Ch 2.1-2.2

Introduction

• Chapter Overview: Introduction to Chapter Two, bridging the gap between basic chemistry and human biology.

• Reference to Chapter One: While the smallest unit of life is the cell, biological function is rooted in the interactions of atoms and molecules.

• Focus: Detailed discussion of chemistry and the critical role chemical structures play in anatomy and physiology. Understanding these concepts is vital for understanding metabolism and homeostatic mechanisms.

• Challenge Level: Chapters Two and Three are foundational; mastery of this prerequisite chemistry is essential for success in later physiological systems.

Definition of Matter

• Matter: Defined as anything that has mass and occupies space.

• Mass vs. Weight:

  • Mass: The amount of matter in an object, which remains constant regardless of gravity.

  • Weight: The force exerted by gravity on an object's mass.

• Forms of Matter: Matter exists in three primary states in the human body:

  • Solid: Definite shape and volume (e.g., bones, teeth).

  • Liquid: Definite volume but takes the shape of its container (e.g., blood plasma, interstitial fluid).

  • Gas: Neither definite shape nor volume (e.g., $O<em>{2}$ and $CO</em>{2}$ in the lungs).

• Additional Form: Plasma is a high-energy state of matter not found within the human body under normal physiological conditions.

Atoms and Elements

• Elements: Pure substances that cannot be broken down into simpler substances by ordinary chemical means (e.g., gold $(Au)$, iron $(Fe)$, and oxygen $(O)$).

• Atoms: The smallest units of an element that retain its unique properties.

  • Hydrogen: The simplest and smallest element, consisting of one proton and one electron.

  • Uranium: One of the heaviest naturally occurring elements.

Composition of the Human Body

• Major Elements (Account for over $98.5\%$ of total body mass):

  • Oxygen ($O$): $65.0\%$ — Essential for cellular respiration and a component of water.

  • Carbon ($C$): $18.5\%$ — The "backbone" of all organic molecules (carbohydrates, lipids, proteins).

  • Hydrogen ($H$): $9.5\%$ — Found in water and all organic molecules; influences $\text{pH}$.

  • Nitrogen ($N$): $3.2\%$ — Component of proteins and nucleic acids ($DNA/RNA$).

  • Calcium ($Ca$) and Phosphorus ($P$) also significant for bone structure and energy transfer.

• Minor Elements (Less than $1\%$ of body weight):

  • Includes Sulfur ($S$: $0.25\%$), Potassium ($K$: $0.2\%$), Sodium ($Na$: $0.15\%$), Chlorine ($Cl$: $0.15\%$), Magnesium ($Mg$: $0.05\%$), and Iron ($Fe$: $0.006\%$).

• Trace Elements: Present in amounts less than $0.01\%$, but vital for health (e.g., Iodine for thyroid function, Zinc for enzyme activity).

Atomic Structure

• Atoms: Composed of three primary subatomic particles:

  • Protons: Positively charged ($+1$), located in the central nucleus. The number of protons determines the element's identity (Atomic Number).

  • Neutrons: Electrically neutral (no charge), located in the nucleus. They contribute to atomic mass and stability.

  • Electrons: Negatively charged ($-1$), found in constant motion within orbitals/shells surrounding the nucleus.

• Mass Relationship: Protons and neutrons are significantly heavier than electrons; an electron's mass is approximately $\frac{1}{1800}$ the mass of a proton.

• Nucleus: The dense central core containing protons and neutrons; it carries an overall positive charge.

Atomic Mass Unit (AMU)

• AMU (Dalton): A standard unit for measuring atomic weight.

  • Protons and Neutrons each weigh approximately $1 \text{ AMU}$.

  • Electrons have a mass so small ($0.0005 \text{ AMU}$) that it is considered negligible for basic calculations.

Electron Configuration and the Periodic Table

• Electron Shells: Regions around the nucleus representing different energy levels.

  • First Shell (Valence Shell 1): Can hold a maximum of $2$ electrons.

