The Chemistry of Life: Part 1 - Interactions of Atoms

Overview of Chemistry and Fundamental Particles

• Definition of Chemistry: Chemistry is the scientific study of the structure, composition, and interactions of elemental matter.

• Matter:

  • Defined as anything that occupies space and has mass.

  • Comprises all living and nonliving structures in the universe.

  • Exists in three primary physical states: solid, liquid, and gas.

• Mass versus Weight:

  • Mass: Represents the inherent energy contained in a physical object at rest, described mathematically by Einstein's mass-energy equivalence equation:         $E = m c^2$

  • Weight: The variable force exerted on matter by gravity. Unlike mass, weight changes depending on gravitational strength.

• Fundamental Subatomic Particles:

  • All physical matter in the universe is constructed from three fundamental particles: electrons, up quarks, and down quarks.

  • Up quarks and down quarks combine in specific arrangements to construct subatomic nucleons (protons and neutrons).

  • Protons, neutrons, and electrons assemble to form fundamental chemical units known as atoms.

Chemical Elements and Cosmic Origin

• Chemical Elements:

  • Matter is composed of unique fundamental chemical substances called elements.

  • There are 118 known chemical elements.

  • Each element is designated by a standard chemical symbol consisting of one or two letters derived from its name (e.g., $\text{H}$ for Hydrogen, $\text{Ca}$ for Calcium, $\text{Mg}$ for Magnesium).

  • Elements are systematically organized in the Periodic Table according to their atomic number.

• Natural versus Synthetic Occurrence:

  • Elements with atomic numbers 1 through 98 occur naturally within the universe.

  • Elements heavier than atomic number 98 are synthetic and have been isolated exclusively through particle accelerator bombardment.

• Radioactivity and Nuclear Decay:

  • All chemical elements heavier than Bismuth (atomic number 83) possess unstable nuclei and are naturally radioactive.

  • Radioactive isotopes decay by emitting protons, electrons, and nuclear energy to convert into lighter, more stable elements.

  • Example: Uranium-235 ($^{235}\text{U}$) undergoes radioactive decay into Lead ($\text{Pb}$) with a physical half-life of $700 \times 10^6\,\text{years}$. Measuring predictable decay rates allows scientists to date ancient geological and biological specimens.

• Cosmic Origin and Abundance:

  • Hydrogen ($\text{H}$, approximately $74\%$) and Helium ($\text{He}$, approximately $24\%$) together account for roughly $99\%$ of all elemental matter in the universe.

  • Bare Hydrogen and Helium atomic nuclei formed within seconds to minutes following the Big Bang; neutral atoms assembled approximately $380,000\,\text{years}$ later.

  • All elements heavier than Helium are synthesized inside stars through nuclear fusion.

  • Because subatomic particles and atomic nuclei are fundamentally conserved, elements are virtually indestructible ("immortal"), making all living organisms composed of ancient cosmic matter.

Composition of the Human Body and Biological Minerals

• Elemental Composition of the Human Body:

  • Out of 118 total elements, exactly 24 possess established biological roles in human anatomy and physiology.

  • Major Elements (Total $98.5\%$ of Body Weight):

    • Oxygen ($\text{O}$): $65.0\%$ of body weight

    • Carbon ($\text{C}$): $18.0\%$ of body weight

    • Hydrogen ($\text{H}$): $10.0\%$ of body weight

    • Nitrogen ($\text{N}$): $3.0\%$ of body weight

    • Calcium ($\text{Ca}$): $1.5\%$ of body weight

    • Phosphorus ($\text{P}$): $1.0\%$ of body weight

  • Lesser Elements (Total $0.8\%$ of Body Weight):

