Chapter 1: Review of General Chemistry

What Is Organic Chemistry?

Organic chemistry is the study of compounds primarily based on carbon. Historically, organic compounds were thought to originate only from living organisms, while inorganic compounds came from minerals. Modern definitions focus on carbon-containing compounds, regardless of their source.

• Early definition: Organic compounds were believed to possess a 'vital force' from living organisms.

• Current definition: Organic compounds are those based on carbon atoms.

• Example: The synthesis of urea from ammonium cyanate demonstrated that organic compounds can be made from inorganic sources.

What Makes Carbon So Special?

Carbon is unique in its ability to form stable covalent bonds by sharing electrons, unlike elements to its left (which give up electrons) or right (which accept electrons) in the periodic table.

• Electron sharing: Carbon forms four covalent bonds, allowing for diverse molecular structures.

• Versatility: Carbon's ability to bond with itself and other elements is foundational to organic chemistry.

The Structure of an Atom

Atoms consist of a nucleus (protons and neutrons) surrounded by an electron cloud. The atomic number indicates the number of protons, which determines the element's identity.

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles in orbitals around the nucleus.

• Example: Carbon has atomic number 6, so it has 6 protons and, in a neutral atom, 6 electrons.

Isotopes

Isotopes are atoms of the same element (same atomic number) with different mass numbers due to varying numbers of neutrons.

• Example: Carbon has three common isotopes: 12C, 13C, and 14C.

• Application: Isotopes are used in radiocarbon dating and tracing mechanisms in organic reactions.

The Electronic Configurations of the First 11 Elements

Electron configuration describes the arrangement of electrons in atomic orbitals, following the Aufbau principle, Pauli exclusion principle, and Hund’s rule.

• Aufbau principle: Electrons fill the lowest energy orbitals first.

• Pauli exclusion principle: No more than two electrons can occupy the same orbital.

• Hund’s rule: Electrons fill empty orbitals before pairing up.

Atoms in the First Column of the Periodic Table Lose an Electron

Elements like lithium and sodium achieve a filled outer shell by losing their single valence electron, forming positively charged ions.

• Stability: Atoms are most stable with a filled or empty outer shell.

• Example:

Atoms on the Right Side of the Periodic Table Gain an Electron

Elements such as fluorine and chlorine achieve a filled outer shell by gaining an electron, forming negatively charged ions.

• Example:

• Example:

A Hydrogen Atom Can Lose or Gain an Electron

Hydrogen can achieve stability by either losing its electron (forming a proton) or gaining an electron (forming a hydride ion).

• Example: (proton)

• Example: (hydride ion)

Achieving a Filled Outer Shell by Sharing Electrons

Atoms can achieve stable electron configurations by sharing electrons, forming covalent bonds. This is the basis for molecular formation in organic chemistry.

• Covalent bond: A bond formed by the sharing of electrons between atoms.

• Example: Two fluorine atoms share electrons to form .

How Many Bonds Does an Atom Form?

The number of bonds an atom forms is determined by its need to achieve a stable electron configuration. Carbon forms four bonds, nitrogen three, oxygen two, and hydrogen one.

• Example: Water (H2O), ammonia (NH3), methane (CH4).

Phosphorus and Sulfur Atoms

Phosphorus and sulfur can have expanded octets due to the presence of d orbitals, allowing them to be surrounded by more than eight electrons.

• Example: Phosphorus can have 10 electrons, sulfur can have 10 or 12 electrons in its valence shell.

Nonpolar and Polar Covalent Bonds

Covalent bonds can be classified as nonpolar or polar based on the difference in electronegativity between the bonded atoms.

• Nonpolar covalent bond: Atoms have similar electronegativities; electrons are shared equally.

• Polar covalent bond: Atoms have different electronegativities; electrons are shared unequally.

• Direction of polarity: Indicated by an arrow pointing toward the more electronegative atom.

The Greater the Difference in Electronegativity, the More Polar the Bond

Bond polarity increases with greater differences in electronegativity. Ionic bonds form when electrons are not shared, and atoms are held together by electrostatic attraction.

• Electronegativity difference: Determines bond type (nonpolar, polar, ionic).

Dipole Moment of a Bond

The dipole moment is a measure of bond polarity, depending on the difference in electronegativity and the distance between charges.

