Small Molecules and the Chemistry of Life

Overview

This section introduces the foundational chemical principles underlying biological systems, focusing on atomic structure, chemical bonding, molecular interactions, and the unique properties of water essential for life.

Atomic Structure

Atoms and Subatomic Particles

Atoms are the smallest units of matter that retain the properties of an element. All matter is composed of atoms, which consist of three main subatomic particles: protons, neutrons, and electrons.

• Protons: Positively charged particles located in the nucleus; each has a mass of approximately 1 atomic mass unit (amu) or 1 dalton.

• Neutrons: Neutral particles (no charge) also found in the nucleus; mass is about 1 amu.

• Electrons: Negatively charged particles orbiting the nucleus in electron clouds; mass is negligible compared to protons and neutrons (about 1/1836th of a proton's mass).

Nucleus: The dense central part of an atom, containing protons and neutrons.

Element: A pure substance composed of only one kind of atom (e.g., Fe, O, H).

Atomic Number and Mass Number

• Atomic Number (Z): The number of protons in the nucleus; unique to each element and determines its identity.

• Mass Number (A): The total number of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

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

Atomic Mass Unit and Atomic Weight

• Atomic Mass Unit (amu): Standard unit for atomic and molecular mass; 1 amu = 1 dalton = grams.

• Atomic Weight: The average mass of an element's atoms, accounting for the relative abundance of isotopes.

The Periodic Table

Organization and Key Elements

The periodic table arranges elements by increasing atomic number. Elements essential for life include carbon (C), hydrogen (H), oxygen (O), nitrogen (N), phosphorus (P), and sulfur (S).

• Groups: Columns with elements sharing similar chemical properties.

• Periods: Rows indicating increasing energy levels (electron shells).

Example: All atoms with 6 protons are carbon; all with 8 protons are oxygen.

Electron Configuration and Chemical Reactivity

Electron Shells and Orbitals

Electrons occupy specific energy levels (shells) around the nucleus, organized into orbitals.

• First shell: 1 s orbital (holds up to 2 electrons).

• Second shell: 1 s orbital and 3 p orbitals (holds up to 8 electrons).

• Valence shell: The outermost shell; determines chemical reactivity.

Pauli Exclusion Principle: Each orbital holds a maximum of two electrons with opposite spins.

Aufbau Principle: Electrons fill the lowest available energy levels first.

Stability and Reactivity

• Atoms are most stable when their valence shell is full (e.g., noble gases).

• Atoms with incomplete valence shells are reactive and tend to gain, lose, or share electrons to achieve stability (octet rule).

Chemical Bonds and Molecules

Types of Chemical Bonds

• Covalent Bonds: Atoms share one or more pairs of electrons. Strongest type of bond in biological molecules.

• Ionic Bonds: Formed when one atom donates an electron to another, resulting in oppositely charged ions that attract each other.

• Hydrogen Bonds: Weak attractions between a hydrogen atom covalently bonded to an electronegative atom (like O or N) and another electronegative atom.

Covalent Bonding and Molecular Geometry

• Single, double, and triple covalent bonds involve sharing one, two, or three pairs of electrons, respectively.

• Single bonds allow free rotation; double and triple bonds restrict rotation, affecting molecular shape.

• Bond strength: Triple > Double > Single.

Example: Oxygen molecule (O2) has a double bond; nitrogen molecule (N2) has a triple bond.

Polarity and Electronegativity

• Electronegativity: The ability of an atom to attract electrons in a bond.

• Nonpolar covalent bond: Electrons are shared equally (e.g., H2).

• Polar covalent bond: Electrons are shared unequally, creating partial charges (e.g., H2O).

Example: In water, oxygen is more electronegative than hydrogen, resulting in a partial negative charge on oxygen and partial positive charges on hydrogens.

Ionic Bonds and Ion Formation

• When the difference in electronegativity is large, electrons are transferred, forming ions.

• Cation: Positively charged ion (loses electrons).

• Anion: Negatively charged ion (gains electrons).

Example: Sodium (Na) donates an electron to chlorine (Cl), forming Na+ and Cl-.

Hydrogen Bonds and Biological Importance

• Hydrogen bonds stabilize the structures of proteins and nucleic acids (e.g., DNA double helix).

• They are weaker than covalent and ionic bonds but crucial for molecular interactions.

Chemical Reactions

Types of Chemical Reactions

• Redox (Oxidation-Reduction) Reactions: Involve the transfer of electrons between molecules.

• Oxidation: Loss of electrons or hydrogen atoms.

• Reduction: Gain of electrons or hydrogen atoms.

Example: Cellular respiration is a redox process: glucose is oxidized, oxygen is reduced.

Conservation Laws

• Law of Conservation of Mass: Matter is neither created nor destroyed in chemical reactions.

• Law of Conservation of Energy: Energy is neither created nor destroyed, only transformed.

The Importance of Water for Life

Properties of Water

• Polarity: Water is a polar molecule, allowing it to form hydrogen bonds.

• Cohesion and Adhesion: Water molecules stick to each other (cohesion) and to other substances (adhesion).

• High Specific Heat: Water can absorb or release large amounts of heat with little temperature change.

• Solvent Abilities: Water dissolves many substances, making it the "universal solvent" for biological systems.

Hydrophilic and Hydrophobic Substances

• Hydrophilic: "Water-loving"; substances that dissolve easily in water (e.g., salts, sugars).

• Hydrophobic: "Water-fearing"; substances that do not dissolve in water (e.g., oils, fats).

Solutions, Moles, and pH

Mole Concept and Molarity

• Mole: The amount of substance containing particles (Avogadro's number).

• Molarity (M): Number of moles of solute per liter of solution.

Example: To make a 1M solution of glucose (C6H12O6), dissolve 180 g in water to make 1 liter.

Acids, Bases, and pH

• Acids: Substances that release hydrogen ions (H+) in solution.

• Bases: Substances that accept H+ or release hydroxide ions (OH-).

• pH: Measure of hydrogen ion concentration; .

• pH Scale: Ranges from 0 (very acidic) to 14 (very basic); 7 is neutral.

Example: Pure water has [H+] = M, so pH = 7.

Ionization of Water and Buffering

• Water can ionize slightly:

• At 25°C, (the ion product constant, ).

• Biological systems use buffers to minimize changes in pH and maintain homeostasis.

Table: pH Values of Some Familiar Substances

Summary Table: Types of Chemical Bonds

Key Equations

• Atomic Number:

• Mass Number:

• pH:

• Molarity:

• Ion Product of Water: at 25°C

Conclusion

Understanding atomic structure, chemical bonding, and the properties of water is essential for comprehending the molecular basis of life. These principles form the foundation for more advanced topics in biology, including metabolism, genetics, and cellular structure.