    • Sulfur ($\text{S}$): $0.25\%$ of body weight

    • Potassium ($\text{K}$): $0.20\%$ of body weight

    • Sodium ($\text{Na}$): $0.15\%$ of body weight

    • Chlorine ($\text{Cl}$): $0.15\%$ of body weight

    • Magnesium ($\text{Mg}$): $0.05\%$ of body weight

    • Iron ($\text{Fe}$): $0.006\%$ of body weight

  • Trace Elements (Total $0.7\%$ of Body Weight): Present in minute concentrations, including Chromium ($\text{Cr}$), Cobalt ($\text{Co}$), Copper ($\text{Cu}$), Fluorine ($\text{F}$), Iodine ($\text{I}$), Manganese ($\text{Mn}$), Molybdenum ($\text{Mo}$), Selenium ($\text{Se}$), Silicon ($\text{Si}$), Tin ($\text{Sn}$), Vanadium ($\text{V}$), and Zinc ($\text{Zn}$).

• Minerals:

  • Inorganic chemical elements extracted directly from soil by plants and subsequently passed up the food chain to human consumers.

  • Include essential mineral elements such as $\text{Ca}$, $\text{P}$, $\text{Cl}$, $\text{Mg}$, $\text{K}$, $\text{Na}$, $\text{I}$, $\text{Fe}$, $\text{Zn}$, $\text{Cu}$, and $\text{S}$.

  • Account for approximately $4\%$ of total human body weight.

  • Provide essential structural integrity (e.g., calcium and phosphorus in teeth and bones) and act as vital enzyme cofactors.

Atomic Structure and Subatomic Properties

• The Atom:

  • The smallest fundamental unit of an element that retains all chemical properties and characteristics of that specific element.

  • Subatomic building blocks (protons, neutrons, electrons) are non-unique; all electrons display identical behavior regardless of the element they belong to.

  • Atoms are extremely small: approximately $200,000$ individual atoms placed side-by-side fit across the period printed at the end of a sentence.

• Central Nucleus:

  • Occupies the geometric center of the atom and contains nucleons:

    • Protons ($p^+$): Possess a positive electric charge ($+1$) and a physical mass of $1\,\text{amu}$ (atomic mass unit).

    • Neutrons ($n^0$): Electrically neutral ($0$ charge) with a physical mass of $1\,\text{amu}$.

• Electron Cloud:

  • Electrons orbit in space surrounding the dense central nucleus.

  • Electrons ($e^-$): Possess a negative electric charge ($-1$) and an negligible mass ($\approx 0\,\text{amu}$). An electron is approximately $2000$ times lighter than a single proton.

• Atomic Number and Mass Number:

  • Atomic Number: Equal to the exact number of protons in an atom's nucleus. Defines elemental identity.

  • Atomic Mass (Mass Number): Equal to the total sum of protons plus neutrons in the nucleus.

• Electrical Neutrality:

  • In an uncharged atom, the total number of negative electrons equals the total number of positive protons, yielding a net electrical charge of zero.

Electron Shells, Sublevels, and Valence Configurations

• Electron Shells:

  • Defined concentric volumes of space surrounding the atomic nucleus where electrons reside.

  • Shells fill progressively starting from the innermost shell closest to the nucleus outward.

  • Maximum Shell Capacities:

    • 1st shell (innermost): maximum 2\text{e}^-$\n * 2nd shell: maximum 8\text{e}^-$

    • 3rd shell: maximum 18\text{e}^-$\n * 4th shell: maximum 32\text{e}^-$

    • 5th shell: maximum 32\text{e}^-$\n * 6th shell: maximum 18\text{e}^-$

    • 7th shell: maximum 8\text{e}^-$\n\n\n\n* **Sublevels and Orbitals:**\n * Electron shells contain sublevels (s, p, d, f), which are subdivided into atomic orbitals capable of holding a maximum of 2 paired electrons each:\n * **s$sublevel:** 1 orbital$\rightarrow$maximum$2\text{e}^-$

    • $p$ sublevel: 3 orbitals $\rightarrow$ maximum 6\text{e}^-$\n * **d$sublevel:** 5 orbitals$\rightarrow$maximum$10\text{e}^-$

