Chapter 2 - Water(1).pptx
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CHAPTER 2: WATER Week 2 CHEM311 Chapter Outline 2-1 Water and Polarity 2-2 Hydrogen Bonds 2-3 Acids, Bases and pH 2-4 Titration Curves 2-5 Buffers REVIEW Polarity The polarity of a molecule depends on two things: 1. The geometry of the molecule 2. The polarity of the bonds holding the atoms together...
CHAPTER 2: WATER Week 2 CHEM311 Chapter Outline 2-1 Water and Polarity 2-2 Hydrogen Bonds 2-3 Acids, Bases and pH 2-4 Titration Curves 2-5 Buffers REVIEW Polarity The polarity of a molecule depends on two things: 1. The geometry of the molecule 2. The polarity of the bonds holding the atoms together Remember that the polarity of the bond depends on the electronegativity of the atoms involved. Copyright © 2012, The McGraw-Hill Compaies, Inc. Permission required for reproduction or display. 3 REVIEW Geometry and Polarity Electron Domain Geometry Basic review from Gen. Chemistry – don’t memorize it… REVIEW Electronegativity and Polarity Electronegativity Ionic and covalent bonds are simply the extremes in a spectrum of bonding. Bonds that fall between these two extremes are polar, meaning that electrons are shared but are not shared equally. Such bonds are referred to as polar covalent bonds. Copyright © 2012, The McGraw-Hill Compaies, Inc. Permission required for reproduction or display. 5 REVIEW Electronegativity and Polarity Electronegativity Copyright © 2012, The McGraw-Hill Compaies, Inc. Permission required for reproduction or display. 6 REVIEW Electronegativity and Polarity Electronegativity is the ability of an atom in a compound to draw electrons to itself. Copyright © 2012, The McGraw-Hill Compaies, Inc. Permission required for reproduction or display. 7 REVIEW Electronegativity and Polarity As a general rule: A bond between atoms whose electronegativities differ by less than 0.5 is generally considered purely covalent or nonpolar. A bond between atoms whose electronegativities differ by the range of 0.5 to 2.0 is generally considered polar covalent. A bond between atoms whose electronegativities differ by 2.0 or more is generally considered ionic. Copyright © 2012, The McGraw-Hill Compaies, Inc. Permission required for reproduction or display. 8 REVIEW Electronegativity and Polarity Dipole Moment and Partial Charges In polar covalent bonds, because electrons are not being shared equally, you effectively get a partial charge forming on the different atoms… two charges separated by a given distance constitutes a dipole. The shift of electron density in a polar bond is symbolized by a crossed arrow (a dipole arrow). The charge separation can also be represented like this: Copyright © 2012, The McGraw-Hill Compaies, Inc. Permission required for reproduction or display. 9 Polar Bonds & Molecules Molecules such as CO2 have polar bonds but, given their geometry, are nonpolar molecules; that is, they have a zero dipole moments What makes water polar Difference in electronegativity of atoms involved in bond (O (3.5) – H (2.1) = 1.4; polar covalent) and its geometry (net dipole in one direction) Different types of interactions (some definitions) Ionic Bonds: Held together by positive and negative ions, e.g. NaCl Salt Bridge: Attraction that occurs when oppositely charged molecules are in close proximity, E.g. charged amino acid side-chains in a protein Ion-dipole interactions: When ions in solution interact with molecules with dipoles, e.g., KCl dissolved in H2O Dipole-dipole interactions: Forces that occur between molecules with dipoles, one positive and one negative Dipole induced-dipole interactions ( + induced dipoleinduced dipole = also known as London dispersion forces): weak and generally do not lead to solubility in water van der Waals Forces: between dipoles or induced dipoles Ion-dipole and Dipole-dipole Interactions: (permanent charge) Hydration shell Ion-dipole and dipole-dipole interactions help ionic and polar compounds dissolve in water Dipole-Induced Dipole Interaction (induced charge) Induced Dipole-Induced Dipole Interactions