Chapter 2: The Chemical Foundation of Life

Atoms, Isotopes, Ions, & Molecules

• The Chemical Foundation of Life explains how atoms and their interactions shape biology.

• Key concepts:

  • Elements: substances that cannot be broken down by chemical reactions.
  • Atoms: the smallest unit of an element that retains its properties.
  • Subatomic structure:
  • Nucleus contains Protons (+) and Neutrons (neutral).
  • Electron cloud/outside region contains Electrons (−).
  • An element is defined by its atomic number Z, the number of protons.
  • Atomic mass is the sum of protons and neutrons; electrons contribute negligibly to mass (~1/2000 Dalton).
  • Mass number A = Z + N (N = number of neutrons).
  • Atomic Mass Unit (amu or Dalton, Da): approximately the mass of a proton or neutron.
  • Isotopes: variants of the same element with different neutrons (same Z, different N).
  • Radioactive isotopes emit particles/energy as they decay.
  • Electrons, energy, and orbitals define chemical behavior.
  • Valence electrons are in the outermost shell and determine an atom’s reactivity and bonding.
  • Periodic trends (brief): as you move across a period, nuclear charge increases, pulling electrons closer and often increasing electronegativity; down a group, atomic radius increases and electronegativity generally decreases.

• Life’s most abundant elements (in living matter): Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N) – together ~96% of living matter.

  • Oxygen ~65%; Carbon ~18-19%; Hydrogen ~9-10%; Nitrogen ~3-4% (approximate values shown in the transcript).
  • Table of major elements in the human body (percent by mass, including water):
  • Oxygen (O) – 65.0%
  • Carbon (C) – 18.5%
  • Hydrogen (H) – 9.5%
  • Nitrogen (N) – 3.3%
  • Calcium (Ca) – 1.5%
  • Phosphorus (P) – 1.0%
  • Potassium (K) – 0.4%
  • Sulfur (S) – 0.3%
  • Sodium (Na) – 0.2%
  • Chlorine (Cl) – 0.2%
  • Magnesium (Mg) – 0.1%
  • Trace elements: B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V, Zn (all <0.01%)

• Subatomic particles (recap):

  • Nucleus: Protons (+) and Neutrons (neutral)
  • Outer region: Electrons (−)

• Defining an Element and basic quantities:

  • Atomic Number Z = number of protons in the nucleus.
  • Atomic Mass A ≈ sum of protons and neutrons; ~1 Dalton per nucleon.
  • Mass Number A = Z + N.
  • 1 Dalton (Da) ≈ 1 atomic mass unit (amu).
  • Electron mass is ~1/2000 of a proton, hence negligible for atomic mass.

• Isotopes:

  • Isotopes differ in neutron number; still the same element (same Z).
  • Radioactive isotopes emit particles and energy, making them useful in biology and medicine (e.g., tracing, dating).

• Bonds and chemical behavior (Overview):

  • The electron configuration and distribution determine chemical behavior.
  • Covalent bonds: sharing of valence electrons to fill outer shells.
  • Ionic bonds: transfer of electrons creating ions that are held together by electrical attraction.
  • Emergent properties depend on the type of bonds and molecular structure.
  • Bond strength varies by bond type; reactivity often arises from unpaired electrons in the valence shell.
  • Molecule vs. Compound:
  • Molecule = two or more atoms covalently bonded.
  • Compound = two or more different elements bonded in a fixed ratio.
  • Structural vs. molecular representations:
  • Molecular formula indicates the number and type of atoms (e.g., ext{H}_2).
  • Structural formula represents shared electron pairs with lines (e.g., H–H for H2; ext{O}= ext{O} for O2).
  • Emergent example: Sodium chloride (NaCl) forms a compound with distinct properties from its elements.

• The Covalent Bond and electron sharing

  • Covalent bonds complete valence shells and produce stable, relatively unreactive molecules.
  • Molecule = 2 or more atoms covalently bonded.
  • Compound = 2+ different elements bonded in a fixed ratio.
  • The way atoms share electrons (single, double, or triple bonds) affects bond strength and molecule geometry.

• Electronegativity and covalent bonds

  • Electronegativity is the attraction of an atom for electrons in a bond.
  • Three main factors:
  • Number of protons in the nucleus: more protons → greater electronegativity.
  • Number of electrons: more electrons in the valence region → greater electronegativity.
  • Distance of outer electrons from the nucleus: greater distance → lower electronegativity.
  • Bond types based on electronegativity differences:
  • Nonpolar covalent: similar electronegativities; electrons shared evenly (e.g., ext{H}2, ext{O}2, ext{CH}_4).
  • Polar covalent: different electronegativities; electrons unequally shared (e.g., ext{H}2 ext{O}, ext{NH}3).
  • Ionic: large electronegativity differences; electron transfer forms ions (e.g., NaCl).

