Chemistry & Physiology: Atoms, Bonds, Water, and Metabolism — Study Notes
Imaging Techniques in Medicine
Review of topics from last week; sign-in options: QR code or PIN.
Medical imaging techniques discussed:
X-ray: good for visualizing bones; still involves ionizing radiation and can be harmful.
PET scan: uses radiopharmaceuticals to visualize physiological functions in tissues; involves radioactivity.
MRI (Magnetic Resonance Imaging): uses a strong magnet to obtain highly detailed images; no ionizing radiation.
Ultrasound: uses sound waves to look inside the body; generally safe for the fetus; commonly used for pregnancy imaging.
Summary takeaway: each imaging modality has different strengths, safety considerations, and best-use scenarios.
Subatomic Particles and Atomic Structure
A substance cannot be broken down into other substances by chemical reaction is defined as an element.
Proton: subatomic particle with a positive charge; located in the nucleus.
Electron: subatomic particle with a negative charge; located in the electron cloud (outside the nucleus).
Neutron: subatomic particle with no charge (neutral).
Atom: smallest unit of matter; has mass and takes up space; consists of nucleus (protons + neutrons) and surrounding electrons.
Key charges: proton +, electron −, neutron 0.
Proton number = atomic number Z; defines the element. Example: element with Z = 7 is nitrogen.
Mass number A = Z + N (protons + neutrons).
When two atoms of different elements bond, they form a compound; if the same element bonds, it forms a molecule.
Isotopes: atoms with the same number of protons but different numbers of neutrons.
Basics of chemical reactivity: determined by electrons in the valence shell (outermost shell).
Four major elements making up most of the body: Carbon (C), Hydrogen (H), Oxygen (O), Nitrogen (N) – collectively CHON.
Periodic Table and Electron Distribution
Periods = rows; groups = columns.
Periods tell you how many electron shells an atom has.
First period: maximum of 2 electrons in the first shell.
Second period: maximum of 8 electrons in the second shell.
Third period: can have up to 18 electrons in the third shell (an exception noted for very small sets of elements).
Group number corresponds to valence shell electrons (outermost electrons). Example:
Group 1: 1 valence electron
Group 2: 2 valence electrons
Group 3: 3 valence electrons
…
Group 8: 8 valence electrons (noble gases) — typically nonreactive.
Octet rule: atoms tend to have a full outer shell with 8 electrons (valence electrons) to achieve stability.
Special note on hydrogen: often considered an exception; hydrogen in group 1 has 1 valence electron but seeks 1 more electron to fill its first shell (2 electrons for the first shell).
Examples of valence electron counts:
Oxygen (group 6): 6 valence electrons; wants 2 more to reach 8.
Fluorine (group 7): 7 valence electrons; wants 1 more to reach 8.
Carbon (group 4): 4 valence electrons; wants 4 more to reach 8.
Concept: atoms “fill their holes” by gaining, losing, or sharing electrons to reach stability.
Ions and Electrical Charge
Ions: charged atoms formed by gaining or losing electrons.
Hydrogen example:
Neutral hydrogen: 1 proton, 1 electron (Z = 1); mass ~1.01; mass number A ≈ 1.
Gaining an electron: becomes negatively charged (cation vs anion confusion). More precisely:
Gaining electron: becomes an anion (negative charge).
Losing electron: becomes a cation (positive charge).
Definitions:
Anion: negatively charged ion.
Cation: positively charged ion.
Practical examples:
Sodium (Na) in group 1 tends to lose 1 electron to form Na⁺ (cation).
Magnesium (Mg) in group 2 tends to lose 2 electrons to form Mg²⁺ (cation).
Fluorine (F) in group 7 tends to gain 1 electron to form F⁻ (anion).
Electronegativity: tendency of an atom to attract electrons. Oxygen and nitrogen are highly electronegative and pull electrons toward themselves.
