1pH Acids, bases, pK Conjugate acid-base pairs Calculate pK from a titration curveBuffersHanderson-Hasselbalch EquationPractice some problems
2Ionization of Water: Quantitative Treatment Concentrations of participating species in an equilibrium processare 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 25C.[H2O] can be determined from water density, it is 55.5 M.Ionic product of water:In pure water [H+] = [OH–] = 10–7 M
3What is pH?pH is defined as the negative logarithm of the hydrogen ion concentrationSimplifies equationsThe pH and pOH must always add to 14In neutral solution, [H+] = [OH–] and the pH is 7pH can be negative ([H+] = 6 M)pH = -log[H+]
5pH of Some Common Liquids FIGURE 2-15 The pH of some aqueous fluids.
6TitrationThe process of gradually adding known amounts of reagent to a solution with which the reagent reacts while monitoring the results is called a titration.
7Equilibrium constant=ionization constant= dissociation constant Each acid has a characteristic tendency to lose its protons in an aqueous solution.The stronger the acid the greater the tendency.The tendency of any acid (HA) to lose a proton and form its conjugate base (A+) is defined by the equilibrium constant K for the reversible rxn.
8Equilibrium constant=ionization constant= dissociation constant HA-----> H+A K= [H][A]/[HA]The relative strengths of weak acids and bases are expressed as their dissociation constant, which expresses the tendency to ionize.
9pKa measures aciditypKa = –log Ka (strong acid large Ka small pKa)FIGURE 2-16 Conjugate acid-base pairs consist of a proton donor and a proton acceptor. Some compounds, such as acetic acid and ammonium ion, are monoprotic; they can give up only one proton. Others are diprotic (carbonic acid and glycine) or triprotic (phosphoric acid). The dissociation reactions for each pair are shown where they occur along a pH gradient. The equilibrium or dissociation constant (Ka) and its negative logarithm, the pKa, are shown for each reaction. *For an explanation of apparent discrepancies in pKa values for carbonic acid (H2CO3), see p. 63.
10Buffers are mixtures of weak acids and their anions (conjugate base) Buffers resist change in pHAt pH = pKa, there is a 50:50 mixture of acid and anion forms of the compoundBuffering capacity of acid/anion system is greatest at pH = pKaBuffering capacity is lost when the pH differs from pKa by more than 1 pH unit
11FIGURE 2-17 The titration curve of acetic acid FIGURE 2-17 The titration curve of acetic acid. After addition of each increment of NaOH to the acetic acid solution, the pH of the mixture is measured. This value is plotted against the amount of NaOH added, expressed as a fraction of the total NaOH required to convert all the acetic acid (CH3COOH) to its deprotonated form, acetate (CH3COO–). The points so obtained yield the titration curve. Shown in the boxes are the predominant ionic forms at the points designated. At the midpoint of the titration, the concentrations of the proton donor and proton acceptor are equal, and the pH is numerically equal to the pKa. The shaded zone is the useful region of buffering power, generally between 10% and 90% titration of the weak acid.
12If you change the labeling of axis, what happens to the titration curve?
13Weak acids have different pKas FIGURE 2–18 Comparison of the titration curves of three weak acids. Shown here are the titration curves for CH3COOH, H2PO-4, and NH+4 . The predominant ionic forms at designated points in the titration are given in boxes. The regions of buffering capacity are indicated at the right. Conjugate acid-base pairs are effective buffers between approximately 10% and 90% neutralization of the proton-donor species.
15More Acids and Bases Acid Conjugate Base HA > A + H+HA > A- + H+HA > A2- + H+Note: In some cases, the conjugate base has a – charge but in others it does not!IMPORTANT POINT: The conjugate base ALWAYS has one less + charge than the acid
16Amino acids titration curves An amino acid can act as a base or an acid. Such substances are called to be amphoteric, and are referred to as ampholyte.A crystalline amino acid dissolved in water is ionized, and can act as a weak acid or base.2 titratable groups: -COOH and -NH3Thus, amino acids have 2 dissociation constants and plots with 2 stages.Depending on the medium’s pH, an amino acid can have a (+), (-) and a net “0” charge.
