Human Anatomy and Physiology: Chemistry, pH, and Homeostasis Review Flashcards

Fundamentals of Chemistry and Matter

• Chemistry Definition: The scientific discipline that deals with the structure, properties, and interactions of matter.
• Matter Definition: Anything that occupies space and has mass.
• States of Matter: Matter exists in three physical states:
  • Solid
  • Liquid
  • Gas
• Chemical Level of Organization: Represents the foundation of biological hierarchy:
  • Subatomic particles $\rightarrow$ Atoms $\rightarrow$ Molecules $\rightarrow$ Organelles $\rightarrow$ Cells $\rightarrow$ Tissues $\rightarrow$ Organs $\rightarrow$ Organ Systems $\rightarrow$ Organisms.
• Biological Importance: Understanding chemistry explains the molecular behaviors that govern cellular processes and clarifies structural-functional relationships at higher levels of physiological organization.

Atomic Structure and Subatomic Particles

• Atom: The smallest stable unit of matter.
• Subatomic Particles: Atoms are composed of three primary subatomic particles:
  • Protons ($\mathbf{p^+}$):
    • Charge: Positive ($+1$
    • Mass: Approximately $1\,\text{atomic mass unit (amu)}$
    • Location: Situated within the central atomic nucleus.
  • Neutrons ($\mathbf{n^0}$):
    • Charge: Neutral / uncharged ($0$
    • Mass: Approximately $1\,\text{atomic mass unit (amu)}$ (similar to protons)
    • Location: Situated within the central atomic nucleus.
  • Electrons ($\mathbf{e^-}$):
    • Charge: Negative ($-1$
    • Mass: Extremely light (negligible mass compared to protons/neutrons)
    • Location: Orbit around the nucleus within designated electron shells or electron clouds.

• Electron Shells and Energy Levels:
  • Regions surrounding the atomic nucleus where electrons exist.
  • Capacity Rules:
    • $1\text{st shell}$ (closest to nucleus): Holds up to $2\text{ electrons}$.
    • $2\text{nd shell}$: Holds up to $8\text{ electrons}$.
    • $3\text{rd shell}$: Holds up to $8\text{ electrons}$.
  • Valence Shell: The outermost electron shell of an atom.
  • Chemical Stability Rule: The number of electrons in the valence shell dictates chemical reactivity. If an atom's valence shell is not full, the atom is unstable and will react with other atoms to gain, lose, or share electrons.

• Quantitative Atomic Terms:
  • Atomic Number: The total number of protons located within the nucleus of an atom. Unique to each specific element.
  • Mass Number (Atomic Mass): The combined total number of protons plus neutrons in the nucleus of an atom. $1\text{ proton}$ or $1\text{ neutron}$ weighs $1\,\text{amu}$.
    • Equation: $\text{Mass Number} = \text{Protons} + \text{Neutrons}$
  • Atomic Weight: The weighted average mass of an element's naturally occurring atoms, accounting for the relative abundances of all its isotopes.

Classification of Atoms, Elements, and Isotopes

• Element: A pure substance composed entirely of atoms of one kind, which cannot be broken down into simpler substances by ordinary physical or chemical means.
  • Total Naturally Occurring Elements: $92\text{ elements}$.
  • Elements in Human Body: $27\text{ total elements}$ ($13\text{ main elements}$ and $14\text{ trace elements}$ present in microscopic quantities).
  • Primary Four Human Body Elements: Oxygen ($\text{O}$), Carbon ($\text{C}$), Hydrogen ($\text{H}$), and Nitrogen ($\text{N}$).

• Isotopes: Alternate structural versions of the exact same element that contain identical numbers of protons (same atomic number) but different numbers of neutrons, resulting in distinct mass numbers.
  • Carbon Isotopes Example:
    • Carbon-12 ($\mathbf{^{12}C}$): Contains $6\text{ protons}$ and $6\text{ neutrons}$. Mass number = $12\,\text{amu}$. Represents $98.9\%$ of natural carbon.
    • Carbon-13 ($\mathbf{^{13}C}$): Contains $6\text{ protons}$ and $7\text{ neutrons}$. Mass number = $13\,\text{amu}$. Represents $1.1\%$ of natural carbon.
    • Carbon-14 ($\mathbf{^{14}C}$): Contains $6\text{ protons}$ and $8\text{ neutrons}$. Mass number = $14\,\text{amu}$. Represents $<0.1\%$ of natural carbon.

