Chapter 2 Notes: Atoms, Bonding, and Biomolecules

Background context
- The four major atoms that make up over 95% of the human body are carbon (C), hydrogen (H), nitrogen (N), and oxygen (O).
- Everything important in biochemistry (proteins, fats, nucleic acids, carbohydrates) is built from combinations of these atoms.
- The smallest building block is the atom; how atoms combine and bond determines the structure and function of molecules.

Subatomic Particles

Subatomic particles and charges
- Protons: positively charged (+1)
- Neutrons: neutral (no charge; add mass)
- Electrons: negatively charged (-1)
- In chemical bonding, focus is on protons and electrons because of their charge interactions.
Basic particle locations
- Protons and neutrons cluster in the nucleus (the atom’s core).
- Electrons orbit around the nucleus in regions called orbitals or energy levels.
- Nucleus analogy: like the sun; electrons orbit around it, with rules governing how many can occupy each shell.

Atomic Structure and the Nucleus

Nuclear composition and location
- Nucleus contains protons and neutrons.
- Protons + neutrons in the nucleus; electrons orbiting outside.
Atomic number (Z)
- Defined as the number of protons in the atom: Z = \text{# protons}.
- In a neutral atom, the number of electrons equals the number of protons: \text{#electrons} = Z\,.$
Examples of atomic numbers and their basic structure
- Hydrogen: Z = 1 → 1 proton; neutral atom has 1 electron; electron occupies the first orbital.
- Helium: Z = 2 → 2 protons, 2 electrons; both electrons reside in the first orbital.
- Neon: Z = 10 → 2 in the first orbital, 8 in the second orbital; the outer shell is full.

Electron Orbitals, Shells, and the Octet Rule

Orbitals and energy levels
- Electrons inhabit orbitals/shells around the nucleus, not a literal planetary path but a similar idea.
- The first shell (closest to the nucleus) is most strongly attracted to the nucleus; electrons here are tightly bound.
Octet rule (capacity of electron shells)
- First shell can hold up to 2 electrons.
- All shells beyond the first can hold up to 8 electrons (the octet for stability, though in very large atoms more complex configurations exist).
- In this course, we mainly consider up to two or three shells away from the nucleus (large atoms beyond chlorine are not the focus here).
Valence electrons
- The electrons in the outermost shell are called valence electrons.
- These determine how atoms bond with others.
- If the outermost shell is full, the atom is usually inert (not reactive).
- If there are empty spots in the outermost shell, the atom is reactive and can form bonds to fill that shell.
Inert vs reactive examples
- Helium: Z = 2; first shell full (2 electrons) → inert; tends not to react.
- Neon: Z = 10; first shell full (2), second shell full (8) → inert.
Carbon example to illustrate distribution
- Carbon: Z = 6; distribution: 2 electrons in the first shell, 4 in the second shell.
- This places carbon with 4 valence electrons, which strongly influences its bonding versatility.

Building Molecules: Covalent and Ionic Bonds

Covalent bonds (sharing electrons)
- Covalent bonds involve sharing electrons between atoms.
- Bonding leads to molecules composed of multiple atoms.
- In drawings, a single covalent bond is represented by a single line and corresponds to one pair of electrons (i.e., 2 electrons total).
Ionic bonds (transfer of electrons)
- Ionic bonds arise when electrons are transferred from one atom to another.
- This transfer creates charged species: a cation (positively charged) and an anion (negatively charged).
- Opposite charges attract and hold the ions together.
- Example sequence (sodium and chlorine)
- Sodium (Na): Z = 11; loses one electron from the outer shell to achieve a stable configuration (outer shell becomes filled up to 2 in the first shell and 8 in the second, with the remaining electron moved to the third shell).
- Chlorine (Cl): Z = 17; gains one electron to complete its outer shell (now it has a full outer shell in the third level).
- Resulting charges: Na becomes \text{Na}^+ $; Cl becomes$ \text{Cl}^-.
- The electrostatic attraction between Na+ and Cl− forms the ionic bond, producing NaCl.
Sharing vs transfer in context
- Covalent bonds can be nonpolar or polar depending on electron sharing symmetry (see below).
- Ionic bonds involve complete electron transfer and formation of ions with charges that attract.

Covalent Bonding: Nonpolar vs Polar

Nonpolar covalent bonds
- Electrons are shared equally between atoms.
- Example: carbon dioxide, \mathrm{CO_2}, where electrons are shared to form bonds that are effectively equal along a straight line; this results in a nonpolar molecule.
Polar covalent bonds
- Electrons are shared but not equally; electron density is drawn toward one atom more than the other.
- Example: water, \mathrm{H_2O}, where oxygen attracts electron density more than hydrogen.
- Partial charges are indicated with delta notation: \delta^- $on the more electronegative atom (oxygen) and$ \delta^+ on the others (hydrogens).
- Resulting partial charges create dipoles and can influence intermolecular interactions.

Hydrogen Bonds and Their Significance

What is a hydrogen bond?
- A hydrogen bond is a weak attraction between the partial negative charge on one molecule (e.g., a lone pair on oxygen) and the partial positive charge on a hydrogen atom bonded to a highly electronegative atom in another molecule.
- Although weaker than covalent or ionic bonds, hydrogen bonds are cumulatively strong and crucial in biology.
Why they matter
- Water’s properties (e.g., surface tension, cohesion) are greatly influenced by hydrogen bonding among H2O molecules, enough to support small organisms (e.g., water striders) on the surface.
- Hydrogen bonds stabilize the DNA double helix by linking complementary bases across strands, enabling the structure essential for genetic information storage and replication.

Connections to Biomolecules and Real-World Relevance

How bonding shapes biomolecules
- The four major atoms (C, H, N, O) form the backbone and functionality of proteins, fats, nucleic acids, and carbohydrates through covalent and noncovalent interactions.
- Bonding patterns determine the three-dimensional shapes, reactivity, and interactions of biomolecules.
Foundational principles tied to broader chemistry and biology
- Atomic number and neutrality govern the electron configuration and potential for bonding.
- The outermost electrons (valence electrons) guide chemical reactivity and compound formation.
- Electronegativity differences drive covalent bond polarity and the formation of hydrogen bonds, which underpin many physiological processes.

Quick Reference: Key Formulas and Concepts

Atomic number: Z = \text{# protons}
Neutral atom electron count: \text{#electrons} = Z
First shell capacity: 2 electrons
Other shells capacity (octet rule): up to 8 electrons
Covalent bond (single): two electrons shared between two atoms ⇒ represented by one line
Ionic bond: transfer of electrons creates ions; opposite charges attract
Nonpolar covalent bond: equal sharing of electrons
Polar covalent bond: unequal sharing; partial charges \delta^- $on the more electronegative atom and$ \delta^+ on the other
Hydrogen bond: interaction between partial charges across molecules; important for water properties and DNA structure

Carbon example recap

Carbon with Z = 6$$: distribution 2 in the first shell, and 4 in the second shell; this outer-shell configuration contributes to carbon’s versatile bonding patterns (forming diverse molecules essential for life)

Practice observations to prepare for class

Identify whether a bond is covalent or ionic based on electron transfer vs sharing
Determine whether a molecule is polar or nonpolar based on electron sharing and electronegativity differences
Recognize the role of hydrogen bonds in biological structures (e.g., DNA, proteins, water-dependent processes)