Weak Interactions in Aqueous Systems PDF

Weak Interactions in Aqueous Systems- Structure and Ionization of Water - Acids&Bases&pH Types of Chemical Bonds IONIC BOND COVALENT BOND H BOND The strength of chemical bonds Bond energy (E): the amount of energy required to break apart a mole of molecules into its...

Weak Interactions in Aqueous Systems- Structure and Ionization of Water - Acids&Bases&pH Types of Chemical Bonds IONIC BOND COVALENT BOND H BOND The strength of chemical bonds Bond energy (E): the amount of energy required to break apart a mole of molecules into its component atoms. Size of the atom Electronegativity Bond length The Importance of Noncovalent Interactions in Biochemistry Marks’ basic medical biochemistry : a clinical approach / Michael Lieberman, Allan Marks, Alisa Peet ; illustrations by Matthew Chansky. — 4th ed. The Importance of Noncovalent Interactions in Biochemistry Molecular Polarity Electronegativity, is the tendency for an atom of a given chemical element to attract shared electrons (or electron density) when forming a chemical bond Molecular Polarity A polar molecule is a molecule in which one end of the molecule is slightly positive, while the other end is slightly negative. Molecular polarity depends on the shape of the molecule as well as the presence of polar covalent bonds and lone The Nature of Noncovalent Interactions Van der Waals Interactions Van der Waals forces are driven by induced electrical interactions between two or more atoms or molecules that are very close to each other. Van der Waals interaction is the weakest of all intermolecular attractions between molecules. Van der Waals Interactions The van der Waals radii of the carbon atoms (1.7 Å) dictate that the planes of the two stacked benzene rings cannot be closer than 3.4 Å. Hydrogen Bond A hydrogen bond is an interaction between a hydrogen atom covalently bonded to another atom and a pair of nonbonded electrons on a separate atom (in biomolecules, this is most often an O or a N atom). Hydrogen Bond The atom to which the hydrogen is covalently bonded is called the hydrogen-bond donor, and the atom with the nonbonded electron pair is called the hydrogen-bond acceptor. Energies of some noncovalent interactions in biomolecules An input of about 350 kJ of energy is required to break a mole of C—C single bonds, and about 410 kJ to break a mole of C—H bonds. The Structure and Properties of Water ֍ Polarity ֍ Ability to dissolve many polar and ionic substances ֍ High heat capacity ֍ High heat of vaporization ֍ Cohesive and adhesive properties ֍ Water is less dense as a solid than as a liquid ֍ Higher melting point, boiling point, and heat of vaporization Water as a Solvent Water is an excellent solvent because of its hydrogen bonding potential and its polar nature. It readily dissolves most biomolecules, which are generally charged or polar compounds; compounds that dissolve easily in water are hydrophilic (Greek, “water-loving”). Hydrophobic Molecules in Aqueous Solution Hydrophobicity is the physical property of a molecule that is seemingly repelled from a mass of water (known as a hydrophobe) Amphipathic Molecules in Aqueous Solution Amphipathic molecules has a “head” group that is strongly hydrophilic, coupled to a hydrophobic “tail” usually a hydrocarbon. Some Examples of Polar, Nonpolar, and Amphipathic Biomolecules (Shown as Ionic Forms at pH 7) Acids and Bases: Proton Donors and Acceptors Bronsted-Lowry Concept:  Acid can donate a proton  Base can accept a proton Strong&Weak Acids A strong acid dissociates almost completely into a proton and a weak conjugate base. Strongacids Strong acidshave have weak weak conjugate conjugate base base Weak Weakacids acidshave have strong strongconjugate conjugate base base Ionization of Water Although water is essentially a neutral molecule, it does have a slight tendency to ionize; in fact, it can act as both a very weak acid and a very weak base. Molecules or ions which can either donate or accept a proton, depending on their circumstances, are called amphiprotic species. Kw, the ion product The equilibrium described above can be expressed in terms of Kw, called the ion product of water, which is 10- 14 at 25 °C. The higher the H+ of a solution describes In pure water, all the an acidic solution. H+ and -OH ions must come from the dissociation of the water itself. Under A low H+ must be these circumstances, accompanied by a the concentrations high -OH, describes a of H+ and - OH must be equal; basic solution. thus, at 25 °C, and the solution is said to Worked Example-1 What is the concentration of H+ in a solution of 0.1M NaOH? Solution: We begin with the equation for the ion product of water: With [OH-]= 0.1 M, solving for [H+] gives The pH Scale and the Physiological pH Range To avoid working with negative powers of 10, we almost always express hydrogen ion concentration in terms of pH, defined as: “The negative logarithm to the base 10 of the hydrogen ion concentration” Note that most body fluids have pH values in the range 6.5–8.0, which is often referred to as the physiological pH range. Worked Example-2 What is the concentration of OH- in a solution with an H+ concentration of 1.3x 10- 4 M? Solution: We begin with the equation for the ion product of water: Kw=[H+][OH-] With [H+]=1.3 x 10-4 M, solving for [OH-] gives: Worked Example-3 What is the pH of a 0.02M sodium hydroxide solution? 