Chapter 2: Chemistry Comes Alive Notes (BIO 201)

Matter and Elements

  • Matter is anything that has mass and occupies space; in the human body it exists as solid, liquid, and gas.
  • All matter is composed of atoms; an atom is the smallest particle that exhibits the chemical properties of an element.
  • Elements: 92 naturally occurring; hydrogen is the smallest, uranium the largest and heaviest.
  • Periodic table: elements organized by atomic number; living systems categorize elements into major, minor, and trace based on weight percent in the body.
    • Major elements compose ~98% of body weight.
    • Minor elements compose <1%.
    • Most abundant in humans: O, C, H, N.
  • Only 12 elements occur in living organisms in greater than trace amounts:
    • O, C, H, N, Ca, P, S, K, Na, Cl, Mg, Fe.
  • Atoms are composed of three subatomic particles: protons, neutrons, and electrons.
  • Key properties: mass and charge differentiate subatomic particles.
  • Atomic mass unit (amu) expresses atomic mass; neutrons and protons each have mass ~1 amu; neutrons are uncharged, protons are positively charged.
  • Electrons carry negative charge and have ~1/800th the mass of a proton or neutron; located in orbitals around the nucleus (electron cloud or discrete energy levels).
  • The periodic table shows an element’s symbol, atomic number, and average atomic mass; each element has a unique chemical symbol (e.g., C for carbon).
  • Atomic number indicates the number of protons in an atom; average atomic mass reflects the total mass of protons and neutrons in the nucleus.
  • Neutrons can be estimated by: n=A−Zn = A - Z where AA is atomic mass and ZZ is atomic number.
  • In neutral atoms, electrons = protons = atomic number.

Atoms

  • Subatomic particles and masses/charges:
    • Protons: positive charge, mass ~1 amu, located in the nucleus.
    • Neutrons: neutral, mass ~1 amu, located in the nucleus.
    • Electrons: negative charge, mass ~/800 of a proton/neutron; located in orbitals around the nucleus.
  • Nucleus contains protons and neutrons; electrons form an electron cloud around the nucleus.
  • The nucleus defines the element; electrons define chemical behavior.

Electron Shells and Isotopes

  • Electron shells surround the nucleus and have specific energy levels.
    • Innermost shell holds up to 2 electrons.
    • Second shell holds up to 8 electrons.
    • All subsequent shells typically hold up to 8 electrons.
  • Some higher-level shells may hold more than eight electrons (note from slides).
  • Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons; they have essentially identical chemical characteristics but different atomic masses.
    • Example: carbon isotopes C-12, C-13, C-14:
    • Protons: 6 in each.
    • Neutrons: 6, 7, 8 respectively (actual masses: 12, 13, 14).
    • Radioisotopes are unstable isotopes that decay by emitting high-energy radiation (alpha, beta, gamma) to reach stability.
    • Half-life: the time required for 50% of a radioisotope to decay to a stable form.
  • Valence electrons drive chemical behavior and the periodic table organization.

Periodic Table and Valence

  • The periodic table is organized by atomic number (left-to-right) and valence shell electron count (column groups I to VIIIA).
  • Elements in the same column (group) have the same number of valence electrons; e.g., Group IIA elements have 2 valence electrons.
  • The octet rule: atoms tend to fill their valence (outermost) shell to eight electrons to achieve stability.
  • Inert (non-reactive) atoms have completely filled valence shells.
  • Atoms that do not have a filled valence shell tend to lose, gain, or share electrons to achieve stability.

Ions and Ionic Bonds

  • Ions are atoms (or groups) with positive or negative charges due to loss or gain of electrons.
    • Cation: positively charged (loss of electrons).
    • Anion: negatively charged (gain of electrons).
  • In general:
    • Left side of the periodic table tends to lose electrons and form cations.
    • Right side tends to gain electrons and form anions.
  • Many elements do not simply lose or gain electrons; they share electrons to achieve stability (covalent bonding).
  • Ionic bonds form when oppositely charged ions attract electrostatically, creating an ionic compound (salt).
    • Not a molecule (ionic compounds are not discrete molecules).
    • Example: NaCl – Na donates one electron to Cl; Na becomes Na⁺, Cl becomes Cl⁻ (chloride).

