Lecture 2 Water and Buffers.pdf

Water and Buffers Lecture 2 Learning Outcomes At the end of this lecture, students should be able to: Explain the importance of water as a universal solvent Recognize water structure and describe hydrogen bonding Describe water interactions with biological molecules. Explain (b...

Water and Buffers Lecture 2 Learning Outcomes At the end of this lecture, students should be able to: Explain the importance of water as a universal solvent Recognize water structure and describe hydrogen bonding Describe water interactions with biological molecules. Explain (briefly) water ionization constant, pH scale, weak acids and bases dissociation constants and buffering systems  Water is the medium for life Organisms typically contain 70–90% water. Dissolve nutrients and easy to transport Chemical reactions occur in aqueous environment Water is a critical determinant of the structure and function of proteins, nucleic acids, and membranes  Structure of the Water Molecule There are four electron pairs around an oxygen atom in water Two of these pairs covalently link two hydrogen atoms to a central oxygen atom The two remaining pairs remain nonbonding (lone pairs)  Structure of the Water Molecule The water molecule is a polar molecule: the opposite ends have opposite charges Polarity allows water molecules to form hydrogen bonds with each other The electronegativity of the oxygen atom induces a net dipole moment – O is more electronegative than H, shared electrons are pulled more towards O. – Because of the dipole moment, water can serve as both a hydrogen bond donor and acceptor  Hydrogen Bonding in Water Strong dipole-dipole or charge-dipole interaction that arises between an acid (proton donor) and a base (proton acceptor) Typically involves two electronegative atoms (frequently nitrogen and oxygen) Hydrogen bonds are strongest when the bonded molecules are oriented to maximize electrostatic interaction Ideally the three atoms involved are in a line  Hydrogen Bonding Gives Water Its Unusual Properties Water can serve as both – an H donor – an H acceptor Up to four H-bonds per water molecule gives water its. – unusually high boiling point – unusually high melting point – unusually large surface tension Hydrogen bonding in water is weak but cooperative (one strengthen the other) Hydrogen bonds between neighboring In liquid: not all water molecules are weak (20 kJ/mol) relative molecules make four H-bonds to the H–O covalent bonds (420 kJ/mol)  Water Forms Hydrogen Bonds with Polar Solutes Hydrogen bonds are not unique to water. They form between an electronegative atom (H acceptor, usually O or N) and a H atom covalently bonded to another electronegative atom (the H donor) in the same or another molecule Importance of Hydrogen Bonds: - Source of unique properties of water - Structures and functions of proteins, DNA, polysaccharides - Binding of substrates to enzymes - Binding of hormones to receptors - Matching of mRNA and tRNA Biological Relevance of Hydrogen Bonds  Weak Interactions in Aqueous Systems Water is a good solvent for charged and polar substances - dissolve readily in water because they can replace water-water interactions with more energetically favorable water-solute interactions – amino acids and peptides – small alcohols – carbohydrates Water is a poor solvent for nonpolar substances because they interfere with water-water interactions but are unable to form water-solute interactions. In aqueous solutions, nonpolar molecules tend to cluster together – nonpolar gases – aromatic moieties – aliphatic chains Hydrophilic Hydrophobic  Water dissolves many salts Strong electrostatic interactions between the solvated ions and water molecules lower the energy of the system NaCl Solubility of Polar and Nonpolar Solutes Why are nonpolar molecules poorly soluble in water?  