# formic acid and naoh titration

Because this value is less than 5% of 0.0500, our assumption is correct. Again, because the concentration of HF is so small, we will consider the initial $\left[{\text{H}}_{3}{\text{O}}^{\text{+}}\right]$ to be 1 $\times$ 10−7M from the ionization of water. Thus, pick an indicator that changes color in the acidic range and brackets the pH at the equivalence point. (17-29) ___5.3635 pOH of solution = 14 - pH = 14 -8.6364 Answer: 5.3635 Each 'hump' has its own half-equivalence point. Figure 3 shows us that methyl orange would be completely useless as an indicator for the CH3CO2H titration. Question: Were going to titrate formic acid with the strong basic,{eq}NaOH {/eq}. Acid-Base Indicators. Figure 2. Substances such as phenolphthalein, which can be used to determine the pH of a solution, are called acid-base indicators. Think about where in the titration this brings you. At this point, the only hydronium ions left are those from the autoionization of water, and there are no OH− particles to neutralize them. $\text{pOH}=\text{-log}\left(5.3\times {10}^{-6}\right)=5.28$ The color change intervals of three indicators are shown in Figure 3. In the example, we calculated pH at four points during a titration. Therefore, $\left[{\text{H}}_{3}{\text{O}}^{\text{+}}\right]$ = 9.8 $\times$ 10−5M: pH = −log(9.8 $\times$ 10−5) = 4.009 = 4.01; mol OH− = M $\times$ V = (0.100 M) $\times$ (0.039 L) = 0.00390 mol, $\begin{array}{l}\\ \\ \left[\text{HA}\right]=\frac{0.00010\text{mol}}{0.0790\text{L}}=0.00127M\\ \left[{\text{A}}^{\text{-}}\right]=\frac{0.00390\text{mol}}{0.0790\text{L}}=0.0494M\end{array}$. There is initially 100 mL of 0.50 M formic acid and the concentration of {eq}NaOH {/eq} is 1.0 M.. A. Use [CO2H-la, instead of [H3O+]eg for the calculation. Table 1 shows data for the titration of a 25.0-mL sample of 0.100 M hydrochloric acid with 0.100 M sodium hydroxide. The pH range between 3.1 (red) and 4.4 (yellow) is the color-change interval of methyl orange; the pronounced color change takes place between these pH values. We use Kw to calculate the concentration. The ${\text{H}}_{3}{\text{O}}^{\text{+}}$ concentration in a 1.0 $\times$ 10−7M HF solution is: ${\text{H}}_{3}{\text{O}}^{\text{+}}$ = 1.0 $\times$ 10−7 + x = 1.0 $\times$ 10−7 + 0.9995 $\times$ 10−7 = 1.999 $\times$ 10−7M. | Assuming that the dissociated amount is small compared to 0.100, After 25.00 mL of NaOH are added, the number of moles of NaOH and CH, In (1), 25.00 mL of the NaOH solution was added, and so practically all the CH, After 37.50 mL of NaOH is added, the amount of NaOH is 0.03750 L $\times$ 0.100. The initial concentration of ${\text{H}}_{3}{\text{O}}^{\text{+}}$ is ${\left[{\text{H}}_{3}{\text{O}}^{\text{+}}\right]}_{0}=0.100M$. 0.00 mL, pH = 2.37. 8.364 1 50mmol Salt = 150mL = 3 log pH solution:ż [pKCO + pKa +logC]orz[14 + 3.75 + logą Answer: 8.6364 8) What is the pH after adding 100 mL of NaOH? The equivalence points of both the titration of the strong acid and of the weak acid are located in the color-change interval of phenolphthalein. The point of inflection (located at the midpoint of the vertical part of the curve) is the equivalence point for the titration. Using the assumption that x is small compared to 0.0500 M, ${K}_{\text{b}}=\frac{{x}^{\text{2}}}{0.0500M}$, and then: $x=\left[{\text{OH}}^{-}\right]=5.3\times {10}^{-6}$ This is the equivalence point, where the moles of base added equal the moles of acid present initially. Indicators for Strong Acid - Weak Base Titrations An aqueous solution of hydrochloric acid, HCl(aq), is a strong acid. Therefore, $\left[{\text{H}}_{3}{\text{O}}^{\text{+}}\right]$ = 3.13 $\times$ 10−3M: pH = −log(3.13 $\times$ 10−3) = 2.504 = 2.50; mol OH− = M $\times$ V = (0.100 M) $\times$ (0.020 L) = 0.00200 mol. initially 100. mL of 0.50 M formic acid and the concentration of Calculate pH at the equivalence point of formic acid titration with NaOH, assuming both titrant and titrated acid concentrations are 0.1 M. pK a = 3.75.