• Ionic bonds, anions, and cations

  • Ionic bonds arise when electrons are transferred to complete valence shells.
  • Ions: atoms with more or fewer electrons than protons, carrying a charge.
  • Anion: negatively charged ion; Cation: positively charged ion.
  • Electrical attraction between opposite charges stabilizes ionic compounds.

• Weak chemical bonds / Intermolecular forces

  • Hydrogen bonds: a hydrogen atom covalently bound to a highly electronegative atom (e.g., O or N) is attracted to another electronegative atom in a different molecule.
  • Common in water and DNA; dynamic and reversible.

• Molecular shape and function

  • Shape is determined by the arrangement of atoms’ valence orbitals.
  • Shape governs how biological molecules recognize and interact with each other.
  • Shapes enable self-assembly and specific binding (e.g., receptors recognizing ligands).
  • Example: Endorphin vs. morphine binding to receptors; structural similarity enables morphine to fit endorphin receptors.

• Carbon and the backbone of life

  • Carbon: essential due to 4 valence electrons and the ability to form up to 4 covalent bonds with different atoms.
  • Carbon forms the backbone of the four major biomolecule classes: proteins, carbohydrates, lipids, nucleic acids.

• Hydrocarbons and organic chemistry in biology

  • Hydrocarbons: nonpolar covalent bonds store energy; not typically found in pure form in living organisms.
  • Often part of larger macromolecules; many biomolecules contain long chains of C and H with occasional heteroatoms.
  • Hydrocarbon rings (aromatics) are common (e.g., benzene rings in some amino acids and cholesterol).
  • Nitrogen can substitute for carbon in heterocyclic structures or side chains.

• Isomers and stereochemistry

  • Structural isomers: same atoms, different covalent connectivity.
  • Geometric (cis/trans) isomers: different arrangement around a double bond.
  • Enantiomers (optical isomers): non-superimposable mirror images around a chiral center.
  • Chiral centers: carbon atoms bonded to four different groups; biological interactions often depend on chirality.

• Functional groups and biochemistry

  • Functional groups replace hydrogens on hydrocarbon chains and confer reactivity and polarity.
  • Major biological functional groups and their typical roles:
  • Hydroxyl (–OH): in alcohols and sugars; polar and increases solubility.
  • Carbonyl (C=O): in carbohydrates (aldehydes in terminals; ketones internal).
  • Carboxyl (–COOH): in fatty acids and amino acids; acidic; polar and hydrophilic.
  • Amino (–NH2): in every amino acid; polar; can accept H+.
  • Sulfhydryl (–SH): in amino acids; can form disulfide bonds; polar.
  • Phosphate (–OPO3^{2-}): in nucleotides and phospholipids; hydrophilic and negatively charged.
  • Methyl (–CH3): nonpolar, hydrophobic; plays a role in gene expression regulation (e.g., methylation).
  • Functional groups enable solubility in water and enable hydrogen bonding, influencing molecular interactions in aqueous environments.

• Carbohydrates, proteins, lipids, nucleic acids (context for functional groups)

  • Carbohydrates: rely on Hydroxyl and Carbonyl groups; monosaccharides contain both; aldehydes or ketones classify sugars as aldoses or ketoses.
  • Proteins: common functional groups include Sulfhydryl and Amino groups; disulfide bonds stabilize protein structure.
  • Lipids and Nucleic Acids: Phosphate and Methyl groups are common; phosphate groups are central to nucleotides and phospholipids; methyl groups influence gene expression and lipid signaling.

• Practical notes on life chemistry

  • Most biological chemistry occurs in aqueous environments.
  • Water’s polarity and hydrogen bonding drive solubility, molecular interactions, and macromolecule structure.
  • Small structural changes (e.g., by adding a functional group) can have large effects on a molecule’s function in water.

