Introduction to Organic Chemistry Practice Flashcards

Historical Background of Organic Chemistry

• Definition and Scope: Organic chemistry is the study of carbon and its seemingly unlimited number of compounds. These compounds significantly impact daily life in fields such as medicine, agriculture, and general biological processes.

• Theoretical Origins (The Big Bang): According to Oparin ($1923$), organic chemistry may have originated with the Big Bang. Components such as ammonia, nitrogen, carbon dioxide, and methane combined to form amino acids. This hypothesis was verified in a laboratory setting by Miller in ($1950$).

• Ancient Usage: Romans and Egyptians utilized organic chemicals from natural sources as dyes, medicines, and poisons, although they were unaware of the specific chemical compositions of these substances.

• Evolution of Experimental Organic Chemistry:

  • $1769$: Scheele isolated organic compounds from nature in their pure state.

  • $1784$: Lavoisier developed analytical methods for determining the elemental composition of these substances.

  • $1807$: Berzelius proposed the "Vital Force" theory, suggesting organic chemicals found in nature possessed a special force directing their synthesis, making laboratory synthesis impossible.

• The Fall of Vitalism: In ($1828$), Frederich Wöhler synthesized urea (a natural component of urine) by heating ammonium cyanate ($NH_4OCN$). This discovery proved that a "vital force" was not required and paved the way for modern synthetic organic chemistry.

• Structural Developments:

  • $1858$: Kekule and Couper developed valence theory, providing the foundation for understanding chemical bonding structures.

  • Late 19th Century: Organic chemistry expanded into biological systems, including the study of proteins and DNA.

The Chemical Bond and Atomic Theory

• Quantum Mechanics: Gained acceptance around ($1926$) following mathematical solutions for electronic energy levels provided by Heisenberg and Schroedinger.

• Electron Distribution:

  • Energy Levels (Shells): Electrons exist in shells surrounding the nucleus; energy increases as shells move further away from the nucleus.

  • Subshells and Orbitals: Within shells are orbitals, each containing up to two electrons. The four types are s, p, d, and f.

• Orbital Capacities by Shell:

  • Shell 1: Contains $1$ s-orbital ($1s$); total capacity: $2$ electrons.

  • Shell 2: Contains $1$ s-orbital ($2s$) and $3$ p-orbitals ($2p_x, 2p_y, 2p_z$); total capacity: $8$ electrons.

  • Shell 3: Contains $1$ s-orbital, $3$ p-orbitals, and $5$ d-orbitals; total capacity: $18$ electrons.

  • Shell 4: Contains $1$ s-orbital, $3$ p-orbitals, $5$ d-orbitals, and $7$ f-orbitals; total capacity: $32$ electrons.

• Shapes of Orbitals:

  • s-orbitals: Spherical shape.

  • p-orbitals: Barbell shape, aligned along the x, y, and z axes ($p_x, p_y, p_z$).

• Aufbau Principle: Electrons fill lower energy levels first. An element possesses a number of electrons equal to its atomic number.

• Electron Configurations of First and Second Row Elements:

  • $H (1): 1s^1$

  • $He (2): 1s^2$

  • $Li (3): 1s^2, 2s^1$

  • $Be (4): 1s^2, 2s^2$

  • $B (5): 1s^2, 2s^2, 2p^1$

  • $C (6): 1s^2, 2s^2, 2p^2$

  • $N (7): 1s^2, 2s^2, 2p^3$

  • $O (8): 1s^2, 2s^2, 2p^4$

  • $F (9): 1s^2, 2s^2, 2p^5$

  • $Ne (10): 1s^2, 2s^2, 2p^6$ (Inert gas with completely filled shell).

Electronegativity and Chemical Bonding Types

• Electronegativity: The ability of an atom to attract electrons to itself.

  • Periodic Trend: Increases from left to right across a row and from bottom to top in a column.

  • Row Trend: Li < Be < B < C < N < O < F

  • Column Trend: I < Br < Cl < F

• Electropositive Elements: Elements that easily lose electrons to form positive charges (e.g., Alkali metals).

• Ionic Bonding: Occurs between atoms with large differences in electronegativity. One atom transfers electrons to another.

  • Example (Lithium Fluoride, LiF): Lithium ($1s^2, 2s^1$) transfers its $2s^1$ electron to the Fluorine ($1s^2, 2s^2, 2p^5$) orbital. This forms a $Li^+$ cation and an $F^-$ anion, both achieving noble gas configurations.

• Covalent Bonding: Formed by the sharing of two electrons between two atoms.

  • Example (Hydrogen, H2): Two hydrogen atoms share their $1s^1$ electrons to form a bond.

• Polar Covalent Bonding: Occurs when electrons are shared unequally due to electronegativity differences.

  • Example (Hydrogen Fluoride, HF): The fluorine atom is more electronegative, drawing shared electrons closer and creating partial charges expressed by the Greek delta ($\delta$) symbol ($H^{\delta^+} - F^{\delta^-}$).

