Chemistry Review for Physiology: Atoms, Ions, and Molecules
Matter and the Basis of Biology
Matter is anything that has mass and occupies space.
The three forms of matter we commonly refer to are:
Solid
Liquid
Gas
In physiology, there are many bodily solids (e.g., bone) and liquids (e.g., blood).
Gases in life: oxygen (O₂) and carbon dioxide (CO₂).
We start from the smallest unit that has chemical properties of an element: the atom.
An atom is the smallest particle exhibiting the chemical properties of an element.
Subatomic particles include:
Protons: mass ≈ 1 AMU, charge = +1
Neutrons: mass ≈ 1 AMU, charge = 0
Electrons: mass ≈ 1/1836 AMU, charge = −1
The nucleus contains protons and neutrons; electrons orbit the nucleus in regions called electron orbitals (often simplified as shells).
In brief, atoms form matter and build up to molecules and cells, which we’ll tie back to cellular organization later.
Atoms, Elements, and the Periodic Table
An element is a pure chemical substance consisting of only one type of atom.
There are 92 naturally occurring elements that make up matter, organized in the Periodic Table.
Element symbols can be simple (e.g., H, C, O) or derived from Latin or other languages (e.g., Fe for iron from ferrum; Na for sodium from natrium).
The Periodic Table is organized by:
Electric structure influences: the arrangement of electrons and their shells.
Columns (groups) reflect elements with related properties, particularly valence electron configurations.
Rows (periods) reflect increasing atomic number across the table.
A top-number reference (as discussed in the lecture) is shown above the symbol and relates to trends such as atomic size in this teaching context.
Important features highlighted in class:
Electronegativity tends to increase across a period (left to right) and affects how atoms interact.
The outermost shell (valence shell) largely determines chemical behavior.
The first electron shell holds a maximum of 2 electrons; all subsequent shells hold up to 8 electrons.
Major vs minor elements in the human body:
Major elements collectively make up about 99% of body mass (e.g., carbon-based life).
Minor elements collectively compose less than about 1% of body weight but are essential.
Trace elements exist in even smaller amounts (about 0.01%), but are physiologically important.
How we describe an element:
Chemical symbol (often one letter, sometimes two; sometimes from Latin).
Atomic number Z: number of protons. Shown above the symbol on the periodic table.
Atomic mass (often shown below the symbol): total mass of protons and neutrons.
Mass number A ≈ Z + N (N = number of neutrons).
Example: iron (Fe)
Symbol: Fe (from ferrum, Latin)
Atomic number Z = 26
Atomic mass ≈ 55.845
Mass number A ≈ 56 (rounded for some calculations)
If A = 56 and Z = 26, neutrons N = A − Z = 56 − 26 = 30
In a neutral atom, electrons = protons → 26 electrons when Fe is neutral
Isotopes and Atomic Mass
Isotopes have the same number of protons and electrons (same Z) but different numbers of neutrons.
Prefix iso- means "same"; tope means "place"; so isotopes share the same place in the periodic table (same element) but differ in mass.
Chemical properties are nearly identical because Z (and thus electron configuration) is the same.
Mass differs due to different neutron counts.
Common example: carbon isotopes
^12C, ^13C, and ^14C are naturally existing isotopes.
^14C is radioactive; ^12C and ^13C are stable.
Atomic mass on the periodic table is a weighted average of the isotopes’ masses (
the relative abundances).Practical uses of isotopes in physiology and medicine:
Scintigraphy (radiolabeled isotopes) to image metabolically active tissues (e.g., thyroid activity).
PET scans often use radiolabeled glucose to detect areas with high glucose uptake (cancer cells).
Biological half-life: time required for half of the radioactive material (or other tracer) to be eliminated from the body.
This varies widely: some molecules have very short half-lives (seconds) like mycosanoids; others like insulin can have half-lives of minutes to hours.
Ions and Ionic Compounds
Ions are atoms with a charge produced by loss or gain of one or more electrons.
Neutral atoms have equal numbers of protons and electrons.
If an atom gains electrons, it becomes negatively charged (anion).
If an atom loses electrons, it becomes positively charged (cation).
Anions: negative charge (e.g., Cl⁻).
Cations: positive charge (e.g., Na⁺, Mg²⁺).
Basic examples:
Chlorine gains one electron to form Cl⁻ (charge −1).
Sodium loses one electron to form Na⁺ (charge +1).
Magnesium loses two electrons to form Mg²⁺ (charge +2).
Polyatomic ions: ions that contain more than one atom (e.g., bicarbonate HCO₃⁻, phosphate PO₄³⁻).
Ionic bonds and salts:
Ionic bonds are electrostatic attractions between a cation and an anion, formed by transfer (not sharing) of electrons.
Resulting structures are lattices (crystalline salts).
Example salts:
Sodium chloride: — Na⁺ transfers an electron to Cl⁻, forming a crystalline lattice.
Magnesium chloride: — Mg²⁺ pairs with two Cl⁻ ions.
Molecular Compounds and Covalent Bonds
Molecular compounds are stable associations between two or more elements held together by covalent bonds (shared electrons).
Note: Not all molecules with more than one atom are compounds; for example, O₂ is a molecule but not a compound because it has only one element.
