Chapter 2: The Chemical Basis of Life - Water & Carbon

Elemental Composition and Basic Atomic Structure

Living organisms are composed primarily of four main chemical elements: Hydrogen, Oxygen, Carbon, and Nitrogen. Together, these four elements make up approximately $96\,\%$ of all matter in living organisms. A fundamental guiding principle in biological chemistry is that form follows function, meaning that the specific structural arrangement of atoms and molecules directly dictates their physical and chemical behaviors.

Atoms represent the fundamental structural units of chemical elements. Every atom consists of three primary subatomic particles: protons, neutrons, and electrons. Protons are positively charged particles ($+$), neutrons are electrically neutral particles carrying no charge, and electrons are negatively charged particles ($-$). Protons and neutrons are packed densely together within the central core of the atom, known as the nucleus. Electrons inhabit three-dimensional regions of space surrounding the nucleus called orbitals.

Every chemical element is uniquely defined by its atomic number, which corresponds directly to the characteristic number of protons present in the nucleus of its atoms. In a neutral atom, the number of protons is equal to the number of electrons. Atoms that possess the exact same atomic number share identical chemical properties and belong to the same chemical element. For example, a helium atom carries an atomic number of $2$, indicating the presence of $2$ protons in its nucleus.

Isotopes are distinct forms of a single element that possess the exact same number of protons but differ in their total number of neutrons. The mass number of an element corresponds to the combined total count of protons and neutrons present in the nucleus of its most common isotope. For example, in a carbon-12 atom (represented symbolically as $^{12}_{6}C$), the mass number is $12$ and the atomic number is $6$. Calculating the neutron count involves subtracting the atomic number from the mass number: $\text{Number of neutrons} = 12 - 6 = 6$. Thus, carbon-12 contains $6$ protons, $6$ electrons, and $6$ neutrons.

Valence, Chemical Bonding, and Molecular Structure

An atom's outer electron shell contains orbitals that may be filled or unfilled. The number of unpaired electrons located in an atom's outermost shell is defined as its valence. Valence governs an atom's reactivity and bonding capacity. Unfilled electron orbitals allow for the formation of chemical bonds. Atoms achieve maximum structural stability when all of their electron orbitals are completely filled. Molecules are defined as distinct chemical substances held together by chemical bonds.

Elements exhibit specific valences based on their electron configurations:

Valence $0$ elements have completely filled outer shells and include Helium ($He$), Neon ($Ne$), and Argon ($Ar$).

Valence $1$ elements have one unpaired outer electron and include Hydrogen ($H$), Lithium ($Li$), Sodium ($Na$), Fluorine ($F$), and Chlorine ($Cl$).

Valence $2$ elements have two unpaired outer electrons and include Beryllium ($Be$), Magnesium ($Mg$), Oxygen ($O$), and Sulfur ($S$).

Valence $3$ elements have three unpaired outer electrons and include Boron ($B$), Aluminum ($Al$), Nitrogen ($N$), and Phosphorus ($P$).

Valence $4$ elements have four unpaired outer electrons and include Carbon ($C$) and Silicon ($Si$).

Chemical bonds are categorized based on how electrons are shared or transferred between atoms. A covalent bond occurs when unpaired valence electrons are shared between two atomic nuclei, effectively filling the outer orbitals of both participating atoms. For example, when two individual hydrogen atoms ($1$ proton each) bond, they share their single valence electrons to produce a stable molecular hydrogen gas ($H_2$) molecule. When one oxygen atom ($8$ protons, $8$ neutrons) covalently bonds with two hydrogen atoms ($1$ proton each), a water ($H_2O$) molecule is formed. When two oxygen atoms covalently bond with each other, molecular oxygen gas ($O_2$) is produced.

Covalent bonds vary based on the relative electronegativity of the participating atoms. Electronegativity refers to the strength with which an atomic nucleus pulls shared electrons toward itself. Electronegativity increases as the number of protons in the nucleus increases, and decreases as the number of shielding electron shells increases. Among common biological elements, the relative electronegativity scale follows the hierarchy: Oxygen > Nitrogen > Carbon = Hydrogen ($O > N > C = H$).

In a nonpolar covalent bond, electrons are shared equally between two atoms because their electronegativities are equal or nearly equal. Consequently, the atoms carry no partial electrical charges. Examples include molecular hydrogen ($H_2$) and methane ($CH_4$). In a polar covalent bond, electrons are shared unequally because one atom exerts a stronger electronegative pull than the other. This unequal distribution causes partial negative charges (denoted as $\delta-$) to gather around the more electronegative atom, while partial positive charges (denoted as $\delta+$) reside near the less electronegative atom. Examples include water ($H_2O$) and ammonia ($NH_3$).

