Organic Chemistry: A Review of General Chemistry

Organic Chemistry: A Review of General Chemistry

1.1 Introduction to Organic Chemistry

• Definition of Organic Chemistry:
  • The study of carbon-containing molecules and their reactions.
• What occurs during a reaction?:
  • Molecules collide.
  • Bonds are broken and bonds are formed.
• Why do reactions occur?:
  • This question will require at least 2 semesters to answer comprehensively.
  • Emphasis on understanding electrons, as they are key players in reactions.

1.2 Constitutional Isomerism

• Historical Context:
  • Proposed in the mid-1800s that substances are defined by specific arrangements of atoms.
• Insufficiency of Molecular Formula:
  • A compound’s formula alone is inadequate for its definition.
  • Compounds can differ based on how atoms are bonded together.
• Definition of Constitutional Isomers:
  • Compounds that share the same molecular formula but possess different structural configurations.

1.3 Covalent Bonding / Common Bonds

• Common Bonded Atoms to Carbon:
  • Common atoms that bond with carbon include nitrogen (N), oxygen (O), hydrogen (H), and halides (F, Cl, Br, I).
• Elemental Bonding Patterns:
  • Generally, each element forms a specific number of bonds with other atoms.
• Exercise:
  • Practice with SkillBuilder 1.1 by drawing constitutional isomers.

1.3 Atomic Structure

• Summary of General Chemistry Concepts:
  • Protons: positively charged (+1) particles located in the nucleus.
  • Neutrons: neutral particles, also in the nucleus.
  • Electrons: negatively charged (−1) particles in orbitals outside the nucleus.
  • Valence Electrons:
  • Valence electrons are those in the outermost shell.
  • Using carbon as an example, valence electrons play a crucial role in bonding.

1.3 Counting Valence Electrons

• Method of Calculation:
  • The number of valence electrons can be calculated through electron configuration analysis.
  • For Group A elements, simply refer to the group number on the periodic table, which indicates the number of valence electrons.
• Exercise:
  • Practice with step 1 of SkillBuilder 1.2 to draw a Lewis structure.

1.4 Formal Charge / Definition

• Concept of Formal Charge:
  • A formal charge occurs when there is an imbalance in the number of valence electrons an atom owns compared to what it needs to be neutral.
• Types of Charges:
  • Anions: negatively charged atoms.
  • Cations: positively charged atoms.
• Determining Formal Charge:
  • Compare the number of valence electrons owned (based on bonding) to those needed for neutrality.

1.4 Formal Charge / Contribution of Carbon

• Neutrality of Carbon:
  • Carbon requires four valence electrons to be neutral as it belongs to Group IV.
  • In certain structures, carbon may be surrounded by eight electrons but only owns four (one from each bond).
  • Hence, with four electrons owned and four needed, carbon has no formal charge.

1.4 Formal Charge / Oxygen Example

• Formal Charge of Oxygen:
  • Oxygen exists in Group VI and needs six valence electrons to be neutral.
  • Surrounded by eight electrons, but only owns seven (one from a bond and three from lone pairs).
  • As it owns seven electrons but needs six, it carries a formal charge of −1.
• Exercise:
  • Practice with SkillBuilder 1.3 by calculating formal charges.

1.5 Electronegativity and Polar Bonds

• Definition of Electronegativity:
  • Electronegativity measures how strongly an atom attracts shared electrons.
  • Fluorine (F) is recognized as the most electronegative element, serving as a reference point for other atoms within the same group or period on the periodic table.

1.5 Types of Bonds Based on Electronegativity

• Covalent Bond:
  • Electrons are shared between two atoms when the difference in electronegativity is less than 0.5.
• Polar Covalent Bond:
  • Electrons are shared between two atoms when the difference in electronegativity is between 0.5 and 1.7.
• Ionic Bond:
  • Electrons are not truly shared; differences in electronegativity exceed 1.7, with the more electronegative atom fully owning the electrons.

1.5 Polar Bonds and Electron Shifting

• Behavior of Electrons:
  • Electrons will tend to shift away from atoms with lower electronegativity towards those with higher electronegativity.
  • The larger the difference in electronegativity, the more polar the bond.