Covalent Bonds and Molecules

  • Covalent bonds form when atoms share electrons to satisfy valence needs.
    • Occurs when participating atoms need electrons in their outer shells (4–7 electrons in outer shell).
    • Common in major body elements: O, C, H, N.
  • Covalent bonds can be single, double, or triple, depending on how many electron pairs are shared.
  • Carbon as a key example: outer shell contains 4 electrons; carbon tends to form 4 covalent bonds, enabling carbon-based skeletons.
  • Carbon skeletons can be straight, branched, or ring-shaped.
  • If bonded atoms are the same element (e.g., O2), the molecule is not a compound; otherwise, it is a molecular compound.
  • Structural formulas depict how atoms are arranged in a molecule and help differentiate isomers (same formula, different arrangement).

Polar vs Nonpolar Covalent Bonds; Electronegativity; Amphipathic Molecules

  • Covalent bonds may be nonpolar or polar depending on how equally electrons are shared, driven by electronegativity.
    • Nonpolar covalent: electrons shared equally (similar electronegativity).
    • Polar covalent: electrons shared unequally (different electronegativities).
  • Amphipathic molecules have both polar and nonpolar regions (e.g., phospholipids in membranes).
  • Hydrogen bonds: a weak attraction between a partial positive hydrogen in one molecule and a partial negative atom in another; crucial in biology.
  • Van der Waals forces: momentary, weak attractions from transient dipoles in nonpolar molecules; contribute to intermolecular attractions.
  • Hydrophobic interactions: nonpolar molecules cluster away from water (water is polar).

Intermolecular Attractions

  • Hydrogen bonds are especially important in water and biological macromolecules.
  • Van der Waals forces contribute to the stability of large molecules and molecular interactions.
  • Hydrophobic interactions help drive the folding of macromolecules and the formation of cellular membranes.

Water: Structure, Properties, and Roles

  • Water has several key functions: transport, lubrication, cushioning, and waste excretion; ~two-thirds of body weight.
  • Water is a polar molecule: one oxygen atom bonded to two hydrogens; polarity arises because O is more electronegative than H, pulling electrons toward itself.
  • Each water molecule can form up to four hydrogen bonds with neighboring water molecules.
  • Phases of water depend on temperature: gas (water vapor), liquid (water), solid (ice).
  • Cohesion: water–water attraction via hydrogen bonding; surface tension arises from cohesive forces at the surface.
  • Adhesion: attraction between water molecules and other substances via hydrogen bonding.
  • Water has a high specific heat (high energy required to raise temperature) because hydrogen bonds must be broken.
  • Water has a high heat of vaporization (energy required to convert liquid to gas by breaking hydrogen bonds).
  • Water is the universal solvent:
    • Hydrophilic substances dissolve in water (polar or charged).
    • Hydrophobic substances do not dissolve in water (nonpolar).
    • Amphipathic molecules partially dissolve due to polar regions interacting with water and nonpolar regions avoiding water.
  • Water dissociation and pH:
    • Water can autoionize: extH2extO<br/>ightleftharpoonsextH++extOH−ext{H}_2 ext{O} <br /> ightleftharpoons ext{H}^+ + ext{OH}^-
    • In biology, the hydronium ion is represented as extH3extO+ext{H}_3 ext{O}^+; hydroxide as extOH−ext{OH}^-.
    • Water is neutral (pH = 7) with equal concentrations of H⁺ and OH⁻.
  • Acids and bases:
    • Acid: dissociates in water to produce H⁺ (and an anion) – a proton donor.
    • Base: accepts H⁺ in solution – a proton acceptor.
  • pH and buffers:
    • pH is a measure of the relative concentration of H⁺ in solution: extpH=−<br/>log⁡[extH+]ext{pH} = -<br /> \log [ ext{H}^+] (0–14 scale).
    • Neutralization: returning an acidic or basic solution to pH 7.
    • Buffers: substances that help resist pH changes when excess acid or base is added.
  • Water mixtures:
    • Suspensions: larger-than-100 nm particles do not stay mixed without motion (e.g., blood cells in plasma).
    • Emulsions: a type of suspension where water mixes with a nonpolar liquid (e.g., oil in water).
    • Colloids: protein particles in water (1–100 nm).
    • Solutions: homogeneous mixtures where particles are <1 nm; solutes dissolve in water.
  • Concentration expressions:
    • Mass/volume: grams solute per liter solution.
    • Mass/volume percent: grams solute per 100 mL solution.
    • Molarity (M): moles per liter of solution: M = rac{n}{V} where nn is moles and VV is liters.
    • Molality (m): moles per kilogram of solvent.
    • Osmol/ osmole: number of osmotic active particles in solution; relates to whether a substance dissociates.
    • Osmolarity: osmoles per liter of solution ( ext{Osmolarity} = rac{ ext{osmoles}}{1~ ext{L}}).
    • Osmolality: osmoles per kilogram of water ( ext{Osmolality} = rac{ ext{osmoles}}{1~ ext{kg}}).
  • A mole is the amount of substance containing 6.02imes10236.02 imes 10^{23} entities (Avogadro's number): NA=6.02imes1023N_A = 6.02 imes 10^{23}
    • One mole of any substance has a mass in grams equal to its atomic (or molecular) mass.