Ion formation in physiology: sodium and chloride ions illustrate how body chemistry relies on gaining or losing electrons to form ions.
Key terms:
Electrons are lost or gained to form ions; opposite charged ions bond via ionic bonds.
Chemical Bonding: Covalent, Ionic, and Hydrogen Bonds
Bonding is the attraction that holds atoms together; largely determined by valence electrons.
Three main bond types: covalent, ionic, and hydrogen bonds.
Covalent bonds:
Electrons are shared between two atoms.
Can be single (one pair shared) or multiple (e.g., double bonds share two pairs).
Example: H–H (H2) is a covalent bond.
Polar covalent vs nonpolar covalent:
Polar covalent bond: unequal sharing of electrons (e.g., H2O; oxygen is more electronegative, pulling shared electrons closer to itself).
Nonpolar covalent bond: equal sharing of electrons (e.g., O=O, or C–H in hydrocarbons generally nonpolar).
Ionic bonds:
Formation between oppositely charged ions (cations and anions).
Example: NaCl (sodium chloride) formed from Na⁺ and Cl⁻.
Hydrogen bonds:
Very weak interactions between a partially positive hydrogen and a partially negative atom in another molecule (or another part of the same molecule).
In water, hydrogen bonds occur between separate water molecules, giving water its characteristic properties.
Molecules vs compounds (revisited):
Molecule: two or more atoms bonded together.
Compound: two or more atoms of different elements bonded together.
Special notes:
Covalent bonds can be nonpolar or polar depending on electronegativity differences.
Hydrogen bonding is essential for water's behavior and many biological interactions.
Water: Structure, Properties, and Role in Biology
Water molecule structure: H₂O with one oxygen atom covalently bonded to two hydrogens.
Oxygen has 6 valence electrons; it tends to form two bonds to complete its shell (two hydrogen atoms).
Bond type between O and H: polar covalent (electrons not shared equally; oxygen pulls electrons closer).
Partial charges: oxygen is partially negative; hydrogens are partially positive (denoted with δ− and δ+).
Key properties of water in biology:
Lubricant and cushion: reduces friction in joints; protects organs from physical trauma.
Heat sink: large capacity to absorb and release heat with minimal temperature change; stabilizes body temperature.
Solvent system: dissolves many substances; forms hydration shells around dissolved ions.
Why water moderates temperature:
Hydrogen bonding must be broken to heat water, which requires energy; this slows temperature rise.
Evaporation of water (sweat cooling) releases energy and cools the body.
Hydration shells and solubility:
When salts dissolve (e.g., NaCl), water molecules surround ions forming hydration shells: partial negative oxygen around Na⁺, partial positive hydrogens around Cl⁻.
Hydration shells are essential for dissolving salts and many polar or charged solutes.
Hydrophilic vs hydrophobic:
Hydrophilic: water-loving; substances that dissolve in water (polar molecules and ions).
Hydrophobic: water-fearing; substances that do not dissolve in water (nonpolar; e.g., oils, fats).
Salt dissolution: water as solvent dissolves salts by surrounding ions and stabilizing them in solution.
Examples:
Salt is NaCl; in water, Na⁺ is stabilized by the partial negative charge of water's oxygen, while Cl⁻ is stabilized by the partial positive charges of water's hydrogens.
Answerable concepts:
Hydrophilic substances are typically polar or charged; hydrophobic substances are nonpolar.
Water's solvent properties underpin many biological processes.
Solutions: Solvent, Solute, and Hydration
Solution: a homogeneous mixture of solute(s) dissolved in a solvent.
Solvent: the component present in the greater amount that does the dissolving (often water in biological systems).
Solute: the substance dissolved in the solvent (e.g., salt, sugars).
Hydration shells: layers of water molecules organized around dissolved ions or molecules.
Example: Dissolving NaCl in water forms Na⁺(aq) and Cl⁻(aq) with hydration shells.