Noncovalent Interactions No sharing a pair of electrons. Ionic (Coulombic) interactions – electrostatic interactions between permanently charged species Hydrogen bonding – electrostatic interactions between uncharged but polar molecules Hydrophobic Effect – Ordering of water molecules around nonpolar substances van der Waals interactions – weak interactions between all atoms, regardless of polarity Examples of Noncovalent Interactions  Water, buffer and pH Water is both the solvent in which metabolic reactions occur and a reactant in many biochemical processes, including hydrolysis, condensation, and oxidation- reduction reactions  Ionization of Pure Water → H+ + OH- H2O  O-H bonds are polar and can dissociate heterolytically Products are a proton (H+) and a hydroxide ion (OH–) Dissociation of water is a rapid reversible process Most water molecules remain un-ionized, thus pure water has very low electrical conductivity The equilibrium is strongly to the left Extent of dissociation depends on the temperature  Ionization of Water Concentrations of participating species in an equilibrium process are not independent but are related via the equilibrium constant: [H +] [OH-] H2O  → H+ + OH- Keq = ———— [H2O] Keq can be determined experimentally, it is 1.8 10–16 M at 25C. [H2O] can be determined from water density, it is 55.5 M. Ionic product of water: K w = K eq  [H 2 O] = [H + ][OH - ] = 1 10 −14 M 2 In pure water [H+] = [OH–] = 10–7 M  Ionization of Water and pH Concentrations of H+ and OH– are equal in pure water Adding certain solutes, called acids and bases, modifies the concentrations of H+ and OH– Biologists use something called the pH scale to describe whether a solution is acidic or basic (the opposite of acidic)  What is pH? pH is defined as the negative logarithm of the hydrogen ion pH = -log[H+] concentration The pH and pOH must always add to 14 K w = [H + ][OH - ] = 1 10 −14 M 2 In neutral solution, [H+] = [OH–] and the pH is 7 − log[ H + ] − log[ OH - ] = +14 pH + pOH = 14  Dissociation of Weak Electrolytes: Principle Weak electrolytes dissociate only O Keq O partially in water. H3C + H2O H3C + H3O+ OH O- Extent of dissociation is determined K a = K eq  [H 2 O] by the acid dissociation constant Ka. [H + ][CH 3COO - ] Ka = = 1.74 10 −5 M pKa is determined by titration [CH 3COOH] curves pKa = - log Ka [CH 3COOH ] [H + ] = Ka  [CH 3COO − ] Henderson–Hasselbalch Equation: Derivation [H + ][A - ] HA  → H+ + A- Ka = [HA] [HA] [H ] = K a + [A - ] [HA] - log[H +] = -logK a − log [A-] - [A ] pH = pK a + log [HA]  Henderson–Hasselbalch Equation: Uses Estimating the pH of a buffer solution Finding the equilibrium pH in an acid-base reaction To determine the amount of acid and conjugate base needed to make a buffer solution of a certain pH.  Titration Curves Reveal the pKa of Weak Acids The titration curve of acetic acid  Buffers are mixtures of weak acids and their anions (conjugate base) Buffers resist change in pH At pH = pKa, there is a 50:50 mixture of acid and anion forms of the compound Buffering capacity of acid/anion system is greatest at pH = pKa Buffering capacity is lost when the pH differs from pKa by more than 1 pH unit  Biological Buffer Systems Maintenance of intracellular pH is vital to all cells – Enzyme-catalyzed reactions have optimal pH – Solubility of polar molecules depends on H-bond donors and acceptors – Equilibrium between CO2 gas and dissolved HCO3– depends on pH Buffer systems in vivo are mainly based on – Phosphate, buffers pH of cytoplasm, other ICF and urine – bicarbonate, important for blood plasma – histidine, efficient buffer at neutral pH – Proteins (Hg in RBCs cytoplasm)  pH Scale pH of Some 0 1 Common Battery acid Increasingly Acidic 2 Gastric juice, lemon juice H+ Liquids [H+] > [OH−] H+ − H+ Vinegar, wine, H+ OH 3 − + OH H H+ cola + H H+ Acidic 4 Tomato juice Beer solution Black coffee 5 Rainwater Acidic solutions have 6 Urine OH− Saliva OH− Neutral pH values less than 7 H+ H+ OH OH− OH − H+ − [H+] = [OH−] 7 Pure water Human blood, tears H+ Basic solutions have pH H+ Neutral 8 Seawater Inside of small intestine values greater than 7 solution Increasingly Basic 9 Most biological fluids [H+] < [OH−] 10 have pH values in OH− OH− Milk of magnesia 11 the range of 6 to 8 OH− H+ OH− OH− OH + − OH− Household ammonia H 12 Basic solution Household 13 bleach Oven cleaner 14