aka. London Dispersion Forces (induced charge) Biochemically-Relevant Bond Energies Van der Waals interactions occur between species that don’t have any full formal charge (i.e. dipoles, induced dipoles)… Hydrogen bonds are a special case Hydrogen Bonds Hydrogen bond: the attractive interaction between dipoles when: positive end of one dipole is a hydrogen atom bonded to an atom of high electronegativity, most commonly O or N, and the negative end of the other dipole is an atom with a lone pair of electrons, most commonly O or N Hydrogen bond is non-covalent Nonlinear bonds are weaker than bonds in which all three atoms lie in a straight line. Hydrogen Bonding Between Polar Groups and Water Interesting and Unique Properties of Water Each water molecule can be involved in 4 hydrogen bonds: 2 as donor, and 2 as acceptor Due to the tetrahedral arrangement of the water molecule. Hydrogen Bonding Even though hydrogen bonds are weaker than covalent bonds, they have a significant effect on the physical properties of hydrogen-bonded compounds Other Biologically Important Hydrogen bonds Hydrogen bonding is important in stabilization of 3-D structures of biological molecules such as: DNA, RNA, proteins. Hydrogen bonds in Biological Molecules Hydrogen bonds in Proteins Hydrogen bonds in Nucleic Acids Solvent Properties of H2O Hydrophilic: water-loving tend to dissolve in water Hydrophobic: water-fearing tend not to dissolve in water Amphipathic: has characteristics of both properties molecules that contain one or more hydrophobic and one or more hydrophilic regions (e.g. membrane lipids) Amphipathic molecules both polar and nonpolar character Interaction between nonpolar molecules is very weak – called van der Waals interactions Micelle formation by amphipathic molecules Micelle: a spherical arrangement of organic molecules in water solution clustered so that their hydrophobic parts are buried inside the sphere their hydrophilic parts are on the surface of the sphere and in contact with the water environment formation depends on the attraction between temporary induced dipoles Acids & Bases A Brønsted-Lowry acid is a proton (hydrogen ion) donor. A Brønsted-Lowry base is a proton (hydrogen ion) acceptor. Acids and Bases A strong acid dissociates (or ionizes) completely in aqueous solution to form hydronium ions (H3O+) A weak acid does not dissociate completely in aqueous solution to form hydronium ion A strong base dissociates completely in aqueous solution to form hydroxide ions A weak base does not dissociate completely in aqueous solution to form hydroxide ions Weak acids and bases; Ka Weak acids and weak bases always exist as conjugate acid-base pairs in an aqueous solution as represented below. Strength of an acid (amount of hydrogen ion released when a given amount of acid is dissolved in water) is described by Ka: Water is always part of the reaction, so written correctly, Ionization of H2O (Kw) and pH Lets quantitatively examine the dissociation of water: Molar concentration of water is 55.5 M, and so Kw is called the ionization constant for water at 25oC: Must define a quantity to express hydrogen ion concentrations…pH [10-7] = 7 Henderson-Hasselbalch (weak acids) Equation to connect K to pH of solution containing a both acid and base. Henderson-Hasselbalch (Cont’d) Henderson-Hasselbalch equation From this equation, we see that when the concentrations of weak acid and its conjugate base are equal, the pH of the solution equals the pKa of the weak acid when pH < pK , the weak acid predominates a when pH > pK , the conjugate base predominates a pKa and Ka and strength of acid pH calculations pH calculations Buffers: solutions resistant to pH change Buffer: a solution containing either a weak acid and its salt (conjugate base) or a weak base and its salt (conjugate acid), which is resistant to changes in pH. Examples of acid-base buffers are solutions containing (weak acid and its salt or conjugate base) CH3COOH and CH3COONa H2CO3 and NaHCO3 NaH2PO4 and Na2HPO4 How buffers work: conjugate base If a weak acid (HA) is mixed