This document provides an overview of key chemical and biochemical principles needed to understand cellular processes at the molecular level, including: - Covalent bonding principles and the most common elements and bonds in cells. - Properties of important biomolecules like water, amino acids, nucleotides, carbohydrates, and lipids. - Noncovalent interactions and how they contribute to molecular recognition and structure. - Chemical equilibrium, dissociation constants, and how pH affects biomolecules. - Free energy and how it relates to the spontaneity of biochemical reactions.

1. Chap. 2. Chemical Foundations
Topics
• Chemical Bonds & Noncovalent Interactions
• Chemical Building Blocks of Cells
• Chemical Equilibrium
• Energetics
Goal
Review chemical and
biochemical principles
essential to the
understanding of cell
processes at the
molecular level.
Fig. 2.1. Key concepts

2. Review of Covalent Bonding
The most abundant elements in cells are H>O>C>N>P>S. The
number of covalent bonds formed by these elements is shown in
Table 2.1. Note that oxygen and nitrogen have unshared pairs
of electrons in bonding orbitals. The most common bonding
orbitals for carbon (sp3, tetrahedral; sp2, trigonal planar) are
shown in Fig. 2.3 (right). When 4 different substituents are
bonded to sp3 carbon, the carbon is asymmetrical. These
carbons are chiral and optically active.

4. Properties of Water Molecules
Water is a polar solvent. It readily dissolves polar and ionic
compounds, but not nonpolar hydrocarbons. Water molecules are
polar because hydrogen and oxygen atoms have substantially
different electronegativities (affinities for electrons) (Fig. 2.5).
Because electrons are shared unequally, the -O-H bonds are
dipolar and partial positive and negative charges occur on H and
O. This feature plus the fact that 2 unshared pairs of electrons
are located at the top of these sp3 hybridized molecules creates
a net dipole moment. Water dipoles interact well with dissolved
polar and ionic solutes.

5. Noncovalent Interactions (Bonding)
Noncovalent interactions are weak electrical bonds between
molecules. Types of noncovalent interactions are 1) ionic
(electrostatic) bonds, 2) H-bonds, and 3) van der Waals
interactions. Noncovalent interactions (1-5 kcal/mol) are typically
~100-fold weaker than covalent bonds (Fig. 2.6). Their stability
is only slightly greater than thermal energy in biological systems.
Nonetheless, noncovalent interactions play important roles in
protein and nucleic acid stabilization because they are "collectively
strong." Note that the hydrophobic effect drives molecular
interactions, but is not a noncovalent bond per se.

6. Ionic Interactions
Ionic interactions occur between cations and anions. These bonds
are non-directional, and strength depends on the distance of
separation (r) according to 1/r2. Strength also depends on the
medium (dielectric constant), and is less in polar than nonpolar
solvents. Ionic compounds such as NaCl are readily dissolved n
water (Fig. 2.7). Solvation spheres of water molecules surround
ions in solutions. Water molecules orient so that the negative
ends of their dipoles contact cations and the positive ends
contact anions in solution.

7. H-bonds
H-bonds are noncovalent interactions occurring between the H
atom of a dipolar molecule such as water, and the unshared
electron pair of another atom (i.e., O or N). These bonds
represent the primary way in which water molecules interact
with themselves and many types of biomolecules (Fig. 2.8). H-
bonds are highly directional in that strength depends on the
proper alignment of the interacting atoms. Directionality
confers bonding specificity as with the Watson-Crick H-bonds
between the bases of double helical DNA.

8. van der Waals Interactions
van der Waals interactions are bonds between fluctuating,
induced dipoles within the electron clouds of interacting
molecules. These bonds can occur between nonpolar or polar
molecules. van der Waals bonds are extremely dependent on
the distance of separation between molecules, and are
significant only when the electron clouds of the molecules are
just touching. van der Waals interactions are demonstrated
for two O2 molecules in Fig. 2.10. The covalent and van der
Waals radii are shown.

9. The Hydrophobic Effect
The hydrophobic effect refers to the entropy-driven aggregation
of nonpolar molecules in aqueous solution that occurs to minimize
the ordering of water molecules with which they are in contact.
This is not an attractive force, but rather a thermodynamically
driven process. Fig. 2.11
shows the cage-like
structures formed by
water molecules
surrounding a nonpolar
solute. The hydrophobic
effect drives the
formation of membranes
and contributes to the
folding of proteins and
the formation of double
helical DNA.