• Quick references to key equations and constants

  • Water dissociation product (autoionization):
  • [ ext{H}^+][ ext{OH}^-] = 10^{-14} at 25°C
  • pH definition: pH is the negative base-10 logarithm of the hydrogen ion concentration:
  • ext{pH} = -\, ext{log}_{10}([ ext{H}^+])
  • Relationship between acidity/alkalinity and proton concentration: more \text{H}^+ means lower pH; fewer \text{H}^+ means higher pH.
  • The concentration of hydronium/hydroxide ions is inversely related: increasing one decreases the other.
  • Avogadro’s number: N_A = 6.023 \times 10^{23} \text{ molecules/mol}
  • 1 mole = amount containing N_A = 6.023 \times 10^{23} entities; molar mass (g/mol) is the mass of 1 mole of that substance.
  • 1 mole of carbon has mass: M_{ ext{C}} = 12\ ext{g/mol}
  • Molarity definition: 1\ \text{M} = 1\ \text{mol solute} / \text{L solution}

• Example calculation: preparing a 2 M NaCl solution in 1 L

  • Step 1: Find molar mass of NaCl: M_{ ext{NaCl}} \approx 58.44\ \text{g/mol}
  • Step 2: Moles needed = 2 mol
  • Step 3: Mass required = 2\ \text{mol} \times 58.44\ \text{g/mol} = 116.88\ \text{g}
  • Therefore, dissolve 116.88 g of NaCl in water to make 1 L of a 2 M NaCl solution.

• States and properties of water

  • Water has three states (liquid, solid, gas) with hydrogen bonding patterns changing accordingly:
  • Liquid: hydrogen bonds constantly break and reform; high cohesion and surface tension.
  • Solid (ice): stable crystalline lattice; less dense than liquid, causing ice to float.
  • Gas: hydrogen bonds are largely broken; molecules move freely with high kinetic energy.
  • Emergent properties of water (4):
  • Cohesion (water–water hydrogen bonding) and Adhesion (water to other surfaces)
  • High surface tension and capillary action
  • Moderation of temperature (high heat capacity and high heat of vaporization)
  • Versatility as a solvent (polar molecule interacts with ions and polar solutes)

• Hydrophilic vs. Hydrophobic

  • Hydrophilic: affinity for water; typically ions and polar molecules (e.g., substances with O–H or N–H bonds).
  • Hydrophobic: no affinity for water; nonpolar and noncharged molecules (e.g., fats, oils, O2, CO2).

• Buffers and homeostasis

  • Buffers are weak acid–base pairs that resist pH changes by absorbing or releasing H+ as needed.
  • Common buffering system: bicarbonate (HCO3−)/carbonic acid (H2CO3).
  • Buffers help maintain near-neutral pH in biological systems.

• pH in the human body and fluids (examples from tissue and organ contexts)

  • Saliva ~7.11; Gastric secretions ~1.0–3.5; Pancreatic secretions ~8.0–8.3;
  • Bile ~7.81; Urine ~4.5–8.0; Blood/plasma ~7.35–7.45 (arterial ~7.4, venous ~7.3);
  • Small intestine secretion ~7.5–8.0; Bone ~7.4; Many body fluids hover near 7.4 to support enzyme activity and stability.
  • Note: pH values are context-specific and essential for enzymatic function and chemical equilibria.

• Acids and bases (overview)

  • Acids donate H+ into solution (Arrhenius concept): e.g., HCl → H+ + Cl−
  • Bases donate OH− or accept H+ to reduce H+ concentration, e.g., NaOH → Na+ + OH−; NH3 + H+ → NH4+.

• Dissociation and solubility in water

  • Ionic compounds dissociate in water due to interactions with the polar water molecules, forming hydrated ions.
  • Hydration shells (sphere of hydration) form around ions, stabilizing them in solution.

• Activity connections

  • The structure of water and its ionic equilibria influence the conformation and interactions of proteins, nucleic acids, carbohydrates, and lipids.

• Carbon-based chemistry and functional groups (preview for Part 3)

  • The presence and arrangement of functional groups determine solubility, reactivity, and interactions with other biomolecules in aqueous environments.

• Dopamine and benzene ring (example from carbon section)

  • Dopamine contains a benzene ring (aromatic hydrocarbon) and functions as a neurotransmitter involved in reward-motivated behavior.

• Notes on notation and units used in biology labs

  • Use ext{H}2 ext{O} for water, ext{H}2 for molecular hydrogen, ext{O}_2 for molecular oxygen.
  • Use 1 M = 1 mol/L for solution concentration; remember SI units for mass (g, kg) and volume (L, mL, μL).

• Carbon: the backbone of life

  • Carbon’s tetravalence enables forming diverse and complex organic molecules essential for life.
  • Carbon-based molecules underpin proteins, carbohydrates, lipids, and nucleic acids.
  • Isomerism and stereochemistry add another layer of diversity to biological chemistry.

• Summary takeaways

  • Life depends on the chemistry of atoms, bonds, and molecules in water.
  • Covalent vs ionic bonds create the diversity of molecular interactions.
  • Water’s properties (hydrogen bonding, solvent capabilities, pH) drive biological processes.
  • Carbon’s versatility as a backbone allows the vast complexity of biomolecules and metabolic pathways.