Detailed Notes on Hybridization, Electronegativity, Formal Charge, and Resonance

Hybridization and Bond Strength

  • Increasing s character (SP3 to SP2 to SP): Increasing the amount of s character in a hybrid orbital.
    • SP3: 25% s character.
    • SP2: 33% s character.
    • SP: 50% s character.
  • Effect on Bond Length: Greater s character leads to shorter bonds due to reduced directionality (more like a sphere than a peanut shape).
  • Bond Strength: Shorter bonds are stronger and harder to break.
  • Carbon bond breaking: Breaking the carbon bonds will be discussed in chapter 12.

Electronegativity Trends

  • Comparison of Oxygen and Chlorine:
    • Oxygen (O) is slightly more electronegative than Chlorine(Cl), with a small difference (0.2).
  • Iodine vs. Carbon:
    • Iodine (I) has an electronegativity of 2.7, which is close to that of Carbon (C).
    • Iodine is not as electronegative as other halogens due to its position on the periodic table (further down).
  • Carbon-Hydrogen Bond Polarity:
    • Electronegativity difference between Carbon (2.5) and Hydrogen (2.2) is small (0.3).
    • Carbon-Hydrogen (C-H) bonds are considered relatively nonpolar due to this small difference.
  • Carbon-Oxygen Bond Polarity:
    • Electronegativity difference between Carbon (2.5) and Oxygen (3.4) is significant (0.9).
    • Carbon-Oxygen (C-O) bonds are polar.
  • General Rule:
    • Molecules containing atoms other than Carbon and Hydrogen tend to have polar regions.
    • Hydrocarbon regions (C-H) are generally nonpolar.

Lipids and Intermolecular Forces

  • Lipids (Fats):
    • Long chain hydrocarbons with a polar group at one end.
    • Amphipathic: Having both polar and nonpolar regions.
    • Nonpolar tails aggregate together, forming membranes.
  • Intermolecular Forces:
    • Dispersion Forces (London Dispersion Forces): Weakest intermolecular force, relevant to nonpolar molecules.
      • Temporary dipoles.
    • Organic molecules composed of primarily Carbon and Hydrogen are weakly held together.
    • Low melting and boiling points.
    • Volatile: Evaporate easily (e.g., acetone).

Electronegativity Definition

  • Definition: The attraction an atom has for electrons in a bond (tug of war analogy).
  • Distinction from Electron Affinity:
    • Electronegativity: Attraction in a bond.
    • Electron Affinity: Ability to accept electrons (relevant for nonmetals forming ionic compounds).
    • Ionic compounds: Electron transfer occurs
      • Nonmetals become negative.
      • Metals become positive.
  • Key Difference:
    • Electronegativity: Occurs within a bond.
    • Electron Affinity: Occurs before bond formation.

Formal Charge Calculation

  • Formula:
    • Formal Charge = (Number of Valence Electrons) - (Number of Lines/Bonds) - (Number of Dots/Nonbonding Electrons)
  • Valence Electrons:
    • Determined by the group number on the periodic table.
      • Not the atomic number. Using the atomic number will assign incorrect formal charges.
  • Preferred Bonding Patterns (Valencies):
    • Carbon: Prefers four bonds.
    • Hydrogen: Prefers one bond.
      • Hydride (H-): Hydrogen with an extra electron, carrying a -1 charge (discussed in chapter 12).
    • Nitrogen: Prefers three bonds and one lone pair.
      • Can form coordinate covalent bond with four bonds, resulting in a positive charge (e.g., ammonium ion).
    • Oxygen: Prefers two bonds and two lone pairs.

Examples: Ammonia and Ammonium Ion

  • Ammonia (NH3):
    • Hydrogen: Group 1, one bond, formal charge = 0.
    • Nitrogen: Group 5, three bonds, two nonbonding electrons, formal charge = 5 - 3 - 2 = 0.
  • Ammonium Ion (NH4+):
    • Hydrogen: Formal charge = 0.
    • Nitrogen: Group 5, four bonds, zero nonbonding electrons, formal charge = 5 - 4 - 0 = +1.
  • Conjugate Acid-Base Pairs:
    • Ammonia (NH3): Base.
    • Ammonium (NH4+): Conjugate acid (differs by H+).