An ionic bond forms when electrons are completely transferred from one atom to another rather than shared, resulting in fully charged atoms known as ions. A cation is a positively charged ion formed when an atom loses one or more electrons. An anion is a negatively charged ion formed when an atom gains one or more electrons. The electrostatic attraction between oppositely charged cations and anions creates an ionic bond. For example, sodium ($Na$) loses an electron to become a sodium cation ($Na^+$), while chlorine ($Cl$) gains that electron to become a chloride anion ($Cl^-$). In the absence of water, these ions arrange into a crystalline solid table salt ($NaCl$) lattice.

Molecules can contain single, double, or triple covalent bonds depending on the number of shared electron pairs. Single covalent bonds involve one shared pair of electrons, as seen in water ($H_2O$), ammonia ($NH_3$), and methane ($CH_4$). Double covalent bonds involve two shared pairs of electrons, as seen in carbon dioxide ($CO_2$). Triple covalent bonds involve three shared pairs of electrons, as seen in molecular nitrogen ($N_2$).

Scientists represent molecular structures using four standardized models: molecular formulas (e.g., $CH_4$, $NH_3$, $H_2O$, $O_2$), structural formulas (showing standard bond lines), ball-and-stick models (showing three-dimensional spatial orientation and bond angles), and space-filling models (showing actual van der Waals atomic radii and molecular shape).

Unique Properties of Water and Aqueous Solutions

Water serves as the biological solvent for all known life forms. The molecular architecture of water features several key elements: a small overall size, a distinct bent molecular geometry, highly polar covalent bonds, and an overall dipole moment. Because oxygen is significantly more electronegative than hydrogen ($O > H$), the electrons within water's polar covalent bonds are pulled strongly toward the oxygen nucleus. This produces a partial negative charge ($\delta-$) on the oxygen atom and a partial positive charge ($\delta+$) on each of the two hydrogen atoms.

The polar nature of water allows adjacent water molecules to form hydrogen bonds. A hydrogen bond is an electrostatic attraction between the partial positive charge ($\delta+$) of a hydrogen atom on one molecule and the partial negative charge ($\delta-$) of an electronegative atom (such as oxygen) on an adjacent molecule.

Hydrophilic ("water-loving") substances consist of ions and polar molecules that interact favorably with water's partial charges, allowing them to remain suspended or dissolved in solution. Hydrogen bonding makes it possible for virtually any charged or polar molecule to dissolve readily in water. When solid table salt ($NaCl$) dissolves in water, water molecules surround the individual ions to form spheres of hydration: partial positive hydrogen atoms orient toward chloride anions ($Cl^-$), while partial negative oxygen atoms orient toward sodium cations ($Na^+$). Conversely, hydrophobic ("water-fearing") substances are uncharged, nonpolar compounds. Because hydrophobic molecules cannot form favorable hydrogen bonds with water, they do not dissolve in aqueous solutions.

Water exhibits extraordinary physical and chemical properties driven largely by its capacity for hydrogen bonding:

Cohesion is the attraction and binding between like molecules. Water molecules hydrogen-bond extensively with one another, resulting in high surface tension. Surface tension causes the surface layer of liquid water to resist elastic deformation or penetration, preventing light objects from sinking.

Adhesion is the attraction and binding between unlike molecules, such as water molecules binding to glass surfaces, plastic containers, or biological epithelial layers.

Capillary action is the movement of water through narrow spaces resulting from the combination of adhesion and cohesion. When water is placed in a glass cylinder, adhesion causes water molecules to adhere to the glass wall and pull upward at the perimeter, while cohesion causes the internal water molecules to resist this upward pull, generating a curved surface called a meniscus.

Density anomaly as a solid: Water is less dense as a solid than as a liquid. As liquid water freezes into ice, its hydrogen bonds stabilize into a rigid, open crystal lattice structure that expands the volume. Consequently, ice floats on liquid water.

Thermal absorption capacity: Water possesses an extraordinarily high specific heat and high heat of vaporization, enabling it to absorb large amounts of heat energy without undergoing rapid shifts in temperature.