0.02M 0.02M 0.02M pOH=−log[OH−]=−log(0.02) pOH=1.70 pH+pOH=14→pH=14 −pOH 14−1.70=12.3 pH scale is logarithmic Change in just one unit of scale = tenfold change in H+ concentration. “The patient’s blood pH has changed by 0.3 pH unit” means it has doubled (or halved) in value. pH values seen in clinical practice Weak Acid and Base Equilibria: Ka and pKa The equilibrium constant for A larger value the dissociation of a weak acid of Ka indicates a stronger acid Because pKa is the negative logarithm of The strength of acids is Ka, a numerically usually expressed in smaller value of pKa terms of the pKa value corresponds to a stronger acid, and a larger value Some weak acids and their conjugate bases The Henderson–Hasselbalch Equation A weak acid [H+]= Ka [HA] dissociate [A-] s as shown -log [H+] = -logKa – log [HA] [A-] pH=pKa – log [HA] [A-] pH=pKa – log [A-] [HA] Titration of Weak Acids ֍ Titration is used to determine the amount of an acid in a given solution. ֍ A measured volume of the acid is titrated with a solution of a strong base, usually sodium hydroxide (NaOH), of known concentration. ֍ The NaOH is added in small increments until the acid is consumed (neutralized), as determined with an indicator dye or a pH meter. ֍ The concentration of the acid in the original solution can be calculated from the volume and concentration of NaOH added. David L. Nelson, Michael M. Cox - Lehninger Principles of Biochemistry, W.H. Freeman (2012) Buffer Solutions A buffer solution is a solution which resists changes in pH when a small amount of acid or base is added. Typically a mixture of a weak acid and a salt of its conjugate base or weak base and a salt of its conjugate acid. David L. Nelson, Michael M. Cox - Lehninger Principles of Biochemistry, W.H. Freeman (2012) Worked Example-4 A buffer solution is prepared by dissolving 0.1 moles of acetic acid (CH₃COOH) and 0.1 moles of sodium acetate (CH₃COONa) in enough water to make 1 liter of solution. The pKa of acetic acid is 4.76. What is the pH of this buffer solution? pKa of acetic acid = 4.76 Moles of [A⁻] (acetate ion) = 0.1 pH is the pH of the solution. moles (from sodium acetate) pKa is the acid dissociation Moles of [HA] (acetic acid) = 0.1 constant of the weak acid. moles (from acetic acid) [A⁻] is the concentration of the Volume of solution = 1 liter conjugate base (acetate ion, Since both the acetic acid and acetate CH₃COO⁻). [HA] is the concentration of the are dissolved in 1 liter, their weak acid (acetic acid, CH₃COOH). concentrations will be: [A⁻] = 0.1 M (since 0.1 moles / 1 liter = 0.1 M) [HA] = 0.1 M The Body & pH Homeostasis of pH is tightly Why is it important to controlled Extracellular fluid pH= 7.4 maintain pH of blood within Arterial Blood pH= 7.35 – 7.45 normal range? (OUR NORMAL RANGE) Venous blood is more acidic Most enzymes function only with than arterial? narrow pH ranges Because it contains more CO2  Acid-base balance can affect than arterial blood. electrolytes (Na+, K+, Cl, Ca++) < 6.8 or > 8.0 death occurs  pH affect hormones. Acidosis : below 7.35  To maintain normal function of Alkalosis : above 7.45 synapses The Body & pH Acids produced by metabolism of lipids and proteins Cellular metabolism produces CO2 Volatile CO2 15,000 mmol H+ Lungs equivalents per day Organic acids Primarily ketones and Several thousand Primarily liver lactate mmol per day Inorganic acids Primarily sulfate and 1.5 mmol/kg per day Primarily renal phosphate Morris CG. Anaesthesia 2008;63:294- Body Defence Against Changes in pH Acid-Base Homeostasis First Second Third Line Line Line Bicarbonat Physiological Buffer e Buffer System Renal Respirator System Chemic y Mechanis Mechanis al Phosphate m m Buffer Buffer System System Protein Buffer System Buffering against pH Changes in Biological Systems First line of defence Two most common chemical buffer groups Bicarbonate Non-bicarbonate (Hb,protein,phosphate) Blood buffer systems act instantaneously Regulate pH by binding or releasing The Bicarbonate Buffer System Blood plasma is buffered in part by the bicarbonate system, consisting of carbonic acid (H2CO3) as proton donor and bicarbonate (HCO-3 ) as proton acceptor The Bicarbonate Buffer System Lactic acid and pH ֍ Lactic acid and pH homeostasis by the bicarbonate buffer system. ֍ The bicarbonate buffer system removes protons [H+] generated during anaerobic glycolysis. ֍ The protons are disposed of as water while the CO2 evolved is expired via the lungs. The phosphate buffer system The phosphate buffer system is maximally effective at a pH close to its pKa of 6.86 and thus tends to resist pH changes in the range between about 5.9 and 7.9. It is therefore an effective buffer in biological fluids; in mammals, for example, extracellular fluids and most cytoplasmic compartments have a pH in the range of 6.9 to 7.4. Protein Buffer Systems The cytoplasm of most cells contains high concentrations of proteins, and these proteins contain many amino acids with functional groups that are weak acids or weak bases. For example, the side chain of histidine has a pKa of 6.0 and thus can exist in either the protonated or unprotonated form near neutral pH. Proteins containing histidine residues therefore buffer effectively near neutral pH