Water
In liquid water at room temperature and atmospheric pressure, water molecules are disorganized
and in continuous motion. In ice, on the other hand, each water molecule is fixed in space.
H2Osolid->H2O liquid ΔH= +5.9 kJ/mol
H2O liquid->H2O gas ΔH= +44.0 kJ/mol
During melting or evaporating, the entropy of the aqueous system increases. At room temperature,
both the melting and evaporation occur spontaneously. The free-energy change ΔG must have a
negative value for a process to occur spontaneously. ΔG=ΔH-TΔS, ΔH-enthalpy change from making
and breaking bonds, ΔS-entropy.
H-O-H bond angle is 104.5°, because of the crowding by the nonbonding orbitals of the oxygen atom.
Hydrogen bond ʎ=0.177 nm (23kJ/mol dissociation energy)
Covalent bond (O-H) ʎ=0.0965 nm (470kJ/mol dissociation energy)
Hydrogen bond: nearly tetrahedral arrangement. 1 water molecule can form multiple hydrogen
bonds with other water molecules: 3.4 bonds in water, 4 bonds in ice. Relatively weak bonds. Great
amount of hydrogen bonds responsible for cohesion of liquid water. Strongest when donor-
hydrogen-acceptor are in a straight line. Between the hydroxyl group of an alcohol and water, the
carbonyl group of a ketone and water, peptide groups in polypeptides, complementary bases of DNA.
H-atom covalently bonded to C-atom do not participate in hydrogen bonding, because C is slightly
more electronegative, C-H very weakly polar.
Hydrophilic: easily dissolved in water, polar compounds.
Hydrophobic: difficult to be dissolved in water, nonpolar compounds.
Hydrophobic effect: the nonpolar regions of the molecules cluster together to present the smallest
hydrophobic area to the aqueous solvent, and the polar regions are arranged to maximize their
interaction with the solvent.
Compounds with both hydrophobic and hydrophilic groups are called amphipathic compounds and
can have important biological functions. Many biomolecules are amphipathic: proteins, pigments,
certain vitamins, and the sterols and phospholipids of membranes.
Amphipatic compounds (e.g. lipids) can form micelles.
Van der Waals interactions (also called London forces): Weak interatomic attractions: when two
uncharged atoms are brought very close together, their surrounding electron clouds influence each
other. Single van der Waals force is very weak (4kJ to break it) but MANY of these simultaneously can
have significant impact on protein-protein interactions.
Ionic interactions: Water dissolves salts by hydrating and stabilizing ions.
Noncovalent (weak) interactions among biomolecules in aqueous solvent: Hydrogen bonds; ionic
interactions; hydrophobic interactions; van der Waals interactions.
Equilibrium constants (Keq) describe the position of any chemical reaction.
, Water itself is a slightly ionized (Water is a weak acid or a weak base). In pure water at 25°, [H2O] is
55.5M, and the concentration of H+ and OH- is extremely low, negligible. Kw ( the ion product of
water) is constant, so if [H+] is cery high, [OH-] muste be very low. This is the basis for the pH scale!
[55.5M]*Keq=[H+]*[OH-]=Kw
Keq of pure water at 25° C=1.8*10-16M
Kw=[55.5M]*(1.8*10-16M)=[H+]*[OH-]
Kw=1.0*10-14M2=[H+]2=[OH-]2
[H+]=[OH-]=1.0*10-7M at neutral pH
pH is a denotation of H+ concentration . p simply mens “negative logarithm of”
1
pH =log +¿ pH scale is logarithmic, this means that a difference in a single pH unit
H =−log ¿ ¿
indicates a 10 fold difference in H+ concentration.
The equilibrium constant defines the tendency to lose protons. The stronger the acid is, the higher its
Ka and the lower its pKa are. The stronger the acid, the greater its tendency to lose a proton. The
tendency of any acid (HA) to lose a proton and form its conjugated base (A -) is defined by the
equilibrium constant for the reversible reaction. Equilibrium constant for ionization reactions are
called ionization constants or acid dissociation constants Ka.
1
K eq =¿ ¿=Ka Acid dissociation constant p K a=log =−log [ K a ]
[K a ]
Titration curve: a plot of pH vs the amount of base (NaOH) added until acid is neutralized
pH=pKa [Ac-]=[HAc] HAc↔H+ + Ac- Buffers tend to buffer at their pKa ± 1 pH unit
acetic acid >> dihydrogen phosphate >> ammonium ion The titration curve of a weak acid shows
you that a weak acid and its anion (the conjugated acid-base pair!) can act as buffer.
Buffers: Intacellular and extracellular fluids have a typical and constant pH 7; High concentration of
proteins in cytosol, many amino acid functional groups as weak acids/bases can buffer near pH 7.
Henderson-Hasselbalch equation relates pH, pKa and buffer concentration: 1) Solve Ka equation for H
[ HA ] / 3) Use pH and pK pH=pK + log
2) Take neg log /–log[H+]=-logKa-log a ¿¿
¿¿
Amino acids
One general structure. 20 common amino acids: different side chains. 1806-Asparagine. Amino acids
asre stereoisomers: Ca is the chiral center (4 different groups); enantiomers (mirror images). Only L-
In liquid water at room temperature and atmospheric pressure, water molecules are disorganized
and in continuous motion. In ice, on the other hand, each water molecule is fixed in space.
