1014NSC Biochemistry Complete Lecture Notes PDF

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Fundamentals of Biochemistry 1014NSC

Trimester 2, 2021

Lecture Notes Week 1 – Water and Weak interactions Water 1. Water plays a central role in the chemistry of all life 2. Proteins, polysaccharides, nucleic acids, and membranes all assume their characteristic shapes in response to water 3. The chemical properties of water are related to the functions of biomolecules, entire cells, an organisms. Weak interactions in aqueous solutions 1. 2. 3. 4.

Hydrogen bonding Ionic interactions Hydrophobic interactions van der Waals interactions

Weak interactions are crucial to macromolecular structure and function The structure of H2O -

Important properties of water arise from its angled shape. Angle of 104.5º between two covalent bonds. bonding orbital - sp3. Polar O-H bonds are due to uneven distribution of charge The oxygen nucleus attracts electrons more strongly than does the hydrogen nucleus — the electrons are more often in the vicinity of the oxygen atom (2δ - ) than the hydrogen (δ +). Angled arrangement of polar bonds creates a permanent dipole for a water molecule.

Hydrogen bonding -

Water molecules attract each other due to their polarity. A hydrogen bond is formed when a partially positive hydrogen atom attracts the partially negative oxygen atom of a second water molecule. Hydrogen bonds can form between electronegative atoms and a hydrogen attached to another electronegative atom.

Hydrogen bonding between water molecules -

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A water molecule can form up to four hydrogen bonds. In liquid water at room temperature and atmospheric pressure water molecules are: o disorganized (randon) o in continuous motion Each molecules forms hydrogen bonds with an average of only 3.4 other molecules.

Hydrogen bonding of ice -

In ice, each water molecule forms the maximum of four hydrogen bonds, creating a regular crystal lattice. This crystal lattice of ice makes it less dense than liquid water, and thus ice floats on liquid water.

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Fundamentals of Biochemistry 1014NSC

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Water as a solvent -

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Water is a good solvent for charged and polar substances: o Amino acids and peptides o Small alcohols o Carbohydrates Water is a poor solvent for nonpolar substances: o Nonpolar gases o Aromatic moieties o Aliphatic chains

Polar, non-polar and amphipathic biomolecules

Examples of hydrogen bond acceptors and donors H-bond acceptor: electronegative atoms such as O or N with an ion pair of electrons H-bond donor: Hydrogen atoms covalently bonded to another electronegative atom

Some biologically important hydrogen bonds -

Source of unique properties of water Structure and function of proteins Structure and function of DNA Structure and function of polysaccharides Binding of substrates to enzymes Binding of hormones to receptors Matching of mRNA and tRNA 2|Page

Fundamentals of Biochemistry 1014NSC

Trimester 2, 2021

Some biologically important hydrogen bonds

Examples of hydrogen bond acceptors and donors -

Hydrogen bonds are strongest when the bonded molecules oriented to maximize electrostatic interaction, which occurs when the hydrogen atom and the two atoms that share it are in a straight line - that is, when the acceptor atom is in line with the covalent bond between the donor atom and H - holding two hydrogen bonded molecules or groups in a specific geometric arrangement.

Ionic interactions -

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Ionic (Coulombic) interactions o electrostatic interactions between permanently charged species, or between the ion and a permanent dipole Water as solvent. Water dissolves many crystalline salts by hydrating their component ions. The NaCl crystal lattice is disrupted as water molecules cluster about the Cl- and Na+ ions. The ionic charges are partially neutralized, and the electrostatic attractions necessary for lattice formation are weakened.

Hydrophobic interactions -

The biologically important gases CO2 , O2 , and N2 are nonpolar. Nonpolar gases are poorly soluble in water. The movement of molecules - from the disordered gas phase into aqueous solution constrains their motion and the motion of water molecules and therefore represents a decrease in entropy.

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Fundamentals of Biochemistry 1014NSC

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Effects of nonpolar molecules in water -

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Nonpolar Amphipathic compounds force energetically unfavourable changes in the structure of water. Eg: Long-chain fatty acids have very hydrophobic alkyl chains, each of which is surrounded by a layer of highly ordered water molecules. Release of ordered water favours formation of an enzymesubstrate complexes. While separate, both enzyme and substrate force neighbouring water molecules into an ordered shell. Binding of substrate to enzyme releases some of the ordered water, and the resulting increase in entropy provides a thermodynamic push toward formation of the enzymesubstrate complex.