10. Molecular Complementarity
Structurally complementary molecules can bind to each other
(Fig. 2.12). Binding often is mediated solely by noncovalent
interactions. The extent of molecular complementarity sets
binding affinity (how
strongly molecules
interact) and
specificity (which
molecules interact).
Affinity and
specificity are very
important for
molecular function.
Intramolecular
complementarity
establishes the 3D
shapes of proteins,
etc.

11. Intro to Biological Macromolecules
As shown in Fig. 2.13, biological macromolecules are polymers
that are assembled from "building block" monomer units. In all
cases, polymerization occurs via condensation reactions in which
water is eliminated.

12. Eukaryotic Cell Membranes
Eukaryotic cell membranes are composed of molecules such as
phospholipids (Fig. 2.13), cholesterol, and proteins. Phospholipid
bilayers are noncovalent assemblies of phospholipids. The
hydrophobic fatty acyl chains of phospholipids aggregate together
in bilayers. The two layers of a bilayer are called leaflets.
Bilayers are mostly impermeable to polar and ionic molecules, but
allow nonpolar compounds to pass through. Biomembranes typically
contain other lipids such as sterols and glycolipids, and proteins.

13. Amino Acids
There are 20 "standard" amino acids that are specified by the
genetic code and polymerized into proteins by ribosomal
translation. Amino acids contain an a-carbon, to which typically
4 different substituent groups are attached (Fig. 2.4). These
groups are the a-amino group, the a-carboxyl group, hydrogen,
and the variable R-group (side-chain). The a-amino and a-
carboxyl groups are charged at neutral pH.
There are two possible
configurations for these groups-
-the "D" and "L" stereoisomers,
which are mirror images of each
other (enantiomers). The
standard amino acids have the
L-configuration. Amino acids are
classified based on the
characteristics of their R-groups
(Fig. 2.14, next three slides).

14. The Hydrophobic Amino Acids
With the exception of
tyrosine, these amino
acids lack polar
functional groups within
their side-chains.
Hydrophobic amino
acids often are found in
the interior of proteins.
The hydroxyl group of
tyrosine residues
exposed at the protein
surface can be
phosphorylated in
receptors, etc.

15. The Hydrophilic Amino Acids
These amino acids contain charged or H-bonding functional
groups within their side-chains. For this reason, they often
are found at the surface of proteins. The R-groups of lysine
and arginine are positively charged at neutral pH. The R-
groups of aspartate and glutamate are negatively charged. The
histidine R-group may or may not be positively charged
depending on its environment. Serine and threonine can be
covalently modified by phosphorylation or glycosylation.

16. Other Amino Acids
These amino acids carry out special functions in proteins. Glycine
and proline often are found in "turns" (bends) within proteins.
Cysteine can be modified with lipid groups or covalently bound to
another cysteine in a "disulfide bond". Disulfide bonds (bottom
right) contribute to protein folding stability.
Disulfide bond

17. Properties of Nucleotides
Nucleotides are the building blocks of nucleic acids (DNA &
RNA). Nucleotides are composed of a base, a sugar, and one or
more phosphates (Fig. 2.16a). A nucleoside contains just the
base plus the sugar. The sugar is ribose in RNA and 2-
deoxyribose in DNA (Fig. 2.16b). The most common bases are
the purines, adenine and guanine, and the pyrimidines, uracil,
thymine, and cytosine (Fig. 2.17). The naming of nucleosides and
nucleotides is described in Table 2.3.

18. Properties of Carbohydrates
The most common monosaccharides in cells contain 5 (pentoses) or
6 (hexoses) carbons. Monosaccharides are classified as aldoses or
ketoses, depending on whether they contain an aldehyde or
ketone group. Pentoses and hexoses typically form rings, as
illustrated for the most common monosaccharide, glucose (Fig.
2.18a). Aldohexoses occur in 16 stereoisomeric forms. Dimers of
monosaccharides are called disaccharides, and polymers are called
polysaccharides. Common disaccharides are lactose (galactose-
glucose) and sucrose (glucose-fructose). Common polysaccharides
formed from glucose are glycogen in animals, and starches and
cellulose in plants. The linkages between monosaccharide units are
called glycosidic bonds.

