BIO 171 A&P I Week 3 Flashcards

Chemical & Biological Buffer Systems

• Acids and Bases

  • Acids: Defined as proton donors. They donate hydrogen ions ($\text{H}^+$), which increases the concentration of $\text{H}^+$ in a solution. A higher concentration of $\text{H}^+$ results in greater acidity.

  • Bases: Defined as proton acceptors. They accept hydrogen ions ($\text{H}^+$), which decreases the concentration of $\text{H}^+$ in a solution. A lower concentration of $\text{H}^+$ results in a less acidic (more basic or alkaline) solution.

  • Hydrogen Ion ($\text{H}^+$): A neutral hydrogen atom becomes a hydrogen ion ($\text{H}^+$) when it loses an electron. Because protons and electrons normally balance out to neutral charge in an element, losing its single negative electron leaves only a single positive proton in the nucleus, rendering the ion positively charged.

• The pH Scale

  • Definition: pH stands for "Potential of Hydrogen".

  • Acidic Range: pH values from $0$ to $6.99$. Characterized by a greater concentration of hydrogen ions ($\text{H}^+$) relative to hydroxide ions ($\text{OH}^-$).

  • Basic/Alkaline Range: pH values from $7.01$ to $14$. Characterized by a lower concentration of hydrogen ions ($\text{H}^+$) relative to hydroxide ions ($\text{OH}^-$).

• Carbonic Acid–Bicarbonate Buffer Equation

  • Origin of Reactants: Carbon dioxide ($\text{CO}_2$) and water ($\text{H}_2\text{O}$) are produced continuously in the human body as waste products of cellular respiration.

  • Chemical Reaction: When $\text{CO}_2$ and $\text{H}_2\text{O}$ react, they form carbonic acid ($\text{H}_2\text{CO}_3$):         $\text{CO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{CO}_3$

  • Dissociation: Carbonic acid ($\text{H}_2\text{CO}_3$) is an unstable weak acid. When placed in water, it dissociates into a hydrogen ion ($\text{H}^+$) and a bicarbonate ion ($\text{HCO}_3^-$):         $\text{H}_2\text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^-$

  • Complete Reversible Buffer Reaction:         $\text{CO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^-$

• Buffer Dynamics & Homeostasis

  • Rightward Shift: Moving to the right adds free $\text{H}^+$ to the solution, causing pH to decrease and acidity to increase. Example: Holding one's breath causes a buildup of $\text{CO}_2$, driving the reaction to the right.

  • Leftward Shift: Moving to the left adds bicarbonate ions ($\text{HCO}_3^-$) to bind free $\text{H}^+$, removing $\text{H}^+$ from solution. This causes pH to increase and acidity to decrease. Example: Blood approaching acidosis triggers a shift to the left.

  • Normal Blood pH Range: Normal blood pH is tightly regulated between $7.35$ and $7.45$

  • Acidosis: Clinical condition where blood pH drops below $7.35$

  • Alkalosis: Clinical condition where blood pH rises above $7.45$

  • Physiological Response to Acidosis: The body introduces bicarbonate ions ($\text{HCO}_3^-$) into the bloodstream to bind extra $\text{H}^+$, neutralizing excess acidity and restoring homeostasis within the strict $7.35 - 7.45$ blood pH window.

Levels of Structural Organization

• Hierarchical Levels of Human Body Organization

  • 1. Chemical Level: The lowest level of structural organization. Atoms (such as Hydrogen and Oxygen) combine via chemical bonding to form molecules with three-dimensional structures (such as Water molecules).

  • 2. Cellular Level: A variety of complex molecules combine to form the fluid and organelles of a functional body cell (e.g., smooth muscle cells, organelles, and intracellular fluid).

  • 3. Tissue Level: A community of similar cells working together to perform a specialized function forms a body tissue (e.g., smooth muscle tissue, skeletal muscle tissue).

  • 4. Organ Level: Two or more different tissue types combine to form a discrete structure with specific functions (e.g., the urinary bladder).

  • 5. Organ System Level: Two or more organs work in close coordination to perform complex systemic functions (e.g., the urinary tract system, consisting of the kidneys, ureters, urinary bladder, and urethra).

  • 6. Organismal Level: Many distinct organ systems work harmoniously together to maintain life and support the functions of an independent living organism.