Van der Waals interactions -

van der Waals interactions arise when two uncharged atoms are brought very close together, their surrounding electron clouds influence each other. Random variations in the positions of the electrons around one nucleus may create a transient electric dipole, which induce a transient, opposite electric dipole in the nearby atom. The two dipoles weakly attract each other, bringing the two nuclei together.

Ionization of water, weak acids, and bases -

Pure water is slightly ionized. The ionization of water is expressed by an equilibrium constant. The pH scale designates the H+ and OH- concentrations. Weak acids and bases have characteristic dissociation constants. Titration curves reveal the pKa of weak acids.

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Fundamentals of Biochemistry 1014NSC

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Ionization of water -

Water is a neutral molecule with a very slight tendency to ionize. H2O ↔ H+ + OH-

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Free protons (H+) don’t actually exist, rather they exist as hydronium ions, H3O+. For simplicity, however, we often represent these ions as H+. H2O ↔ H+ + OH

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The ionization (dissociation) of water is described by the expression

H +¿ ¿ OH −¿ ¿ ¿ K =¿ -

where K is the dissociation constant Because of the undissociated [H2O] is much larger than the concentrations of the component ions, it can be considered constant (unchanging), and incorporated into K to yield an expression for the ionization of water Kw. Kw = [H+] [OH- ]

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The value of Kw, the ionization constant of water is 10-14 at 25°C. If the ionization constant of water is 10-14 then the concentration of H+ in solution is 10-7 mol/L with an equivalent concentration of OH- ions.

H +¿ ¿ OH −¿ ¿ OH −¿ ¿ H +¿ ¿ ¿ K w=¿ -

Since these values are very low and involve negative powers of 10, the pH scale can be used.

H +¿ ¿ H +¿ ¿ pH=−log 10 ¿ -

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Eg: Human blood plasma has a [H+] of ~ 0.4 x 10-7 mol/L or 10-7.4 mol/L which gives a pH of 7.4 The value where an equal amount of H+ and OHions are present is termed neutrality: at 25°C the pH of pure water is 7.0. At this temperature pH values below 7.0 are acidic and above pH 7.0 are alkaline. Neutral solutions change with temperature, due to enhanced dissociation of water with increasing temperature. 5|Page

Fundamentals of Biochemistry 1014NSC -

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Always remember that the pH scale is logarithmic, and not a linear one. Thus, a solution of pH 3.0 is not twice as acidic as a solution at pH 6.0 but 1000 times more acidic (ie: contains 1000 times more H+ ions).

Proton movement between water molecules -

The proton of a hydronium ion can jump rapidly from one water molecule to another Therefore, the mobility of H+ and OH ions in solution are much higher than for other ions. Proton jumping is the reason for acid-base reactions being among the fastest.

Acid-base chemistry -

Acid – a compound that acts as a proton donor in an aqueous solution. Base – a compound that acts as a proton acceptor in an aqueous solution. Conjugate pair – an acid together with its corresponding base. By the above definitions, an acid-base reaction can be written as HA + H2O ↔ H3O+ + A An acid (HA) reacts with a base (H2O) to form a conjugate base of the acid (A- ) and the conjugate acid (H3O+). Eg: acetate ion (CH3COO- ) is the conjugate base of acetic acid (CH3COOH) and the ammonium ion (NH4 +) is the conjugate acid of NH3 . The acid-base reaction is usually abbreviated to HA ↔ H+ + A

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The participation of water in the reaction is implied. An alternative expression for a basic solution is HB+ ↔ H+ + B

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The strength of an acid is specified by its dissociation constant, or its efficiency as a proton donor. The equilibrium constant for an acid-base reaction is expressed as a dissociation constant.

H 3O +¿ ¿ A−¿ ¿ ¿ [ products] =¿ K= [reactants] -

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In dilute solutions, the water concentration is essentially constant 55.5 M. Therefore, the term [H2O] is customarily combined with the dissociation constant, to take the form.

