Lecture notes Protein Science
HC 1 Introduction 6-9-2023
Exam 20-10-2023: lectures and research papers discussed in the lectures.
Atoms contain a nucleus and a cloud of electrons, orbiting around that nucleus. Atoms differ
by the number of positively charged protons within the nucleus and an equal number of
negatively charged electrons. Electrons do not occupy circular orbits as shown here for
simplicity. Instead, they are distributed into energy levels or shells, and within shells again
into orbital. Orbitals have complicated shapes that are described by quantum mechanical
wave formulas.
Within molecules, atoms are connected by covalent bonds. A single covalent bond is created
by sharing one pair of electrons. By sharing electrons, atoms strive to get their utmost
orbitals filled. The outermost orbital can accommodate 8 electrons.
In double bonds, four instead of two electrons are shared. This has consequences for the physical
properties of the molecule, e.g. free rotation is impossible around double bonds.
Carbon constitutes the most abundant element in biology (60% of our body’s dry weight). Carbon has
4 electrons in its outer shell, so it has to lose or acquire 4 electrons for outer shell
saturation. The respective bonds show tetrahedral orientation.
Suppose a carbon atom has bound 4 different substituents. In that case,
asymmetry dictates the existence of two stereoisomers, an L and a D form (e.g.
amino acids in proteins are exclusively found in L-configuration).
Carbon has a unique role in the cell because of its ability to form strong covalent
bonds with other carbon atoms.
Hydrogen is the smallest atom in chemistry as well as biology. Hydrogen atoms contain one electron
and one proton.
Nitrogen represents the second to most abundant element in biological dry body mass (11%). With 5
electrons in its outer shell, nitrogen mostly accepts three electrons to form 3 bonds.
Oxygen represents 9% of our body’s dry mass. With two electrons missing in its outer shell, O mostly
accepts 2 electrons to form 2 bonds.
Many C-N and C-O compounds are building blocks, also called functional groups or just functions.
They constitute recurring themes in biochemistry. Shown here is the amine (or amino) group, which
can form amide bonds with acidic groups.
Within a covalent bond, some atoms are more able to attract electrons than others. This is
determined by electronegativity. Electronegativity can invoke polarity, i.e. an uneven charge
distribution within molecules. The magnitude of this effect is dependent on the differences in
electronegativity between the atoms actually participating in a chemical bond. One
prominent example of polarity is H2O.
In extreme cases, differences in electronegativity can even allow the generation of ions,
which may still attract each other in ionic bonds.
When bound to atoms of higher electronegativity, hydrogens can dissociate and cause ion pairs, as
shown here for the carboxyl group in acetic acid. When the released proton subsequently binds to
water, hydronium ions (H3O+) are formed. Amino groups show the opposite tendency, they can
accept protons and thereby become positively charged. In chemistry, protonation (or hydronation) is
,the adding of a proton (or hydron, or hydrogen cation), usually denoted by H+, to an atom, molecule,
or ion, forming a conjugate acid. Deprotonation is the removal/transfer of a proton from an acid in an
acid-base reaction. The concentration of hydronium ions (in mol/l, M) is expressed as pH:
pH = - log [H3O+] [H3O+] = 10-pH M
- Strong acid/base: virtually complete (de)protonation;
- Weak acid/base: equilibrium between protonation and deprotonation.
To quantify the tendency to acquire or lose protons, the pKa value indicates at which pH a compound
is half protonated.
- Ka=¿ ¿ equilibrium dissociation constant.
- pKa=−log10 Ka logarithm, easier to handle. pH value at which 50% of protons is
released. So half is protonated and the other half is deprotonated.
Interactions between chemical groups:
- Noncovalent interactions involving charges:
o Short-range repulsions
o Electrostatic forces – point charges
o Van der Waals interactions
o Hydrogen bonds
- Hydrophobic effect
Short-range repulsions occur upon narrowing the distance between atoms. As
e- orbitals start to overlap, the Pauli principle starts to rule. Energy increases
with the 12th power of distance (treatment as distinct border possible). Van
der Waals radii may vary responding to environmental influence. Atoms are
visualized as spheres with impenetrable volume defined by van der Waals
radius.
Electrostatic forces – point charges and salt bridges
Energy difference for approaching two atoms A and B with Z charge from
infinite distance to distance r (Coulomb’s law). Value of ɛ depends on
environment/medium. Energies of interaction vary from high values in vacuum
down to a few kJ/mol.
