Study Notes on Biochemistry Concepts

Partial Charges and Hydrogen Bonds

• Hydrogen does not form hydrogen bonds with carbon.

• Full charges arise from the loss or gain of electrons, leading to ionic bonds.

Van der Waals Forces

• Caused by transient charges due to electron distribution.

• Occur in all molecules; stronger in polar, weaker in non-polar.

• Gecko climbing example illustrates Van der Waals forces in action.

Hydrophobic Interactions

• Nonpolar molecules cluster in water, leading to energy stability and increased entropy.

• Separates oil and water, promoting lower energy configurations.

Amphipathic Molecules

• Have both hydrophilic and hydrophobic properties; interactions lead to protein folding and cellular structures.

Functional Groups

• Important functional groups: hydrocarbons, alcohols, aldehydes, ketones, carboxyls, amines, and phosphates.

• Alcohols, amines, and certain groups can form hydrogen bonds.

pH and Biological Importance

• pH measures hydrogen ion concentration; crucial for enzyme activity and protein folding.

• Blood pH around 7.4 is vital for health.

Autoionization of Water

• Water can ionize to form H+ and OH−; equilibrium is key for maintaining proper pH.

Acids and Bases

• Strong acids completely ionize in water; weak acids partially ionize to form conjugate bases.

• pKa is the pH at which the concentrations of acid and conjugate base are equal.

Buffer Systems

• Buffers resist pH changes within certain limits; works best near pKa values.

• Example: carbonic acid and bicarbonate in blood maintain pH.

Partial Charges and Hydrogen Bonds

• Partial charges ($\delta+$ and $\delta-$) are created when electrons are shared unequally between atoms with different electronegativity values.

• Hydrogen bonds are a specific type of dipole-dipole interaction involving a hydrogen atom covalently bonded to a highly electronegative atom ($N$, $O$, or $F$) which is then attracted to the lone pair of another electronegative atom.

• Hydrogen does not form hydrogen bonds with carbon because the electronegativity difference between them ($C \approx 2.5$, $H \approx 2.1$) is too small to polarize the bond significantly.

• Full charges arise from the complete gain or loss of electrons ($e^-$), leading to ionic bonds and the formation of distinct cations and anions.

Van der Waals Forces

• These represent weak, short-range intermolecular attractions caused by transient (temporary) dipoles as electron distributions fluctuate around nuclei.

• They occur in all molecules, whether polar or non-polar; while weak individually, their cumulative effect in large numbers can be substantial.

• Example: Geckos utilize Van der Waals forces between the millions of microscopic hairs (setae) on their feet and the molecules of a surface to walk on walls and ceilings.

Hydrophobic Interactions

• Nonpolar molecules (hydrophobes) do not interact favorably with water and tend to cluster together to minimize their surface area in contact with the polar solvent.

• This clustering is driven by thermodynamics, specifically an increase in global entropy ($\Delta S$). When nonpolar molecules aggregate, fewer water molecules are forced into rigid, "orderly" cage-like structures (clathrates) around them.

• This effect is responsible for the separation of oil and water and is a primary driver for energy-stable configurations in biological systems.

Amphipathic Molecules

• These molecules possess both a hydrophilic (polar/charged) "head" and a hydrophobic (nonpolar) "tail."

• In water, they spontaneously assemble into structures such as micelles or lipid bilayers, which form the basis of all cellular membranes.

• Amphipathic interactions are critical in protein folding, as they drive hydrophobic amino acids to the interior of the protein away from water.

Functional Groups

• Functional groups determine the chemical properties and reactivity of organic molecules. Key groups include:

  • Hydrocarbons: $-CH_3$ (methyl), which are non-polar and hydrophobic.

  • Alcohols: $-OH$ (hydroxyl), which are polar and capable of forming hydrogen bonds.

  • Carbonyls: Found in aldehydes ($-CHO$) and ketones ($C=O$).

  • Carboxyls: $-COOH$, which can release a proton ($H^+$) to act as an acid.

  • Amines: $-NH_2$, which can accept a proton ($H^+$) to act as a base.

  • Phosphates: $-PO_4^{3-}$, which are highly charged and central to energy transfer molecules like ATP.

pH and Biological Importance

• pH is a logarithmic scale used to specify the acidity or basicity of an aqueous solution: $pH = -\log_{10}[H^+]$.

• The scale typically ranges from $0$ to $14$. Most biological processes occur near neutrality ($pH \approx 7$).

• Human Blood: Maintenance of blood pH around $7.4$ is critical; even small deviations can lead to medical emergencies as enzymes become denatured.

Autoionization of Water

• Water molecules occasionally dissociate into hydrogen and hydroxide ions: $H_2O \rightleftharpoons H^+ + OH^-$.

• The equilibrium constant for this reaction at $25^\circ C$ is the ion product constant, $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$.

• In pure water, $[H^+] = [OH^-] = 10^{-7} M$, resulting in a neutral $pH$ of $7.0$.

Acids and Bases

• Strong Acids: Dissociate completely in water (e.g., $HCl \rightarrow H^+ + Cl^-$).

• Weak Acids: Only partially ionize in water, establishing an equilibrium between the acid ($HA$) and its conjugate base ($A^-$).

• pKa: This value is the negative log of the acid dissociation constant ($pK<em>a = -\log</em>{10} K_a$). It represents the pH at which the concentrations of the acid and conjugate base are exactly equal ($[HA] = [A^-]$).

Buffer Systems

• Buffers are solutions that resist changes in pH when small amounts of acid or base are added; they work best when the pH is close to the compound's $pK_a$.

• The Henderson-Hasselbalch equation defines this relationship: $pH = pK<em>a + \log</em>{10}\left(\frac{[A^-]}{[HA]}\right)$.

• Biological Buffer: The bicarbonate system ($H<em>2CO</em>3 / HCO_3^-$) in human blood is a key physiological buffer that prevents rapid pH changes.