19. Properties of Fatty Acids
Fatty acids contain a polar carboxylic acid group and a
hydrocarbon tail of variable length (14-20 carbons in animals)
(Table 2.4). Fatty acids are further classified as saturated or
unsaturated depending on their content of double bonds (usually
cis in animals). cis double bonds introduce kinks into the tail (see
Fig. 2.21), perturb packing interactions, and increase the fluidity
of membrane lipids. The carboxyl group is joined to glycerol via
ester (acyl) linkages in triacylglycerols and phosphoglycerides.

20. Phospholipids
Phospholipids are the most
prevalent components within
biomembranes. In
phosphoglycerides, two fatty
acids are linked to glycerol at
carbons-1 and -2. A phosphate
and polar group (such as choline)
are linked to carbon-3 (Fig.
2.20). The phosphate and polar
group constitute the "head
group.” The head groups of
phospholipids commonly found in
eukaryotic cells are shown in
Table 2.5.

21. Equilibrium Constants
A reaction, aA  bB, is at equilibrium when the rate of the
forward (vf = kf[A]a) and reverse (vr = kr[B]b) reactions are the
same. The equilibrium constant, Keq, is defined as
Keq = kf/kr = [B]b/[A]a
Note that the steady-state concentrations and ratios of the
components of a reaction occurring within a cell often are very
different from the concentrations and ratios of the components
when the reaction is at equilibrium (Fig. 2.23).
X C

22. Dissociation Constants
Many proteins bind to molecules
called ligands. Ligands can be
small molecules, proteins, or
other biopolymers, such as DNA
or polysaccharides (Fig. 2.24).
Dissociation constants (Kds) are
used to measure the affinity of
protein-protein and protein-
DNA interactions, for example.
Kds are equivalent to the inverse
of the Keq for a reaction, as
illustrated in the following
equation.
Protein (P) + DNA (D)  Protein-DNA (P.D)
Kd = [P] [D] / [P.D]

23. Chap. 2 Meaning of the Kd
Example: Protein (P) binding to DNA (D)
P + D  P.D
Kd = [P][D]/[P.D]
What is the ratio of [D]/[P.D] for different values of [P]?
[P] = 0.1 x Kd [D]/[P.D] = 10/1
[P] = Kd [D]/[P.D] = 1/1
[P] = 10 x Kd [D]/[P.D] = 1/10
This shows that the DNA binding site is about 10% occupied when the
concentration of [P] is 10-fold lower than the Kd, 50% occupied when
[P] is the same as the Kd, and 90% occupied when [P] is 10-fold
greater than the Kd.

24. Solution pH
Water has a small but measurable tendency to ionize
H2O  H+ + OH-
In pure water, [H+].[OH-] = 1 x 10-14 M2. Thus, in a neutral
solution [H+] = [OH-] = 1 x 10-7 M.
pH is defined as -log [H+].
Thus, pH = 7.0 for a neutral
solution. Acidic solutions have
pH < 7.0 and basic solutions
have pH > 7.0. Remember, [H+]
changes by 10-fold for a 1 unit
change in pH. pH affects the
properties of nearly all
biomolecules. The pHs of a
number of solutions is shown in
Fig. 2.25.

25. Ionization of Weak Acids
Biomolecules containing carboxylic acid groups are weak acids.
Unlike strong acids (e.g., HCl), weak acids are << 100%
dissociated in water. The dissociation of a weak acid is written as
HA  H+ + A-
for which the acid dissociation constant is
Ka = Keq = [H+] [A-] / [HA]
Henderson-Hasselbach Equation
Another form for the equilibrium equation is
pH = pKa + log [A-]/[HA]
Examination of this equation shows that
[A-]/[HA] = 10/1 when pH = pKa + 1
[A-]/[HA] = 1/1 when pH = pKa
[A-][HA] = 1/10 when pH = pKa - 1

26. Relationship Between pH, pKa, and Acid
Dissociation
The percentages of the conjugate acid (H2CO3) and conjugate
base (HCO3
-) forms of carbonic acid are plotted as a function of
pH in Fig. 2.26. At pH extremes, one or the other species
makes up essentially 100% of the solution. At the pKa, a 50/50
ratio of the two forms is present. The same type of information
can be derived from a titration curve (next slide).