• Radioisotopes and Radioactive Decay:
  • Radioisotopes: Isotopes possessing unstable, radioactive nuclei.
  • Radioactive Decay: The continuous spontaneous breakdown of an unstable nucleus into a more stable configuration.
  • Medical Applications:
    • Diagnostic procedures: Diagnostic imaging and molecular tracing.
    • Therapeutic procedures: High-potency radiation emissions intentionally applied to destroy targeted malignant cells (e.g., cancer treatment).

• Ions and Free Radicals:
  • Ion: An atom or group of atoms carrying an electrical charge formed by gaining or losing valence electrons.
    • Cation: A positively charged ion formed when an atom loses one or more electrons.
    • Anion: A negatively charged ion formed when an atom gains one or more electrons.
  • Free Radical: An extremely reactive atom, ion, or molecule containing an unpaired electron in its outermost valence shell.

Chemical Bonds, Compounds, and Electrolytes

• Bonding Concepts:

  • Chemical Reactions: Processes that allow reactive atoms to achieve electron stability by gaining, losing, or sharing valence electrons.
  • Chemical Bond: Attractive forces that hold participating atoms together following chemical reactions.
  • Molecule: Two or more atoms bonded together via shared electrons (covalent) or ionic interactions. Can consist of identical elements ($\text{H}_2$) or different elements ($\text{H}_2\text{O}$).
  • Compound: A substance composed of two or more atoms of different elements bound together in fixed, definite mass proportions.

• Ionic Bonds:

  • Bonds formed when one or more valence electrons are completely transferred from one atom to another, producing distinct cations and anions that attract each other electrostatically.
  • Example: Sodium ($\text{Na}$) transfers an electron to Chlorine ($\text{Cl}$), yielding a sodium cation ($\text{Na}^+$) and a chloride anion ($\text{Cl}^-$) to form Sodium Chloride ($\text{NaCl}$).

• Salts and Electrolytes:
  • Compounds composed of cations and anions bound by ionic interactions form crystalline structures known as salts.
  • When placed in aqueous solutions, salts dissociate (separate) into free ions.
  • Electrolytes: Soluble inorganic compounds whose ions conduct electrical currents in solution.
  • Physiological Importance: Essential for regulating body fluid distribution, facilitating muscle tissue contraction, and propagating nerve impulses.

• Key Biological Electrolyte Dissociations:

  • $\text{NaCl} \rightarrow \text{Na}^+ + \text{Cl}^-$
  • $\text{KCl} \rightarrow \text{K}^+ + \text{Cl}^-$
  • $\text{CaPO}_4 \rightarrow \text{Ca}^{2+} + \text{PO}_4^{2-}$
  • $\text{NaHCO}_3 \rightarrow \text{Na}^+ + \text{HCO}_3^-$
  • $\text{MgCl}_2 \rightarrow \text{Mg}^{2+} + 2\text{Cl}^-$
  • $\text{Na}_2\text{HPO}_4 \rightarrow 2\text{Na}^+ + \text{HPO}_4^{2-}$
  • $\text{Na}_2\text{SO}_4 \rightarrow 2\text{Na}^+ + \text{SO}_4^{2-}$

• Covalent Bonds:

  • The strongest category of chemical bonds, formed when two or more atoms share valence electrons.
  • Bond Orders:
    • Single Covalent Bond: Sharing of $1\text{ pair}$ of electrons ($2\text{ electrons total}$).
    • Double Covalent Bond: Sharing of $2\text{ pairs}$ of electrons ($4\text{ electrons total}$).
    • Triple Covalent Bond: Sharing of $3\text{ pairs}$ of electrons ($6\text{ electrons total}$).
  • Polarity Classifications:
    • Nonpolar Covalent Bonds: Occur between atoms with similar electronegativities; electrons are shared equally. Outer surrounding atoms are typically identical. Examples: Hydrogen gas ($\text{H}_2$), Oxygen gas ($\text{O}_2$), Carbon dioxide ($\text{CO}_2$), Nitrogen gas ($\text{N}_2$), Methane ($\text{CH}_4$).
    • Polar Covalent Bonds: Occur between atoms with significantly different electronegativities; electrons are shared unequally. The strongly electronegative atom attracts electrons closer, acquiring a partial negative charge ($\delta^-$ or $2\delta^-$), while the other atom acquires a partial positive charge ($\delta^+$). Examples: Nitric oxide ($\text{NO}$), Water ($\text{H}_2\text{O}$).

• Hydrogen Bonds:
  • Weak electrostatic attractions occurring between a partial positive hydrogen atom ($\delta^+$) in a polar covalent bond and a partial negative atom ($\delta^-$) of another polar covalent bond.
  • Functional Significance:
    • Generates cohesion between neighboring water molecules, producing high surface tension.
    • Empowers water to absorb and distribute thermal energy without rapid temperature swings.
    • Stabilizes three-dimensional folded conformations of giant bio-macromolecules such as $\text{DNA}$ and proteins.

Chemical and Physical Properties of Water

• Abundance and Fundamental Importance:

  • Water ($\text{H}_2\text{O}$) is the single most vital inorganic chemical compound in living organisms.
  • Accounts for up to $\approx 60\%$ of total adult human body mass.

• Four Major Chemical Properties of Water:

  1. Universal Solvent:
     • Dissolves a broad range of biological molecules and compounds.
     • Essential for nutrient absorption/transport, metabolic chemical reactions, and cellular waste removal.
     • Solution Vocabulary:
       • Solution: A homogenous fluid mixture of two or more distinct substances.
       • Solute: The dissolved material; present in smaller quantitative amounts.
       • Solvent: The liquid medium in which solutes are dissolved; present in the greatest quantitative amount.
       • Aqueous Solution: Any fluid solution where water functions as the primary solvent.
     • Aqueous Heterogeneous Mixtures:
       • Colloid: A fluid solution containing large molecular solutes (such as proteins) that remain permanently distributed without settling out over time. Examples: Blood plasma (water + solutes + proteins), Jello.
       • Suspension: A mixture containing large particulate components that will physically settle out of solution if left undisturbed. Examples: Whole blood (plasma + red/white blood cells + platelets), sand mixed into water.
     • Molecular Water Interactions:
       • Hydrophilic ("water-loving"): Molecules that interact readily with water molecules. Includes charged ions and polar compounds.
       • Hydrophobic ("water-fearing"): Molecules that do not interact readily with water. Includes nonpolar molecules, fats, lipids, and oils.
  2. Reactivity:
     • Serves as the medium where cellular chemical reactions take place.
     • Actively participates directly as a reactant or product in fundamental metabolic reactions, including dehydration synthesis (building complex molecules by removing water) and hydrolysis (cleaving chemical bonds by adding water).
  3. High Heat Capacity:
     • Heat Capacity Definition: The quantity of thermal energy required to elevate the temperature of a unit mass of a substance by exactly $1^\circ\text{C}$.
     • Water possesses an unusually high heat capacity due to its extensive intermolecular hydrogen bonding network.
     • Physiological Benefits:
       • Maintains liquid state over a wide temperature range (freezing and boiling points are far apart).
       • Absorbs and dissipates substantial thermal energy during evaporation (sweating cools the body).
       • Resists rapid temperature shifts (a property called thermal inertia), helping stabilize body core temperature.
  4. Lubrication:
     • Dramatically reduces mechanical friction between opposing tissue surfaces.
     • Protects organs and facilitates smooth biological movements.
     • Example: Serous aqueous fluid within closed ventral body cavities coats internal organs, preventing abrasive friction against cavity walls.

