Principles of Biochemistry 1-23

The ionization of water, weak acids, and weak bases.

Higher melting point, boiling point, and heat of vaporization compared to most other common solvents.

When the acceptor atom is in line with the covalent bond between the donor atom and hydrogen.

An aqueous solution of a weak acid and its salt.

0.11 nm.

About 0.154 nm (0.077 nm + 0.077 nm).

It exposes the smallest possible hydrophobic surface area to the water, requiring fewer water molecules in the shell of ordered water.

It causes an ordering of water molecules, but the energetic effect is less significant than with nonpolar solutes.

They are unable to undergo energetically favorable interactions with water molecules and interfere with hydrogen bonding among water molecules.

The structure and function of biomolecules.

Hydrogen bonds, ionic interactions, hydrophobic interactions, and van der Waals interactions.

Water replaces solute-solute hydrogen bonds with solute-water hydrogen bonds, screening the electrostatic interactions between solute molecules.

It is disrupted as water molecules cluster about the Cl- and Na+ ions.

Hydrogen bonds are about 10% covalent due to overlaps in the bonding orbitals and about 90% electrostatic.

The highly localized dielectric constant, such as in a hydrophobic pocket of a protein

ΔG must have a negative value.

Their electron clouds begin to repel each other, and at the point where net attraction is maximal, the nuclei are said to be at van der Waals contact distance.

Two phases form; neither liquid is soluble in the other.

1 to 20 picoseconds.

Each hydrogen atom shares an electron pair with the central oxygen atom.

It puts the positive charge of the hydrogen ion directly between the two partial negative charges.

117 °C

Oxygen or nitrogen.

4.6

The increase in entropy (randomness) of the system

ΔG = ΔH − TΔS, where ΔH represents the enthalpy change and ΔS represents the change in randomness.

The formation of lipid bilayers in biological membranes.

It may result in a small gain of enthalpy due to the breaking of hydrogen bonds between water molecules.

Heat is taken up by the system.

The crystal lattice structure of ice makes it less dense than liquid water.

By hydrating and stabilizing the Na+ and Cl- ions, weakening the electrostatic interactions between them.

Conditions that produce blood pH outside the range of 7.3 to 7.5 are life-threatening in humans.

0.018 g/L.

By the properties of the aqueous medium in which life began.

Flickering clusters.

Water weakens the electrostatic interactions between ions, counteracting their tendency to associate in a crystalline lattice.

Noncovalent interactions, influenced by water's properties as a solvent, are responsible for the strength and specificity of recognition among biomolecules.

Because of the stabilizing effect of hydrogen bonds between the hydroxyl groups or carbonyl oxygen of the sugar and the polar water molecules.

In sizes proportional to their van der Waals radii.

10 to 40 nm, depending on the electrolyte concentration

The number of ordered water molecules, and thus the magnitude of the entropy decrease, is proportional to the surface area of the hydrophobic solute.

Weak interatomic attractions that occur when two uncharged atoms are brought very close together, influencing each other's electron clouds.

The arrows represent electric dipoles; there is a partial negative charge (−) at the head of the arrow and a partial positive charge (+) at the tail.

To break a sufficient proportion of hydrogen bonds.

Another hydrogen bond forms, with the same partner or a new one, within 0.1 picoseconds.

The shapes of the outer electron orbitals of the oxygen atom.

They profoundly influence the structure, self-assembly, and properties of all cellular components, including proteins, nucleic acids, and lipids.

The combined effects of hydrogen bonds, ionic interactions, and the hydrophobic effect.

Because there is less screening of charges by the nonpolar solvent.

0.077 nm.

Because their movement from the disordered gas phase into aqueous solution constrains their motion and the motion of water

Because ΔH is positive for melting and evaporation, the increase in entropy (ΔS) makes ΔG negative and drives these changes.

They form structures such as micelles and lipid bilayers.

Carbon dioxide forms carbonic acid (H2CO3) and is transported as the bicarbonate ion (HCO3-), either free or bound to hemoglobin.

The physical and chemical properties of water.

Because ice is less dense than liquid water.

Water readily dissolves charged biomolecules, including compounds with ionized carboxylic acids, protonated amines, and phosphate esters or anhydrides.

Enzymes have evolved to function optimally at near-neutral (physiological) pH.

By one or more equilibrium constants.

When the three atoms involved in the bond lie in a straight line.

The magnitude of the charges (Q), the distance between the charged groups (r), and the dielectric constant (ε) of the solvent.

Stable structures of amphipathic compounds in water, where nonpolar regions cluster together and polar regions interact with water.

Because carbon is only slightly more electronegative than hydrogen, making the C—H bond only very weakly polar.

900 g/L.

A regular crystal lattice.

Four hydrogen bonds, forming a regular lattice structure.

Ionized carboxylic acids (—COO⁻), protonated amines (—NH₃⁺), and phosphate esters or anhydrides.