−¿ +¿+ A¿ For an acid ¿ HA=H

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H +¿ ¿ A−¿ At equilibrium ¿ ¿ Ka= K [ H 2 O ] =¿ Thus, the stronger the acid the greater the value of Ka. Because the acid dissociation constants, like [H+] values are cumbersome to work with, they are transformed into pK values by the formula. pK = -logK analogous to pH = -log [H+] = log 1 [H+]

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Acids can be classified according to their relative strengths, that is, their ability to transfer a proton to water. Weak acids K < 1

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strong acids K >> 1

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Virtually all the acid-base reactions in biological systems involve H3O+, (OH- ) and weak acids (and their conjugate bases). The pH of a solution is determined by the relative concentration of acids and bases. The relationship between the pH of a solution and the concentrations of an acid and its conjugate base can easily be derived.

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The Henderson-Hasselbach Equation

A−¿ ¿ ¿ ¿ pH = pKa+ log ¿ -

Relates the extent of ionisation of a weak acid (and base) to the pH of the solution.

pH = pKa+log -

[conjugate base] [ acid]

This equation is of fundamental importance in preparing buffer solutions to control pH during biochemical reactions. The ionisation characteristics of some of these compounds/groups actually controls the pH in cells and physiological fluids so pH varies only over a narrow range. In Blood the pH stays in the range 7.0 – 7.8 The concentration of the acid in the original solution can be calculated from the volume and concentration of NaOH added and a titration curve plotted. At the midpoint of the titration, at which exactly 0.5 equivalent of NaOH has been added, [HA]=[A- ] and pH=pKa. 7|Page

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The titration curves of these acids have the same shape, they are displaced along the pH axis because the three acids have different strengths. Acetic acid, with the highest Ka (lowest pKa) of the three, is the strongest (loses its proton most readily); it is already half dissociated at pH 4.76. Dihydrogen phosphate loses a proton less readily, being half dissociated at pH 6.86. Ammonium ion is the weakest acid of the three and does not become half dissociated until pH 9.25. Buffer are mixtures of weak acids and their conjugate bases. Buffers are aqueous systems that tend to resist changes in pH when small amounts of acids or base added. A mixture of equal concentration of acetic acid and acetate ion, found at the midpoint of the titration curve is a buffer system.

Buffering in biological solutions -

Buffers are mixtures of weak acids and their conjugate bases. A simple expression relates pH, pKa, and buffer concentration (Henderson-Hasselbalch equation). Weak acids or bases buffer cells and tissues against pH changes.

Week 3 – Amino Acids and Peptide bonds -

Amino acid: α-amino-substituted carboxylic acids, the building blocks of proteins Peptide bond: A substituted amide linkage between the α-amino group of one amino acid and the α-carboxylic group of another.

Amino acids -

All amino acids have a carboxyl and amino group bonded to an α-carbon. The α-carbon also has a hydrogen and a side chain (R-group). Each amino acid differs in the side chain which vary in charge, size, structure, water solubility and chemical properties.

Chirality of amino acids -

All amino acids with the exception of glycine are chiral and thus exist in L and D isomers. Only L-enantiomers are found in proteins. The formation of stable repeating substructures in proteins generally require their constituent amino acids to be of one stereochemical series. Cells are able to specifically synthesize the L isomer of amino acids because the active site of enzymes is asymmetric, causing the reactions they catalyse to be stereospecific.

Non-polar aliphatic amino acids -

Aliphatic amino acids tend to be hydrophobic. Become more hydrophobic as the side chain increases in length. Hydrophobic amino acids are usually buried in proteins for protection against aqueous environments.

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Tend to be clustered together within proteins, stabilizing structure by means of hydrophobic interactions. Methionine is one of only 2 amino acids that contain sulfur. It contains a nonpolar thioether group in its side chain.

Proline -

Proline is a cyclic amino acid. Shares many of the properties of aliphatic amino acids. The rigid nature of proline makes the folding of proline into proteins difficult. Generally introduces a kink into polypeptide chain.

Polar, uncharged amino acids -

R groups of these amino acids are more soluble in water and more hydrophilic than the nonpolar amino acids. Side chains contain functional groups that are able to form hydrogen bonds in water. Polarity is due to the hydroxyl groups of serine and threonine, the sulfhydryl group of cysteine and the amide groups of asparagine and glutamine. Asparagine and glutamine are the amide derivatives of aspartate and glutamate, respectively.