  • Outer Shells: Can hold up to $8$ electrons (for the elements relevant to biology).

• Valence Electrons: Electrons in the outermost shell that determine the chemical reactivity and bonding behavior of the atom.

• Periodic Table Organization:

  • Chemical Symbol: A one- or two-letter abbreviation (e.g., $Na$ for Sodium, derived from Natrium).

  • Atomic Number ($Z$): Number of protons in the nucleus.

  • Atomic Mass ($A$): The weighted average of the mass of all naturally occurring isotopes.

Calculating Subatomic Particles

• Proton Count: Equals the Atomic Number.

• Electron Count: In a neutral atom, electrons equal the number of protons.

• Neutron Count: Calculated as $\text{Mass Number} - \text{Atomic Number}$.

  • Example: Sodium ($Na$) with a mass of $23$ and atomic number of $11$: $23 - 11 = 12 \text{ Neutrons}$.

Isotopes and Radioactivity

• Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons (e.g., $^{12}C$, $^{13}C$, and $^{14}C$).

• Radioisotopes: Unstable isotopes that decay over time, releasing radiation.

  • Physical Half-Life: Time required for $50\%$ of a radioisotope to decay into a stable form.

  • Biological Half-Life: Time required for the body to eliminate $50\%$ of a radioactive substance through excretion.

• Medical Uses: Radioisotopes are used in imaging (e.g., Iodine-131 for thyroid scans) and localized cancer treatments.

Chemical Bonding and Molecules

• Compounds: Substances composed of two or more different elements in fixed proportions (e.g., $H_{2}O$).

• Octet Rule: Atoms tend to react in ways that leave them with a full outer shell (usually $8$ electrons) to achieve maximum stability.

Ionic vs. Covalent Bonding

• Ionic Bonds: Formed when one atom transfers electrons to another, resulting in electrostatic attraction between oppositely charged ions.

  • Ions: Charged particles.

  • Cations: Positively charged ions (formed by losing electrons; e.g., $Na^+$). Mnemonic: "The 't' in cation looks like a plus sign."

  • Anions: Negatively charged ions (formed by gaining electrons; e.g., $Cl^-$).

  • Example: $NaCl$ (Table Salt). Sodium loses an electron to Chlorine.

• Covalent Bonds: Formed when atoms share electrons.

  • Nonpolar Covalent: Electrons are shared equally (e.g., $O<em>{2}$, $CH</em>{4}$).

  • Polar Covalent: Electrons are shared unequally due to differences in electronegativity, creating partial charges (e.g., water, where Oxygen is $\delta-$ and Hydrogen is $\delta+$).

• Polyatomic Ions: Groups of covalently bonded atoms that carry a net charge (e.g., Bicarbonate $HCO<em>{3}^-$ or Phosphate $PO</em>{4}^{3-}$).

Summary of Common Ions in Physiology

• Common Cations:

  • Sodium ($Na^+$): Principal extracellular cation; critical for water balance and nerve impulses.

  • Potassium ($K^+$): Principal intracellular cation; essential for muscle contraction and heart rhythm.

  • Calcium ($Ca^{2+}$): Vital for bone health, blood clotting, and neurotransmitter release.

  • Magnesium ($Mg^{2+}$): Required for many enzymatic reactions.

  • Hydrogen ($H^+$): Concentration determines the acidity ($\text{pH}$) of body fluids.

• Common Anions:

  • Chloride ($Cl^-$): Most abundant extracellular anion; tracks with sodium.

  • Bicarbonate ($HCO_{3}^-$): Acts as a buffer to maintain blood $\text{pH}$.

  • Phosphate ($PO_{4}^{3-}$): Found in bones, $ATP$ (energy), and the sugar-phosphate backbone of $DNA$.

Conclusion

• All physiological processes are chemical in nature. Atoms bond to achieve stability via the octet rule, leading to the formation of molecules and ionic salts like $NaCl$.

• The balance of ions (electrolytes) and the structure of molecules are fundamental to cellular function and overall human health.