1.5 Bonds and Electronegativity Values

• Ambiguity in Bond Classification:
  • Certain bonds may be classified as either covalent or ionic based on electronegativity differences, as seen in an example where the difference is 1.5, suggesting a borderline between polar covalent and ionic bonds.
  • Pointers about the variability in treatment of electronegativity differences.
• Exercise:
  • Practice with SkillBuilder 1.4 by locating partial charges.

1.7 Atomic Orbitals / Definition

• Introduction to Quantum Mechanics:
  • Established in the 1920s, quantum mechanics explains the wave behaviors of electrons.
  • Solutions to wave equations yield wave functions.
  • The three-dimensional representation of these wave functions squared reveals atomic orbitals.

1.7 Electron Density and Probability

• Characteristics of Orbitals:
  • Each type of orbital is characterized by its shape (e.g., s, p).
  • Electron Density Defined:
  • Describes the probability of locating an electron; the shape indicates 90-95% of the area where an electron is likely to be found.
• Visualization of Orbitals:
  • Atomic orbitals can be conceptualized as clouds of electron density.

1.7 Phases of Atomic Orbitals

• Dual Nature of Electrons:
  • Electrons demonstrate both particle-like and wave-like properties, although the theory of quantum mechanics is not fully complete.
  • The theory confirms experimental data and maintains predictive capabilities.
• Wave Function Values:
  • Similar to how waves behave on a lake, electron wavefunctions can have positive (+), negative (−), or zero values (nodes).

1.7 Significance of Signs and Nodes in Orbitals

• Mathematical Derivation of Orbitals:
  • The signs of wave functions are mathematically derived and are not indicative of electrical charge.
  • The presence of a nodal plane in a p-orbital is significant for understanding orbital overlap in bond formation.

1.7 Properties of s Orbitals

• Stability of Electrons in Orbitals:
  • Electrons achieve maximum stability (lowest energy state) in the 1s orbital.
  • The 1s orbital accommodates up to two electrons.
  • For atoms with more electrons, they fill the 2s and 2p orbitals subsequently. The 2p orbitals exhibit equal energy levels, referred to as degenerate orbitals.

1.8 Valence Bond Theory / Introduction

• Concept of Bond Formation:
  • A bond forms when atomic orbitals overlap, akin to the overlapping of waves.
  • Bonding occurs only through constructive interference of wave functions.

1.8 Sigma Bond Formation

• Mechanism of Sigma Bonding:
  • Example: The bond in H$_{2}$ (hydrogen molecule) arises from constructive interference of overlapping atomic orbitals, leading to sigma (σ) bond formation.
  • The bonded electrons are primarily localized in the region of overlapping orbitals, represented as a sigma bond.

1.9 Molecular Orbital Theory / Introduction

• Understanding Molecular Orbitals (MOs):
  • Molecular orbitals are formed when atomic orbital wavefunctions overlap, creating MOs that span the entire molecule.
  • MOs provide a more thorough analysis of bonds, including both constructive and destructive interferences.
  • Preservation of quantity: The number of MOs produced equals the number of AOs that contribute to them.
• Visualizing MOs in H$_{2}$:
  • Visualization of molecular structure reveals the bonding environment.

1.9 Antibonding Orbitals

• Properties of Antibonding MOs:
  • Antibonding MOs have higher energy due to the presence of one node.
  • Electrons preferentially occupy bonding MOs to reach a lower energy state when atomic orbitals overlap.

1.10 Hybridized Atomic Orbitals

• Hybridization Requirement for Carbon:
  • The ground state electron configuration of carbon does not adequately explain the formation of four bonds.
  • While considering the excited state also fails to create four equivalent bonds, as only two orbitals have unpaired electrons.

1.10 Hybridized State of Atomic Orbitals

• Mechanism of Hybridization:
  • Hybridization enables carbon to create four equal atomic orbitals that exhibit symmetrical geometry.
  • These atomic orbitals need to have equivalent energy levels to create C–H bonds of equal energy and symmetry.
• Characterization of sp$^{3}$ Orbitals:
  • The shape of an sp$^{3}$ hybrid orbital is determined by 25% s-character and 75% p-character.

