Lecture2

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II Water cont. D. Properties of water (cont.) 2. Water has a high melting point, boiling point and heat of vaporization as compared to most solvents a-the amount of H-bonds that one water molecule is capable of forming gives it a great amount of internal cohesion - each H2O molecule can form 4 H-bonds 1) The oxygen from the H2O molecule can accept two hydrogens from two other H2O molecules for a total of two H-bonds 2) Each hydrogen of the water molecule can form an H-bond for a total of two more giving a grand total of four

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b- Although water can form four H-bonds maximum it does not always do so: - When would a single water molecule most stably form four H-bonds? - What happens in liquid water?

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c- this is significant because water’s stability is a major reason why cells exist in an aqueous environment d- to melt ice or evaporate water takes some thermal energy H2O (solid) H20 (liquid) ∆H= +5.9 kJ/mol H2O (liquid) H2O (gas) ∆H= +44.0 kJ/mol - both processes are still spontaneous with a -∆G at room temperature. Why? e- water can absorb large amounts of heat making it a good insulator for the earth as well as our bodies -Sweating- water absorbs heat from our bodies and takes it to the outer surface to be dissipated

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3. Water is more dense as a liquid then a solid a- this is why ice floats - if it didn’t then lakes would freeze from the bottom to the top and life would not sustain itself b- solid water forms maximum H-bonds more stably so the water molecules are spaced apart more while in liquid we have the “flickering clusters” so that overall more water molecules can be packed into the same space making it more dense

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The crystal structure of water is less dense because H-bonds stabilize a tetrahedral configuration

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4. Water has a high surface tension a- the cohesion of H-bonds allows for high surface tension b- this also helps with stability of environment and sustaining life c- the cohesive forces between molecules down into a liquid are shared with all neighboring atoms. Those on the surface have no neighboring atoms above, and exhibit stronger attractive forces upon their nearest neighbors on the surface. This enhancement of the intermolecular attractive forces at the surface is called surface tension.

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E. Beyond H-bonds there are other noncovalent interactions that are of biological significance 1. Ionic interactions 2. Hydrophobic interactions 3. Van der Waals forces- a weak interaction that occurs when two atoms are in close proximity to each other a- as electrons of one atom are randomly moving around its nucleus a transient dipole can be created inducing the opposite transient dipole in a neighboring atom. the electrostatic attraction brings the nuclei together to a point but their electron clouds will repel each other if they get too close the balanced point is known a Van der Waals contact

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Why is there a difference between covalent and van der waals radii?

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4. All of these noncovalent interactions are biologically significant a- remember that alone each of these types of interactions are weak when compared to covalent bonds b- the strength comes in the number of these interactions that are possible c- to separate two molecules from each other that share a lot of noncovalent interactions means that you must disrupt all of those interactions at the same time - the probability of this occurring spontaneously is not that likely and generally will require energy in standard conditions - therefore for most biological molecules its native form is the one that maximizes its noncovalent interactions

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Hemoglobin- when it is crystallized water molecules are often attached so tightly that they crystallize as well Water- red circles Heme group- red lines α- subunit β-subunit

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III Acids and Bases A. The pH scale 1. pure water is slightly ionized H2O H+ + OH- a- pure water has a tendency to dissociate and donate a proton to a solution b- this proton does not usually exist freely in pure water but instead forms a hydronium ion (H3O+) H O H O H O+ H + OH- H H H

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c- this ionization can be measured through its electrical conductivity - pure water carries electrical current very efficiently d- movement of hydrogen and hydroxide ions is very fast as compared to other ions - this is why acid/base reactions in aqueous solutions occur so quickly - this quickness can be accounted for by proton hopping

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2. Equilibrium Constant ( Keq) of water a- the equilibrium constant of any reaction is the concentration of products divided by the concentration of reactant at equilibrium for a given temperature A + B C + D Keq = {C}{D} {A}{B} - this value is defined in terms of molarity (M) since we are measuring concentrations

