General Chemistry Review - Organic Chemistry

1.1: Introduction to Organic Chemistry

  • Organic chemistry is the study of carbon-containing compounds.
  • Inorganic chemistry deals with compounds lacking carbon.
  • Organic chemistry is crucial in medicine, chemical engineering, and chemistry.
  • It contributes to household chemicals, foods, plastics, drugs, and more.
  • Our bodies are mainly composed of organic compounds like DNA, RNA, and proteins.
  • Vitalism, the early belief that organic compounds could only come from living organisms, was disproven in 1828 when Friedrich Wöhler synthesized urea from ammonium cyanate.

1.2: The Structural Theory of Matter

  • In the mid-19th century, Kekulé, Couper, and Butlerov proposed that substances are defined by the specific arrangement of atoms.
  • Isomers are compounds with the same molecular formula but different structures.
  • Constitutional isomers have the same atoms but differ in their connectivity, leading to different chemical and physical properties.
  • Each element forms a predictable number of bonds (HONC: H-1, O-2, N-3, C-4).

1.3: Electrons, Bonds, and Lewis Structures

  • Atoms have a nucleus with neutrons and positively charged protons.
  • Electrons are confined to shells (numbered 1, 2, 3, etc.) around the nucleus.
  • Shells are divided into subshells (s, p, d, f), which contain orbitals.
  • An orbital can hold up to two electrons.
  • Valence electrons, found in the outermost shell (valence shell), are involved in bonding and reactions.
  • Lewis structures (electron-dot structures) show valence electrons as dots around the element symbol.
  • Atoms react to achieve the electronic configuration of a noble gas, typically an octet of electrons in the valence shell (octet rule), except for hydrogen, which needs two electrons.
  • Covalent bonds form through the sharing of electrons between atoms with similar electronegativities.
  • Molecules consist of atoms joined by covalent bonds.

Drawing Lewis Structures:

  1. Determine the number of valence electrons in the molecule or ion.
    • Add 1 electron for each negative charge (anion).
    • Subtract 1 electron for each positive charge (cation).
  2. Arrange the atoms, typically with H and halogens (F, Cl, Br, I) on the ends.
  3. Connect atoms with single bonds (one line represents two electrons).
  4. Arrange the remaining electrons in pairs to give each atom a full valence shell, starting with the most electronegative atoms.
    • Aim for 8 electrons (octet rule) except for H, which needs 2.
  5. Show nonbonding electron pairs (lone pairs) as Lewis dots (:).
  6. Use multiple bonds (double or triple) if needed to satisfy the octet rule and eliminate lone pairs.
    • Double bonds: share two pairs of electrons (two parallel lines).
    • Triple bonds: share three pairs of electrons (three parallel lines).
  • Examples: HCN, CO2, NH2^{-}, [CH3CO2]^-

1.4: Identifying Formal Charges

  • Formal charge is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity.
  • Calculated by: Formal Charge = (# of valence electrons) - (# of non-bonding electrons) - (1/2 # of bonding electrons).
  • Examples: Ammonium ion (NH4^+$), water (H2O).

1.5: Induction and Polar Covalent Bonds

  • Electronegativity: An atom's attraction for electrons in a chemical bond.
  • Pauling scale: F is 4.0; other elements are relative to F.
  • Electronegativity increases from left to right across a period and from bottom to top within a group.
    • Difference < 0.5: covalent bond (nonpolar covalent bond).
    • Difference between 0.5 and 1.7: polar covalent bond.
    • Difference > 1.7: ionic bond (electrons are transferred).
  • Induction: The more electronegative atom attracts electrons more strongly.
  • Polar covalent bonds: The more electronegative atom has a partial negative charge (\delta^-), and the less electronegative atom has a partial positive charge (\delta^+).
  • Dipole: Separation of charge, represented by an arrow pointing to the negative end, with a cross at the positive end.

1.6: Reading Bond Line Structures

  • Bond-line structures simplify drawing organic molecules.

Rules for Reading Bond-Line Structures:

  1. Each line represents a bond.
  2. Each bend or terminus represents a carbon atom unless another group is shown explicitly.
  3. Carbon atoms are not labeled with "C" (can optionally write CH_3 at terminus).
  4. Hydrogen atoms bonded to carbon atoms are not shown (unless needed to infer 3D structure).
  5. The number of hydrogen atoms bonded to a carbon atom is inferred to satisfy the octet rule.
  6. Atoms other than C or H are shown with their element symbol.
  7. Hydrogen atoms bonded to atoms other than C are shown (e.g., –NH_2, –OH).
  8. Double bonds are shown as two lines; triple bonds are shown as three lines.

