Chapter 1: Structure and Bonding – Introduction and Review

Chapter 1: Structure and Bonding – Introduction and Review

Atoms and Ions

This section introduces the basic building blocks of matter: atoms, ions, and molecules, which are foundational to understanding organic chemistry.

• Atoms: Consist of a nucleus (protons and neutrons) surrounded by electrons.

• Molecules: Groups of atoms held together by covalent bonds (e.g., H2O, CH4).

• Ions: Atoms or molecules that have gained or lost electrons, resulting in a net charge.

  • Cation: Positively charged ion (loss of electrons).

  • Anion: Negatively charged ion (gain of electrons).

• Alkali and alkaline earth metals tend to form cations; halogens tend to form anions.

Ionic and Covalent Bonds

Chemical bonds are the forces that hold atoms together in compounds. The two main types are ionic and covalent bonds.

• Ionic Bonds: Electrostatic attraction between oppositely charged ions (e.g., Na+Cl-).

• Covalent Bonds: Sharing of electron pairs between atoms (e.g., H2, H2O).

• Polar Covalent Bonds: Electrons are shared unequally due to differences in electronegativity, resulting in partial charges.

Electronegativity and Bond Polarity

• Electronegativity: The tendency of an atom to attract electrons in a bond. Pauling scale is commonly used.

• If two atoms have different electronegativities, the bond is polar (e.g., O–H in water).

• Dipole Moment: A measure of bond polarity, defined as $\mu = q \times d$ (charge × distance).

Element	Electronegativity
H	2.1
C	2.5
N	3.0
O	3.5
F	4.0
Cl	3.0
Br	2.8
I	2.5

Additional info: Table values are typical Pauling electronegativities for elements common in organic chemistry.

Electronic Configuration

Electronic configuration describes the arrangement of electrons in atomic orbitals, which determines chemical properties and bonding.

• Electrons fill orbitals in order of increasing energy: $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s$ etc.

• Example (Carbon): 6 electrons: $1s^2 2s^2 2p^2$

• Example (Nitrogen): 7 electrons: $1s^2 2s^2 2p^3$

Lewis Structures

Lewis structures are diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule.

• Each bond is represented by a line (–), and lone pairs are shown as dots.

• Atoms (except H) tend to satisfy the octet rule (8 electrons in the valence shell).

• Examples:

  • Methane (CH4): All atoms have a complete octet (except H).

  • Ethylene (C2H4): Double bond between carbons.

  • Hydrogen cyanide (HCN): Triple bond between C and N.

Formal Charges

Formal charge is the charge assigned to an atom in a molecule, calculated by:

$\text{Formal charge} = \text{valence electrons} - \text{nonbonding electrons} - \frac{1}{2}(\text{shared electrons})$

• Helps identify charged atoms in molecules and resonance structures.

• Example: In NH4+, N has a formal charge of +1.

Resonance Structures

Resonance structures are different Lewis structures for the same molecule, showing delocalization of electrons.

• Resonance forms do not exist independently; the actual molecule is a hybrid.

• Curved arrows indicate movement of electron pairs.

• Guidelines for major contributors:

  • Structures with full octets and minimal charge separation are favored.

  • Negative charges should be on more electronegative atoms.

Structural Formulas

There are several ways to represent organic molecules:

Lewis Structure	Condensed Structure	Line-Angle Structure
Shows all atoms and bonds explicitly	Groups atoms together (e.g., CH3CH2OH)	Each vertex/line end is a carbon; hydrogens on carbons are implied

Example: Ethanol can be written as:

• Lewis: H–C–C–O–H (all atoms shown)

• Condensed: CH3CH2OH

• Line-angle: two lines with an endpoint for the OH group

Molecular Orbitals and Bonding

Molecular orbitals (MOs) are formed when atomic orbitals combine. The number of MOs equals the number of atomic orbitals combined.

• Sigma (σ) bond: Electron density is concentrated along the axis joining two nuclei (single bonds).

• Pi (π) bond: Electron density is above and below the plane of the nuclei (double/triple bonds).

• Bonding and antibonding MOs are formed; bonding MOs are lower in energy.

Hybridization and Molecular Shapes

Hybridization explains the observed shapes of molecules by mixing atomic orbitals to form new, equivalent hybrid orbitals.

• sp3 hybridization: 4 electron domains, tetrahedral shape, 109.5° bond angles (e.g., methane).

• sp2 hybridization: 3 electron domains, trigonal planar shape, 120° bond angles (e.g., ethylene).

• sp hybridization: 2 electron domains, linear shape, 180° bond angles (e.g., acetylene).

Number of Electron Domains	Hybridization
2	sp
3	sp2
4	sp3

Isomerism

Isomers are compounds with the same molecular formula but different structures or spatial arrangements.

• Constitutional (structural) isomers: Differ in the connectivity of atoms.

• Stereoisomers: Same connectivity, different spatial arrangement.

  • Geometric (cis-trans) isomers: Differ in arrangement around a double bond or ring.

  • Enantiomers: Non-superimposable mirror images.

  • Diastereomers: Not mirror images, differ at some but not all stereocenters.

• Geometric isomerism requires two different groups on each carbon of a double bond.

Example: C5H12 has three constitutional isomers (n-pentane, isopentane, neopentane).

Additional info: Stereoisomerism and the E/Z system for alkenes are covered in more detail in later chapters.