1.10 Example of sp$^{3}$ Hybridization

• Formation of CH$_{4}$:
  • In methane (CH$_{4}$), the 1s atomic orbitals from four hydrogen atoms overlap with the sp$^{3}$ hybrid orbitals of carbon.

1.10 Hybridization in Ethene

• Use of Hybridized Orbitals:
  • In ethene, each carbon atom bonds with three others, necessitating only the formation of three hybridized atomic orbitals.
  • Carbon adopts sp$^{2}$ hybridization, yielding three equal-energy sp$^{2}$ orbitals and one unhybridized p orbital.
  • The sp$^{2}$ orbital shape is defined by a 33% s-character and 67% p-character distribution.

1.10 Pi Bonding in Ethene

• Pi Bond Formation:
  • The sp$^{2}$ orbital overlaps to generate sigma (σ) bonds.
  • Unhybridized p orbitals engage in side-by-side overlap to form pi (π) bonds.
  • The electron density of the pi bond is distributed above and below the molecular plane.
  • Pi bonds are intrinsically weaker than sigma bonds.

1.10 Pi Bond Description under MO Theory

• Visualization of the Pi Bond:
  • Pi bonds can also be described using molecular orbital theory, emphasizing the role of wave function phase.

1.10 Hybridization State in Ethyne

• Bonding in Ethyne:
  • Ethyne requires each carbon to bond with two others, resulting in the formation of two hybridized atomic orbitals.
  • Overlapping sp hybridized orbitals create sigma (σ) bonds, while unhybridized p orbitals overlap to form pi bonds.
• Exercise:
  • Practice with SkillBuilder 1.6 by identifying hybridization states.

1.10 Bond Strength and Length

• Comparison of Bond Strengths:
  • Sigma bonds are stronger than pi bonds, typically requiring nearly double the energy to break.
• Bond Lengths According to s-Character:
  • The nature of overlap (sp$^{3}$-sp$^{3}$ versus sp-sp) influences bond lengths—more s-character correlates with shorter bonds.
  • Observed order of bond lengths: sp$^{3}$ (longest) > sp$^{2}$ > sp (shortest).

1.11 Molecular Geometry / VSEPR Theory

• Valence Shell Electron Pair Repulsion (VSEPR) Theory:
  • VSEPR posits that valence electrons (both shared and lone pairs) repel one another.
• Molecular Geometry Prediction:
  • Initiate geometry determination with the steric number, providing a preliminary prediction.
  • Steric Number and Hybridization:
  • Steric number = 4 results in sp$^{3}$ hybridization.
  • Steric number = 3 implies sp$^{2}$ hybridization.
  • Steric number = 2 indicates sp hybridization.
• Note on Further Factors:
  • Additional considerations affecting hybridization/molecular geometry will be covered in Chapter 2.

1.11 Molecular Geometry – sp$^{3}$ Configuration

• Geometry for sp$^{3}$ Atoms:
  • Atoms with sp$^{3}$ hybridization present tetrahedral electron group geometry with four valence electron pairs.
  • Methane has four equivalent bonds, yielding equal bond angles.
  • Ammonia’s bond angles are slightly less than tetrahedral, while water has bond angles that are even smaller.

1.11 Molecular Geometry Overview

• Describing Molecular Geometry:

• Molecular geometry refers purely to atoms bonded to the central atom, while electron group geometry includes lone pairs.

• Example Calculation for CH$_{2}$O:
  • The electron pairs in sp$^{2}$ hybridized orbitals give trigonal planar electron group geometry (steric number = 3 = trigonal planar).
• Example Calculation for CO$_{2}$:
  • When steric number = 2, the geometry is linear, and the atom is sp-hybridized.

1.12 Molecular Polarity & Dipoles

• Formation of Polar Bonds:
  • Electronegativity differences result in the creation of polar bonds.
• Induction and Dipole Moment:
  • Induction refers to the shifting of electrons within an orbital to create a dipole moment.
  • Dipole moment equation:
  • Dipole moment = (amount of partial charge) × (distance of separation between δ+ and δ−).