    • $f$ sublevel: 7 orbitals $\rightarrow$ maximum 14\text{e}^-$\n\n\n\n* **Aufbau Filling Principle:**\n * Electrons fill available atomic sublevels strictly in order of increasing energy:\n        1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p \rightarrow 5s \rightarrow 4d \rightarrow 5p \rightarrow 6s \rightarrow 4f \rightarrow 5d \rightarrow 6p \rightarrow 7s \rightarrow 5f \rightarrow 6d \rightarrow 7p \rightarrow 8s\n * *Example (Lead, \text{Pb}, Atomic Number 82):*\n * Full electron configuration:\n            1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6 6s^2 4f^{14} 5d^{10} 6p^2\n * The outermost energy shell (shell 6) contains 6s^2$and$6p^2, resulting in a total of 4 valence electrons.\n* **Valence Electrons and Octet Rule:**\n * **Valence Shell:** The outermost occupied electron shell of an atom.\n * **Valence Electrons:** Electrons inhabiting the valence shell. These electrons participate in chemical reactions.\n * **The Octet Rule:** Atoms tend to react, gain, lose, or share electrons until their valence shell reaches a stable state of 8 electrons (or 2 electrons for elements possessing only a single electron shell, such as Hydrogen and Helium).\n\n\n\n* **Chemical Reactivity Classification:**\n * **Chemically Inert Elements:** Possess a completely filled valence shell. They do not readily gain, lose, or share electrons, making them unreactive noble gases (e.g., Helium, Neon).\n * **Chemically Reactive Elements:** Possess incomplete valence shells. They undergo chemical bonding to achieve electronic stability (e.g., Hydrogen, Carbon, Oxygen, Sodium).\n\n# Isotopes and Atomic Mass Variations\n\n* **Isotopes:**\n * Varieties of a single element that differ exclusively in their number of nuclear neutrons, resulting in different overall atomic mass numbers.\n * Because isotopes retain identical proton and electron counts, they possess identical valence configurations and behave identically in chemical reactions.\n\n\n\n* **Isotopes of Hydrogen:**\n * **Hydrogen-1 (Protium, ^1\text{H}$):**$1p^+$,$0n^0$,$1e^-\n * **Hydrogen-2 (Deuterium, ^2\text{H}$):**$1p^+$,$1n^0$,$1e^-\n * **Hydrogen-3 (Tritium, ^3\text{H}$):**$1p^+$,$2n^0$,$1e^-\n* **Relative Atomic Mass:**\n * Reflects the weighted average atomic mass of all naturally occurring isotopes of an element.\n * Accounted for by the non-integer values on the periodic table; for example, Hydrogen has a relative atomic mass of 1.008\,\text{amu}$rather than exactly$1\,\text{amu}.\n\n# Ions, Ionization, and Physiological Electrolytes\n\n* **Ionization:** The chemical process in which an atom transfers one or more electrons to another atom, breaking electrical neutrality and producing charged particles called ions.\n\n\n\n* **Types of Ions:**\n * **Anion:** A negatively charged ion formed when an atom gains electrons (\text{electrons} > \text{protons}$). *Example:* Chloride ion ($\text{Cl}^-).\n * **Cation:** A positively charged ion formed when an atom loses or donates electrons (\text{protons} > \text{electrons}$). *Example:* Sodium ion ($\text{Na}^+).\n* **Electrolytes:**\n * Substances that ionize in water and conduct electrical currents.\n * **Electrolyte Balance:** The tight homeostatic regulation of specific ion concentrations in blood and bodily fluids required for optimal physiological function.\n * **Electrolyte Imbalance:** Severe elevation or depletion of blood ion levels that disrupts cellular function and can prove fatal.\n * *Practical Application:* Sports beverages like Gatorade contain small quantities of K^+ to replace potassium ions lost through perspiration during exercise.\n * *Extreme Application:* High-dose concentrations of potassium chloride (\text{KCl}) are administered in lethal injections to halt cardiac muscle contractions.