Solution Chemistry, pH, Acids, Bases, and Buffers

• The Molar Concept:

  • Mole: A standardized quantitative unit possessing a weight in grams equal to that element's atomic weight.
  • Avogadro's Principle: One mole of any element contains the exact same number of individual atoms as one mole of any other element.
  • Comparative Molar Examples:
    • $1\text{ mole}$ of oxygen atoms weigh significantly more than $1\text{ mole}$ of hydrogen atoms ($\approx 16\,\text{g}$ vs $\approx 1\,\text{g}$).
    • $1\text{ mole}$ of oxygen atoms contains the exact same total count of individual atoms as $1\text{ mole}$ of hydrogen atoms ($6.022 \times 10^{23}\text{ atoms}$).

• pH Fundamentals:

  • Definition: Measures the absolute concentration of hydrogen ions ($\text{H}^+$) in a solution.
  • Mathematical Formula: Defined as the negative logarithm of the hydrogen ion concentration measured in moles per liter:         $\text{pH} = -\log[\text{H}^+]$
  • Scale Limits: Ranges from $0.00$ to $14.00$.
  • Inverse Logarithmic Behavior: Lower pH values indicate higher concentrations of $\text{H}^+$ (acidic); higher pH values indicate lower concentrations of $\text{H}^+$ (basic/alkaline). Each single unit change on the pH scale represents a $10\text{-fold}$ ($10\times$) shift in $\text{H}^+$ concentration.
  • pH Ranges:
    • Neutral pH: Equal concentrations of $\text{H}^+$ and $\text{OH}^-$ ions ($[\text{H}^+] = [\text{OH}^-]$). Pure water at equilibrium has a neutral pH of $7.00$
    • Acidic pH: pH values between $0$ and $7$. High $\text{H}^+$ concentration, low $\text{OH}^-$ concentration.
    • Basic (Alkaline) pH: pH values between $7$ and $14$. Low $\text{H}^+$ concentration, high $\text{OH}^-$ concentration.
  • Human Blood Homeostatic Range: Arterial blood plasma pH is strictly maintained between $\text{pH } 7.35$ and $7.45$.

• pH Spectrum Values of Reference Substances:

  • $1\,\text{mol/L}\text{ Hydrochloric acid (HCl)}$: $\text{pH } 0.00$ ($[\text{H}^+] = 1 \times 10^0\,\text{M}$
  • Stomach acid: $\text{pH } 1.00$ ($[\text{H}^+] = 1 \times 10^{-1}\,\text{M}$
  • Beer, vinegar, wine, pickles: $\text{pH } 3.00$ ($[\text{H}^+] = 1 \times 10^{-3}\,\text{M}$
  • Tomatoes, grapes: $\text{pH } 4.00$ ($[\text{H}^+] = 1 \times 10^{-4}\,\text{M}$
  • Saliva, milk: $\text{pH } 6.00$ ($[\text{H}^+] = 1 \times 10^{-6}\,\text{M}$
  • Pure water: $\text{pH } 7.00$ ($[\text{H}^+] = 1 \times 10^{-7}\,\text{M}$
  • Urine: Fluctuates across a physiological range of $\text{pH } 4.50$ to $8.00$
  • Human blood plasma: $\text{pH } 7.35 - 7.45$
  • Eggs: $\text{pH } 8.00$ ($[\text{H}^+] = 1 \times 10^{-8}\,\text{M}$
  • Ocean water: $\text{pH } 8.00 - 9.00$
  • Household bleach: $\text{pH } 9.50 - 10.00$
  • Household ammonia: $\text{pH } 11.00 - 12.00$ ($[\text{H}^+] = 1 \times 10^{-11} - 1 \times 10^{-12}\,\text{M}$
  • Oven cleaner: $\text{pH } 13.50$
  • $1\,\text{mol/L}\text{ Sodium hydroxide (NaOH)}$: $\text{pH } 14.00$ ($[\text{H}^+] = 1 \times 10^{-14}\,\text{M}$

• Acids, Bases, and Salts:

