Biochemistry: An Evolving Science PDF

Biochemistry: An Evolving Science Andrew Jakymiw, PhD Office: BSB 230C Phone: 843-792-2551 Email: [email protected] Learning objectives: 1. Compare and contrast the unity of biology at the biochemical level 2. Know the difference between covalent and noncovalent bonds 3. Understand the four t...

Biochemistry: An Evolving Science Andrew Jakymiw, PhD Office: BSB 230C Phone: 843-792-2551 Email: [email protected] Learning objectives: 1. Compare and contrast the unity of biology at the biochemical level 2. Know the difference between covalent and noncovalent bonds 3. Understand the four types of noncovalent bonds and their importance to macromolecular structures 4. Understand the properties of water 5. Understand the concepts of free energy and entropy 6. Understand how the laws of thermodynamics govern biochemical reactions 7. Understand the principles of acid-base reactions and know how buffers regulate pH Biochemical unity underlies biological diversity The biological world is magnificently diverse, yet a key feature underlies this diversity:  The construction of animals, plants, and microorganisms from cells implies that these diverse organisms might have more in common than their outward appearance would suggest In fact, all organisms have many common features at the biochemical level: 1) Similar classes of molecules (macromolecules, such as proteins and nucleic acids) 2) Common metabolic processes (metabolites, such as glucose and glycerol) The observations made at the biochemical level suggest that all living things on Earth have evolved from a common ancestor: Based on biochemical characteristics, the diverse organisms of the modern world can be divided into three groups/domains: 1) Eukarya (eukaryotes) 2) Bacteria 3) Archaea Concepts from chemistry help explain the properties of biological molecules 1. 2. 3. 4. Chemical bonds Properties of water Laws of Thermodynamics Acid-base chemistry 1. Chemical bonds I. Covalent bonds Strongest bonds that hold atoms together Formed by the sharing of a pair of electrons between adjacent atoms Require considerable amounts of energy to break them A typical C-C covalent bond has a bond energy of 85 kcal/mol and 1.5Å bond length • More than one electron pair can be shared between two atoms (multiple covalent bond) • • • • • For some molecules, more than one pattern of covalent bonding can be depicted (resonance structures) Adenine (A) • Critical aspect of carbon chemistry macromolecules in living systems permit the formation of 1. Chemical bonds continued… II. Noncovalent bonds • Weaker than covalent bonds • Allow for reversible molecular interactions • Consist of four bond types: A. Ionic Interactions B. Hydrogen bonds C. van der Waals Interactions D. Hydrophobic Interactions A. Ionic/Electrostatic Interactions • Charged group on one molecule can attract an oppositely charged group on the same or another molecule • An attractive interaction has a negative energy • Strength is greatly affected by the presence of the solvent and its dielectric constant  Example: Interactions between two ions bearing single opposite charges in water have an energy of -1.4 kcal/mol, whereas in hexane, it is -55 kcal/mol • Examples: Na+Cl- or Na+CH3CO2- 1. Chemical bonds continued… B. Hydrogen bonds • Fundamentally electrostatic interactions (responsible for specific base-pair formations in the DNA double helix) • Electronegative atom to which the hydrogen atom is covalently bonded pulls electron density away from the hydrogen atom (hydrogen-bond donor), developing a positively charged hydrogen atom (δ+) that can then interact with an atom having a partial negative charge (δ-; hydrogen-bond acceptor) through an ionic interaction • Weaker (1-5 kcal/mol) and longer than covalent bonds Hydrogen bonds are depicted by dashed green lines. The positions of the partial charges are shown. Watson–Crick base pairs. Adenine pairs with thymine (A–T) and guanine with cytosine (G–C). The dashed green lines represent hydrogen bonds. 1. Chemical bonds continued… C. van der Waals interactions • Distribution of electronic charge around an atom fluctuates with time resulting in a transient asymmetry in the electronic charge around an atom • This asymmetry induces a complementary asymmetry in the electron distribution of neighboring atoms – thus causing attraction between the atoms • Attraction increases as two atoms come closer to each other, until they are separated by van der Waals contact distance • At distances shorter than van der Waals contact distance very strong repulsive forces become dominant • Weak bonds (0.5-1 kcal/mol), however, the net effect of a large number of atoms in van der Waals contact, can be substantial Energy of a van der Waals interaction as two atoms approach each other. The energy is most favorable at the van der Waals contact distance. Owing to electron– electron repulsion, the energy rises rapidly as the distance between the atoms becomes shorter than the contact distance. 