\n* **Physiological Roles of Key Electrolytes:**\n * **Sodium Ion (\text{Na}^+):** The primary extracellular cation; essential for nerve impulse conduction, muscle contraction, and controlling the osmotic movement of water.\n * **Potassium Ion (\text{K}^+):** The primary intracellular cation; vital for membrane potential regulation, nerve signaling, and muscle contraction.\n * **Calcium Ion (\text{Ca}^{2+}):** Provides structural hardness to bone matrix and teeth; triggers muscle contraction mechanisms and stimulates neurotransmitter exocytosis.\n * **Hydrogen Ion (\text{H}^+$):** Free concentration directly dictates the$\text{pH} of blood, extracellular fluids, and intracellular compartments.\n * **Chloride Ion (\text{Cl}^-$):** Most abundant extracellular anion; regulates neuronal excitability and serves as a major constituent of stomach gastric acid ($\text{HCl}).\n * **Bicarbonate Ion (\text{HCO}3^-$):** Acts as the primary transport vehicle for carbon dioxide ($\text{CO}_2) in blood plasma and functions as a key physiological buffer.\n * **Phosphate Ion (\text{PO}_4^{3-}$):** The most abundant intracellular anion; hardens bone mineral structure and serves as a core structural component of cell membrane phospholipids, nucleic acids ($\text{DNA/RNA}), and ATP.\n\n# Molecules, Compounds, and Molecular Weight\n\n* **Molecules and Compounds:**\n * **Molecule:** A stable chemical particle consisting of two or more atoms held together by chemical bonds.\n * **Compound:** A molecule composed specifically of two or more *different* chemical elements (e.g., Water \text{H}_2\text{O}$, Methane$\text{CH}_4$). All compounds are molecules, but not all molecules are compounds (e.g.,$\text{O}_2 is a molecule, not a compound).\n* **Representing Chemical Structures:**\n * **Molecular Formula:** Identifies the exact elemental types and quantitative ratios of atoms present in a molecule (e.g., \text{H}_2\text{O}$,$\text{C}_6\text{H}{12}\text{O}6).\n * **Structural Formula:** Illustrates the precise spatial arrangement and relative positions of chemical bonds joining atoms together.\n\n\n\n* **Isomers:** Chemical compounds that share identical molecular formulas but exhibit different structural arrangements and distinct physical/chemical properties (e.g., Ethanol \text{CH}_3\text{CH}_2\text{OH}$vs. Ethyl ether$\text{CH}_3\text{OCH}_3$, both sharing the molecular formula$\text{C}_2\text{H}_6\text{O}).\n* **Molecular Weight Calculation:**\n * The molecular weight (MW) of a molecule is calculated by summing the atomic weights of all constituent atoms.\n * *Calculation Example (Glucose, \text{C}_6\text{H}{12}\text{O}_6):*\n * 6 Carbon atoms \times 12\,\text{amu} = 72\,\text{amu}\n * 12 Hydrogen atoms \times 1\,\text{amu} = 12\,\text{amu}\n * 6 Oxygen atoms \times 16\,\text{amu} = 96\,\text{amu}\n * \text{Total Molecular Weight} = 72 + 12 + 96 = 180\,\text{amu}\n\n# Chemical Bonds\n\n* **Chemical Bonds:** Forces of attraction that link atoms together to form molecules or attract adjacent molecules to one another.\n\n| Bond Type | Description | Strength |\n| :--- | :--- | :--- |\n| **Ionic Bond** | Electrostatic attraction between a cation and an anion resulting from complete electron transfer | Moderately strong attraction |\n| **Nonpolar Covalent Bond** | Equal sharing of valence electron pairs between atomic nuclei | Strongest of all chemical bonds |\n| **Polar Covalent Bond** | Unequal sharing of valence electron pairs; electrons spend more time around the more electronegative nucleus | Strong attraction |\n| **Hydrogen Bond** | Weak electrostatic attraction between a slightly positive hydrogen atom (\delta+$) and a slightly negative atom ($\delta-) | Relatively weak attraction |\n\n* **Covalent Bonds:**\n * Formed when two adjacent atoms share pairs of valence electrons to fulfill the octet rule.