1.12 Evaluating Net Dipole Moment

• Handling Multiple Polar Bonds:
  • In molecules with several polar bonds, the net dipole moment is the vector sum of all individual bond dipoles.
  • Accurate assessment of molecular polarity requires knowledge of molecular geometry.
  • Failure to represent the molecule in correct geometry may lead to incorrect polarity evaluations.

1.13 Intermolecular Forces

• Impact on Properties:
  • Various properties—solubility, boiling point, density, state of matter, melting point—are influenced by inter-molecular attractions.
• Types of Intermolecular Attractions:
  • Dipole-dipole interactions.
  • Hydrogen bonding.
  • Dispersion forces (also known as London forces or fleeting dipole-dipole forces).

1.13 Understanding Dipole-Dipole Attractions

• Mechanism of Dipole-Dipole Forces:
  • These forces arise when polar molecules align their opposing charges.
  • Example: Acetone’s permanent dipole is caused by differing electronegativity between carbon and oxygen.
• Effects on Physical Properties:
  • Dipole-dipole interactions raise the boiling and melting points of acetone.
  • In contrast, isobutylene lacks significant dipole moments, resulting in lower mp and bp.

1.13 Comparative Examples of Boiling Points

• Dissimilarities in Properties:
  • Isobutylene and acetone exhibit markedly different melting and boiling points due to the presence or absence of dipole-dipole interactions.
  • Isobutylene is less polar and experiences weaker dipole-dipole attractions, thus has a lower boiling point compared to acetone, which is more polar.

1.13 Hydrogen Bonding / Introduction

• Strength of Hydrogen Bonds:
  • Hydrogen bonds represent a strong type of dipole-dipole attraction.
  • These bonds are potent due to large partial positive and negative charges.
• Formation of Hydrogen Bonds:
  • Hydrogen bonding occurs between H atoms covalently bonded to electronegative atoms (N, O, F) and the lone pairs on other electronegative atoms.

1.13 Hydrogen Bonds and Electronegativity

• Role of Electronegativity:
  • Hydrogen must be sharing electrons with highly electronegative elements (O, N, or F) to manifest a large partial positive charge.
  • The significant δ+ charge on H can attract large δ− charges from other molecules, which explains variations in boiling points among isomers based on their hydrogen bonding capability.

1.13 Protic and Aprotic Solvents

• Solvent Classification:
  • Protic solvents engage in hydrogen bonding.
  • Aprotic solvents do not engage in hydrogen bonding.

1.13 London Dispersion Forces

• Attraction in Nonpolar Molecules:
  • Nonpolar molecules (with negligible dipoles) experience attractions due to transient dipole moments, referred to as London dispersion forces.
• Structure of Nonpolar Molecules:
  • These molecules typically exhibit an even electron distribution that balances positive charges, but electron motion can produce temporary dipoles.

1.13 Definition and Mechanism of London Dispersion Forces

• Formation of Temporary Dipoles:
  • The random motion of electrons within a molecule can yield an uneven electronic distribution, creating a temporary dipole.
  • This, in turn, can induce temporary dipoles in proximate molecules, leading to fleeting attractions.
  • Despite their weakness, numerous such interactions can accumulate and become significant.

1.13 Correlation Between Mass and Boiling Points

• Effect of Mass on Boiling Points:
  • London dispersion forces explain the trend where heavier molecules tend to have higher boiling points.
• Molecular Structure and Surface Area:
  • Increased branching within a molecule decreases its surface area, leading to weaker London dispersion forces.

1.14 Solubility / Polar vs Apolar Compounds

• Principle of Solubility:
  • The principle of “like dissolves like” holds true.
• Compatibility of Polar Compounds:
  • Polar compounds generally mix well with other polar compounds due to H-bonding and strong dipole-dipole interactions.
• Compatibility of Nonpolar Compounds:
  • Nonpolar compounds effectively mix with other nonpolar compounds, as weak attractions yield low resistance to mixing.