Fundamentals of Atomic Structure & The Periodic Table

• The Periodic Table of Elements

  • Element: A pure substance that cannot be broken down into simpler substances by ordinary chemical means. Each box on the periodic table represents one unique element.

  • Chemical Compounds: Substances like water ($\text{H}_2\text{O}$) result from a chemical reaction and chemical bonding between distinct elements (Hydrogen and Oxygen).

  • Chemical Symbol: Represents the element's name using a single capital letter or a capital letter combined with a second lowercase letter.

  • Atomic Number: Displayed in the top corner of an element's entry; indicates the exact number of protons in an atom's nucleus. The periodic table is organized sequentially by atomic number.

  • Atomic Mass: Displayed at the bottom of an element's entry; indicates the combined mass of both protons and neutrons located in the atomic nucleus, reflecting the overall heaviness of an atom relative to others.

  • Columns (Groups): Elements in the same vertical column behave similarly chemically and perform similar functions because they share the same number of electrons in their outermost valence shells (denoted by Roman numerals).

  • Rows (Periods): Horizontal rows represent energy levels being filled by electrons.

  • Major Elements of the Human Body: Oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus.

• Atomic Structure

  • Atom: The smallest particle that can exist as a distinct element; cannot be further sub-divided by chemical means.

  • Nucleus: Solid core of the atom housing positively charged protons ($+$) and uncharged neutrons ($0$), which together constitute the atomic mass.

  • Electrons ($-$): Small subatomic particles with a negative charge located in shells/orbitals surrounding the nucleus. They spin rapidly, act as waves of energy, and mediate chemical reactions.

  • Atomic Mass Units (AMUs): The measurement unit used to quantify atomic mass.

  • Electrical Neutrality: In any standard element on the periodic table, the number of electrons equals the number of protons, rendering the overall atom electrically neutral.

  • Calculating Neutrons: To determine the number of neutrons in an atom, subtract the atomic number (protons) from the rounded atomic mass:         $\text{Number of Neutrons} = \text{Atomic Mass} - \text{Atomic Number}$

Isotopes & Radioactivity

• Isotopes

  • Definition: Atoms of the same element that possess the exact same atomic number (same number of protons) but different atomic mass numbers (different number of neutrons).

  • Elemental Identity: The atomic number dictates proton count and element identity; altering proton count creates a different element entirely, whereas altering neutron count creates an isotope.

  • Carbon Isotopes:

    • Carbon-12 ($^{12}\text{C}$): Contains $6$ protons and $6$ neutrons. Non-radioactive.

    • Carbon-13 ($^{13}\text{C}$): Contains $6$ protons and $7$ neutrons. Non-radioactive.

    • Carbon-14 ($^{14}\text{C}$): Contains $6$ protons and $8$ neutrons. Typical radioactive isotope of carbon.

• Radioactive Decay & Emissions

  • Radioactive Isotopes: Unstable isotopes (such as Carbon-14) whose nuclei spontaneously decay, emitting subatomic particles and energy over time to achieve stability.

  • Alpha Particles ($\alpha$): Composed of $2$ protons and $2$ neutrons ejected from the nucleus (essentially a Helium nucleus, $^4_2\text{He}$). Harmless under normal exposure conditions.

  • Beta Particles ($\beta$): High-speed subatomic electrons ejected from the atom. Neutrons are conceptualized as being composed of equal parts proton and electron to maintain neutral charge. Beta particles can be blocked and shielded using aluminum foil.

  • Gamma Rays ($\gamma$): High-energy electromagnetic radiation emitted from the nucleus. Poses severe biological risk; high-energy radiation (including intense solar radiation) can damage cellular DNA, inducing mutations and tumor formation.

• Half-Life Concepts

  • Half-Life (General): The duration of time required for half of the radioactive atoms in a sample nucleus to decay into a stable state; measures the rate of decay.

  • Physical Half-Life ($t_{1/2}$): The specific time needed for half of the radioactive nuclei in a sample to undergo physical radioactive decay.

  • Biological Half-Life: The time required for the human body to eliminate half of a given substance through natural biological excretion pathways.

  • Effective Half-Life: The overall time required for radioactivity inside the body to be reduced by half; represents a combined parameter of physical half-life and biological half-life.