with its conjugate base (A-), both will remain in solution. The acid and conjugate base may react with one another HA + A- → A- + HA But they just trade places, and the concentrations of [HA] and [A-] remain the same How buffers work If a strong base is added to a buffer: the weak acid will give up its H+ to neutralize the base (OH-) into water (H2O). Since the added OH- is consumed by this reaction, the pH will change only slightly. If a strong acid is added to a buffer: the weak base will react with the H+ from the strong acid to form the weak acid HA. The H+ gets absorbed by the A- instead of reacting with water to form H3O+ (H+), so the pH changes only slightly. Buffers Buffers are used to maintain a stable pH in a solution, because they can neutralize small quantities of additional acid of base. Buffers are useful over specific pH ranges. For example, pH range of common buffering agents: Buffer citric acid acetic acid KH2PO4 borate CHES pKa 3.13., 4.76, 6.40 4.8 7.2 9.24 9.3 Useful pH range 2.1 to 7.4 3.8 to 5.8 6.2 to 8.2 8.25 to 10.25 8.3 to 10.3 A buffer is effective in a range of about +/- 1 pH unit of the pKa of the weak acid How buffers work (e.g. acetic acid, pKa 4.8) Consider a buffer solution made by dissolving sodium acetate (CH3COONa) into acetic acid (CH3COOH). Acetic acid is the acid, while the sodium acetate dissociates in solution to yield the conjugate base (acetate ion: CH3COO-). Note Na is not shown because it is a spectacle ion. A strong base would be ‘neutralized’ by acetic acid. CH3COOH(aq) + OH-(aq) ⇆ CH3COO-(aq) + H2O(aq) A strong acid would be ‘neutralized’ by the acetate ion. CH3COO-(aq) + H+(aq) ⇆ CH3COOH(aq) Titration Curves Acid-base titrations are monitored by the change of pH as titration progresses to reach an equivalence point. Equivalence point: the point in an acid-base titration at which enough acid has been added to exactly neutralize the base (or vice versa) Titration Curves acids-bases (how its done) titrant: solution of a known concentration of an acid or base. This is used to determine the unknown concentration of base or acid. analyte: unknown concentration of acid or base equivalence point: moles of base = moles of acid Titrant: NaOH Analyte: acetic acid How Do We Choose a Buffer? By understanding the relationship between pH and buffering capacity for the ionic environment needed for the biological samples being studied Phosphate buffer, pka = 7.2 Selecting a Buffer The following criteria are typical suitable pKa no interference with the reaction or detection of the assay suitable ionic strength suitable solubility Making Buffers in the Laboratory: Buffer Capacity Buffering capacity is related to the concentrations of the weak acid and its conjugate base the greater the concentration of the weak acid and its conjugate base, the greater the buffer capacity Naturally Occurring Buffers H PO -/HPO 2- is the principal buffer in cells 2 4 4 H CO /HCO - is an important (but not the only) buffer 2 3 3 in blood hyperventilation can result in increased blood pH hypoventilation can result in decreased blood pH Importance of regulating blood pH Buffer Example What is the new pH if add 1.5mL of 2.0M HCl to (a) 10.0mL water or (b) 10.0mL of a mixture of 0.55M carbonic acid and 0.55M sodium hydrogen carbonate (carbonic acid pKa is 6.37)? (a) Assuming 100% dissociation of strong acid: New M(HCl) = (0.0015L*2.0M)/0.0115L = 0.261M pH = -log[H3O+] = 0.583 ~ 0.58 Buffer Example What is the new pH if add 1.5mL of 2.0M HCl to (a) 10.0mL water or (b) 10.0mL of a mixture of 0.55M carbonic acid and 0.55M sodium hydrogen carbonate (carbonic acid pKa is 6.37)? (b) n(HCl) = 0.0015L*2.0M = 0.003mol This HCl will react with an equivalent number of moles of sodium hydrogen carbonate to produce carbonic acid old: n(HCO3–) = n(H2CO3) = 0.0100L*0.55M = 0.0055mol new: n(HCO3–) = 0.0055mol-0.003mol = 0.0025mol new: n(H2CO3) = 0.0055mol+0.003mol = 0.0085mol 𝑝𝐻 =𝑝 𝐾 𝑎 + log ([ − 3 [ 𝐻𝐶 𝑂 ] 𝐻2 𝐶 𝑂3] ) =6.37 +log =5.84 ( 0.0025 0.0085 ) Note: the ratio of the moles is the same as the ratio of the molarities.