Solutions are categorized on the pH spectrum based on their relative hydrogen ion concentration relative to pure water:

Acidic solutions (pH < 7) include stomach acid, lemon juice, vinegar, soft drinks, beer, wine, tomatoes, black coffee, urine, and milk.

Neutral solutions (pH = 7) include pure water and human blood.

Basic solutions (pH > 7) include seawater, baking soda, milk of magnesia, household ammonia, household bleach, and oven cleaner.

Thermodynamics, Chemical Energy, and Reaction Spontaneity

Chemical reactions involve the breaking and forming of chemical bonds, transforming starting substances (reactants) into final substances (products). A representative aqueous reaction is the combination of carbon dioxide gas and liquid water to yield carbonic acid:

$CO_2(g) + H_2O(l) \rightleftharpoons H_2CO_3(aq)$

Chemical equilibrium is established when the forward reaction rate equals the reverse reaction rate. At equilibrium, the relative quantities of reactants and products remain constant over time. Reactions that absorb heat from their surroundings to proceed are defined as endothermic. Reactions that release heat into their surroundings are defined as exothermic.

Energy is defined as the capacity to do work or supply heat. Energy exists in two main forms: potential energy (stored energy) and kinetic energy (energy of active spatial motion or thermal energy). Chemical energy is a specific form of potential energy stored within the chemical bonds of molecules.

The amount of potential energy held within a covalent bond depends on electron position. Electrons held loosely in nonpolar covalent bonds between atoms with equal electronegativities possess high potential energy. Electrons held tightly by highly electronegative atoms in polar covalent bonds possess low potential energy.

The spontaneity of a chemical reaction is determined by two thermodynamic factors:

1. The change in potential energy: Reactions tend to be spontaneous when the products have lower potential energy than the reactants.

2. The change in molecular order (entropy): Reactions tend to be spontaneous when the products are less ordered (possess higher entropy) than the reactants.

For example, when molecular hydrogen gas and oxygen gas react, high-potential-energy reactants convert into lower-potential-energy water products, releasing energy as heat and light that vaporizes the resulting water:

$2H_2(g) + O_2(g) \rightarrow 2H_2O(g) + \text{released heat and light}$

Similarly, cellular respiration and wood combustion involve the reaction of glucose and oxygen to form carbon dioxide, water, and heat:

$C_6H_{12}O_6(s) + 6O_2(g) \rightarrow 6CO_2(g) + 6H_2O(g) + \text{released heat}$

In this system, glucose ($C_6H_{12}O_6$) and oxygen ($O_2$) represent reactants with high potential energy and a high degree of structural order (low entropy). The resulting products—carbon dioxide ($CO_2$) and water ($H_2O$)—possess low potential energy and a low degree of structural order (high entropy).

Chemical Evolution Theory and Prebiotic Synthesis Experiments

The Chemical Evolution Theory is the leading scientific explanation for the origin of life on Earth. The theory consists of a pattern and a process:

The pattern dictates that, in addition to simple molecules, complex carbon-containing molecules exist and are required for life.

The process dictates that early in Earth's history, simple chemical compounds combined via spontaneous reactions to form increasingly complex carbon-containing substances prior to the evolution of biological life.

Chemical evolution is hypothesized to have initiated in one of two ancient environments:

1. The primordial atmosphere: Volcanic gases ejected into the atmosphere consisted of water vapor ($H_2O$), carbon dioxide ($CO_2$), molecular nitrogen ($N_2$), and potentially molecular hydrogen ($H_2$) and carbon monoxide ($CO$).

2. Deep-sea hydrothermal vents: High-temperature oceanic environments supplied with gases such as $CO_2$ and $H_2$, alongside reactive metal catalysts such as nickel ($Ni$) and iron ($Fe$).

Two models describe prebiotic synthesis pathways:

In the atmospheric model, simple molecules present in ancient Earth's atmosphere were exposed to sunlight energy. Solar kinetic energy drove reactions among atmospheric gases to synthesize precursor molecules (such as formaldehyde, hydrogen cyanide, ribose, glycine, and acetaldehyde), which then condensed into early oceans.

In the hydrothermal vent model, simple molecules dissolved in early oceans and hydrothermal vents underwent spontaneous reactions catalyzed by minerals on ocean crust walls. Stimulated by heat and concentration, these precursors formed complex molecules such as acetic acid, methane, ribose, and organic polymers.

