Chapter 1: Structure and Bonding - Vocabulary Flashcards

Electronic Structure and Bonding

• Organic chemistry focuses on carbon compounds; living organisms are organic by traditional definition; organic chemistry studies compounds derived from living organisms and natural products.

• Electronic Structure of the Atom:

  • An atom has a dense, positively charged nucleus surrounded by a cloud of electrons.

  • Electron density is highest at the nucleus and drops off exponentially with distance from the nucleus in any direction.

  • Visual representations can be shown in two dimensions and three dimensions (cutaway views).

  • The 2s and 2p orbitals exist around the nucleus; electron density and nodes define orbital shapes.

The 2p Orbitals

• There are three 2p orbitals oriented at right angles to each other.

• Each p orbital consists of two lobes.

• Each orbital is labeled according to its orientation along the x, y, or z axis (px, py, pz).

Isotopes

• Isotopes are atoms with the same number of protons but different numbers of neutrons.

• Mass number is the sum of protons and neutrons in an atom: A = Z + N.

• Example: Carbon-12 (C-12) has 6 protons and 6 neutrons; Carbon-14 (C-14) has 6 protons and 8 neutrons.

Electronic Configurations of Atoms

• Valence electrons are electrons in the outermost shell of the atom.

• Aufbau principle: fill the lowest energy orbitals first.

• Hund’s rule: when there are two or more orbitals of the same energy (degenerate), electrons enter different orbitals rather than pairing in the same orbital.

Bond Formation: Key Rules

• Valence electrons: number of electrons in the outermost shell.

• Duet rule: H seeks 2 electrons to complete its outer shell.

• Octet rule: C, N, O, and other second-row elements seek 8 electrons in their valence shell.

• Types of bonding: Ionic bonding or Covalent bonding.

Ionic Bonding

• Atoms may transfer electrons to achieve noble gas configurations (full valence shells).

• Resulting opposite charges (cation and anion) attract each other, forming an ionic bond.

Covalent Bonding

• Electrons are shared between atoms to complete the octet.

• Nonpolar covalent (pure covalent): electrons are shared evenly between atoms.

• Polar covalent: electrons are shared unevenly, creating partial charges.

Electronegativity and Bond Polarity

• Electronegativity differences help predict whether a bond is polar and the direction of the dipole moment.

• If electronegativities are similar (e.g., C–H), bonds are considered nonpolar.

Dipole Moment

• Definition: μ = q × r, where q is the charge separation and r is the distance over which charge separation occurs.

• Electrostatic Potential Map (EPM) illustrates charge distribution: red indicates partially negative regions; blue indicates partially positive regions.

Pauling Electronegativities

• Electronegativity values can predict polarity and dipole direction.

• C–H bonds are typically nonpolar due to similar electronegativities.

Lewis Structures

• Examples: CH4, NH3, H2O, Cl2 illustrate the distribution of valence electrons.

• Lewis diagrams show valence electrons as dots; bonds as lines.

• Carbon typically forms 4 bonds; nitrogen 3 or 4; oxygen 2 or 1 with charges; halogens 1.

• Example electron counts for common atoms (valence electrons): B = 3, C = 4, N = 5, halogens = 7.

Bonding Patterns

• Common Bonding Patterns in Organic Compounds (Uncharged):

  • Carbon: valence electrons = 4

  • Nitrogen: valence electrons = 5

  • Oxygen: valence electrons = 6

  • Hydrogen: valence electrons = 1

  • Halogens: valence electrons = 7

• This table helps recognize typical bonding patterns and charges in organic structures.

Nonbonding Electrons

• Also called lone pairs.

• Nonbonding electrons are valence-shell electrons not shared between atoms.

Multiple Bonding

• Double bond: sharing two pairs of electrons.

• Triple bond: sharing three pairs of electrons.

Formal Charges

• Formal charge = (group number, i.e., valence electrons) – (nonbonding electrons) – 1/2 × (shared electrons)

• Example usage: ammonia borane (borazane) H3N—BH3; formal charges help assess electron bookkeeping.

• Note: Formal charges are bookkeeping devices and may not reflect actual charges in all resonance structures.

Solved Problem: Formal Charge in Ammonia Borane

• Example: Determine FC on each atom in H3N—BH3.

• Approach: count valence electrons for each atom, count nonbonding electrons, count shared electrons, apply FC formula.

Condensed Structural Formulas and Bonding Patterns

• Condensed structural formulas omit explicit bonds; list groups attached to a central atom after the center (e.g., CH3CH3 instead of H3C-CH3).

• If there are identical groups, parentheses with subscripts are used: (CH3)2CH for isopropyl-like fragments.

• Examples:

  • Ethane: CH3CH3

  • Isobutane: (CH3)3CH or (CH3)2CHCH3 depending on representation

  • n-Hexane: CH3(CH2)4CH3 or CH3(CH2)4CH3 (depending on how shown)

• Condensed formulas for alkenes, nitriles, aldehydes, ketones, carboxylic acids, etc., can be written with or without showing all bonds.

• Table 1-2 provides examples of condensed structural formulas for several compounds.

Line-Angle Drawings (Skeletal Formulas)

• Bonds are drawn as lines; carbons are implied at line ends and intersections.

• Hydrogens attached to carbon are not shown.

• Nitrogen, oxygen, and halides must be shown explicitly.

• Double and triple bonds are shown explicitly.

• The convention emphasizes the carbon framework while reducing clutter of hydrogens.

Line-Angle Formulas (Continued)

• Examples include hexane, hex-2-ene, hexan-3-ol, cyclohex-2-en-1-one, 2-methylcyclohexan-1-ol, and nicotinic acid (niacin).

• Condensed structures may be converted to line-angle forms and vice versa.