Hydrophilic vs Hydrophobic: Practical Implications
Hydrophilic substances: dissolve in water due to polar bonds or ionic charges.
Hydrophobic substances: do not dissolve in water due to nonpolar bonds; examples include fats and oils.
Practical relevance: lipid portions of cell membranes are hydrophobic and influence membrane properties and transport.
Metabolism: Reactions, Energy, and Coupling
Metabolism: sum total of all chemical reactions in the body.
Catabolic reactions (decomposition): break down complex molecules into simpler ones; energy released (exergonic).
Anabolic reactions: build up larger molecules from smaller ones; energy required (endergonic).
Energy in metabolism is often stored as ATP and used to drive anabolic reactions.
Reactions can be reversible; equilibrium depends on body needs.
Exchange reactions: involve swapping components between reactants (e.g., HCl and NaHCO₃ forming NaCl and H₂CO₃).
Energy concepts:
Kinetic energy: energy of motion; related to temperature and molecular movement.
Potential energy: stored energy; in chemistry, chemical energy stored in bonds.
Important terminology:
Exergonic reactions: release energy.
Endergonic (ender) reactions: require/absorb energy.
Interconnection:
Catabolic reactions release energy that is used to drive anabolic reactions; metabolism is a coupled system.
Inorganic vs Organic Compounds
Organic compounds: contain carbon and hydrogen.
Inorganic compounds: do not contain both carbon and hydrogen (or lack one of them).
In biology, three key inorganic compounds are emphasized:
Water (H₂O)
Salts (e.g., NaCl)
Acids and bases and their roles in the body
Note: Organic chemistry is typically built around carbon-containing compounds; this course focuses on inorganic basics first.
Quick Concepts and Practice Points
Atomic number Z defines the element; mass number A defines total protons and neutrons: A = Z + N.
Neutral atoms have equal numbers of protons and electrons.
Ions form when electrons are gained or lost; charges are shown as superscripts (e.g., Na⁺, Cl⁻).
Valence shell electrons determine chemical reactivity; Group number ≈ number of valence electrons (for main-group elements).
Noble gases (Group 8) have full valence shells and are generally nonreactive.
Water’s properties (lubrication, heat capacity, solvent) support life processes and homeostasis.
Hydration shells enable ions to dissolve in water; polarization of water around ions drives solubility.
Bond types summary:
Covalent bonds: shared electrons; can be polar (unequal sharing) or nonpolar (equal sharing).
Ionic bonds: attraction between oppositely charged ions.
Hydrogen bonds: weak attractions between partially positive hydrogen and another electronegative atom; crucial for water and biological interactions.
Practical Implications and Real-World Relevance
Understanding bonding and molecular structure explains why water behaves the way it does in physiology and why life relies on water’s unique properties.
The concepts of ions, electronegativity, and hydration shells underpin nerve signaling, muscle contraction, and osmoregulation in biology.
Energy flow in metabolism (catabolic to anabolic coupling) is fundamental to cellular energetics and ATP production.
Distinctions between hydrophilic and hydrophobic substances influence nutrient transport, membrane structure, and many diagnostic or therapeutic approaches.
Quick Reference Formulas and Key Notations
Atomic/mass relations:
Z = ext{number of protons}
N = ext{number of neutrons}
A = Z + N
Electron configuration basics: valence electron count approximated by group number (for main-group elements).
Water dissolution example: ext{HCl} + ext{NaHCO}3 ightarrow ext{NaCl} + ext{H}2 ext{CO}_3
Water: polar covalent O–H bonds; partial charges: δ− on O, δ+ on H.
Hydration around ions in solution: water molecules orient with O toward cations and H toward anions.
Solvent/solute definitions:
Solvent: substance doing the dissolving (often water).
Solute: substance being dissolved (e.g., NaCl).
Hydrophilic vs hydrophobic:
Hydrophilic: dissolve well in water (polar/charged).
Hydrophobic: do not dissolve in water (nonpolar).