• Formula: (where is charge, is distance)

Electrostatic Potential Maps

Electrostatic potential maps visually represent the distribution of electron density in molecules, highlighting regions of partial positive and negative charge.

• Application: Used to predict reactivity and intermolecular interactions.

Lewis Structures

Lewis structures show how atoms are bonded and indicate lone pairs and formal charges. They are fundamental for representing organic molecules.

• Lone pairs: Non-bonding pairs of electrons.

• Formal charge: Calculated as:

Carbon, Nitrogen, Oxygen, and Halogen Bonding Patterns

Each element has a characteristic bonding pattern. Deviations from these patterns result in formal charges.

• Carbon: Forms 4 bonds.

• Nitrogen: Forms 3 bonds and has 1 lone pair.

• Oxygen: Forms 2 bonds and has 2 lone pairs.

• Halogens: Form 1 bond and have 3 lone pairs.

• Hydrogen: Forms 1 bond.

The Number of Bonds Plus the Number of Lone Pairs Equals 4

For second-row elements (C, N, O, F), the sum of bonds and lone pairs always equals four, reflecting their valence shell electron configuration.

How to Draw a Lewis Structure

Drawing Lewis structures involves determining the total number of valence electrons, arranging atoms to satisfy octet rules, and checking for formal charges.

• Steps: Count valence electrons, arrange atoms, assign bonds and lone pairs, check octets, and calculate formal charges.

Kekulé, Condensed, and Skeletal Structures

Organic molecules can be represented in various ways: Kekulé structures (showing all bonds), condensed structures (omitting some bonds), and skeletal structures (showing only carbon skeletons).

• Kekulé: Shows all bonds, omits lone pairs.

• Condensed: Omits some or all bonds, shows groups.

• Skeletal: Shows carbon-carbon bonds as lines, omits carbons and hydrogens bonded to carbons.

Atomic Orbitals and Electron Behavior

Atomic orbitals are regions where electrons are most likely to be found. Electrons behave as standing waves, and orbitals have distinct shapes and energies.

• s orbitals: Spherical shape.

• p orbitals: Dumbbell shape, with opposite phases in lobes.

Bond Formation: Sigma and Pi Bonds

Bonds form by the overlap of atomic orbitals. Sigma (σ) bonds result from end-to-end overlap, while pi (π) bonds result from side-to-side overlap of p orbitals.

• Sigma bond: Strongest type of covalent bond, found in single bonds.

• Pi bond: Weaker, found in double and triple bonds.

Hybridization and Molecular Geometry

Hybridization is the mixing of atomic orbitals to form new hybrid orbitals, determining molecular geometry and bond angles.

• sp3 hybridization: Tetrahedral geometry, bond angle ≈ 109.5°.

• sp2 hybridization: Trigonal planar geometry, bond angle ≈ 120°.

• sp hybridization: Linear geometry, bond angle ≈ 180°.

Bond Strength, Bond Length, and Bond Order

The strength and length of a bond depend on the number of bonds (bond order) and the hybridization of the orbitals involved.

• Bond order: Single (1), double (2), triple (3).

• Higher bond order: Stronger and shorter bonds.

• More s character: Shorter and stronger bonds, larger bond angles.

Dipole Moments of Molecules

The dipole moment of a molecule depends on the magnitude and direction of individual bond dipoles. Symmetrical molecules do not have a net dipole moment.

• Example: Methane (CH4) is symmetrical and has no dipole moment.

Summary of Key Concepts

• The shorter the bond, the stronger it is.

• The greater the electron density in the region of orbital overlap, the stronger the bond.

• The more s character, the shorter and stronger the bond.

• The more s character, the larger the bond angle.

Learning Objectives

• Write ground-state electronic configurations for elements hydrogen through calcium.

• Describe the relative polarity of bonds and determine the direction of dipoles.

• Represent organic compounds using Lewis, Kekulé, condensed, and skeletal structures.

• Assign lone pairs and calculate formal charges.

• Determine hybridization of C, O, or N atoms from molecular formula.

• Describe how molecular geometry is determined by hybridization.

• Describe how hybridization affects bond strength and length.

• Describe how bond order affects bond length and strength.