The Stanley Miller spark-discharge experiment tested whether kinetic energy added to simple atmospheric gases could drive chemical evolution. The experimental apparatus consisted of a small glass flask containing boiling water (providing heat energy), connected via glass tubing to a large glass flask containing primitive gases ($CH_4$, $NH_3$, and $H_2$). Electrodes introduced spark discharges (simulating lightning) into the gas mixture. A condenser cooled the vapor back into water droplets, which collected in a lower liquid trap.

The experiment tested two opposing hypotheses:

Null hypothesis: Chemical evolution will not occur, even with an input of energy, leaving only starting molecules in the liquid trap.

Alternative hypothesis: Complex organic compounds will form in the liquid water if kinetic energy is supplied.

Analysis of the trap liquid revealed the synthesis of formaldehyde ($H_2CO$), hydrogen cyanide ($HCN$), and complex organic compounds containing carbon-carbon ($C-C$) bonds, including amino acids. This confirmed that chemical evolution occurs readily when simple high-free-energy molecules are exposed to a kinetic energy source.

Subsequent origin-of-life studies adjusted starting compositions to reflect primitive volcanic output ($CO$, $CO_2$, and $H_2$) rather than the highly reduced gases ($CH_4$, $NH_3$) used in Miller's initial apparatus. These studies demonstrated photochemical synthesis driven by solar radiation:

High-energy solar photons bombard molecular hydrogen ($H_2$), breaking bonds to generate highly reactive hydrogen radicals ($H^\bullet$) containing unpaired electrons. High-energy photons also strike carbon dioxide ($CO_2$), producing carbon monoxide radicals ($CO^\bullet$) and oxygen radicals ($O^\bullet$). These reactive radicals combine to form formaldehyde ($CH_2O$) and water:

$CO_2(g) + 2H_2(g) \rightarrow CH_2O(g) + H_2O(g)$

Through these photochemical steps, solar radiant energy was converted and stored as chemical potential energy within the chemical bonds of formaldehyde ($H_2CO$) and hydrogen cyanide ($HCN$).

Carbon Versatility and Organic Functional Groups

Carbon is the most versatile building block on Earth due to its four valence electrons, which allow it to form up to four covalent bonds simultaneously. Carbon-containing molecules can form an almost limitless array of molecular shapes, incorporating single and double bonds. Carbons can link together in linear chains, as seen in octane ($C_8H_{18}$), or in ring structures, as seen in glucose ($C_6H_{12}O_6$). The formation of carbon-carbon ($C-C$) bonds represented a critical milestone in chemical evolution.

The carbon atoms in an organic molecule furnish the skeleton that gives the molecule its overall structural shape. Specific combinations of atoms, known as functional groups, attach to carbon skeletons and determine the chemical behavior and reactivity of organic molecules.

Six functional groups commonly attached to organic carbon skeletons (where $R$ represents the remaining attached portion of the organic molecule) include:

1. Amino Group ($-NH_2$): Family: Amines. Formula: $R-NH_2$ Properties: Acts as a base in aqueous solution, attracting a proton to form $R-NH_3^+$.. Example: Glycine (an amino acid).

2. Carboxyl Group ($-COOH$): Family: Carboxylic acids. Formula: $R-COOH$ Properties: Acts as an acid in aqueous solution, dropping a proton to form $R-COO^-$.. Example: Acetic acid.

3. Carbonyl Group ($-CHO$ or $R_1-CO-R_2$): Family: Aldehydes (when bound to at least one hydrogen, $R-CHO$) and Ketones (when bound to two carbon chains, $R_1-CO-R_2$). Properties: Aldehydes readily react with specific compounds to produce larger organic molecules. Examples: Acetaldehyde (aldehyde) and Acetone (ketone).

4. Hydroxyl Group ($-OH$): Family: Alcohols. Formula: $R-OH$ Properties: Highly polar structure makes compounds significantly more soluble in water via hydrogen bonding; may also act as a weak acid and drop a proton. Example: Ethanol.

5. Phosphate Group ($-PO_4^{2-}$): Family: Organic phosphates. Formula: $R-PO_4^{2-}$ Properties: Molecules containing multiple phosphate groups linked directly together store large amounts of chemical potential energy. Example: $3$-Phosphoglyceric acid.

6. Sulfhydryl Group ($-SH$): Family: Thiols. Formula: $R-SH$ Properties: When present in proteins, sulfhydryl groups can react to form disulfide bonds ($S-S$) that covalently cross-link and stabilize protein tertiary and quaternary structures. Example: Cysteine (an amino acid).