Endocrine Physiology: Thyroid Hormones & Metabolism

• Thyroid Gland & Hormones

  • Anatomy: The thyroid gland is situated directly inferior to the larynx (voice box).

  • Iodine Utilization: Thyroid hormones incorporate iodine, which possesses radioisotopes.

  • Triiodothyronine ($\text{T}_3$): Key thyroid hormone that regulates basal metabolic rate, cardiac function, skeletal muscle control, and brain development.

  • Thyroxine ($\text{T}_4$): The primary hormone secreted by the thyroid gland; transported throughout the body and converted into $\text{T}_3$ (the biologically active form).

  • Hypothyroidism: Clinical state defined by pathologically low circulatory levels of $\text{T}_3$ and $\text{T}_4$

• Cellular Metabolism & Obligate Glucose Aerobes

  • Metabolic Role of Thyroid Hormones: Regulate cellular respiration and adenosine triphosphate ($\text{ATP}$) production in every cell across the human body.

  • Neuronal Energy Demands: Neurons are extremely sensitive to $\text{ATP}$ availability due to high energy requirements fueled by $\text{ATP}$.

  • Obligate Glucose Aerobes: Neurons in the brain and peripheral nervous system are obligate glucose aerobes, meaning they have an absolute physiological requirement for both glucose and oxygen. The brain consumes approximately $20\%$ of total systemic glucose.

  • Glucose Preservation: If systemic glucose levels drop, the body prioritizes and preserves available glucose strictly for brain tissue.

• Symptomatology of Hypothyroidism

  • Low Energy: Reduced $\text{T}_3$ and $\text{T}_4$ impair cellular $\text{ATP}$ synthesis, causing systemic fatigue.

  • Cold Intolerance: Cellular respiration converts energy from food into $\text{ATP}$, but approximately $60\%$ of total energy is released as body heat. Reduced metabolic rate in hypothyroidism decreases heat generation, leaving patients chronically cold.

  • Weight Gain & Adiposity: As long as the pancreas secretes insulin, circulating glucose enters cells. If cells cannot consume this glucose via cellular respiration, the excess glucose is stored as glycogen and subsequently converted into body fat.

• Hypothalamic-Pituitary-Thyroid Axis

  • Thyrotropin-Releasing Hormone ($\text{TRH}$): Secreted by the hypothalamus.

  • Thyroid-Stimulating Hormone ($\text{TSH}$): Secreted by the anterior pituitary gland.

  • Regulatory Loop: $\text{TRH}$ from the hypothalamus stimulates the anterior pituitary to release $\text{TSH}$, which subsequently stimulates the thyroid gland to synthesize and secrete $\text{T}_3$ and $\text{T}_4$

Stoichiometry & The Mole Concept

• The Mole & Avogadro's Number

  • Definition: One mole of any substance contains exactly $6.023 \times 10^{23}$ particles (Avogadro's number).

  • Molar Mass: The mass of one mole of a substance expressed in grams ($g$), equivalent to the element's atomic mass value on the periodic table rounded to the hundredths place ($1/100\text{th}$).

  • Carbon Molar Mass Example: One mole of Carbon contains $6.023 \times 10^{23}$ carbon atoms and has a total mass of $12.01\,g$

• Specific Molar Calculations

  • Boron ($\text{B}$):

    • Protons = $5$; Electrons = $5$

    • Atomic Mass = $10.811$, which rounds to $11$ total nuclear particles.

    • Neutrons = $11 - 5 = 6$ neutrons.

    • One mole of Boron contains $6.023 \times 10^{23}$ particles and has a mass of $10.81\,g$

  • Lithium ($\text{Li}$):

    • Protons = $3$; Electrons = $3$

    • Atomic Mass = $6.941$, which rounds to $7$ total nuclear particles.

    • Neutrons = $7 - 3 = 4$ neutrons.

    • One mole of Lithium contains $6.023 \times 10^{23}$ particles and has a mass of $6.94\,g$

Chemical Bonding & Electronegativity

• Electrons & Energy Storage

  • Valence Shell Stability: Atoms achieve structural stability by completely filling their outermost valence energy shell (requiring $8$ electrons in most outer shells, or $2$ in the first shell).

  • Chemical Energy: The chemical energy found in organic molecules (e.g., glucose) resides in covalent bonds holding carbon atoms together. Covalent bonds represent shared electrons storing potential chemical and electrical energy.