Bonding in Carbon Compounds

• Tetravalency: In all carbon compounds, carbon forms four bonds.

• Hybridization and the Sigma ($\sigma$) Bond:

  • sp3 Hybridization (Single Bonds): One electron from the $2s^2$ orbital is promoted to a vacant $2p$ orbital. The one s and three p orbitals mix to form four equivalent $sp^3$ hybrid orbitals of equal energy.

  • Methane ($CH_4$): Uses four $sp^3$ orbitals to bond with the $1s^1$ orbital of four hydrogens. The shape is tetrahedral with bond angles of $109.5^\circ$.

  • Ethane ($CH_3CH_3$): Features covalent sigma ($\sigma$) bonds between carbon and hydrogen, and a sigma bond between the two carbon atoms.

• The Pi ($\pi$) Bond and the Carbon-Carbon Double Bond:

  • sp2 Hybridization: One s orbital and two p orbitals mix to create three $sp^2$ hybrid orbitals, leaving one unhybridized p orbital with one electron.

  • Ethene ($CH_2=CH_2$): Two carbons join via $sp^2$ orbital overlap (sigma bond) and side-to-side overlap of the unhybridized p orbitals (pi bond).

  • Physical Properties of Ethene: Includes $5$ sigma bonds and $1$ pi bond. The molecule is planar with $120^\circ$ bond angles. The $C=C$ bond length is $1.34\,\text{Å}$ (longer than the $1.1\,\text{Å}$ $C-H$ bond).

• The Carbon-Carbon Triple Bond:

  • sp Hybridization: One s orbital mixes with one p orbital to give two hybrid $sp$ orbitals, leaving two unhybridized p orbitals.

  • Ethyne ($HC \equiv CH$): Carbons are bound by one sigma bond and two pi bonds. The molecule is linear ($180^\circ$ bond angle).

Molecular Polarity and Hydrogen Bonding

• Polarity and Dipole Moments: Bonds between carbon and electronegative atoms (O, N, S, Halogens) are polar. A separation of charge creates a dipole moment.

• Hydrogen Bonding: A strong attraction between a partially positive hydrogen atom (bonded to O, N, or F) and the non-bonding (lone pair) electrons of an electronegative atom in another molecule.

  • Properties: Responsible for high boiling and melting points. For example, water boils at a high temperature relative to its mass because extra energy is needed to break hydrogen bonds.

  • Organic Context: Critical for the solubility and boiling points of alcohols ($R-OH$) and carboxylic acids ($R-COOH$).

Organic Structures and Representations

• Structure Visualization: Carbon must always have four bonds.

• Formula Types:

  • Complete Structure: Shows every atom and every bond.

  • Condensed Structure: Groups atoms together (e.g., $CH_3CH_2CH_2CH_2CH_3$ for pentane).

  • Line (Skeletal) Structure: Ends and vertices represent carbon atoms; hydrogen atoms are implied and not explicitly written.

• Cyclic Compounds: Carbon atoms form rings, such as cyclohexane ($C_6H_{12}$).

• Complex Biological Molecules:

  • Vitamin B1 (Thiamin diphosphate): Essential nutrient found in grains, liver, and pork.

  • Nicotine: Addictive tobacco component.

  • Corilagin: Found in cranberries; prevents bacteria from attaching to kidneys.

  • Sex pheromone of bees: A complex carboxylic acid.

Classification of Organic Functional Groups

• Alkane: $CH_3CH_3$

• Alkene: $CH_2=CH_2$

• Alkyne: $HC \equiv CH$

• Alkyl Halide: $R-\text{halogen}$

• Aromatic: Benzene-related structures.

• Alcohol: $R-OH$

• Phenol: Ar-OH.

• Ether: $R-O-R$

• Aldehyde: $RCH=O$

• Ketone: $R-C(=O)-R$

• Acid (Carboxylic): $RCOOH$

• Ester: $RCOOR$

• Anhydride: $(RCO)_2O$

• Amide: $RCONH_2$

• Nitrile: $RCN$

• Amine: $R-NH_2$

Supplemental Information and Problem Set Details

• Linus Pauling: A two-time Nobel Prize winner noted for contributions to bonding theory.

• Third Row Ions: Common ions include $Na^{+1}$, $Mg^{+2}$, $Al^{+3}$, $Si^{+4}$, $P^{+5}$, $S^{-2}$, and $Cl^{-1}$.

• Dipole Moment Calculations: Based on bond moments such as $H-C$ ($0.4$) and $C-Cl$ ($1.5$).

• Hydrogen Bond Strength: In alcohols and carboxylic acids, the strength is approximately $8-10\,kcal/mole$.

• Chemical Examples for Analysis:

  • Aspirin

  • Freon-11 ($FCCl_3$)

  • Ethane, Propene, and Propyne