They show their greatest activity and stability at those pH values.

Hydrogen bonds are represented by three parallel blue lines.

The hydrogen bond is weaker.

Because they are nonpolar and do not dissolve polar biomolecules easily.

Binding of substrate to enzyme releases some of the ordered water molecules.

By clustering together, minimizing the number of ordered water molecules required to surround hydrophobic portions, and increasing the entropy of the system.

By forming new solute-water interactions, resulting in a generally small net change in enthalpy (ΔH).

In an aqueous environment.

Because they can replace water-water interactions with energetically favorable water-solute interactions.

104.5°, slightly less than the 109.5° of a perfect tetrahedron.

Some biologically important hydrogen bonds.

Hydrogen atoms have localized partial positive charges (δ+), and the oxygen atom has a partial negative charge (δ−).

The bond dissociation energy of hydrogen bonds in liquid water is about 23 kJ/mol, compared to 470 kJ/mol for the covalent O—H bond in water or 350 kJ/mol for a covalent C—C bond.

Because it readily dissolves most biomolecules, which are generally charged or polar compounds.

The phenomenon where nonpolar regions of molecules cluster together to minimize hydrophobic area exposed to water, while polar regions maximize their interaction with water.

Amphipathic biomolecules have both polar and nonpolar surface regions. Examples include proteins, pigments, certain vitamins, sterols, and phospholipids of membranes.

To minimize the number of ordered water molecules required to surround the hydrophobic portions, thereby increasing the entropy of the system.

Ammonia (NH3), nitric oxide (NO), and hydrogen sulfide (H2S).

+5.9 kJ/mol

Because they interfere with water-water interactions but are unable to form water-solute interactions.

Because of crowding by the nonbonding orbitals of the oxygen atom.

It helps in holding molecules or groups in a specific geometric arrangement, which is crucial for their function.

By hydrating their component ions.

Hydrogen bonds stabilize interactions between polar biomolecules and water, enhancing solubility.

Because the joined atoms are pulled together by the shared electron pair.

The entropy increases as the highly ordered arrays of water molecules in ice relax into less orderly hydrogen-bonded arrays in liquid water or into the wholly disordered gaseous state.

The hydrophobic effect favors the aggregation of nonpolar regions and is the most important determinant of structure in biological membranes.

Polar molecules dissolve far better even at low temperatures than nonpolar molecules at relatively high temperatures.

+44.0 kJ/mol

3.4 hydrogen bonds.

A nearly tetrahedral arrangement.

Because it has a polar hydroxyl group and can form intermolecular hydrogen bonds.

A hydrogen bond joins two H2O molecules.

78.5

Hydrophilic compounds.

Because the tendency of water molecules to associate through hydrogen bonds is outweighed by the energetic push toward randomness.

Random variations in the positions of electrons around one nucleus may create a transient electric dipole, inducing a transient, opposite electric dipole in a nearby atom.

70% or more.

Each water molecule hydrogen-bonds with an average of 3.4 other water molecules.

Attractions between adjacent water molecules.

To maintain a narrow pH range in which their macromolecules retain their functional structure, which depends on their ionization state.

It is a measure of how close an atom will allow another to approach.

The number of dipoles in the solvent.

They form a highly ordered cagelike shell to maximize solvent-solvent hydrogen bonding.

The entropy gained by freeing immobilized water molecules stabilizes the micelle.

The entropy of the system increases.

The nonpolar nature of these gases and the decrease in entropy when they enter solution make them very poorly soluble in water.

Each water molecule forms four hydrogen bonds in ice.

The electron structure of the H₂O molecule.

When the bonded molecules are oriented to maximize electrostatic interaction, with the hydrogen atom and the two atoms that share it in a straight line.

Most small intermediates of metabolism, as well as nucleic acids and proteins, are soluble in water. Lipid bilayers, the likely forerunners of biological membranes, form spontaneously in water and are stabilized by their interaction with it.

Alcohols, aldehydes, ketones, and compounds containing N—H bonds.

Because it has a high dielectric constant.

There is a measurable decrease in entropy.

Both enzyme and substrate force neighboring water molecules into an ordered shell.

The aggregation of nonpolar amino acids in protein interiors, driven by the hydrophobic effect, stabilizes the three-dimensional structures of proteins.

Hydrogen bonds between water molecules.

Hydrogen bonds.

Because they are capable of holding two hydrogen-bonded molecules or groups in a specific geometric arrangement.

The presence of many intramolecular hydrogen bonds.

The oxygen nucleus attracts electrons more strongly than the hydrogen nucleus (a proton), making oxygen more electronegative.

ΔG is positive, indicating an unfavorable process.

The release of ordered water molecules increases entropy, providing a thermodynamic push toward formation of the enzyme-substrate complex.

Ordered water molecules surround hydrophobic portions of solute molecules, and clustering minimizes the number of these ordered water molecules.

They facilitate the transport of oxygen.