CH339K Lecture 2

Bonding Covalent Ionic Dipole Interactions Van der Waals Forces Hydrogen Bonds

Covalent Bonds Electrons form new orbitals around multiple atomic nuclei Bond energy results from electrostatic force between redefined electron cloud and nuclei Strong – typically 150 – 400 kJ/mol

Ionic Interactions Energy from non-directional force between ions Biomolecules frequently have large numbers of charged groups Charge-charge interactions stabilize intra- and intermolecular structures Coulomb’s Law: Energy drops off as function of distance between charges

Dipoles Fixed dipoles –Molecules with asymmetric charge distributions form dipoles Induced Dipoles –One dipole can induce a charge in an adjacent molecule

van der Waals Interactions Technically, all induced dipole interactions are van der Waals interactions Biochemists usually mean induced dipole-induced dipole (London Dispersion) forces Any atom will have an uneven distribution of charge at any given instant

Van der Waals (cont.) That temporary dipole will induce a dipole in adjacent atoms This results in a net attractive force between atoms Force is weak -.5 to 2 kJ/mol Net biochemical effect – molecules that FIT together STICK together.

Van der Waals (cont.) If you live in Central Texas, you see van der Waals forces in action every summer night:

Hydrogen Bonds Hydrogen Bonds form between –A hydrogen covalently bound to an electronegative atom –Another electronegative atom

Hydrogen Bonds (cont.) The group to which the hydrogen is covalently bound is the donor. The other group is the acceptor. Donors: –-OH, -NH 2, -SH (lesser donor) Acceptors –-N:, =O:, -O:

Hydrogen Bonds (cont.) Hydrogen bonds are not just electrostatic – partially covalent Therefore, they are directional Intermdiate strength: 5 – 10 kJ/mol

Water Structure

Water Forms Clusters in Solution

Hydrophobic Effect

Water –Has a high specific heat –Has a high heat of vaporization –Is an excellent solvent for polar materials –Is a powerful dielectric –Readily forms hydrogen bonds –Has a strong surface tension –Is less dense when it freezes (i.e. ice floats)

Acids and Bases Definitions –Arrhenius –Bronsted-Lowry –Lewis

Conjugate Pairs Every acid has its conjugate base Every base has its conjugate acid Conjugate AcidConjugate Base H 3 C - COOHH 3 C-COO - NH 4 + NH 3

Acids and bases: pH Water ionizes

Typical pH Values SubstancepH Stomach acid Coca-cola2.5 Human saliva6.5 Human blood7.5 Human urine5 - 8 Oven cleaner14

Acids and Bases Water thus acts as both a weak acid and a weak base (A Strong acid is one that dissociates completely in water; a weak acid is one that doesn’t.) All biochemically significant acids and bases are weak (except for HCl – stomach acid)

Acids and Bases Just like water, a weak acid has an ion product, the Ka For the weak acid HA: Therefore

Acids and Bases Ka’s for weak acids range over several orders of magnitude They are generally small More convenient to define pKa = -log Ka Just like pH = -log[H + ]

Typcal Ka’s and pKa’s AcidKapKa Acetic1.8 x Formic1.7 x Benzoic6.5 x Carbonic4.3 x Imidazole2.8 x Phenol1.3 x

pH for Strong Acids Since a strong acid dissociates completely: pH = -log([Acid]) For a 0.1 M (100 mM) solution of HCl, pH = -log(0.1) = 1

pH for Weak Acids What’s the pH of a 100 mM solution of Acetic Acid? [H+] = M

Shortcut The quadratic solution is a pain, but we can approxmate: [H+] = M

Titrating a Strong Acid 10 ml of an HCL sln Titrate with 0.5 M NaOH OH - + H + –> H 2 O Takes 8.5 ml NaOH to bring solution to neutrality

Titrating a Weak Acid Titrating.1 M Hac Initial pH is 2.88 instead of 1 Little change until large amounts of NaOH have been added Buffering effect Caused by equilibrium that exists between a weak acid and conjugate base.

Henderson-Hasselbach Equation

Predicting pH Let’s make 1 liter of a solution that is 0.1 M in acetic acid ( pKa = 4.74 ) and 0.3 M in sodium acetate.

Buffering Effect Addition of significant amounts of acid or base changes the ratio of conjugate base to conjugate acid pH changes as the log of that ratio Result is resistance to pH change in a buffered solution