The Chemical Basis of Life: Elements, Bonds, and the Properties of Water

The Rationale for Chemistry in Biology

• Fundamental Connection: The basis of all life is rooted in chemistry. This is illustrated by three primary observations:   - Chemical substances constitute the physical bodies of all living organisms as well as their surrounding physical environments.   - The internal functions and biological processes of cells are governed by underlying chemical reactions.   - Life and chemistry are inextricably tied to water ($H_2O$); life originated in water, and every living organism remains dependent upon it.

• Elemental Composition:   - Living organisms are composed of approximately $25$ distinct chemical elements.   - The Bulk Elements: Carbon ($C$), Hydrogen ($H$), Oxygen ($O$), and Nitrogen ($N$) make up the vast majority of all living matter.   - Trace Elements: These are chemical elements essential to life but required only in minute (trace) amounts.

Trace Element Deficiencies and Physiological Effects

• Impact of Deficiency: Dietary deficiencies in essential trace elements can lead to significant physiological conditions.

• Case Study: Goiter:   - A goiter is characterized as an enlarged thyroid gland.   - Cause: This condition is typically caused by a deficiency of the trace element Iodine ($I$) in the diet.

• Prevention: To combat such deficiencies, trace elements are frequently added to essential resources such as food supplies and water (e.g., iodized salt).

Chemical Elements and the Formation of Compounds

• Nature of Compounds: Chemical elements combine in specific, fixed ratios to form compounds.

• Emergent Properties: Compounds often possess characteristics entirely different from the elements that compose them. An example involves the synthesis of Sodium Chloride ($NaCl$):   - Sodium ($Na^+$): Separately, this is a chalky, soft solid.   - Chlorine ($Cl^-$): Separately, this is a toxic gas.   - Sodium Chloride ($NaCl$): When combined in a fixed ratio, they form white, crystalline table salt.

Atomic Structure and Isotopes

• The Atom: Defined as the smallest particle of matter that still retains the specific properties of its chemical element.

• Subatomic Particles:   - Protons: Carry a positive ($+$) charge and are located within the nucleus.   - Neutrons: Carry a neutral ($0$) charge and are located within the nucleus.   - Electrons: Carry a negative ($-$) charge and orbit the nucleus in specific shells.

• Atomic Measurements:   - Atomic Number: Determined by the number of protons in an atom; this number defines the specific element.   - Mass Number: Calculated as the sum of the protons and neutrons within the nucleus ($ext{Protons} + ext{Neutrons}$).

• Isotopes:   - Isotopes are variants of an element that have the same number of protons but a different number of neutrons.   - Some isotopes are radioactive, meaning they decompose spontaneously, giving off particles and energy.   - Applications: Radioactive isotopes are useful as biological tracers to follow molecules through metabolic pathways.   - Risks: While beneficial for medical use, uncontrolled exposure to radioactive isotopes can cause harm to living organisms.   - Measurement: The decay of radioactive isotopes is measured by their half-life.

Applications in Medical Diagnostics

• Radioactive Tracers: These substances are frequently used in medical diagnosis.

• Imaging: Tracers are utilized in conjunction with sophisticated imaging instruments (e.g., PET scans) to visualize internal biological structures and processes.

Principles of Chemical Bonding

• Ionic Bonds: These are attractions formed between ions of opposite charges.   - Ions are created when atoms either gain or lose electrons.   - Example: The attraction between $Na^+$ and $Cl^-$ creates an ionic bond.

• Covalent Bonds: These bonds join atoms together into molecules through the process of electron sharing.   - Atoms may share one or more pairs of their outer-shell (valence) electrons.   - Methane ($CH_4$): A molecule formed by covalent bonds where one Carbon atom shares electrons with four Hydrogen atoms.

• Covalent Bond Types:   - Non-polar Covalent Bonds: Occur when covalently bonded atoms share their electrons equally. An example is Carbon Dioxide ($CO_2$).   - Polar Covalent Bonds: Occur when electrons are shared unequally between atoms, leading to a polar molecule with partial charges. An example is Water ($H_2O$).

• Hydrogen Bonds: These are weak but vital bonds in the chemistry of life.   - They occur when the charged regions of water molecules are attracted to the oppositely charged regions on nearby water molecules.