Cysteine -

The side chain can ionize at moderately high pH. It can form a covalently linked dimeric amino acid called a cystine in which two cysteine residues are joined to form a disulfide bond. Play a special role in structures by covalently links between different parts of a protein or different polypeptides. Disulfide linkages are strongly hydrophobic (nonpolar).

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Fundamentals of Biochemistry 1014NSC -

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The R groups carry positive charges and are strongly polar at pH 7.0. They are very hydrophilic and usually found on the exterior of proteins. Lysine has a secondary amino group on its aliphatic side chain. Arginine has a positively charged guanidino group.

Histidine -

Histidine contains a imidazole group. Is the only common amino acid having an ionizable with a pKa near neutrality. In enzyme catalysed reactions a his residue facilitates the reaction by serving as a proton donor/acceptor.

Negatively charged amino acids -

R groups carry a net negative charge at pH 7.0. Like the positively charged amino acids they are very hydrophilic. Both aspartate and glutamate have a second carboxylic acid group.

Aromatic amino acids -

The R groups in this class are relatively nonpolar and participate in hydrophobic interactions. The hydroxyl group of tyrosine is also important because of its ability to form H-bonds, and it is an important functional group in some enzymes. Tyrosine and tryptophan are significantly more polar than phenylalanine because of the tyrosine hydroxyl group and the tryptophan indole ring. Tryptophan, tyrosine and to a lesser extent phenylalanine absorb ultraviolet light.

Modified amino acids -

There are a number of modified amino acids which are formed from the common amino acids. Examples: o 4-hydroxyproline (cell walls, collagen) o 5-hydroxylysine (collagen) o 6-N-methyllysine (myosin) o γ-carboxyglutamate (prothrombin) o selenocysteine

4-hydroxyproline -

Example: collagen, a fibrous protein of connective tissue.

5-hydroxylysine -

Example: collagen, a fibrous protein of connective tissue

6-N-methyllysine -

Example: a constituent of myosin, a contractile protein of the muscle.

γ-carboxyglutamate -

Example: found in the blood clotting protein prothrombin and in certain Ca2+ binding proteins.

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Fundamentals of Biochemistry 1014NSC

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Selenocysteine -

Is introduced during protein synthesis rather than created through a postsynthetic modification.

Other properties of amino acids -

Amino acids are ionised in aqueous solutions. At pH 7.0, the amino group is largely protonated (NH3+) and the carboxyl group largely deprotonated (COO-) forming a zwitterion.

Amino acids titration curves -

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Two distinct stages corresponding to the deprotonation of carboxy group and the amino group. At low pH glycine is fully protonated. At the midpoint of this part of the titration (point of inflection) the pH = pKa (pK1). As the titration continues another inflection point is reached. At this point removal of the first proton is complete. The second half of the titration corresponds to the removal of the proton from the amino. It gives a quantitative measure of the pKa of each of the ionizing groups. There are two regions of buffering power, extending for approximately 1 pH unit either side of the pKa. The relationship between net electric charge and the pH of the solution can also be derived from the titration curve. So at pH 5.97, glycine is present in its dipolar form, fully ionized but with no net electrical charge. This point is called the Isoelectric point or Isoelectric pH. So at a pH < pI glycine has a net + charge & at a pH > pI glycine has a net – charge

Ionization of amino acids -

The side chains of some amino acids are also ionisible and each group has a specific pKa. At pH > pKa the proton tends to be off. At pH < pKa the proton tends to be on. The average pKa for an amino acid is called the isoelectric point (pI). As pH increases above the pI the net charge becomes negative. As pH decreases below the pI the net charge becomes positive. When the amino acids are incorporated into proteins only the charges on the side chains remain.

Peptide bond -

A peptide bond is formed when two amino acids are covalently linked between the amino group of one amino acid and the carboxyl group of another. The result is a dipeptide.

Properties of peptide bonds -

Peptide bonds are planar and prefer to be in the trans form. The peptide bond exhibits a partial double bond character. Oligo and polypeptides have unreacted amino group at one terminus (the amino N terminus) and an unreacted carboxyl group at the other end (carboxy C terminus).
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