  • Acid (Proton Donor): A solute that releases hydrogen ions ($\text{H}^+$) into solution.
    • Strong Acids: Dissociate completely in solution (e.g., $\text{HCl} \rightarrow \text{H}^+ + \text{Cl}^-$).
    • Weak Acids: Dissociate incompletely in solution; achieve equilibrium without full dissociation.
  • Base (Proton Acceptor): A solute that removes hydrogen ions ($\text{H}^+$) from solution (frequently releasing hydroxide ions $\text{OH}^-$).
    • Strong Bases: Dissociate completely in solution (e.g., $\text{NaOH} \rightarrow \text{Na}^+ + \text{OH}^-$).
    • Weak Bases: Dissociate incompletely in solution.
  • Salt: An ionic compound consisting of cations other than $\text{H}^+$ and anions other than $\text{OH}^-$. Produced when acids react with bases; does not directly alter solution $\text{H}^+$ or $\text{OH}^-$ balance.

• Buffers and Buffer Systems:

  • Buffer Definition: Chemical compounds that stabilize solution pH by neutralizing added acids or bases, thereby resisting pH shifts.
  • Composition: A functional buffer system requires two components: a weak acid and a weak base.
  • Experimental Demonstration of Buffer Action:
    • Adding $1\,\text{mL}$ of $0.1\,\text{M}\text{ HCl}$ to unbuffered water drops pH sharply from $7.0$ to $3.0$. Adding $1\,\text{mL}$ of $0.1\,\text{M}\text{ HCl}$ to a buffered solution preserves a stable pH of $6.9$
    • Adding $1\,\text{mL}$ of $0.1\,\text{M}\text{ NaOH}$ to unbuffered water spikes pH sharply from $7.0$ to $11.0$. Adding $1\,\text{mL}$ of $0.1\,\text{M}\text{ NaOH}$ to a buffered solution preserves a stable pH of $7.1$

• Three Major Human Body Buffers:
  1. Carbonic Acid–Bicarbonate Buffer System: Primary buffer system stabilizing extracellular fluid (ECF) pH; regulated via renal and respiratory activity.
  2. Phosphate Buffer System: Critical intracellular fluid (ICF) buffer system; also stabilizes urine pH.
  3. Protein Buffer Systems: Plasma and intracellular proteins that accept or release $\text{H}^+$ ions; e.g., Hemoglobin proteins inside red blood cells.

Homeostasis and the Internal Environment

• Homeostasis Definition: The continuous physiological process of establishing and maintaining a relatively stable internal environment within living cells, tissues, organs, and organ systems.
• Dynamic Equilibrium: Physiological systems continuously adapt and adjust to changing environmental conditions. Regulated variables fluctuate within an acceptable normal range around a set point rather than remaining completely static.
• Homeostatic Imbalance: Severe disruption of normal physiological regulation resulting from physical trauma, clinical illness, genetic mutations, or aging processes. Causes systemic pathology and potentially death.
• The Internal Environment (Extracellular Fluid - ECF):
  • The extracellular fluid that directly bathes and surrounds tissue cells.
  • Three Primary ECF Compartments:
    1. Interstitial Fluid: Fluid filling the intercellular spaces surrounding cells.
    2. Plasma: Liquid portion of circulating blood carrying nutrients and blood cells.
    3. Cerebrospinal Fluid (CSF): Specialized fluid secreted within brain ventricles that circulates across the brain surface and spinal cord central canal.

• Key Physiological Variables Regulated Homeostatically:
  • Core body temperature
  • Fluid concentrations (dissolved nutrients, oxygen, dissolved gases, metabolic toxins)
  • Fluid volume (balance between water absorption and excretory loss)
  • pH (concentration of free hydrogen ions $\text{H}^+$)
  • Blood pressure
  • Ion and waste concentrations (electrolytes such as $\text{Na}^+$, $\text{Cl}^-$, $\text{K}^+$, $\text{Ca}^{2+}$, and metabolic waste products)

Homeostatic Regulatory Mechanisms and Feedback Systems

• Three Major Components of Homeostatic Regulation:
  1. Receptor (Sensor):
     • Specialized sensory cells sensitive to a specific environmental stimulus or change.
     • Action: Detects deviations and sends input signals to control centers.
  2. Control Center (Integration Center):
     • Cells that receive, process, and analyze sensory input from receptors.
     • Action: Determines if deviation exceeds normal limits and sends instructions to effectors via electrical (nervous system) or chemical (endocrine system hormones) signals.
  3. Effector (Target):
     • Target cells, tissues, or organs responding to control center instructions.
     • Action: Executes actions to counteract or amplify the stimulus to restore balance. Typically muscle tissue or glandular tissue.

