Chemistry AP Study Notes 7 (Bonding and Geometry)

Thus far:

1. The Basics
2. Basic Chemical Reactions
3. Reactions in Solution
4. Gases
5. Thermodynamics
6. Electron Orbitals

7. Bonding

• Ionic bonds trade electrons. Covalent bonds share electrons.
• Ionic bonds are a metal and a non-metal. Use the criss-cross method to figure out subscripts. These are polar bonds.
• Use Lewis dot structures to draw shared covalent bonds. These tend to be more non-polar, although they can be partially polar (polar covalent bonds).
• Electronegativity is the measure of the tendency to hold or gain electrons. It increases from bottom left to top right on the periodic table.
• The N - A = S rule is 1) N -- what is the ideal filled valence electron number? (2 or 8); 2) A -- what is the total number of available electrons; 3) S is the total number of electrons that must therefore be shared and S/2 tells the total number of bonds.
• The idea of formal charge helps you determine which of more than one possible structure is the most likely. For each possible structure, take the number of valence electrons for an atom. Subtract the number of electrons that aren't bonded and add half the number of bonded electrons. Make sure the total of all the "formal charges" on each atom add up to the actual charge on the ion.
• The preferred structure is the one with the most zeros, especially on the most electronegative atom, without any like charges next to each other.
• There are a few circumstances where the octet rule does not work.

Molecular Geometry

• The VSEPR theory predicts the shape a molecule will take (valence shell electron pair repulsion). Basically, electron pairs try to move as far away from each other as they can.
• 1) Write the Lewis-dot structure, 2) how many electron pairs are there (count double and triple bonds as a single group, 3) what shape maximizes the distance (this is the geometry of the electron groups), 4) now for drawing purposes, pretend that the non-binding electron groups aren't there and draw the molecular geometry.
• Here are the possibilities. With only two bonding pairs, we have linear geometry (like CO₂).
• With three total electron pairs, we have trigonal planar (all three used, BF₃) and bent (only two used NO₂)
• With four total electron pairs, we have tetrahedral (with all four used, CH₄), trigonal pyramidal (with only three used, NH₃), or bent (with only two used, H₂O).
• With five total electron pairs we have trigonal bipyrimidal (with all five used, PF₅), seesaw (with four used, SF₄), T shaped (with three used, ClF₃), and linear (with two used, XeF₂).
• Finally, with six total electron pairs, we have octohedral (with all six used, SF6), square pyrimidal (with 5 used, ClF₅) and square planar (with 4 used, XeF₄).

Valence Bond Theory

• Explains geometry by hybrid orbitals. 
• Linear is one s and one p orbital (sp hybridization).
• sp² hybridization is an s with two p orbitals. This is trigonal planar.
• sp³ hybridization is an s with three p orbitals. This is tetrahedral.
• sp³d hybridization is an s with three p and one d orbital. This is trigonal bipyrimidal.
• sp³d² hybridization is an s with three p and two d orbitals. This is octohedral.
• Sigma bonds are the straight bonds between atoms. Pi bonds are the second and third bonds in double and triple bonds.

Molecular Orbital Theory

• A theory of covalent bonds that sees the electrons as assigned to the whole molecule rather than the individual atoms.
• Speaks of bonding orbitals and antibonding orbitals.
• A concept called "bond order" is half the bonding orbitals minus the antibonding orbitals.
• The higher the bond order, the shorter and stronger the bond.

In addition

• Resonance structures are instances when, say, a double bond isn't just in one location but moves around, so to speak (e.g., NO₃^-).
• Paramagnetism is an attraction to a magnetic field due to unpaired electrons.
• Diamagnetism is a slight repulsion from a magnetic field due to the presence of paired electrons.