Strictly, Coulomb’s law holds true only for point charges. Since ions in proteins are of finite size, the
equation is valid only at distances significantly greater than atomic dimensions. Salt bridges in
proteins occur between ionizable side chains. Due to processes like resonance
and protonation, often components of hydrogen bonding are involved.
- Point charges are an approximation in the Coulombs equation, so it
has no size and doesn’t really exist, whereas dipoles are the two
charges in an bond of a molecule/group, a positive and negative
charge, that is separated by a small difference of the bond.
,Electrostatic forces – dipoles
Different electron attraction by atoms comprising a group or side chain may induce an uneven
distribution of electrons.
µD = Zd
The dipole moment µ equals the product of the magnitude of the separated excess charge Z times
the separation distance d. Dipole moments are vectors with magnitude and direction.
Types of dipoles:
- Permanent dipoles – may exert attractive or repulsive forces, dependent on
spatial arrangement;
- Induced dipoles – emerge by forces originating from neighboring permanent
dipoles (always attractive);
- (London) Dispersion forces – a charge Separation emerging transiently due to
Electron motions in one group may induce a dipole in another group, leading
to attraction. So a spontaneous dipole is formed in a non-polar group which
induces an induced dipole in another group. Other name: van der Waals
interaction.
EXAM QUESTION
Van der Waals interactions
London or dispersion forces lead to van der Waals interactions, with
distance dependence d-6. Since they occur amongst all kinds of
molecules, they are more abundant than other dipolar interactions.
But contribution of single events is small (<1 kJ/mol).
- When the electron clouds overlaps > the atoms will repulse.
- Delta E is the difference in energy/interaction without entropy
influence.
- Attraction (setting free negative energy): interaction of
spontaneous and induced dipoles the two dipoles now
attract each other. (London dispersion forces, last for a short
amount of time)
Hydrogen bonds occur as two electronegative atoms compete for the same hydrogen atom.
Hydrogen can be covalently attached to one atom, but still interact with another one due to its small
size and low electronegativity (high charge). The lengths and strengths of hydrogen bonds depend on
the electronegativities of the participating atoms. Typical lengths vary around 3 Å, typical energies
around 1-10 kJ/mol at r.t. Water and amino acid side chains are often involved in hydrogen bonding.
Hydrophobic effect forces nonpolar solutes to interact with each other instead of the solvent,
excluding them from water. As a result, hydrophobic side chains are forced into the protein’s interior.
This effect comprises a major force ruling in protein folding.
The free energy change combines the changes of Enthalpy and Entropy. Nonpolar molecules prefer
nonpolar environments.
- ΔH: enthalpy change, reflects the difference of energy between all bonds
(noncovalent/bipolar interactions, for example, H-bonds); how much energy do we need to
put in to break the bond? The stronger the bond, the more energy we need to put in.
, - ΔS: entropy change, reflects the difference in disorder/probability of a certain state (water
cage);
o If we don’t do anything, things tend to get disordered.
- ΔG = free enthalpy, in which direction the reaction will occur.
o Protein fold is influenced
hydrophobic effect (less water is
trapped on folded protein surface >
bulk water has more freedom
increasing the entropy) and
conformational entropy (lost during
folding)
- T = temperature
- ΔG = ΔH – TΔS
o Negative ΔG = energy is set free
The second law of thermodynamics: ΔSuniverse = ΔSsystem+ ΔSsurrounding > 0 meaning that the disorder of
the universe always increases.
A cage of water molecules (calatherate, which look more like the organization of water in ice where it
also is fixed) surrounding the non-polar molecule. This cage has more structure than the surrounding
bulk media. To minimize the structure of water the hydrophobic molecules cluster together
minimizing the surface area/relatively smaller cage of water around the non-polar cluster, increasing
the entropy of this state.
The hydrophobic effect is the most important interaction, but it is hard to quantify for single events.
Every interaction contains energy:
- Peptide bond ca 500 kJ/mol
- Disulfide bond ca 330 kJ/mol
- Salt bridge up to 84 kJ/mol
- H-bond 2-10 kJ/mol
- Van der Waals attraction (side chain) ~10 kJ/mol
The amount of atoms/molecules is measured in Mole (mol). 1 mol = 6.022*1023 particles.
The concentration of molecules is measured in Molar (M). 1 M = 1mol/l.
Equilibrium is dependent to ΔG. Small changes in ΔG can have major effects
to the equilibrium (logarithmic dependence). Equilibrium dissociation
constants (KD): frequently used value to quantify interaction. For strong
(effective) interaction, KD should be in nM or lower µM range. Unit: mol/l
(M). The stronger the interaction, the lower the equilibrium dissociation. µM
range block 50% of your ligand.
