Small Molecules and the Chemistry of Life

Overview

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

Atomic Structure

The Helium Atom and Subatomic Particles

All matter is composed of atoms, which are the basic building blocks of elements and molecules. Understanding atomic structure is crucial for explaining the properties and behaviors of biological molecules.

• Atom: The smallest unit of matter that retains the properties of an element.

• Subatomic particles:

  • Protons: Positively charged particles with mass, located in the nucleus.

  • Neutrons: Neutral particles with mass, also in the nucleus.

  • Electrons: Negatively charged particles with negligible mass, orbiting the nucleus in energy levels (electron clouds).

• Nucleus: Dense central part of the atom containing protons and neutrons.

Additional info: The mass of an electron is about 1/1836th that of a proton or neutron and is usually ignored in atomic mass calculations.

Atomic Number, Mass Number, and Isotopes

Atoms are identified by their number of protons, which determines the element. The mass number includes both protons and neutrons.

• Atomic number (Z): Number of protons in the nucleus; unique to each element.

• Mass number (A): Total number of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element with different numbers of neutrons (and thus different mass numbers).

Example: Carbon has three common isotopes: 12C (6 protons, 6 neutrons), 13C (6 protons, 7 neutrons), 14C (6 protons, 8 neutrons).

Atomic Mass Unit and Dalton

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

• Electron mass: grams (usually ignored).

Atoms Bond to Form Molecules

Chemical Bonds and Stability

Atoms interact to achieve stable electron configurations, often by filling their outermost (valence) electron shells. This leads to the formation of chemical bonds.

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve a full outer shell of 8 electrons.

• Chemical bond: Attractive force linking atoms together in molecules.

Types of Chemical Bonds

• Ionic bonds: Formed when one atom donates an electron to another, resulting in oppositely charged ions (cations and anions) that attract each other.

• Covalent bonds: Formed when two atoms share one or more pairs of electrons.

• Hydrogen bonds: Weak attractions between a hydrogen atom covalently bonded to an electronegative atom (like oxygen or nitrogen) and another electronegative atom.

Covalent Bonding and Molecular Structure

• Single bond: Sharing one pair of electrons (e.g., C–H).

• Double bond: Sharing two pairs of electrons (e.g., C=C).

• Triple bond: Sharing three pairs of electrons (e.g., N≡N).

• Bond orientation: The geometry of covalent bonds determines molecular shape and function.

Example: Methane (CH4) has a tetrahedral geometry due to four single covalent bonds between carbon and hydrogen.

Polar and Nonpolar Covalent Bonds

• Electronegativity: Measure of an atom's ability to attract electrons in a bond.

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

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

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

Ionic Bonds and Ion Formation

• Cation: Atom that loses electrons and becomes positively charged.

• Anion: Atom that gains electrons and becomes negatively charged.

• Ionic compound: Formed by electrostatic attraction between cations and anions (e.g., NaCl).

Atoms Change Partners in Chemical Reactions

Chemical Reactions and Energy

Chemical reactions involve the rearrangement of atoms and the making or breaking of chemical bonds. These processes are fundamental to metabolism and energy flow in biological systems.

• Reactants: Substances that undergo change during a reaction.

• Products: Substances formed as a result of a reaction.

• Redox reactions: Involve the transfer of electrons; oxidation is loss of electrons, reduction is gain of electrons.

Example: Combustion of propane:

Conservation Laws

• Law of conservation of mass: Matter is neither created nor destroyed in chemical reactions.

• First law of thermodynamics: Energy is conserved; it can change forms but is not created or destroyed.

The Importance of Water for Life

Water's Unique Properties

Water is essential for life due to its chemical and physical properties, which arise from its molecular structure and ability to form hydrogen bonds.

• Polarity: Water is a polar molecule, with partial negative charge on oxygen and partial positive charges on hydrogens.

• Hydrogen bonding: Water molecules form hydrogen bonds, leading to high cohesion, surface tension, and solvent capabilities.

• Solvent properties: Water dissolves many polar and ionic substances, facilitating biochemical reactions.

Hydrophilic and Hydrophobic Interactions

• Hydrophilic substances: "Water-loving"; typically polar or charged, dissolve well in water.

• Hydrophobic substances: "Water-fearing"; typically nonpolar, do not dissolve in water and tend to aggregate.

Acids, Bases, and pH

Acids and bases are important in biological systems for maintaining pH balance, which affects molecular structure and function.

• Acid: Substance that releases hydrogen ions (H+) in water.

• Base: Substance that accepts hydrogen ions or releases hydroxide ions (OH-).

• pH: Measure of hydrogen ion concentration;

• Neutral pH: pH = 7 (pure water)

• Acidic: pH < 7; Basic (alkaline): pH > 7

Example: 1 M HCl solution has [H+] = 1 M, so pH = 0 (strongly acidic).

Moles, Molar Mass, and Concentration

• Mole: Standard unit for amount of substance; 1 mole = particles (Avogadro's number).

• Molar mass: Mass of one mole of a substance, in grams per mole (g/mol).

• Molarity (M): Concentration of solute in solution; moles of solute per liter of solution.

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

Periodic Table and Biologically Important Elements

Key Elements in Biology

The periodic table organizes elements by atomic number. The most biologically important elements are carbon (C), hydrogen (H), oxygen (O), nitrogen (N), phosphorus (P), and sulfur (S).

• Atomic number: Number of protons; determines element identity.

• Usual covalent bonding capabilities:

  • Hydrogen (H): 1 bond

  • Oxygen (O): 2 bonds

  • Nitrogen (N): 3 bonds

  • Carbon (C): 4 bonds

  • Phosphorus (P): 5 bonds

  • Sulfur (S): 2 bonds

Table: Biologically Important Elements and Their Covalent Bonding Capabilities

Summary

• Atoms are the fundamental units of matter, composed of protons, neutrons, and electrons.

• Chemical bonds (ionic, covalent, hydrogen) enable atoms to form molecules and compounds essential for life.

• Water's unique properties, including polarity and hydrogen bonding, make it vital for biological processes.

• Acids, bases, and pH regulation are crucial for maintaining cellular function.

• The periodic table highlights the elements most important for biological molecules.