Life-Supporting Properties of Water

1. Cohesion and Adhesion:     - Cohesion: Hydrogen bonds make liquid water cohesive, meaning water molecules stick to one another.     - Surface Tension: Created by cohesion, this allows certain insects to walk on the surface of water.     - Biological Transport: The combination of adhesion and cohesion allows water to travel upwards against gravity from the roots of plants to their leaves.

2. Temperature Moderation:     - Water can absorb significant amounts of heat without a large rise in temperature because much of the energy is used to disrupt hydrogen bonds rather than increasing molecular motion.     - Conversely, as water cools, a slight drop in temperature releases a large amount of heat as hydrogen bonds form.     - Evaporative Cooling: As water molecules evaporate, they carry energy away with them, cooling the surface left behind.

3. Density and Phase Change:     - Ice (solid water) is less dense than liquid water because hydrogen bonds hold the molecules further apart in a stable crystal lattice.     - In liquid water, hydrogen bonds constantly break and re-form.     - Environmental Impact: Because ice floats, it protects lakes and oceans from freezing solid, insulating the liquid water and life beneath it.

4. The Solvent of Life:     - Water dissolves more solutes than any other known solvent.     - Polar or charged solutes dissolve when water molecules surround them, forming aqueous solutions.     - This follows the principle of "like dissolves like."

Acid-Base Chemistry and the pH Scale

• Acids: Compounds that release Hydrogen ions ($H^+$) into a solution.

• Bases: Compounds that accept $H^+$ ions or release Hydroxide ions ($OH^-$) in a solution.

• The pH Scale:   - This is a factor-of-10 (logarithmic) scale used to measure acidity and alkalinity.   - The scale ranges from $0$ (most acidic) to $14$ (most basic or alkaline).   - A pH of $7$ is considered neutral.

• Biological Sensitivity: The chemistry of life is highly sensitive to acidic and basic conditions. Most cells maintain an internal pH close to $7$.

Buffers and Physiological Homeostasis

• Buffers: Substances that resist changes in pH by accepting or donating $H^+$ ions as needed.

• Bicarbonate System: This is the body’s natural buffer system used to regulate blood pH.   - Optimal Blood pH: Must stay in a very narrow range, with $7.4$ being optimal.   - Survival Range: Death may occur if the blood pH falls below $7.0$ or rises above $7.8$.

• Mechanism of the Bicarbonate/Carbonic Acid Buffer:   - The reaction is governed by the following equilibrium: $H_2CO_3 \rightleftharpoons HCO_3^- + H^+$.   - If pH is too high (too basic): Carbonic acid ($H_2CO_3$) dissociates to release $H^+$ ions.   - If pH is too low (too acidic): Bicarbonate ions ($HCO_3^-$) pick up stray $H^+$ ions to form carbonic acid.

Environmental Threats of Acidic Conditions

• Impact on Marine Environments:   - Approximately $25\%$ of atmospheric $CO_2$ is absorbed by the oceans.   - High $CO_2$ levels produce carbonic acid in the water, which releases excess $H^+$ ions.   - These $H^+$ ions combine with carbonate ($CO_3^{2-}$) to form bicarbonate ($HCO_3^-$).   - This process depletes the carbonate supply required by corals and shell-building animals to produce Calcium Carbonate ($CaCO_3$) shells, weakening reef structures.

• Impact on Terrestrial Environments:   - Acid Precipitation: Formed when air pollutants from fossil fuel combustion combine with atmospheric water vapor, creating sulfuric and nitric acids.   - These acids damage buildings, kill trees, and alter the pH of water supplies for land animals.

Chemical Reactions and Biochemistry

• Definition: Chemical reactions change the composition of matter.

• Process: Reactants interact, atoms are rearranged, and products result from the transformation.

• Biochemistry: This field study involves the thousands of chemical reactions carried out by living cells to rearrange matter in significant ways.

Questions & Discussion

• Q: Why start biology with chemistry?

• A: Because the basis of all life is chemistry:   - Chemicals compose the bodies of living organisms.   - Chemical reactions drive cell functions.   - Water is the medium where life began and continues to depend upon.   - Living organisms are made of about $25$ chemical elements, primarily $C$, $H$, $O$, and $N$.   - Trace elements are required in minute amounts for survival.