  • Electricity: Defined physically as the flow of moving electrons.

• Covalent Bonds

  • Definition: Strong chemical bonds formed when two atoms share valence electrons through the overlap of their atomic orbitals.

  • Energy Dynamics:

    • Bond Formation: Anabolic processes (such as dehydration synthesis) require energy input to build covalent bonds.

    • Bond Cleavage: Catabolic processes (such as hydrolysis and cellular respiration) break covalent bonds to release stored potential energy.

  • Group Tendency: Covalent bonds are characteristic of elements in Groups III, IV, V, and VI of the periodic table.

  • Nonpolar Covalent Bonds: Occur when electrons are shared equally between two bonding atoms. This happens when two identical atoms bond (sharing identical electronegativities) or atoms with similar pull bond.

    • Examples: Hydrogen gas ($\text{H}_2$) and Methane gas ($\text{CH}_4$).

    • Representation: A single line represents $1$ covalent bond containing $2$ shared electrons.

  • Polar Covalent Bonds: Occur when electrons are shared unequally between atoms because one atom exerts a stronger attraction on shared electrons.

    • Example: Water ($\text{H}_2\text{O}$).

  • Double Bonds: Carbon dioxide ($\text{O}=\text{C}=\text{O}$) features double covalent bonds between carbon and each oxygen atom. Two lines represent a double bond consisting of $4$ shared electrons total per double bond ($2$ shared electrons per line).

• Ionic Bonds

  • Definition: Electrostatic attractions formed between oppositely charged ions resulting from the complete transfer of one or more valence electrons from one atom to another.

  • Stability: Ionic bonds are weaker and less stable than covalent bonds.

  • Cation ($+$): An atom that has lost one or more electrons, acquiring a net positive charge.

  • Anion ($-$): An atom that has gained one or more electrons, acquiring a net negative charge.

  • Electrolytes: Dissociated ions floating in aqueous solution that conduct electrical current.

  • Sodium Chloride ($\text{NaCl}$) Example:

    • Sodium ($\text{Na}$): $11$ protons, $12$ neutrons, $1$ electron in valence shell (Group 1). Highly unstable; seeks to dump its $1$ valence electron to form a cation ($\text{Na}^+$).

    • Chlorine ($\text{Cl}$): $17$ protons, $19$ neutrons, $7$ electrons in valence shell (Group 7). Seeks to gain $1$ electron to complete its octet, forming an anion ($\text{Cl}^-$).

    • The transfer of $1$ electron from $\text{Na}$ to $\text{Cl}$ creates an ionic bond.

  • Group 2 & Group 7 Trends:

    • Group 2 elements dump $2$ valence electrons to form divalent cations (e.g., Calcium in Calcium Chloride: $\text{Ca}^{2+} + 2\text{Cl}^- \rightarrow \text{CaCl}_2$).

    • Group 7 elements possess $7$ valence electrons and acquire $1$ electron to form anions ($-$, gain electron) or form nonpolar covalent bonds.

• Electronegativity

  • Definition: The measure of an atom's relative attraction or pull on shared bonding electrons.

  • Periodic Trend: Electronegativity increases progressively from left to right across the periodic table from Group 1 to Group 7.

  • Group 8: Noble gases are inert, completely stable, and do not form bonds under standard conditions.

  • Oxygen Properties: Oxygen sits in Group 6 and is extremely electronegative ($6$ valence electrons). It forcefully pulls electrons toward itself, oxidizes carbon compounds, and breaks covalent bonds in glucose to harvest energy.

Physical Properties & Biological Functions of Water

• Molecular Geometry & Polarity of Water

  • Structure: Water ($\text{H}_2\text{O}$) possesses a bent molecular geometry.

  • Polar Covalent Nature: Oxygen is significantly more electronegative than hydrogen. Oxygen acts as an electron hog, pulling shared electrons closer to its nucleus.

  • Charge Distribution: Creates a partial negative charge (denoted $\delta^-$) on the oxygen pole, while exposing hydrogen protons to create a partial positive charge (denoted $\delta^+$) on the hydrogen pole.