Biological Macromolecules (Biomolecules)

  • Four major classes of organic biomolecules: lipids, carbohydrates, nucleic acids, and proteins.
  • General features:
    • Always contain carbon and hydrogen; often oxygen; many also contain nitrogen, phosphorus, and sulfur.
    • Carbon is central to macromolecules; carbon skeletons and hydrocarbons are common; functional groups define chemical behavior.
    • Polymers are formed by repeating subunits (monomers) for carbohydrates, proteins, and nucleic acids; lipids do not form true polymers.
  • Monomer/polymer concepts:
    • Dehydration synthesis (condensation): monomers join via loss of water (—H from one monomer and —OH from another) forming a covalent bond and water as a byproduct.
    • Hydrolysis: polymers split into monomers by adding water, breaking covalent bonds.
  • Lipids:
    • Lipids are diverse, nonpolar or amphipathic, and not formed as polymers.
    • Functions: stored nutrients, cell membranes components, hormones.
    • Four primary classes: triglycerides (neutral fats), phospholipids, steroids, eicosanoids.
    • Triglycerides: glycerol + 3 fatty acids; major energy storage; fatty acids typically 14–20 carbons; unsaturated if they contain double bonds, saturated if no double bonds.
    • Phospholipids: amphipathic with polar phosphate head and nonpolar fatty acid tails; key components of membranes.
    • Steroids: four-ring hydrocarbon structure; include cholesterol, steroid hormones (e.g., testosterone, estrogen, progesterone), and bile acids.
    • Eicosanoids: derived from arachidonic acid; include prostaglandins, thromboxanes, leukotrienes, prostacyclins; involved in signaling.
    • Glycolipids: lipids with carbohydrate attached; involved in membrane functions and cell binding.
    • Fat-soluble vitamins A, E, K.
  • Carbohydrates:
    • Monomer: monosaccharide (e.g., glucose, fructose, galactose).
    • Disaccharides formed by two monosaccharides (e.g., sucrose, lactose, maltose).
    • Polysaccharides: glycogen (animal storage), cellulose (plant fiber).
    • Hexose sugars: six-carbon monosaccharides (glucose, fructose, galactose).
    • Pentose sugars: five-carbon (ribose in RNA, deoxyribose in DNA).
  • Nucleic acids:
    • Store and transmit genetic information; two classes: DNA and RNA.
    • Monomer: nucleotide (nitrogenous base, pentose sugar, phosphate group).
    • Pentose sugars: DNA uses deoxyribose; RNA uses ribose.
    • Nitrogenous bases: pyrimidines (single ring) and purines (double ring):
    • Pyrimidines: adenine (A) and thymine (T) are pyrimidines;
      cytosine (C), uracil (U), and guanine (G) are purines.
    • DNA: double-stranded; bases are A, C, G, T; no uracil.
    • RNA: single-stranded; bases are A, C, G, U; no thymine.
    • Adenosine triphosphate (ATP): nucleotide composed of adenine, ribose, and three phosphate groups; provides cellular energy.
    • NAD+ and FAD: nucleotides that participate in ATP production in mitochondria.
  • Proteins:
    • Polymers of amino acids; many cellular functions: enzymes, defense, transport, support, movement, regulation, storage.
    • Twenty different amino acids; each has an amino group, a carboxyl group, and a variable R group (side chain).
    • Peptide bonds: covalent bonds linking amino acids via dehydration synthesis.
    • Dipeptide (2 amino acids), oligopeptide (3–20), polypeptide (21–199), protein (200+).
    • Some proteins covalently bond to carbohydrates to form glycoproteins (e.g., surface markers on red blood cells).
    • Protein structure is critical to function; structure is hierarchical: primary, secondary, tertiary, and quaternary levels.
    • Chaperone proteins assist in proper protein folding.
    • Intermolecular/intramolecular interactions shaping conformation include: hydrophobic exclusion, hydrogen bonding, ionic bonding, and disulfide bonds.
    • Primary structure: linear sequence of amino acids.
    • Secondary structure: regular patterns like alpha-helix and beta-pleated sheet.
    • Alpha-helix: helical coil; provides elasticity in fibrous proteins (e.g., skin, hair).
    • Beta-pleated sheet: planar sheet; contributes to flexibility in many globular proteins.
    • Tertiary structure: final 3D shape of a completed polypeptide; fibrous vs globular proteins.
    • Globular proteins: enzymes, antibodies, some hormones; compact, often spherical.
    • Fibrous proteins: extended linear molecules (ligaments, tendons, muscle proteins).
    • Quaternary structure: present only in proteins with two or more polypeptide chains (e.g., hemoglobin).
    • Denaturation: disruption of a protein’s conformation due to non-optimal environment (pH, temperature); usually inactivates function.

