1.1 – Chemical Fundamentals Elements on the Periodic Table

1. 1.1 – Chemical Fundamentals • Elements on the Periodic Table • isotopes – different atomic mass (i.e. # of neutrons) but same atomic number (i.e. # of protons) • radioisotopes - result from nuclear decay, releasing subatomic particles and radiation • C-14 Dating - C-12:C-14 ratio changes upon death at a predictable rate • radioisotope tracers - follow chemicals through reactions, used in nuclear medicine (i.e. I-131 and thyroid) • Chemical Bonding • orbitals can hold 2 e-, have similar shape at each energy level • e- in an orbital pair up for stability • generally, e-max = 2n2, where n = shell number • n=1 spherical 1s (max 2 e- pair) • n=2  larger spherical 2s + 3 dumbbell-shaped 2p orbitals (max 4 e- pr) • n=3  3s, 3p, 3d (max 9 e- pair) • outermost s + p are valence orbitals  families (columns) •  period (rows)

2. Electronegativity • measure of how strongly an atom attracts electrons of a covalent bond • polar covalent bond – unequally-shared e- results in one atom with a partial negative charge and one with a partial positive charge (e- spends more time around one atom than the other) • If ΔEn = 0nonpolar covalent • If 0 < ΔEn < 1.7  polar covalent • If ΔEn ≥ 1.7  ionic

3. Molecular Shape • function follows form (orientation of bonding e- pair, atom type, and bond type) • formation of covalent bonds causes valence orbits to change orientation (= hybridization)  Valence Shell Electron Pair Repulsion (VSEPR) Theory • valence e- pairs repel each other, will move away as far as possible(see Table 5, P. 15) • most biologically important molecules have up to 4 e- pairs around a central nucleus (non-bonding pairs take up more space, compressing bond angles) • Molecular Polarity • dependent on bond polarity and molecular shape • symmetrical shapes produce non-polar molecules (polarity of bonds has no effect) • asymmetrical shapes produce polar molecules (if at least one bond is non-polar)

4. Water • polar covalent bonds in a symmetrical structure = highly polar molecule • bonds readily with other molecules and ions • intermolecular bonds determine state at given T and P • 3 types of intermolecular bonds: • London dispersion forces (weakest, ubiquitous) are formed by temporary unequal distribution of e- • dipole-dipole forces hold polar molecules to one another (partial +ve, partial –ve) • hydrogen bonds (strongest) form between +ve H and –ve N,O,F • angular shape and H-bonds give H2O its unique characteristics • universal solvent (polar, allows formation of H-bonds) • small, non-polar molecules (O2, CO2) dissolve only slightly as no H-bonds can form (hence the need for hemoglobin as O2, CO2 carriers) • large, non-polar molecules (i.e. lipids) are insoluble in H2O (hydrophobic vs. polar hydrophilic) • “like dissolves like” • polar substances dissolve in other polar substances • non-polar substances dissolve in other non-polar substances

5. Types of Biochemical Reactions • Reduction-oxidation ( aka redox, in which an electron is lost by one reactant [oxidation] and gained by another reactant [reduction], resulting in a transfer of electrons) • Hydrolysis (catabolic reaction in which water molecule is used to break a covalent bond between subunits of a macromolecule) • Condensation (aka dehydration synthesis, an anabolic reaction in which a covalent bond is formed between two subunits by the removal of water to form a macromolecule) • Neutralization (in which an acid and a base combine to form water and a salt)

6. Acids, Bases, and Buffers • ionization of water = transfer of H+ from one H2O molecule to another  H3O+ + OH- (hydronium and hydroxide ion) •  H3O+ gives acidic solutions their characteristic properties • Acids: increase the concentration of H3O+(aq) in solution • contain ≥ 1 ionizable H-atom • e.g. HCl + H2O  H3O+(aq) + Cl-(aq) • concentration of any solute in aqueous solution is expressed in mol/L or mol•L-1 (1 mole = 6.02 x 1023 entities) • for pure water at 25oC: H3O+ concentration [H3O+(aq)] = 1.0 x 10-7 mol/L • pH = -log10 [H3O+(aq)]  ph of pure water = -log10 [10-7 mol/L] = 7 • Strong and Weak Acids and Bases • strength of acids and bases is determined by the degree to which they ionize in water • strong acids and bases ionize completely: • HCl + H2O  H3O+ + Cl- (100%) • NaOH  Na+ + OH- (100%) • weak acids and bases ionize partially: • CH3COOH + H2O ↔ H3O+ + CH3COO- (1.3%) • NH3 + H2O ↔ NH4+ + OH- (10%)

7. reverse reaction occurs more frequently as ion concentration increases • when 1.3% of acetic acid molecules are ionized, the number of forward and reverse reactions are equal (equilibrium) • most organic acids and bases are weak and reach equilibrium Bases OH- gives basic solutions their characteristic properties bases increase the concentration of OH- in solution directly (i.e. NaOH(s)  Na+(aq) + OH-(aq)) indirectly (i.e. NH3 + H2O  NH4+ + OH-) • in a neutralization reaction, equal number of H3O+ and OH- form water • H3O+ + OH-  2H2O or H+ + OH-  H2O • the reaction of an acid and a base produces water and a salt • e.g. HCl + NaOH  H2O + NaCl

8. Conjugate Acids and Bases • reversible acid-base reactions involve the transfer of a proton (H+) • acids are proton donors and bases are proton acceptors In an acetic acid solution: • the forward reaction involves a proton transfer from acetic acid to water •  acetic acid is the acid (donor), water is the base (acceptor) • the reverse reaction involves a proton transfer from a hydronium ion to an acetate ion •  the hydronium ion is the acid (donor), the acetate ion is the base (acceptor) • acetic acid-acetate ion, and hydronium ion-water are conjugate acid-base pairs

9. Acid-Base Buffers • usually conjugate acid-base pairs in equilibrium that resist changes in pH e.g. extra H+ in blood are accepted by the HCO3- to form H2CO3 if a base in the blood removes H+,H2CO3 ionizes to replace H+