1.7-1.9: Atomic Orbitals, MO Theory

  • Orbitals dictate the shape of molecules.
  • We can't pinpoint an electron's exact location but can describe the probability of finding it in a certain place.
  • Atomic orbital: 3D distribution of probabilities for a single atom.
  • Wavefunctions (\psi) are mathematical functions that determine an orbital's shape.
  • Orbitals can be occupied or unoccupied; each holds up to two electrons.
  • For organic chemistry, focus on 1s, 2s, and 2p orbitals.
    • The number "2" indicates a higher energy level with a node.
  • Electrons possess wavelike properties.
  • \psi = + (red) above the average level; \psi = − (blue) below the average level.
    • Nodes: Locations where \psi = 0.
  • Filling Orbitals:
    • Aufbau principle: Fill the lowest energy orbital first.
    • Pauli exclusion principle: Max of two electrons per orbital, with paired spins (↑↓).
    • Hund’s rule: With degenerate orbitals, add one electron to each before pairing.
  • Atomic orbitals combine like waves (constructive vs. destructive interference).
  • Molecular orbital energy level diagrams show the combination of two atomic orbitals to form a molecular orbital.
  • Bonding orbitals from 1s-1s and 2s-2s interactions are cylindrically symmetrical (\sigma orbitals, form \sigma bonds).
  • 2p orbitals can combine end-on or side-on.
    • End-on overlap: Creates \sigma bonds.
    • Side-on overlap: Lacks cylindrical symmetry (π orbitals, form π bonds).

1.10: Hybridized Atomic Orbitals

  • Methane (CH_4) has four C–H bonds, but carbon's electron configuration suggests it can only make two bonds.
  • Hybridization: Combine atomic orbitals to create new, equivalent orbitals.
    • In methane, the 2s and 2p orbitals combine to make four sp^3-hybridized orbitals (1/4 s, 3/4 p).
  • Ethene (H2C=CH2) needs three equivalent bonds from each C atom.
    • Combine the 2s with two 2p orbitals to make three sp^2-hybridized orbitals (1/3 s, 2/3 p).
  • Acetylene (ethyne, HC≡CH):
    • Hybridize 2s and one 2p orbital to form two sp-hybridized orbitals on each C atom.
  • More "s" character leads to shorter and stronger bonds.

1.11: VSEPR Theory

  • Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes: electron pairs (bonding and nonbonding) repel each other and arrange themselves to maximize distance.
  • Steric number: Total number of electron pairs (bonding and nonbonding).
    • Steric number = 4: sp^3
    • Steric number = 3: sp^2
    • Steric number = 2: sp
  • Examples: Methane (CH4) is tetrahedral; Ammonia (NH3) is trigonal pyramidal; Water (H_2O) is bent.
    • Lone pairs repel more strongly than \sigma bonds, reducing bond angles.
  • Boron trifluoride (BF_3): Steric number = 3, sp^2, trigonal planar.
  • Beryllium hydride (BeH_2): Steric number = 2, sp, linear.

1.12: Dipole Moments and Polarity

  • Dipole moment (\mu): Separation of positive and negative charge in a polar covalent bond.
  • Measured experimentally, on the order of 10^{-18} esu cm.
  • 1 debye (D) = 1 \times 10^{-18} esu cm.
  • A molecule with a dipole moment is polar.
  • Molecules with polar bonds may not have a net dipole moment if the bond dipoles cancel due to symmetry (e.g., CCl4 vs. CH3Cl).
  • Unshared electrons significantly contribute to dipole moments (e.g., H2O and NH3$$).

1.13: Intermolecular Forces and Physical Properties

  • Intermolecular forces: Attractive forces between individual molecules, determining physical properties.
    • All intermolecular forces are electrostatic.

Types of Intermolecular Forces:

  • Dipole–Dipole Interactions: Occur between polar molecules with permanent dipole moments.
  • Hydrogen Bonds: Strong dipole-dipole forces between H bonded to small, strongly electronegative atoms (O, N, F) and lone pairs on other electronegative atoms.
  • London Dispersion Forces: Temporary dipoles in nonpolar molecules induce dipoles in neighboring molecules, leading to attractive forces.
    • Larger hydrocarbons have stronger London dispersion forces due to increased surface area.
  • Compactness and rigidity also affect physical properties.

1.14: Solubility

  • "Like dissolves like": polar compounds dissolve in polar solvents, and nonpolar compounds dissolve in nonpolar solvents.

Assigned Homework Problems

  • Practice Problems: 1.32, 1.33, 1.35, 1.39, 1.41, 1.45, 1.47, 1.48, 1.52, 1.56, 1.58, 1.62
  • ACS-Style Problems: All
  • Integrated Problems: 1.73, 1.77
  • Challenge Problems: 1.81