27. Titration of Acetic Acid
In a titration, the conjugate acid form (HA) of a weak acid is
stoichiometrically converted to its conjugate base form (A-) by
the addition of a strong base. The titration curve for acetic
acid (Fig. 2.27) can be used to explain the HH equation. The
ratios of [A_]/[HA] are indicated as a function of pH. The
optimum buffering range of acetic acid is +/- 1 pH unit from
the pKa.
[A-]/[HA] = 1/10
[A-]/[HA] = 1/1
[A-]/[HA] = 10/1

28. Behavior of H3PO4, a Triprotic Weak Acid
H3PO4 has 3 dissociable protons, 3 pKas, and 3 plateaus on its
titration curve (Fig. 2.28). The midpoint of each plateau occurs
where the pH is equivalent to one of the pKas. The cell cytosol is
buffered mostly by phosphate, namely the H2PO4
-/HPO4
2- conjugate
acid/conjugate base pair. The ratio of these species is 1/1 at pKa2.
pKa2

29. Free Energy Changes for Reactions
The free energy of a system or chemical compound is denoted
by "G". The free energy change (∆G) for a reaction (Reactants
 Products) is calculated using the equation
∆G = GProducts - GReactants
If ∆G < 0, the reaction will go to the right.
If ∆G = 0, the reaction is at equilibrium, and goes neither
direction in a net sense.
If ∆G > 0, the reaction will go to the left.
Note that the term ∆G0' is used for reactions occurring under
standard biochemical conditions, where T = 298˚K, P = 1 atm,
and the starting concentrations of all components (other than
H+) is 1 M. For [H+], pH = 7.0. Units of G are kcal/mol.

30. ∆Gs for Reactions (cont.)
Gibbs free energy equation
The ∆G for a reaction is determined by the enthalpy change (∆H)
and entropy change (∆S) for the reaction.
∆G = ∆H - T ∆S
∆H reflects changes in the chemical bonds for all molecules in the
reaction. ∆S reflects changes in the entropy (randomness) of all
components participating in the reaction.
If ∆G < 0, the reaction is
exergonic, and will proceed
spontaneously to the right. If
∆G > 0, the reaction is
endergonic and will not proceed
to the right without an
investment of energy. ∆G is
always negative when ∆H < 0
and ∆S > 0. Conceptual views of
exergonic and endergonic
reactions are shown in Fig. 2.29.

31. Calculation of ∆G Under Any Condition
The calculation of ∆Gs for reactions under any concentration
conditions (e.g., inside cells) is performed using the equation
∆G = ∆G0' + 2.303 RT log Q
where Q = [products]/[reactants] (the mass action ratio) inside
the cell, etc. In other words, ∆G is equivalent to the standard
free energy change (∆G0') plus energy (positive or negative)
resulting from differences in the concentrations of components
from 1 M (standard conditions). Depending on the magnitude of
Q, a reaction can proceed in the opposite direction from that
under standard conditions. Note that R = 1.987 cal/˚K-mol.

32. Relationship Between ∆G0' and Keq
When a reaction is at equilibrium, substitution of ∆G = 0 into
the previous equation shows that
∆G0' = - 2.303 R T log Keq
(Here Q = Keq = [products]/[reactants]). This equation is the
same as
Keq = 10 -∆G0'/2.303RT
An inspection of the equations shows that
∆G0' < 0 when Keq > 1
∆G0' = 0 when Keq = 1
∆G0' > 0 when Keq < 1

33. Energy Coupling in Biological Reactions
Endergonic reactions of biochemistry very often are driven
forward by coupling them to the hydrolysis of ATP. As long as
∆G1 + ∆G2 = ∆Gsum < 0
the reaction will go forward.
Hydrolysis of either of the two
phosphoanhydride bonds in ATP
(Fig. 2.31) releases -7.3
kcal/mol of energy, which is
more than enough energy to
drive most endergonic reactions
forward. Hydrolysis of the
phosphoester linkage between
the a-phosphate and ribose
releases only -2.2 kcal/mol, and
typically is not used for energy
coupling purposes.
a
b
g

34. Mechanisms of Energy Coupling
Assume that ∆G > 0 for reaction B + C  D. Very often such
a reaction can be driven forward via a high energy
intermediate created via ATP phosphorylation.
B + ATP  B-p + ADP
B-p + C  D + p
B + C + ATP  D + ADP + p
B-p serves as a common intermediate in the two reactions
which go forward because ∆Gsum < 0. ATP (or GTP) hydrolysis
also can provide energy for reactions via the creation of a high
energy conformational state in a protein (e.g., myosin) carrying
out an endergonic reaction.