Chapter 2: Water and Carbon – The Chemical Basis of Life

Introduction to Chemical Evolution

Chemical evolution is the leading explanation for the origin of life on Earth. It describes the process by which simple molecules formed increasingly complex carbon-containing substances, eventually leading to molecules capable of self-replication. This transition from chemical to biological evolution marked the beginning of life as we know it.

• Chemical evolution: Formation of complex molecules from simpler ones, leading to the first self-replicating molecules.

• Biological evolution: Took over once molecules could replicate, with natural selection favoring those that were metabolically active and membrane-bound.

• Five characteristics of life: Metabolism, growth, reproduction, response to stimuli, and adaptation through evolution.

Atoms, Ions, and Molecules: The Building Blocks of Chemical Evolution

Elements Essential for Life

Four types of atoms make up 96% of matter in living organisms: hydrogen, carbon, nitrogen, and oxygen. Understanding their structure is fundamental to biology.

• Hydrogen (H), Carbon (C), Nitrogen (N), Oxygen (O): Most abundant elements in biological molecules.

• These atoms form the basis for water, carbon dioxide, and other simple molecules crucial for chemical evolution.

Basic Atomic Structure

Atoms consist of a nucleus (protons and neutrons) surrounded by electrons.

• Protons: Positive charge (+1)

• Neutrons: No charge (neutral)

• Electrons: Negative charge (-1), orbit the nucleus

• Atoms are electrically neutral when protons = electrons.

Example: A hydrogen atom has 1 proton and 1 electron; a carbon atom has 6 protons, 6 neutrons, and 6 electrons.

Atomic Number and Mass Number

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

• Mass number: Sum of protons and neutrons.

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

• Atomic weight: Average mass of all isotopes, weighted by abundance.

• Radioactive isotopes: Unstable, decay over time, used in dating and tracing biological processes.

Electron Arrangement and Chemical Behavior

• Electrons occupy orbitals, grouped into electron shells.

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

• Valence shell: Outermost shell; electrons here are called valence electrons.

• Valence: Number of unpaired electrons in the valence shell; determines bonding capacity.

Chemical Bonds and Molecule Formation

Covalent Bonds

Covalent bonds form when atoms share unpaired valence electrons, resulting in stable molecules.

• Single, double, and triple bonds: Atoms can share one, two, or three pairs of electrons.

• Molecule: Group of atoms held together by covalent bonds.

Example: Two hydrogen atoms share electrons to form H2.

Polar and Nonpolar Covalent Bonds

• Electronegativity: Atom's ability to attract electrons in a bond.

• Nonpolar covalent bond: Electrons shared equally (e.g., C–H bond).

• Polar covalent bond: Electrons shared unequally, creating partial charges (e.g., O–H bond in water).

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

Ionic Bonds and the Electron-Sharing Continuum

• Ionic bond: Electron is transferred from one atom to another, creating ions.

• Cation: Positively charged ion (lost electron).

• Anion: Negatively charged ion (gained electron).

• Bonds exist on a continuum from equal sharing (nonpolar covalent) to complete transfer (ionic).

Geometry and Representation of Molecules

• Molecular shape is determined by the arrangement of bonds and unshared electron pairs.

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

• Molecules can be represented by molecular formulas, structural formulas, ball-and-stick models, and space-filling models.

Properties of Water

Water as a Solvent

Water is essential for life due to its unique properties as a solvent.

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

• Hydrogen bonds: Weak attractions between the partial charges of water molecules.

• Hydrophilic substances: Ions and polar molecules dissolve easily in water.

• Hydrophobic substances: Nonpolar molecules do not dissolve; they cluster together via hydrophobic interactions and van der Waals forces.

Cohesion, Adhesion, and Surface Tension

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

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

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

• These properties enable water to move against gravity in plants and support small organisms on its surface.

Density and Thermal Properties of Water

• Water is denser as a liquid than as a solid; ice floats due to its open crystal structure.

• Specific heat: Water has a high specific heat (4.18 J/g°C), meaning it resists temperature changes.

• Heat of vaporization: Water requires a lot of energy to evaporate, aiding in temperature regulation (e.g., sweating).

Acids, Bases, and pH

Acid-Base Chemistry in Water

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

• In solution, protons associate with water to form hydronium ions ():

• Acids: Donate protons, increasing .

• Bases: Accept protons, decreasing .

Measuring Concentration and pH

• Mole: particles; mass of one mole equals atomic/molecular weight in grams.

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

• pH: Measures hydrogen ion concentration; logarithmic scale:

• Acidic: pH < 7; Basic: pH > 7; Neutral: pH = 7 (typical of living cells).

• Buffers: Substances that minimize changes in pH, maintaining homeostasis.

Chemical Reactions, Energy, and Chemical Evolution

Types of Chemical Reactions

• Chemical reactions involve breaking and forming bonds, converting reactants to products.

• Example: Formation of carbonic acid from CO2 and H2O:

• Reactions can be endothermic (absorb energy) or exothermic (release energy).

Energy in Biological Systems

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

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

• First law of thermodynamics: Energy is conserved; it can be transferred or transformed, but not created or destroyed.

• Second law of thermodynamics: Entropy (disorder) always increases in spontaneous processes.

Spontaneity of Chemical Reactions

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

• Spontaneity 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 diverse molecular structures.

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

• Carbon skeletons can be chains, rings, or branched structures (e.g., octane, glucose).

Functional Groups in Organic Molecules

• Amino group (–NH2): Acts as a base, attracts protons.

• Carboxyl group (–COOH): Acts as an acid, donates protons.

• Carbonyl group (–CO–): Site for linking molecules in larger compounds.

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

• Phosphate group (–PO42−): Carries two negative charges.

• Sulfhydryl group (–SH): Can form disulfide bonds, stabilizing protein structure.

Macromolecules and Polymerization

• Macromolecules: Large molecules made of repeating subunits (monomers).

• Polymerization: Monomers join via condensation (dehydration) reactions, releasing water.

• Hydrolysis: Polymers are broken down by adding water, increasing entropy and favoring monomer formation.

• Examples of biological macromolecules: Proteins, nucleic acids, carbohydrates.

Additional info: These foundational chemical principles are essential for understanding the molecular basis of life and the processes that led to the origin of living systems on Earth.