  • Reversible Re-association: Water undergoes minor internal dissociation ($\text{H}_2\text{O} \rightleftharpoons \text{H}^+ + \text{OH}^-$). Because hydroxide ($\text{OH}^-$) is a strong base, it readily re-attaches to free $\text{H}^+$.

• Hydrogen Bonding

  • Definition: Weak intermolecular electrostatic attractions that occur between separate molecules (not within a single molecule).

  • Mechanism: The partial negative ($\delta^-$) oxygen atom of one water molecule attracts the partial positive ($\delta^+$) hydrogen atom of a neighboring water molecule.

  • Significance: Hydrogen bonding provides water with unique cohesive, thermal, and solvent characteristics necessary to support life.

• Phases of Water

  • Gas (Water Vapor): Composed of low molecular mass substances in a gaseous state.

  • Liquid: Represents almost all water within the human body at standard body temperature.

  • Solid (Ice): Ice is less dense than liquid water because hydrogen bonds expand to create more space between molecules in a rigid matrix; consequently, ice floats.

• Physiological Functions of Liquid Water

  • 1. Transport: Dissolved solutes move easily throughout body fluid compartments.

  • 2. Lubrication: Reduces friction between sliding anatomical structures (e.g., serous membranes).

  • 3. Cushioning: Absorbs sudden physical impacts caused by body movements.

  • 4. Waste Excretion: Soluble waste products dissolve in water for rapid elimination.

• Cohesive and Adhesive Properties

  • Cohesion: The attraction between adjacent water molecules due to hydrogen bonding.

  • Surface Tension: The inward pulling force exerted by cohesive forces at the surface of liquid water. Surface tension causes moist air sacs (alveoli) in the lungs to collapse if uncorrected.

  • Adhesion: The attraction between water molecules and a different non-water substance.

• Thermal Properties

  • Specific Heat: The amount of thermal energy required to raise the temperature of $1\,g$ of a substance by $1^\circ\text{C}$. Water has an extremely high specific heat because significant heat energy must be consumed to break hydrogen bonds before kinetic molecular movement increases. This maintains constant body temperature.

  • Heat of Vaporization: The heat required to transform $1\,g$ of a liquid substance into a gaseous phase. Water has a high heat of vaporization. Sweating cools the body because excess thermal energy is absorbed and dissipated into the environment as water evaporates.

Solution Chemistry & Solute Interactions

• Water as the Universal Solvent

  • Solvent vs. Solute: Water acts as the principal solvent in the body. Solutes are dissolved substances.

  • Universal Solvent: Water is termed the universal solvent because most biological molecules dissolve in it.

  • Solubility Rule: Only polar molecules and charged ions readily dissolve in water.

• Hydrophilic Solutes

  • Hydrophilic: Meaning "water-loving". Refers to polar molecules and charged ions.

  • Hydration Shell: Spheres of water molecules that surround charged or polar particles in solution.

  • Nonelectrolytes: Substances that dissolve in water but remain molecularly intact (e.g., glucose, alcohol). They do not dissociate into ions and cannot conduct electrical currents.

  • Electrolytes: Substances that dissolve and completely dissociate (separate) into charged ions (e.g., salts such as $\text{NaCl} \rightarrow \text{Na}^+ + \text{Cl}^-$, acids like $\text{HCl}$, and bases like $\text{NaOH}$). Dissociated ions conduct electric currents.

• Hydrophobic Solutes

  • Hydrophobic: Meaning "water-fearing". Refers to nonpolar, uncharged molecules.

  • Hydrophobic Exclusion: Cohesive water molecules exert force that pushes nonpolar molecules out of aqueous solution.

  • Hydrophobic Interaction: The clustering behavior of excluded nonpolar molecules together when placed in water.

  • Blood Transport: Hydrophobic substances (such as fats and cholesterol) cannot dissolve in blood water and require specialized carrier proteins to be transported through the bloodstream.

• Amphipathic Molecules

  • Definition: Molecules containing both polar (hydrophilic) and nonpolar (hydrophobic) regions.

  • Behavior in Water: The polar region dissolves in water while the nonpolar region is repelled by water.

  • Phospholipids: Classic amphipathic molecules containing polar heads (which contact water) and nonpolar tails (which group together away from water).

  • Structures Formed: Amphipathic molecules spontaneously assemble into phospholipid bilayers (forming cellular membranes) and micelles (e.g., bile salts).