Linear Combination of Atomic Orbitals (LCAO)

• Bonding arises from combining orbitals between two different atoms (molecular orbitals, MOs).

• Hybridization arises from combining orbitals on the same atom.

• Conservation of orbitals: the number of MOs equals the number of atomic orbitals combined.

• Constructive interference (in phase) increases amplitude; destructive interference (out of phase) cancels amplitude.

The Bonding Region and Electron Distribution

• In a diatomic molecule, electrons in the bonding region help attract both nuclei and reduce repulsion, enabling bond formation.

• EPM visuals show electron density distribution in the bonding region.

Sigma Bonding and MO Theory

• Sigma (σ) bonding: electron density lies along the internuclear axis; can involve s–s, p–p, s–p, or hybridized overlaps.

• Bonding MO is lower in energy than the original atomic orbitals; antibonding MO is higher.

• For H2, σ bond forms from 1s–1s overlap (in-phase).

• Antibonding MO (σ*) forms from out-of-phase overlap.

• Energy diagrams show bonding orbitals at lower energy and antibonding at higher energy.

s–s, p–p, and s–p Overlaps

• s–s overlap yields σ bonding MO.

• s–p overlap yields σ bonding MO and σ* antibonding MO.

• p–p overlap along the internuclear axis yields σ bonding MO; sideways overlap of parallel p orbitals yields π bonding MO.

Pi Bonding and Antibonding

• Sideways overlap of two parallel p orbitals creates a π bond (π MO and π* antibonding MO).

• Pi bonds are generally weaker than sigma bonds due to less effective overlap.

Multiple Bonds

• Double bond consists of one σ bond and one π bond.

• Triple bond consists of one σ bond and two π bonds.

Molecular Shapes and Hybridization

• Molecular shapes are explained by VSEPR theory: valence electron pairs repel and arrange to minimize repulsion.

• Hybridized orbitals are lower in energy when electron pairs are farther apart.

• Hybridization and corresponding geometries:

  • sp: linear geometry, 180° bond angle

  • sp2: trigonal planar geometry, 120° bond angle

  • sp3: tetrahedral geometry, 109.5° bond angle

sp, sp2, sp3 Hybrid Orbitals

• sp hybrids arise from combining one s and one p orbital; two sp orbitals form linear geometry.

• sp2 hybrids arise from combining one s and two p orbitals; three sp2 orbitals form trigonal planar geometry.

• sp3 hybrids arise from combining one s and three p orbitals; four sp3 orbitals form tetrahedral geometry.

Hybridization and Geometry Summary

• Hybridization table:

  • 2: sp — linear, 180°

  • 3: sp2 — trigonal planar, 120°

  • 4: sp3 — tetrahedral, 109.5°

Bonding in Ethylene and Acetylene

• Ethylene (C2H4): five sigma bonds formed by sp2 hybrid orbitals in a trigonal planar geometry.

  • Unhybridized p orbitals on each carbon are perpendicular to the plane of the sp2 hybrids and overlap to form a π bond, located above and below the σ bond.

• Acetylene (C2H2): sp hybridization leads to a triple bond consisting of one σ bond and two π bonds.

Rotation and Isomerism

• Rotation in single bonds:

  • Single bonds allow rotation, yielding various conformations.

• Rotation around double bonds:

  • Double bonds do not rotate; substituent arrangement around the double bond is fixed.

• Isomerism:

  • Constitutional (structural) isomers differ in bonding sequence; same molecular formula but different connectivity.

  • Stereoisomers differ only in spatial arrangement; geometric (cis/trans) is a common example for alkenes.

Constitutional Isomers

• Constitutional isomers have the same formula but different connectivity.

• They typically have different physical and chemical properties.

• The number of possible isomers grows rapidly with the number of carbon atoms.

Geometric Isomers: Cis and Trans

• Stereoisomers have the same order of bonding, but different spatial orientations.

• Cis and trans isomers occur in compounds with restricted rotation about a double bond, leading to distinct isomers where substituents are on the same or opposite sides.

• These are examples of geometric isomers; the lack of free rotation around C=C allows distinct arrangements.

Key Connections to Foundational Concepts

• The octet and duet rules underpin Lewis structures and Lewis electron counting.

• Orbital hybridization explains molecular geometry and the observed bond angles.

• MO theory provides a deeper view of bond formation beyond simple localized bonds (σ and π components).

• Electronegativity and bond polarity connect to the distribution of electron density and dipole moments in molecules.

• Isomerism links to chemical reactivity, physical properties, and separation techniques in synthesis.

Practical and Ethical Implications (General)

• Understanding bonding and molecular structure informs drug design, materials science, and environmental chemistry.

• Ethical considerations include responsible use of chemical knowledge in synthesis, safety, and environmental impact.

Notation and Formulas (LaTeX Examples)

• Dipole moment: oldsymbol{3 } = q \, r

• Formal charge: FC = V - N - \tfrac{1}{2}S

• Bonding geometries: sp (linear) angle = 180^{\circ}, sp2 (trigonal planar) angle = 120^{\circ}, sp3 (tetrahedral) angle = 109.5^{\circ}

Summary Tips for Exam

• Be able to predict bond types (ionic vs covalent) based on electronegativity differences.

• Draw Lewis structures for simple molecules and assign formal charges correctly.

• Identify lone pairs and count valence electrons to satisfy octets where applicable.

• Recognize line-angle drawings and convert between condensed formulas and skeletal structures.

• Apply VSEPR to predict molecular shapes and correlate with hybridization.

• Distinguish between sigma and pi bonds and understand how rotation affects single vs. multiple bonds.

• Understand the concept of isomerism, including constitutional vs geometric isomers, and identify examples in simple hydrocarbons.