\n * **Single Covalent Bond:** The sharing of 1 pair of electrons (represented as \text{H}-\text{H}).\n * **Double Covalent Bond:** The sharing of 2 pairs of electrons (represented as \text{O}=\text{O}).\n * **Triple Covalent Bond:** The sharing of 3 pairs of electrons (represented as \text{N}\equiv\text{N}).\n\n\n\n* **Polarity in Covalent Bonds:**\n * In a **nonpolar covalent bond**, shared electrons distribute symmetrically and spend equal time around both nuclei.\n * In a **polar covalent bond**, shared electrons orbit more frequently around one electronegative nucleus, bestowing a partial negative charge (\delta-$) on that pole and a partial positive charge ($\delta+) on the opposite pole.\n* **Hydrogen Bonds:**\n * Formed between a partially positive hydrogen atom (\delta+$) of one polar molecule and a partially negative atom ($\delta-), usually Oxygen or Nitrogen, on another molecule.\n * Indicated structurally by a dotted or dashed line.\n * Although individually weak, collective hydrogen bonding is vital in establishing three-dimensional protein folding and stabilizing the double-helical structure of \text{DNA}.\n\n# Physical and Chemical Properties of Water\n\n* **Biological Importance:** Water constitutes 50\%$to$75\% of total human body weight. Most body fluids consist of chemicals dissolved or suspended in water.\n* **Solvency:**\n * Water is known as the **Universal Solvent** due to its extreme polarity.\n * **Hydrophilic Substances:** Dissolve readily in water; molecules must be charged or polarized.\n * **Hydrophobic Substances:** Do not dissolve in water; molecules are non-polar or electrically neutral (e.g., lipids, fats).\n\n\n\n* **Hydration Spheres:**\n * When an ionic compound dissolves in water, water molecules orient their partial charges around individual ions (\delta-$Oxygen towards cations like$\text{Na}^+$,$\delta+$Hydrogen towards anions like$\text{Cl}^-).\n * These spherical clusters of water molecules, called **hydration spheres**, isolate the ions and prevent them from re-associating into ionic crystals.\n* **Adhesion:**\n * The tendency of water molecules to cling to surfaces of other distinct substances.\n * Water adheres to living tissue membranes, forming a lubricating fluid layer within serous cavities (e.g., pleura surrounding lungs, pericardium surrounding the heart).\n * Reduces mechanical friction as the heart and lungs expand and contract against thoracic walls.\n\n\n\n* **Cohesion and Surface Tension:**\n * **Cohesion:** The molecular tendency of like water molecules to cling tightly to one another via extensive hydrogen bonding.\n * Cohesive hydrogen bonding at liquid-gas interfaces generates high **surface tension**, allowing light insects (such as water striders) to walk across open water.\n\n\n\n* **Chemical Reactivity:**\n * Water ionizes directly into active hydrogen ions (\text{H}^+$) and hydroxide ions ($\text{OH}^-).\n * Facilitates ionization of dissolved salts (e.g., \text{NaCl} \rightarrow \text{Na}^+ + \text{Cl}^-).\n * Participates directly in metabolic **hydrolysis** (cleaving bonds via water addition) and **dehydration synthesis** (forming bonds via water removal) reactions.\n* **Thermal Stability:**\n * Water possesses a high specific heat capacity, meaning hydrogen bonds restrict molecular movement and prevent rapid temperature swings.\n * Allows the human body to absorb substantial internal metabolic heat without suffering large thermal variations.\n * Functions as an effective evaporative coolant: as liquid water transitions into vapor on skin surfaces, significant thermal energy is dissipated.\n\n# Solutions, Concentration, and Molarity\n\n* **Concentration Principles:**\n * Solution concentration measures the quantity of solute dissolved in a defined volume of solvent.