• Two Regulatory Approaches:

  1. Autoregulation (Intrinsic Regulation):
     • Automatic, localized cellular adjustments occurring directly within a cell, tissue, organ, or organ system in response to local environmental shifts.
     • Example: Oxygen-deprived tissue cells locally release vasodilator chemicals that dilate local blood vessel capillaries, increasing localized blood flow to restore oxygenation without activating central nervous or hormonal systems.
  2. Extrinsic Regulation:
     • Systemic regulatory responses controlled remotely by the nervous system (rapid electrical signals) or endocrine system (slower, sustained hormonal signals).
     • Example: Touching a hot stove triggers pain sensory receptors, sending signals to the central nervous system, which commands specific arm muscles to contract, rapidly removing the hand from danger.

• Feedback Loops:

  1. Negative Feedback Loop:
     • Effector response opposes or negates the initial stimulus direction.
     • Primary mechanism responsible for keeping physiological parameters within a normal range around a designated set point.
     • Thermostat Analogy: Elevated room temperature (stimulus) is detected by a thermometer (receptor) and processed by a thermostat set at $22^\circ\text{C}$ (control center), which commands an air conditioner (effector) to activate, lowering room temperature back to normal.
  2. Positive Feedback Loop:
     • Effector response reinforces or amplifies the initial stimulus direction, enhancing deviation from normal.
     • Uncommon in daily physiology; utilized when potentially dangerous or critical processes must be completed rapidly to re-establish homeostasis.
     • Example: Blood Clotting Cascade:
       1. Damaged cells within an injured blood vessel wall release chemical factors.
       2. Chemicals initiate chain reactions causing blood cells, platelets, and soluble clotting proteins to assemble.
       3. Accelerating clot formation releases additional chemicals at each step.
       4. Escalating loop terminates when a solid blood clot patches the damaged vessel wall and halts bleeding.

Practical Applications, Exercises, and Thought Problems

• Solutions and Water Properties Exercises:

  • IV Drip Scenario: A patient receives an intravenous drip containing saline, electrolytes, and sugar designed to dissolve glucose. The solvent is water/saline, and the solute is glucose.
  • Saline Mixture Scenario: Mixing $1\,\text{tsp}$ of table salt ($\text{NaCl}$) into $3\,\text{cups}$ of water forms a solution. The solute is salt ($\mathbf{NaCl}$), and the solvent is water.
  • Sweating Cooling Mechanism: Sweating cools the body because water possesses a high heat capacity and high heat of vaporization. As surface sweat evaporates, it absorbs and removes significant body heat.
  • Playground Slide vs. Pond Heat Capacity: On a hot day, a metal slide becomes scorching hot while a nearby pond remains cool because water has a significantly higher heat capacity than metal due to hydrogen bonding. This exact property (thermal inertia) prevents human body fluids from overheating rapidly.

• pH Calculations and Evaluations:

  • pH Solution Classifications:
    • $\text{pH } 3$: Acidic
    • $\text{pH } 9$: Basic (Alkaline)
    • $\text{pH } 7$: Neutral
    • $\text{pH } 1$: Strongly Acidic
    • $\text{pH } 8$: Slightly Basic (Alkaline)
  • Relative Hydrogen Concentration Comparison: A solution with a $\text{pH } 3$ contains significantly more $\text{H}^+$ ions than a solution with a $\text{pH } 9$ (exactly $10^6\times$ or $1,000,000\times$ higher $\text{H}^+$ concentration).

• Buffer Mechanism Exercise:

  • Dual-Component Necessity: A functional buffer requires both a weak acid and a weak base so it can resist pH changes in both directions. The weak base neutralizes added $\text{H}^+$ (acid), while the weak acid neutralizes added $\text{OH}^-$ (base).