Water and Carbon: The Chemical Basis of Life

Introduction

This chapter explores the fundamental chemical principles that underlie biological processes, focusing on the roles of water and carbon in the origin and maintenance of life. Understanding the structure of atoms, ions, and molecules, as well as the unique properties of water and the chemistry of carbon, is essential for studying biology at the molecular level.

The Structure of Atoms, Ions, and Molecules

Basic Atomic Structure

• Atoms are the basic units of matter, composed of a nucleus (containing protons and neutrons) and orbiting electrons.

• Protons: Positively charged particles (+1 charge).

• Neutrons: Neutral particles (no charge).

• Electrons: Negatively charged particles (-1 charge) that orbit the nucleus.

• An atom is electrically neutral if it has equal numbers of protons and electrons.

Elements, Atomic Number, and Mass Number

• Element: A substance consisting entirely of one type of atom.

• Atomic number: The number of protons in the nucleus (defines the element).

• Mass number: The sum of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element with different numbers of neutrons (e.g., Carbon-12, Carbon-13, Carbon-14).

• Atomic weight: The average mass of all naturally occurring isotopes of an element, weighted by abundance.

• Radioactive isotopes: Unstable isotopes that decay over time, releasing radiation.

Electron Arrangement and Chemical Behavior

• Electrons occupy orbitals, which are grouped into electron shells (energy levels).

• Each orbital holds up to two electrons; shells are filled from the inside out.

• The valence shell is the outermost shell; electrons here are called valence electrons.

• The number of unpaired valence electrons determines an atom's valence (bonding capacity).

Chemical Bonds and Molecules

Covalent Bonds

• Covalent bonds form when two atoms share pairs of valence electrons, creating a molecule.

• Atoms are most stable when their valence shells are full.

• Single, double, and triple covalent bonds are possible, depending on the number of shared electron pairs.

Electronegativity and Bond Polarity

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

• Electronegativity increases up and to the right on the periodic table (O > N > S ≈ C ≈ H ≈ P).

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

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

Ionic Bonds

• Formed when electrons are transferred from one atom to another, creating ions.

• Cation: Positively charged ion (loses electron).

• Anion: Negatively charged ion (gains electron).

The Electron-Sharing Continuum

• Bonds range from equal sharing (nonpolar covalent) to unequal sharing (polar covalent) to complete transfer (ionic).

Molecular Geometry and Representation

• The shape of a molecule is determined by the arrangement of its bonds and unshared electron pairs.

• Common representations: Molecular formulas, structural formulas, ball-and-stick models, and space-filling models.

• Examples: Methane (CH4) is tetrahedral; water (H2O) is bent.

The Unique Properties of Water

Water as a Solvent

• Water makes up about 75% of a cell's mass and is an excellent solvent due to its polarity.

• Solvent: The substance in which solutes dissolve to form a solution.

• Polar and charged (hydrophilic) substances dissolve easily in water; nonpolar (hydrophobic) substances do not.

Hydrogen Bonding

• Water molecules form hydrogen bonds due to partial charges on oxygen (δ-) and hydrogen (δ+).

• Hydrogen bonds are weak individually but strong collectively, giving water its unique properties.

Cohesion, Adhesion, and Surface Tension

• Cohesion: Attraction between like molecules (water to water).

• Adhesion: Attraction between unlike molecules (water to other substances).

• Surface tension: The cohesive force at the surface of water, making it act like an elastic membrane.

Density and States of Water

• Water is denser as a liquid than as a solid (ice), so ice floats.

• Ice forms an open crystal structure stabilized by hydrogen bonds.

High Specific Heat and Heat of Vaporization

• Specific heat: The energy required to raise the temperature of 1 g of a substance by 1°C.

• Water has a high specific heat and high heat of vaporization due to extensive hydrogen bonding.

• This property helps regulate temperature in organisms and environments.

Acids, Bases, and pH

Acid-Base Chemistry in Water

• Water can dissociate into hydrogen ions (H+) and hydroxide ions (OH-):

• Acids: Substances that donate protons (H+), increasing hydronium ion concentration.

• Bases: Substances that accept protons, decreasing hydronium ion concentration.

pH Scale

• pH measures the concentration of hydrogen ions in solution:

• Acidic solutions: pH < 7; Basic solutions: pH > 7; Neutral: pH = 7.

• Biological systems maintain pH near neutrality (pH ≈ 7).

• Buffers help maintain stable pH by absorbing or releasing H+ as needed.

Buffer Systems Example

• Carbonic acid-bicarbonate buffer system in blood:

Chemical Evolution and the Origin of Life

Chemical Evolution

• Life likely began with the formation of increasingly complex carbon-containing molecules from simple substances.

• Key environments: Early atmosphere (volcanic gases), deep-sea hydrothermal vents.

• Stanley Miller's experiment demonstrated that organic molecules (e.g., amino acids) can form under prebiotic conditions.

Energy in Chemical Reactions

• Potential energy: Stored energy due to position or structure (e.g., chemical bonds).

• Kinetic energy: Energy of motion (e.g., heat, molecular movement).

• First law of thermodynamics: Energy cannot be created or destroyed, only transformed.

• Second law of thermodynamics: Entropy (disorder) always increases in a closed system.

Spontaneity of Chemical Reactions

• Reactions are spontaneous if they proceed without continuous external input.

• Spontaneity is favored when products have lower potential energy and/or higher entropy than reactants.

Carbon: The Backbone of Life

Properties of Carbon

• Carbon forms four covalent bonds, allowing for a variety of stable, complex molecules.

• Organic compounds contain carbon bonded to other elements (H, N, O, P, S).

• Carbon skeletons can form chains, rings, and branched structures.

Functional Groups

• Common functional groups attached to carbon atoms determine the chemical behavior of organic molecules:

• Amino group (-NH2): Acts as a base.

• Carboxyl group (-COOH): Acts as an acid.

• Carbonyl group (C=O): Involved in linking molecules.

• Hydroxyl group (-OH): Acts as a weak acid.

• Phosphate group (-PO42-): Contributes negative charge.

• Sulfhydryl group (-SH): Can form disulfide bonds.

Macromolecules and Polymerization

• Monomers: Small molecular subunits (e.g., amino acids, nucleotides, monosaccharides).

• Polymers: Large molecules made by linking monomers via polymerization.

• Condensation (dehydration) reactions: Join monomers by removing water.

• Hydrolysis reactions: Break polymers into monomers by adding water.

• Polymerization is energetically unfavorable unless monomer concentration is high.

Summary Table: Key Concepts

Example: The Miller-Urey experiment demonstrated that amino acids could form from simple molecules under conditions simulating early Earth, supporting the theory of chemical evolution.