\n * Physiological biological responses depend directly on the total *number* of dissolved solute particles, rather than the absolute physical *weight* of the solute.\n* **Mole and Avogadro's Number:**\n * **Mole:** An amount of any chemical substance whose mass in grams is numerically equal to its molecular weight in atomic mass units (\text{amu}).\n * *Example:* Glucose has a molecular weight of 180\,\text{amu}$; therefore,$180\,\text{g}$of glucose equals exactly$1\,\text{mole} of glucose.\n * **Avogadro's Number:** One mole of any substance contains exactly 6.023 \times 10^{23} individual molecules or formula units.\n* **Molarity (\text{M}):**\n * Defined as the number of moles of solute contained per 1 liter of total solution:\n        \text{Molarity (M)} = \frac{\text{Moles of Solute}}{\text{Liters of Solution}}\n * *Example:* A 1.0\,\text{M}$glucose solution contains$1\,\text{mole}$($180\,\text{g}$) dissolved in$1\,\text{L}$total solution, containing$6.023 \times 10^{23} glucose molecules.\n\n\n\n* **Equimolar Particle Equality:**\n * Solutions sharing identical molarities contain identical numbers of solute particles per unit volume regardless of solute size or molecular weight.\n * A 0.1\,\text{M}$glucose solution ($18\,\text{g/L}$) contains the exact same number of dissolved particles as a$0.1\,\text{M}$sucrose solution ($34\,\text{g/L}).\n\n# Acids, Bases, the pH Scale, and Buffer Systems\n\n* **Acids and Bases:**\n * **Acid:** A proton donor; a chemical molecule that releases hydrogen ions (\text{H}^+) in aqueous solution.\n * **Base:** A proton acceptor; a chemical molecule that binds free hydrogen ions (\text{H}^+) in aqueous solution.\n* **The \text{pH} Scale:**\n * Quantifies solution acidity based on the molar concentration of free hydrogen ions ([\text{H}^+]):\n        \text{pH} = -\log([\text{H}^+])\n * *Mathematical Example:* If a solution has a hydrogen ion concentration of [\text{H}^+] = 0.0000001\,\text{M} = 10^{-7}\,\text{M}:\n        \log([\text{H}^+]) = -7\n        \text{pH} = -(-\log(10^{-7})) = 7\n * **Scale Range:** Spans from 0 to 14.\n * \text{pH} = 7$: Electrically neutral ($[\text{H}^+] = [\text{OH}^-])\n * \text{pH} < 7$: Acidic solution ($[\text{H}^+] > [\text{OH}^-])\n * \text{pH} > 7$: Basic/Alkaline solution ($[\text{H}^+] < [\text{OH}^-])\n * **Logarithmic Nature:** The \text{pH}$scale is logarithmic; a single unit change in$\text{pH}$represents a 10-fold increase or decrease in free$\text{H}^+$concentration (e.g., dropping from$\text{pH } 5$to$\text{pH } 4 increases acidity by 10 times).\n\n\n\n* **Buffers and Homeostasis:**\n * **Buffer:** A chemical system that resists dramatic changes in \text{pH}$by accepting or donating$\text{H}^+ ions as concentrations fluctuate.\n * **Physiological Importance:** Normal human arterial blood \text{pH} must be maintained within a tight range of 7.35 to 7.45. Significant deviations disrupt cellular protein structures, impair drug action, and induce muscle tremors, paralysis, or death.\n * **Carbonic Acid-Bicarbonate Buffer System:** A dynamic, reversible equilibrium system that regulates human blood \text{pH}:\n        \text{H}_2\text{CO}_3 \rightleftharpoons \text{HCO}_3^- + \text{H}^+\n * *Response to a Rise in \text{pH}$(Excessively Basic):* Carbonic acid ($\text{H}_2\text{CO}_3$, weak acid) acts as a proton donor, dissociating into bicarbonate ($\text{HCO}_3^-$) and releasing free$\text{H}^+$ions into solution to lower the$\text{pH}.\n * *Response to a Drop in \text{pH}$(Excessively Acidic):* Bicarbonate ($\text{HCO}_3^-$), acting as a weak base/proton acceptor, binds excess free$\text{H}^+$ions to form carbonic acid ($\text{H}_2\text{CO}_3$), restoring optimal$\text{pH}$$.