Additional Key Concepts and Examples

  • Ionic bonds: Na+ (cation) binds to Cl− (anion) to form table salt (NaCl).
  • Covalent bonds: atoms share electrons; count of bonds relates to outer-shell needs (e.g., O often forms 2 bonds; H forms 1; carbon 4).
  • Carbon-based life rationale: carbon’s ability to form four covalent bonds enables diverse organic frameworks (chains, branches, rings).
  • Hydrolysis vs dehydration synthesis: energy-using reactions that build or break polymers.
  • Hydration and solubility:
    • Hydrophilic substances dissolve in water; hydrophobic substances do not.
    • Amphipathic molecules have regions that interact with water and regions that avoid water (e.g., phospholipid bilayers).
Important Formulas and Constants
  • Avogadro number: NA=6.02×1023N_A = 6.02 \times 10^{23}
  • pH of a solution: pH=−log⁡[H+]\text{pH} = -\log [\mathrm{H^+}]
  • A mole: amount of substance containing 6.02×10236.02 \times 10^{23} entities; mass in grams equals the atomic or molecular mass.
  • Osmolarity and osmolality:
    • Osmolarity=osmoles1 L\text{Osmolarity} = \frac{\text{osmoles}}{1~\text{L}}
    • Osmolality=osmoles1 kg\text{Osmolality} = \frac{\text{osmoles}}{1~\text{kg}}

Connections to Foundational Principles and Real-World Relevance

  • Understanding the chemistry of matter underpins all biology: how atoms bond to form molecules and macromolecules that drive cellular structure, energy metabolism, signaling, and genetic information processing.
  • The properties of water and its behavior as a solvent influence protein folding, enzyme activity, and transport in blood and cells.
  • The concept of electronegativity, polarity, and hydrogen bonding explains the behavior of macromolecules and the formation of cell membranes.
  • The four levels of protein structure relate directly to function; misfolding or denaturation under stress conditions can lead to dysfunction and disease.
  • Lipids, carbohydrates, nucleic acids, and proteins form a coordinated network that sustains energy storage, genetic information flow, catalysis, transport, and structure in living organisms.

Practical Implications and Ethical/Philosophical Considerations

  • Many therapeutic strategies target molecular interactions (e.g., enzyme inhibitors, receptor antagonists) by exploiting our understanding of bonds and molecular structure.
  • The concept of isotopes and radioactivity has medical applications (diagnostics, cancer therapy) but requires careful ethical and safety considerations.
  • Understanding pH and buffers informs clinical decisions (e.g., acid-base balance in patients, cardiac and renal function).
  • Recognizing the distinction between hydrophilic and hydrophobic substances guides drug design and delivery, as well as nutrient absorption.
  • Knowledge of macromolecule folding and chaperone-assisted folding underscores the importance of proper cellular environments for protein function and implications for diseases like prion disorders.

Summary for Exam Readiness

  • Matter, elements, and the structure of atoms establish the building blocks of all biomolecules.
  • The periodic table guides predictions about bonding and reactivity through valence electrons and electronegativity.
  • Ions, ionic bonds, and covalent bonds determine how atoms combine to form salts and molecules.
  • Water’s unique properties (polarity, hydrogen bonding, high specific heat, solvent abilities) are central to biochemistry and physiology.
  • Biomolecules (lipids, carbohydrates, nucleic acids, proteins) form the basis of structure and function in cells, with polymers formed by dehydration synthesis and broken by hydrolysis.
  • Protein structure is hierarchical and essential for function; denaturation disrupts activity.
  • A strong grasp of these concepts supports understanding of metabolism, genetics, physiology, and disease mechanisms.