Oxygen Bonding Preferences

  • Preferred Bonding: Two bonds and two lone pairs.
  • Example: Protonated Oxygen
    • Structure: Oxygen with three bonds to hydrogen.
    • Formal Charge: 6 (valence electrons) - 3 (bonds) - 2 (nonbonding electrons) = +1.
    • Significance: H3O+ is the conjugate acid of water.

Resonance Theory

  • Moving Electrons: Resonance involves creating various pictures of a molecule by moving electrons.
  • Electrons Moved: Only electrons in pi bonds or nonbonding electrons can be moved. Sigma bonds cannot be moved.
  • Stability: The most stable resonance structure is the one with the fewest formal charges.

Formal Charge and Stability

  • Absolute Value: The absolute value is taken, not the signed value. This is done by just looking at the numerical value of the charges.
  • Neutral Molecules: Even if a molecule is neutral, it can still have formal charges.
  • Minimizing Charge: Spreading out electron density stabilizes the molecule.
  • Example: Molecule with +1 and -1 formal charges has two formal charges in total.

Using Preferred Bonding to Identify Formal Charges

  • Nitrogen with two bonds: Probably has a formal charge.
  • Oxygen with three bonds: Likely has a formal charge.

Calculating Formal Charges in Resonance Structures

  • Example 1:
    • Structure: N=N=N
    • Formal Charges:
      • End nitrogens: -1
      • Middle nitrogen: +1
    • Overall Charge: -1
    • Total Formal Charge: 3 (absolute value: 1 + 1 + 1).
  • Example 2:
    • Structure: N-N≡N
    • Formal Charges:
      • Nitrogen with single bond: -2
      • Middle nitrogen: +1
      • Nitrogen with triple bond: 0
    • Total Formal Charge: 3 (absolute value: 2 + 1 + 0).
  • Example 3:
    • Structure: O=C-N
      • Formal Charges:
        • Oxygen: 0, Carbon: 0, Nitrogen: -1
      • Overall Charge: -1
      • Total Formal Charge: 1
  • Example 4:
    • Structure: O-C≡N
      • Formal Charges:
        • Oxygen: -1, Carbon: 0, Nitrogen: 0
      • Overall Charge: -1
      • Total Formal Charge: 1

Movement of Electrons and Curved Arrows

  • Curved Arrows: Used to show the movement of electrons.
    • Tail of arrow: Indicates the electrons that are moving.
    • Head of arrow: Indicates where the electrons are going, forming a new bond.
  • Relationship of Resonance Hybrids: Two pictures that are resonance hybrids are related by a double-headed arrow.
  • Actual Structure: The real structure is a hybrid of the different resonance structures.
  • Resonance is not Equilibrium: They are the same structure but two different resonance hybrids.
  • Concentrated Electron Density: A more concentrated electron density leads to a negative charge and molecule instability.
  • Spreading electron density: Molecule stability results from spreading the cloud out over everybody.

Water Bucket Analogy

  • Concentrated Charge: A charged molecule is akin to having all the water in a five-gallon bucket (concentrated).
  • Spreading Charge: Resonance is like dumping the water on the floor, spreading it out and reducing the concentration in any one spot.
  • More Resonance Structures: The more resonance structures, the more stable the molecule because the charge is more dispersed.

Major and Minor Contributors

  • Major Contributor: Has either the charges on an atom that prefers charge or the one with the least formal charges.

Moving Electrons in Resonance

  • Example 1: Converting one resonance structure to another.
    • Initial Structure:
    • Process: Taking nonbonding electrons and transforming them into a pi bond.
    • Curved Arrows: This is indicated literally by curved arrows.
  • Example 2:
    • Process:
      • Electrons are taken from pair and made a bond, taking a bond and making a bond.
  • Curved Arrows: A curved arrow is used to indicate how you perform each transformation.

Identifying Major Contributors

  • Resonance Structure: Both pictures are resonance structure.
    • Structure 1: -1, +1, -1.
    • Structure 2: -2, 0, +1, 0.
    • The one where is least the charges is the major contributor.

Identifying Resonance Contributors (cont.)

  • Major Contributor: Will either be the structure that charges an atom that prefers it or with less charges overall.
  • The carbon also needs its own charges to be considered okay.

Drawing Organic Molecules: Lewis Structure

  • Lewis Structure: Shows all atoms and bonds.

  • Too much detail: Structure can result in being lost and overwhelmed.

Drawing Organic Molecules: Condensed Formula

  • Condensed Formula: Atoms are shown, but the bonds aren't with special cases being exceptions.
  • Branches: The branches off the main chain are shown using lines.
  • Parentheses: The branches can also be shown without lines, in the parenthese.