H2Osolid->H2O liquid ΔH= +5.9 kJ/mol
H2O liquid->H2O gas ΔH= +44.0 kJ/mol
During melting or evaporating, the entropy of the aqueous system increases. At room temperature,
both the melting and evaporation occur spontaneously. The free-energy change ΔG must have a
negative value for a process to occur spontaneously. ΔG=ΔH-TΔS, ΔH-enthalpy change from making
and breaking bonds, ΔS-entropy.
H-O-H bond angle is 104.5°, because of the crowding by the nonbonding orbitals of the oxygen atom.
Hydrogen bond ʎ=0.177 nm (23kJ/mol dissociation energy)
Covalent bond (O-H) ʎ=0.0965 nm (470kJ/mol dissociation energy)
Hydrogen bond: nearly tetrahedral arrangement. 1 water molecule can form multiple hydrogen
bonds with other water molecules: 3.4 bonds in water, 4 bonds in ice. Relatively weak bonds. Great
amount of hydrogen bonds responsible for cohesion of liquid water. Strongest when donor-
hydrogen-acceptor are in a straight line. Between the hydroxyl group of an alcohol and water, the
carbonyl group of a ketone and water, peptide groups in polypeptides, complementary bases of DNA.
H-atom covalently bonded to C-atom do not participate in hydrogen bonding, because C is slightly
more electronegative, C-H very weakly polar.
Hydrophilic: easily dissolved in water, polar compounds.
Hydrophobic: difficult to be dissolved in water, nonpolar compounds.
Hydrophobic effect: the nonpolar regions of the molecules cluster together to present the smallest
hydrophobic area to the aqueous solvent, and the polar regions are arranged to maximize their
interaction with the solvent.
Compounds with both hydrophobic and hydrophilic groups are called amphipathic compounds and
can have important biological functions. Many biomolecules are amphipathic: proteins, pigments,
certain vitamins, and the sterols and phospholipids of membranes.
Amphipatic compounds (e.g. lipids) can form micelles.
Van der Waals interactions (also called London forces): Weak interatomic attractions: when two
uncharged atoms are brought very close together, their surrounding electron clouds influence each
other. Single van der Waals force is very weak (4kJ to break it) but MANY of these simultaneously can
have significant impact on protein-protein interactions.
Ionic interactions: Water dissolves salts by hydrating and stabilizing ions.
Noncovalent (weak) interactions among biomolecules in aqueous solvent: Hydrogen bonds; ionic
interactions; hydrophobic interactions; van der Waals interactions.
Equilibrium constants (Keq) describe the position of any chemical reaction.
, Water itself is a slightly ionized (Water is a weak acid or a weak base). In pure water at 25°, [H2O] is
55.5M, and the concentration of H+ and OH- is extremely low, negligible. Kw ( the ion product of
water) is constant, so if [H+] is cery high, [OH-] muste be very low. This is the basis for the pH scale!
[55.5M]*Keq=[H+]*[OH-]=Kw
Keq of pure water at 25° C=1.8*10-16M
Kw=[55.5M]*(1.8*10-16M)=[H+]*[OH-]
Kw=1.0*10-14M2=[H+]2=[OH-]2
[H+]=[OH-]=1.0*10-7M at neutral pH
pH is a denotation of H+ concentration . p simply mens “negative logarithm of”
1
pH =log +¿ pH scale is logarithmic, this means that a difference in a single pH unit
H =−log ¿ ¿
indicates a 10 fold difference in H+ concentration.
The equilibrium constant defines the tendency to lose protons. The stronger the acid is, the higher its
Ka and the lower its pKa are. The stronger the acid, the greater its tendency to lose a proton. The
tendency of any acid (HA) to lose a proton and form its conjugated base (A -) is defined by the
equilibrium constant for the reversible reaction. Equilibrium constant for ionization reactions are
called ionization constants or acid dissociation constants Ka.
1
K eq =¿ ¿=Ka Acid dissociation constant p K a=log =−log [ K a ]
[K a ]
Titration curve: a plot of pH vs the amount of base (NaOH) added until acid is neutralized
pH=pKa [Ac-]=[HAc] HAc↔H+ + Ac- Buffers tend to buffer at their pKa ± 1 pH unit
acetic acid >> dihydrogen phosphate >> ammonium ion The titration curve of a weak acid shows
you that a weak acid and its anion (the conjugated acid-base pair!) can act as buffer.
Buffers: Intacellular and extracellular fluids have a typical and constant pH 7; High concentration of
proteins in cytosol, many amino acid functional groups as weak acids/bases can buffer near pH 7.
Henderson-Hasselbalch equation relates pH, pKa and buffer concentration: 1) Solve Ka equation for H
[ HA ] / 3) Use pH and pK pH=pK + log
2) Take neg log /–log[H+]=-logKa-log a ¿¿
¿¿
Amino acids
One general structure. 20 common amino acids: different side chains. 1806-Asparagine. Amino acids
asre stereoisomers: Ca is the chiral center (4 different groups); enantiomers (mirror images). Only L-