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b- to calculate the equilibrium constant of pure water at 25°C we would set up our equation as follows: Keq = {H+}{OH-} {H2O} - the concentration of ionized water molecules is very small at this state and is negligible when compared to pure water molecules - concentration of pure water is 55.5M 1,000 g/L amount of water per liter = 55.5M 18.015 g/mol molecular weight of water

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- therefore: Keq = {H+}{OH-} 55.5M - if you solve for the concentration of hydrogen and hydroxide ions you get: (55.5M)(Keq) = {H+}{OH-} = Kw - Kw is the ion product of water at 25°C - the Keq of pure water can be measured through electrical-conductivity experiments and is 1.8 X 10-16 M thus the Kw for pure water at 25°C is : (55.5M)(1.8 X 10-16 M) = 1.0 X 10 -14 M2 = Kw

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c- in pure water the balance of hydrogen and hydroxide ions are equal at equilibrium ( this is the foundation for neutral pH as we will see in a minute) {H+}{OH-} = 1.0 X 10-14 M2 = (1.0 X 10-7 M )(1.0 X 10-7 M) - the Kw is constant such that any changes in H+ concentration must be balanced by a change in OH- concentration. i.e.: If one goes up the other must do down

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3. The pH scale is a reflection of the hydrogen and hydroxide ions just discussed a- pH reflects the concentration of hydrogen ions in an aqueous solution and thus by default the concentration of hydroxide ions as well - Why? b- the ion product of water (Kw) is the basis for the pH scale c- the scale reflects water ion concentration in a range from 1.0 M H+ to 1.0 M OH- d- the symbol p denotes “ negative logarithm of”. - so for pure water at 25°C pH = log 1 = -log {H+} {H+} pH = log 1 = log 1 + log 107 = 0 + 7 = 7 1X10-7M

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e- thus setting the number 7 as the neutral value in the pH scale is directly derived from the concentration of ion product of pure water at 25° C f- the pH scale is logarithmic so a change in the scale by a value of 1 denotes an increase or a decrease of hydrogen ions by a factor of 10. - thus cola at pH 3 has a hydrogen ion concentration 10,000 times that of blood at pH 7.4 g- pH can be measured using dyes like litmus or phenol red and with a pH meter that uses a glass electrode that is specifically sensitive to hydrogen ion concentration. The pH then compares its reading to a standard to get an accurate value - accurate measurement and control of pH is critical in biological systems or studies as it can have drastic effects on a system

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B. Buffer Systems 1. Weak Acids and Bases a- Acids = proton donors Bases = proton acceptors Conjugate acid-base pair = a proton donor and its corresponding proton acceptor (Acetic Acid) CH3COOH H+ + CH3COO- (Acetate) conjugate acid-base pair

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b- All acids have a tendency to lose their protons in an aqueous solution. Some lose it more readily (strong acid) while others do not (weak acids) - strong acids ionize completely in an aqueous solution as do strong bases - this tendency to ionize can be defined by the Keq for that solution. Consider the following acid HA HA H+ + A- Keq = {H+}{A-} = Ka {HA}

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c. Ka = the dissociation constant or ionization constant - used for ionization reactions - for Ka the higher the value of the stronger the acid i.e.: more of the substance is ionized d. pKa = log 1/Ka = -log Ka - pKa defines the relative strength of a weak acid or base and can be determined experimentally as the pH at the mid-point in a titration curve (we will discuss this in a moment - as we will see pKa is analogous to pH - for pKa the lower the value the stronger the acid

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2. Titration curves a- Titration can be used to determine the amount of an acid in a given solution - take a measured volume of acid solution and titrate it with a strong base (NaOH) - measure the pH as the base is added in increments b- the titration curve is a plot of the pH values against the amount of NaOH added

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c- from the titration curve the pKa of a weak acid can be determined as well as the pH range of the acids buffering power d- For example: The titration of a .1 M solution of acetic acid (HAc) with .1 M NaOH at 25°C was done. H2O H+ + OH- HAc H+ + Ac- - both must conform to their characteristic equilibrium constants Kw = {H+}{OH-} = 1X10-14 M2 Ka = {H+}{Ac-} = 1.74 X 105 M {HAc}