2. Properties of water • Solvent in which most biochemical reactions take place • Two properties are particularly important: i. Water is a polar molecule ii. Water is highly cohesive (i) Polar Molecule • A water molecule is bent, not linear, and so the distribution of charge is asymmetric (thus, an electrically polar structure) (ii) Highly cohesive • Water molecules interact strongly with one another through hydrogen bonds • Water is a versatile solvent able to readily dissolve polar and charged compounds that can participate in the formation of hydrogen bonds and ionic interactions Polar: Cohesive: Ice Structure of ice. Hydrogen bonds (shown as dashed green lines) are formed between water molecules to produce a highly ordered and open structure. 1. Chemical bonds continued… D. Hydrophobic interactions • A manifestation of the properties of water • Nonpolar molecules cannot participate in hydrogen bonding or ionic interactions • Interaction of nonpolar molecules with water molecules is energetically not favorable • Water forms cages around nonpolar molecules • Release of water molecules into solution makes aggregation of nonpolar groups favorable Hydrophobic effect 1. Chemical bonds continued… The DNA double helix is an expression of the rules of chemistry • • • • Phosphate groups in DNA carry a negative charge  So, unfavorable ionic interactions oppose the formation of double helix  But these repulsive ionic interactions are diminished by the high dielectric constant of water and the presence of ions (Na+/Mg++) which partially interact with phosphate groups and neutralize their charges Hydrogen bonds determine the formation of specific base pairs in the double helix Base pairs are parallel and stacked nearly on top of one another, the typical separation distance corresponds to van der Waals contact distance The hydrophobic effect contributes to the favorability of base stacking 3. Laws of Thermodynamics • govern biochemical systems, distinguish between system and its surroundings System = matter within a defined region of space Surroundings = matter in the rest of the universe The 1st Law • The total energy of a system and its surroundings is constant The 2nd Law • The total entropy of a system plus that of its surroundings always increases • Example: the release of water from nonpolar surfaces is favorable because water molecules free in solution are more disordered than when they are associated with nonpolar molecules Heat Gibbs Free Energy Free energy of a system can be considered as the energy available to perform work Free energy has an Enthalpy component (H) and an Entropy component (S), described in the following relationship: G = H - TS Gibbs Free Energy continued… The free energy change of a system that occurs during a reaction can therefore be described by the following equation: ΔG = ΔH - ΔTS If the temperature is constant, the equation can be written: ΔG = ΔH – TΔS Gibbs Free Energy continued… Free energy change, ΔG, is used to describe the energetics of biochemical reactions: ΔG = ΔH - TΔS The reaction may release heat (exothermic, ΔH is negative), absorb heat (endothermic, ΔH is positive), lead to an increase in entropy (ΔS is positive), or a decrease in entropy (ΔS is negative). The spontaneity of a reaction may therefore be determined by taking the entropy change, subtracting it from the change in enthalpy, and thereby determining the sign of ΔG Thus, a reaction is spontaneous when ΔG is negative Gibbs Free Energy continued… (ΔH<0) (ΔS>0) (ΔS<0) (ΔH>0) Spontaneous at all T (ΔG<0) Spontaneous at high T (when TΔS is large) Spontaneous at low T (when TΔS is small) Nonspontaneous at all T (ΔG>0) By looking at ΔH and ΔS, we can tell whether a reaction will be spontaneous or otherwise. 4. Acid-base chemistry: • reactions that are central in many biochemical processes Acid-base reactions = the addition or removal of hydrogen ions Hydrogen ion = H+ = proton = hydronium ion = H3O+ The concentration of H+ in solution is expressed as the pH: pH = -log[H+] Proton dissociation for a substance HA has an equilibrium constant defined by the expression: Ka = [H+][A-]/[HA] The susceptibility of a proton to removal by reaction with a base is described by its pKa value: pKa = -log(Ka) Buffers regulate pH in organisms and in the laboratory Changes in pH can protonate or deprotonate key groups in macromolecules, thus potentially disrupting structures and initiating harmful reactions Systems have evolved to mitigate changes in pH in biological systems Solutions that resist such changes are called buffers The effect of buffers can be analyzed in quantitative terms via the Henderson-Hasselbach equation: pH = pKa + log([A-]/[HA]) [A-]/[HA] = 10pH - pKa CH3COOH Biological relevance in terms of oral health: Adapted from Figure 1 in Perkins S, Wetmore M. Acid induced erosion of teeth. Dentistry Today. 2001; 20:82–87. HCO3 = bicarbonate ion CA VI = carbonic anhydrase VI Questions?