane o Thick Lines- Coming out of the plane towards you o Dashed Lines- Going back into the plane
 You may not carelessly use these.
VSEPR- Rules and Regs
 VSEPR is a model… use it in conjunction with a LDS
 e- pairs around the central atoms count as groups
 Atoms and lone pairs determine the overall geometry of the molecule
 However, only the atoms around the central atom determine the molecular shape.
 Multiple bonds are treated as single bonds in VSEPR
Hybridization (pg 201-203)
 VSEPR gets its geometry via orbital hybridization.
 Orbitals are mathematical constructs—it’s possible to theorize that they can mix in such a way as to create new degenerate, orbital hybrids that help keep e- as far apart as possible.
 CH4, O2 as examples.
Resonance (pg. 189)
 (Given that orbital’s can hybridize it’s also possible to envision multiple LDS structures for a single molecule.
 Resonance is bonding in molecules or ions that cannot be correctly represented by one LDS—in actuality, the molecule exists as the average of its LDS structures – “double arrow.”
 O3, NO3-

Covalent Bond Energies
 In favorable situations, energy is released in bond formation (-E)
 In order to separate the atoms, the same amount of E must be added (+E)—Bond energy (kJ/mol)
 Bond energies vary depending on what the atom is bonded to (as do bond lengths) —reported as averages.
 In order of Increasing E, o C-C< C=C< CΞC
 As the bond energy increases, the distance between atoms gets smaller because the more bonds you have to break the more energy, and the closer they are.
Ionic Bonding
 Ionic compounds are composed of cations and anions held together by electrostatic forces.
 Crystalline solids— three-dimensional network of ions—an array made up of formula units.
 Order lowers PE
 Use Pauling Values to determine whether ionic or covalent.
 Whenever you make an ionic bond, you need to ensure charge balance. o Na and F; Ca and N; Mg and O
 Criss-cross method for charge balance.
 Generally, it’s Group I and II mixing with groups V-VII
 Formula Unit- Simplest ratio of ions that provides electrical neutrality. o Different from a molecule, which is a discrete single entity.
 Electron Dot Notation can be used to describe ionic bonding o Ex) NaCl, CaCl2, Li3N
 Crystal lattices are usually very robust & stable thermodynamically.
 Lattice energy- The energy released when one mole of an ionic crystalline compound is formed from gaseous ions. o NaCl- -787.5 kJ/mole o CaF2- -2634.7 kJ/mol o MgO- -3760 kJ/mol

 Polyatomic ions are ions made up of covalently bonded atoms wherein the entire collection of atoms carries a collective charge. o These can bond with cations and anions as discrete ions themselves.
 Ex) (OH)-, (HCO3)-, (NO3)-
 Na+ vs. (CO3)2-, Ca2+ vs. (NO3)-
Covalent Vs. Ionic Bonds
 Ionic o Formula units have strong attraction to one another o High melting point, boiling point, and very hard, brittle materials.
 Brittle because the charges go close when altered, making them repel. o Electrically conductive in solution or in a molten state, not as a solid.
 Covalent o Molecules have weak attraction to one another, generally speaking. o Lower melting point, boiling point, and generally softer materials (powders, tacky solids), many are gases and liquids at room temperature.
Intermolecular Forces
 Molecules are attracted to one another
 Boiling point is a good metric of how well molecules are attracted to one another. o The higher the boiling point, the closer they are attracted.
 These forces are weaker than intramolecular forces.
 From Weakest to Strongest o London dispersion forces < dipole-dipole < H-bond
 London Dispersion Forces o All molecules have these o Introduced dipoles o Caused by e- motion o More electrons = more mass = stronger LDF’s

 Dipole-Dipole o Strongest intermolecular force o Opposite charges separated by short distance. o Positive ends attract negative ends. o Whether or not a molecule has a dipole is determined by its geometry (VSEPR). o Dipoles can induce other nonpolar molecules to have a dipole. o H- Bond
 Strongest of Dipole forces.
 H-N, H-O, H-F only.
 Large electronegativity difference creates a significant positive charge on H- almost to the level of a free proton.
 Highly charged H is attracted to lone pairs on other N,
O, or F’s.
 Draw a dotted line between elements in two substances

27.11.2012 19:23:00

27.11.2012 19:23:00