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- before any NaOH is added both molecules have ionized to some extant - as NaOH is introduced it will fully ionize and the OH- will bind up the free H+ in the solution to form H2O until the equilibrium relationship with water is met - As the OH- removes the free H+ the HAc will further dissociate to release more H+ to satisfy its own equilibrium requirements -The midpoint of the titration curve when .5 equivalent of NaOH has been added is when half the acetic acid has dissociated and thus the {HAc} = { Ac-} - conjugate acid-base pair - this is the point when the pH = pKa

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- titration ends when the HAc has been fully ionized e- titration curves allow you to graphically see how a weak acid and its conjugate base can act as a buffer - different weak acids or bases would have different strengths and thus different curves and different buffer regions

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3. Buffers- aqueous systems that tend to resist changes in pH when small amounts of acid (H+) or base (OH-) are added. a- buffers consist of a weak acid (the proton donor) and its conjugate base (the proton acceptor) b- the flat area of a titration curve represents its buffering zone - this is the area where an addition of hydrogen or hydroxide ions will not cause a major shift in the pH of the system - the midpoint of the titration curve is the point when the buffering power of a system is maximized when pH = pKa

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c- In a buffer whenever H+ or H- is added the result is a small change in the ratio of the relative concentrations of the weak acid and its anion and thus a small change in pH - realize that the sum of the buffer components does not change only their ratio i.e.: If the concentration of one part goes up the other must go down d- each conjugate acid-base pair has a pH zone in which it is an effective buffer which is generally 1 pH unit above and below its pKa. H2PO4-/HPO4-2 pKa = 6.86 Buffer Zone = pH 5.9-7.9 NH4+/NH3 pKa = 9.25 Buffer Zone = pH 8.3-10.3

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acetic acid-acetate pair as a buffer system

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4. Henderson-Hasselbalch equation a- a useful way to express the dissociation constant of an acid b- it is derived by the following: Ka = {H+}{A-} {HA} 1. Solve for {H+} {H+} = Ka {Ha} {A-} 2. Take the negative logarithm of both sides -log {H+} = -log Ka – log {Ha} {A-}

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3. Substitute in pH and pKa and you have pH = pKa – log {HA} {A-} 4. Invert the negative log to get the Henderson-Hasselbalch equation pH = pKa + log {A-} {HA} Proton acceptor Proton donor c- this equation fits the titration curve of all weak acids and it also explains why pH = pKa at the midpoint of titration because that is when the concentrations of the acid and its conjugate base are equal pH = pKa + log 1 = pKa + 0 = pKa

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d- the H-H equation allows us to calculate the pH given the proper information, pKa given the proper information or the molar ratio of a proton donor and acceptor at a qiven pH and pKA 5. Buffer Systems a- buffering systems are very important in biological systems - most enzymes have a narrow pH range in which they function - blood pH must be tightly monitored - cytoplasm must be regulated

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b- many molecules can assist in buffering - amino acids have functional groups that serve as weak acids or weak bases - ammonia buffers urine c- two key systems used for buffering are the phosphate system and the bicarbonate system to buffer blood pH - remember that the pH you want to maintain must be near the pKa value for effective buffering For example the pKa of H2PO4-/ HPO4-2 is 6.86 thus it works well for cytoplasmic compartments and extracellular fluids which have a pH range of 6.9 to 7.4

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d- the bicarbonate system to buffer blood plasma (carbonic acid) H2CO3 H+ + HCO-3 (bicarbonate) -carbonic acid is formed from dissolved carbon dioxide and water CO2 (d) + H2O H2CO3 - carbon dioxide is normally a gas and thus there is a balance between the gaseous and dissolved form CO2 (g) CO2 (d) -in animals the bicarbonate system is a balance between blood conditions and lung capacity of CO2

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