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IL. Titrating HCISolutlon with NaOH Solution III. Titrating HC2H2O2 Solution with NaOH Solution
II. Titrating HC2H2O2 Solution with NaOHSolution concentration of NaOH solution, M concentration of HC2H3O2 solution. M
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HCl(aq)+H2O(t)→H1O∗( aq )+C (aq) (Fq. 2) Sodium hydroxide is a water-soluble ionic comporand that completely dissociates into sodium ions (Na+)and hydronde ions (OH=. Hence, NaOH is a strong base. We can more accurately represent the neutralization reaction of HCl with an NaOH solution in the complete ionic equation shown in Equation 3. H3O+(aq)+Cl−(aq)+Na+(aq)+OH−(aq)→2H2O(I)+Na+(aq)+Cl−(aq) (Eq. 3) Sodium ions and chloride ions (Cl∘) are spectator ions in this reaction, appearing unchanged as both reactants and products. Therefore, we can cancel them from Equation 3 to create the net ionic equation, Equation 4. H3O+(aq)+OH−(aq)=2H2O(t) (Fq. 4) We can monitor any acid-base titration by following changes in HO+ concentration in solution. This is convenient, because H3O∗ concentration is related to the pH of the solution as defined by Equation 5 . Note that the square brackets represent the molarity (M,mol/L) of H3O∗. pH=−log[H3O+] We monitor pH during a titration using a pH meter equipped with an electrode responsive to H3O" concentration in solution. Initially, the pH of the HCl solution will be low, because there are as many moles of H3O+present as there are moles of HCZ. As we add NaOH solution, the reaction in Equation 4 occurs, decreasing [H, O+] This causes a corresponding increase in solution pH. This neutralization reaction is complete at the equivalence point, when the number of moles of OH−added as titrant equals the number of moles of H3O+originally present in solution. The pH at the equivalence point is established by the components of the titration mixture. For the titration of HCl with NaOH, a strong acid with a strong base, Equation 3 shows that the only species present at the equivalence point are Na+,Cl−, and H2O. Beciuse neither Na+nor Cl−react with water (and are, therefore, netral ions), the pH of the titration mixture at the equivalence point is established by the dissociation of water, shown in Equation 6 . 2H2O(t) wi H3O+(aq)+OH−(aq) (Eq,6) The equilbrium constant expression representing this reaction is shown in Equation 7. At 25∘C, the dissociation constant for water is 1.0×10−14. Ks=[H3O−1OH−]=1.0×10−14 [Eq,7) At the equivalence point of our NaOH and HCl titration, H3O+and OH− concentrations are equal (Equation 6). Therefore, [H3O+]=(f.0×10−14)1/2= 1×10−7M, and the pH of the solution is 7 (Equation 5 ). If we continue to add NaOH solution to the titration mixture after the cquivalence point, the pH will continue to increase.
Figure 1 A fypioal tifnation carre for a strong acid arith 0.100MNaOH If we plot the pH of a titration mixture versus volume of titrant added. we obtain a graph called a titration curve. Figure 1 shows a typical curve for the titration of a strong acid with NaOH solution. We can locate the equivalence point of the titration by drawing a vertical line through the midpoint of the steep portion of the curve. The titrant volume and pH of the solution at the equivalence point correspond to the x and y cootdinates, respectively, of the point where the vertical line intersects the titration curve. II. Thtrating Weak Acid with Most acids are weak. When a weak acid is discolved in water, only a small a Strong Base percent of the acid molecules dissociate to produce Hu, Acetic+acid ( HC1H3O2 ), the pungent component of vinegar, is a typical weak acd. We can represent its dissociation in water by Equation 8 . HC2H3O2(aq)+H2O(a)=H3O+(aq)+C2H3O2−(aq) The equilibrium constant expression for the dissociation of HC2H3O2 is shown in Equation 9. The dissociation constant for H2H7O2 at 25∘C is 18×10−5 Ka=HC1H3O2[H3O+CC2H3O2−1)=1.5×10−5 (Eq.9) The pHH of an HC2H3O2 solution is higher than that of an HCl solution. of equal concentration, because only about 5% of the HC2H3O2 molecules dissociate, whereas essentially all of the HCl molecules do so. If we titrate an HC2H3O2 solution with NaOH solution, the reaction shown in Equation 10 occurs. HC2H3O2(aq)+NaOH(aq)→NaC2H3O2(aq)+H2O(f) (Eq. 10) The complete ionic equation (Equation 11) and net ionic equation (Equation 12) for this reaction are shown below.