Covalent bonds share electrons between atoms, allowing each atom to fill its valence shell.
Common elements forming covalent bonds in physiology:
Hydrogen (H): needs 1 more electron to fill its valence shell; typically forms 1 covalent bond.
Oxygen (O): needs 2 electrons; typically forms 2 covalent bonds.
Nitrogen (N): needs 3 electrons; typically forms 3 covalent bonds.
Carbon (C): needs 4 electrons; typically forms 4 covalent bonds.
Bond types:
Single covalent bond: one pair of electrons shared (depicted as a single line); e.g., H–H.
Double covalent bond: two pairs shared (depicted as a double line); e.g., O═O (and as in CO₂ where carbon shares two electrons with each oxygen).
Triple covalent bond: three pairs shared (depicted as a triple line); e.g., N≡N.
Carbon skeletons:
Carbon can form straight chains, branched chains, or rings.
Carbon often forms four covalent bonds, resulting in various structures: methane (CH₄), carbon dioxide (CO₂), ethanol (C₂H₅OH), formic acid (HCOOH), etc.
Representation conventions:
Line-angle (bond-line) drawings show bonds as lines; vertices/endpoints represent carbon atoms; implicit hydrogens complete the valence.
Polar vs nonpolar covalent bonds:
Electronegativity is the atoms’ attraction for electrons.
Nonpolar covalent bonds: equal sharing of electrons (approximately equal electronegativities).
Polar covalent bonds: unequal sharing due to differences in electronegativity (e.g., O–H, N–H, C–O).
A molecule may be nonpolar even if it contains polar bonds if the bond dipoles cancel (e.g.,
has two polar C=O bonds arranged linearly such that the dipoles cancel).
Examples of polarity:
Water (H₂O) is a polar molecule due to its bent geometry and polar O–H bonds.
Glucose is polar; it dissolves in water.
Oil and triglycerides are largely nonpolar; oil does not mix with water due to nonpolar–polar incompatibility.
Amphipathic molecules:
Contain both polar and nonpolar regions (e.g., phospholipids with a polar head and nonpolar tails).
Phospholipids form the bilayer of cell membranes, with polar heads facing aqueous environments and nonpolar tails in the interior.
Intermolecular and Intramolecular Forces
Intermolecular attractions: weak attractions between molecules that influence structure, shape, and interactions.
Hydrogen bonding: attraction between a partially positive hydrogen atom and a partially negative atom (commonly O or N) in another molecule; e.g., as in water and glucose.
Temporary dipole (induced dipole) interactions can occur in nonpolar molecules due to momentary unequal electron distribution; these are weaker but cumulatively significant.
Hydrophobic interactions: nonpolar molecules aggregate in polar solvents (e.g., oil droplets in water) due to unfavorable interactions between nonpolar groups and water.
Intermolecular forces contribute to properties like surface tension (e.g., hydrogen bonding in water).
Intramolecular forces: forces within the same molecule that influence its shape and stability (e.g., hydrogen bonding within a protein or nucleic acid can influence folding; strong covalent bonds within a molecule determine its core structure).
The balance of these forces helps explain structure–function relationships in biology (structure equals function).
Isomerism
Isomer: molecules with the same molecular formula but different arrangement of atoms in space.
Why it matters: different structures can yield different properties and reactivity; enzyme specificity often distinguishes between isomers.
Common example in physiology: the hexose family — glucose, galactose, and fructose share the same molecular formula but differ in structure.
Glucose and galactose form similar ring structures, with a key difference at one carbon’s orientation.
Fructose is structurally different (a five-membered ring with additional carbon attachments).
Consequences for biology: enzymes recognize specific isomers; transporters may differentiate between closely related sugars; glucose is a key energy currency, while galactose and fructose are often converted into glucose or used in other pathways.
Practical Implications and Applications in Physiology
The chemistry of life underpins biology: understanding atoms, ions, molecules, and their interactions helps explain cellular processes.
Common physiological ions and molecules (examples):
Sodium Na⁺, potassium K⁺, calcium Ca²⁺, chloride Cl⁻, bicarbonate HCO₃⁻, phosphate PO₄³⁻.
Important covalent molecules: water (H₂O), carbon dioxide (CO₂), organic molecules (carbohydrates, lipids, proteins, nucleic acids).
Medical imaging and lab techniques rely on isotopes and chemical behavior:
Radiolabeled isotopes used in scintigraphy and PET scans.
Biological half-life informs how long a tracer remains active in the body.
Quick reference formulas and concepts:
Mass number:
Neutrons:
Neutral atom electron count:
Ionic charge:
Covalent bonds: single, double, triple (1, 2, or 3 electron pairs shared)
Ionic bond: transfer of electrons; electrostatic attraction; lattice structure
Polar covalent bond: unequal sharing due to electronegativity differences; polar molecules arise from net dipoles
Nonpolar covalent bond: equal sharing; may still have polar bonds that cancel in the molecule (e.g., )
Hydrophobic interactions: aggregation of nonpolar regions in polar solvents
Amphipathic molecules: phospholipids with polar heads and nonpolar tails; form membranes
Quick Cross-References to What’s Next
Monday: macromolecules (proteins, nucleic acids, carbohydrates, lipids) and their chemistry.
Later: cellular level organization and physiology applications of these chemical principles.