Drawing Organic Molecules: Line Structure (Skeletal Formula)

  • Line Structure: Only shows carbon-carbon connections. Uses a zigzag to show geometry, putting carbons at zigzag points.
  • Carbons Bonds: Carbons are only at the end of the chain and the apex points of the branches.
  • Tetrahedral Geometry: Tetrahedral geometry is shown by zigzagging the connections.
  • Rings: Are geometric shapes (triangle, diamond, pentagon, hexagon, heptagon, octagon).
  • Cyclocarbon: Always given as a line structure.

Isomers: Definition

  • Isomers: Have the molecular formula, but have different connections.
  • Same: Molecular formula.
  • Different: Connections.

Isomers: Example C5H12

  • Arrangement: Various carbon arrangements.
  • Hydrogens: Carbon arrangement must be arranged to have the same amount of hydrogens.
  • Chain Twisting: Chain twisting does not make another equivalent structure. Chain ends can be thought of as still being attached/connected if they are just move around.
  • Different Isomers: Will require rearrangement that causes different attachment points.
  • Drawing isomers: Take the original then and start removing branches.
  • Draw branches: Different chain arrangements, etc. can require completely removing branches. It boils down to being able to identify carbon locations and that equivalent structures can be moved around.

Isomers: Saturated definition

  • Saturated Structure: Saturated means: the maximum number of hydrogens per carbon.
  • Example given:
    * CnH2n+2 (where n is the number of carbons) is the chemical formula where hydrogens are saturated
  • Rings or Multiple Bonds: Saturated chemical structures DO NOT have rings or multiple bonds.

Line Structure:

  • Drawing the line structure can require less steps or a better idea.
  • Remember: all lines are made of carbons unless identified otherwise. There is NO CARBON in that part.

C6H14 Isomers (Line Note)

  • Remember: The structure can be checked based on molecular weight. Start with drawing 6 in a line and go from there.
  • One can take systematically and count from there (removing from one end to the middle, etc.)
  • Different structures: start with no branch, one branch, two branches.

Presence of Heteroatoms

  • Hetero: Hetero means not carbon, in which something other than carbon is present. Oxygen has more isomers than carbon has.
  • If Oxygen is in the C4H10. Oxygen likes to be the divalent structure:
    * O puts O in between the C and H to make an OH..
  • Increased isomers: Different arrangements where each are completely viable will lead to vastly increased isomers.

Resonance Theory

  • Scenario 1: Carbon Has Positive Charge:
    • Easy to find.
    • Carbon is sp2 with an empty orbital (two buckets, one has 5 gallons of water and the one the does not).
    • Electrons:
      • Electron being moved that it.
    • Where are:
      • Double/pi bond at or near the carbon to swing that gate over.
  • Hinge-Gate:
    • Swing that gate over (close that then open that over. One is no open and one is now open).
  • Positive Charges:
    • The way positive charges can change is due to electrons moving because you removed them.
      • However, the arrows are NOT put to the charges
        #
  • Resonance Structure 2:
    * This looks just like the original. This is very common with resonance.
    * Double Bond:
    * However, the double bonds should be partially all the way over at some point. It's partially a double bond and a positive charge over all the atoms over there.
  • Movies:
    * One day that's going to be a movie or something where people can just see it over there.

Resonance Structure: Scenario 2

  • Carbon: This scenario has carbon a negative charge * Which equals two electrons in a pi orbital
    • Process: Take a pair of dots, make a bond (a new pi bond), then the pi bond makes a pair of dots
      • Whole Pattern: Whole idea is just to keep moving that down
  • Example:
    • Picture: Same as the other one, except except two dots on the carbon.
    • Can't draw this here. I can't move the pi bond over. I will give it so many. Okay. All together. I don't want to have five so I have to have four only
      • Sites with the four electrons (two being hydrogens and 2 be dots)
        * Now we take the pair and make it a pi bond and there are
  • Two Electrons + 1 pi bond
    * Pair or dots on C means it had a negative charge.
    * If it's full with dots and not swing then it must move out.
  • If It's Top:
    * Spreading the positive charge.
  • If It's Bottom:
    * Spreading the negative charge.
    Perfect. The lecture notes for tomorrow will be shorter Chapter 13 and so all I have to do for tomorrow, Chapter 13 and so all I have to do for tomorrow