Figure 2 A fypioal titration curoe for a resak acid with 0.100MNaOH HC2H3O2(aq)+Na+(aq)+OH−(aq) →Na+(aq)+C2H3O2−(aq)+H2O(t)HC2H3O2(aq)+OH−(aq)→C2H3O2−(aq)+H2O(l) As we add NaOH solution to the HC2H3O2 solution, HC2H3O2 is converted into acetate ions (C2H3O2−). The solution pH depends upon the relative concentrations of these two species (Equation 9). As with the titration of HCl solution with NaOH solution, the pH of the HC2H2O2 NaOH titration mixture at the equivalence point is established by the components present at that point C2H3O2−,Na+, and H2O. However, because C2H3O2 - ions are the anions of a weak acid, they react with H2O, or hydrolyze, as shown in Equation 13. C2H3O2−(aq)+H2O(l)=HC2H3O2( aq )+OH−(aq) The equilibrium constant expression, Kh at 25∘C for the hydrolysis reaction is shown in Equation 14. Kh=[C2H3O2−][OH−][HC2H3O2]=5.6×10t00 The other equilibrium that could occur at this equivalence point is the dissociation of water, shown in Equation 6. Because Kb is larger than Kw, the hydrolysis reaction takes precedence. Therefore, the pH of the reaction mixture at the equivalence point is greater than 7.0, due to the presence of OH−produced by the hydrolysis of C2H3O2 ? A typical titration curve for the titration of a weak acid with 0.100M NaOH is shown in Figure 2. Caleulating H3O+and We can calculate the molarity of H3O+and OH−from the pH at any point DH−Concentrations on a titration curve. Suppose we find from our titration curve that the or Points on a solution pH is 4.60 when we have added 10.0 mL of 0.100MNaOH to itration Curve 25.0 mL of an HC2H3O2 solution. To determine [H3O+]and [OH−]for
the mixture at this point, we substitute 4.60 into Equation 5 , and solve for [H3O6]. 4.60−4.60[H3O+]=−log[H3O+]=log∣∣H3O+∣∣=antilog−4.60=2.5×10−4M Using Equation 7, we can calculate OHH−]. ∣∣OH−∣∣=2.5×10−31.0×10−14=4.0×10−111M Using the same process, we can calculate [H3O+]and [OH−]for any point on the titration curve. - Calibrate a pH meter - Fill buret with standardized NaOH solution - Titrate HCl solution with standardized NaOH solution, monitoring the reaction mixture pH as titrant added - Titrate HC2H3O2 solution with standardized NaOH solution, monitoring reaction mixture pH as titrant added CHEMICAI AIRHT 0.100M acetic acid-irritant 0.100M hydrochloric acid-toxic and corroslve 0.100M sodium hydroxide-toxie and corrosive C. A it \& it 8 it Wear departmentally approved sately goggles whilie doing this experiment. Glass electrodes are fragllo and expenslve. Do not bump the glass membrane against anything solid. NOTE: Yout laboratory instructor will deseribe and demonstrate a satisfactory meshod for calibrating the pH mesers avaiable in your laboratory. Depending upon the model of pH meters available, you will perform this celibration using either one or Wwo buifter solutions. Consult your laboratory instructor to make sure your electrode is property vented.