17FIGURE 3-9 Nonionic and zwitterionic forms of amino acids FIGURE 3-9 Nonionic and zwitterionic forms of amino acids. The nonionic form does not occur in significant amounts in aqueous solutions. The zwitterion predominates at neutral pH. A zwitterion can act as either an acid (proton donor) or a base (proton acceptor).
19FIGURE 3-11 Effect of the chemical environment on pKa FIGURE 3-11 Effect of the chemical environment on pKa. The pKa values for the ionizable groups in glycine are lower than those for simple, methyl-substituted amino and carboxyl groups. These downward perturbations of pKa are due to intramolecular interactions. Similar effects can be caused by chemical groups that happen to be positioned nearby—for example, in the active site of an enzyme.
27pIIsoelectric pH or inflection point: The midpoint of the titration curve; the pH at which a molecule has a zero charge (zwitterion or dipolar form)For a simple amino acid with only an a-carboxyl and an a-amino group, the pI is determinedpI=(pK1+pK2)/2However, for an amino acid with three or more ionizable groups, you must avoid the trap of thinking that pI is the average of pKa values:pI=(pKn+pKn+1)/2pKn and pKn+1 are the two pKa values that describe the ionizaton of the species with a zero net charge; that is the the first ionization that adds a proton to the neutral species and gives it a net charge of -1 and the first ionization that remove a proton from the neutral species and gives it a net charge of -1
28FIGURE 3-10 Titration of an amino acid FIGURE 3-10 Titration of an amino acid. Shown here is the titration curve of 0.1 M glycine at 25°C. The ionic species predominating at key points in the titration are shown above the graph. The shaded boxes, centered at about pK1 = 2.34 and pK2 = 9.60, indicate the regions of greatest buffering power. Note that 1 equivalent of OH– = 0.1 M NaOH added.
31At what point(s):glycine will be present predominantly as the species +H3N-CH2-COOH?is the average net charge of glycine +1?is the pH is equal to the pKa of the carboxyl group?does glycine have its maximum buffering capacity?is the average net charge zero?is the predominant species +H3N-CH2-COO- ?is the net charge if Glycine -1?do the predominant species consist of a 50:50 mixture of +H3N-CH2-COOH and +H3N-CH2-COO- ?is the predominant species +H2N-CH2-COO- ?What point corresponds to the pI?Which points have the worst buffering efficiency?
32Acidosis and alkalosis Blood pH< >acidosisBlood pH< > alkalosisRespiratory and metabolicRespiratoryA change in acid-base status induced by altered respirationMetabolicA change in acid-base status induced by metabolic problems (diabetes, alcoholism, poisoning)
33Blood pH 7.35-7.45 Buffers Phosphate buffer (cell cytoplasm) Bicarbonate buffer (plasma)Bicarbonate buffer system is unique in the sense thatH2C03----> CO2(d) + H2OCO2 is a gas under normal conditions
34FIGURE 2-20 The bicarbonate buffer system FIGURE 2-20 The bicarbonate buffer system. CO2 in the air space of the lungs is in equilibrium with the bicarbonate buffer in the blood plasma passing through the lung capillaries. Because the concentration of dissolved CO2 can be adjusted rapidly through changes in the rate of breathing, the bicarbonate buffer system of the blood is in near equilibrium with a large potential reservoir of CO2.
35moreThe pH of a HCO3 system depends on the H2CO3 and HCO3 donor/acceptor concentrations.H2CO3 depends on CO2(d)----> depends on CO2(g)Thus, the pH of a bicarbonate buffer is determined by the [HCO3 ]the solution and partial pressure of CO2 in the gas phase