in titration of 25 mL of the solution containing Cu2+ with the concentration of 0.01 molar , with SCN – (0.01 Molar) , if we use the indicator that changed its color in PSCN= 9. Calculate the error of titration PKSP(CuSCN) =13.4?

Answer 1

just the color change from the indicator. Thus, a careful selection of the indicator will reduce the indicator error. For example, if the equivalence point is at a pH of 8.4, then the phenolphthalein indicator would be used instead of Alizarin Yellow because phenolphthalein would reduce the indicator error. Common indicators, their colors, and the pH range in which they change color are given in the table above. When more precise results are required, or when the reagents are a weak acid and a weak base, a pH meter or a conductance meter are used. For very strong bases, such as organolithium reagent, metal amides, and hydrides, water is generally not a suitable solvent and indicators whose pKa are in the range of aqueous pH changes are of little use. Instead, the titrant and indicator used are much weaker acids, and anhydrous solvents such as THF are used. The approximate pH during titration can be approximated by three kinds of calculations. Before beginning of titration, the concentration of [ H + ] {\displaystyle {\ce {[H+]}}} is calculated in aqueous solution of weak acid before adding any base. When the number of moles of bases added equals the number of moles of initial acid or so called equivalence point, one of hydrolysis and the pH is calculated in the same way that the conjugate bases of the acid titrated was calculated. Between starting and end points, [ H + ] {\displaystyle {\ce {[H+]}}} is obtained from the Henderson-Hasselbalch equation and titration mixture is considered as buffer. In Henderson-Hasselbalch equation the [acid] and [base] are said to be the molarities that would have been present even with dissociation or hydrolysis. In a buffer, [ H + ] {\displaystyle {\ce {[H+]}}} can be calculated exactly but the dissociation of HA, the hydrolysis of A − {\displaystyle {\ce {A-}}} and self-ionization of water must be taken into account. Four independent equations must be used: [ H + ] [ OH − ] = 10 − 14 {\displaystyle [{\ce {H+}}][{\ce {OH-}}]=10^{-14}} [ H + ] = K a [ HA ] [ A − ] {\displaystyle [{\ce {H+}}]=K_{a}{\ce {{\frac {[HA]}{[A^{-}]}}}}} [ HA ] + [ A − ] = ( n A + n B ) V {\displaystyle [{\ce {HA}}]+[{\ce {A-}}]={\frac {(n_{{\ce {A}}}+n_{{\ce {B}}})}{V}}} [ H + ] + n B V = [ A − ] + [ OH − ] {\displaystyle [{\ce {H+}}]+{\frac {n_{{\ce {B}}}}{V}}=[{\ce {A-}}]+[{\ce {OH-}}]} In the equations, n A {\displaystyle n_{{\ce {A}}}} and n B {\displaystyle n_{{\ce {B}}}} are the moles of acid (HA) and salt (XA where X is the cation), respectively, used in the buffer, and the volume of solution is V. The law of mass action is applied to the ionization of water and the dissociation of acid to derived the first and second equations. The mass balance is used in the third equation, where the sum of V [ HA ] {\displaystyle V[{\ce {HA}}]} and V [ A − ] {\displaystyle V[{\ce {A-}}]} must equal to the number of moles of dissolved acid and base, respectively. Charge balance is used in the fourth equation, where the left hand side represents the total charge of the cations and the right hand side represents the total charge of the anions: n B V {\displaystyle {\frac {n_{{\ce {B}}}}{V}}} is the molarity of the cation (e.g. sodium, if sodium salt of the acid or sodium hydroxide is used in making the buffer). Redox titrations are based on a reduction-oxidation reaction between an oxidizing agent and a reducing agent. A potentiometer or a redox indicator is usually used to determine the endpoint of the titration, as when one of the constituents is the oxidizing agent potassium dichromate. The color change of the solution from orange to green is not definite, therefore an indicator such as sodium diphenylamine is used. Analysis of wines for sulfur dioxide requires iodine as an oxidizing agent. In this case, starch is used as an indicator; a blue starch-iodine complex is formed in the presence of excess iodine, signalling the endpoint.Some redox titrations do not require an indicator, due to the intense color of the constituents. For instance, in permanganometry a slight persisting pink color signals the endpoint of the titration because of the color of the excess oxidizing agent potassium permanganate. In iodometry, at sufficiently large concentrations, the disappearance of the deep red-brown triiodide ion can itself be used as an endpoint, though at lower concentrations sensitivity is improved by adding starch indicator, which forms an intensely blue complex with triiodide. Gas phase titrations are titrations done in the gas phase, specifically as methods for determining reactive species by reaction with an excess of some other gas, acting as the titrant. In one common gas phase titration, gaseous ozone is titrated with nitrogen oxide according to the reaction O3 + NO → O2 + NO2.After the reaction is complete, the remaining titrant and product are quantified (e.g., by Fourier transform spectroscopy) (FT-IR); this is used to determine the amount of analyte in the original sample. Gas phase titration has several advantages over simple spectrophotometry. First, the measurement does not depend on path length, because the same path length is used for the measurement of both the excess titrant and the product. Second, the measurement does not depend on a linear change in absorbance as a function of analyte concentration as defined by the Beer-Lambert law. Third, it is useful for samples containing species which interfere at wavelengths typically used for the analyte. Complexometric titrations rely on the formation of a complex between the analyte and the titrant. In general, they require specialized complexometric indicators that form weak complexes with the analyte. The most common example is the use of starch indicator to increase the sensitivity of iodometric titration, the dark blue complex of starch with iodine and iodide being more visible than iodine alone. Other complexometric indicators are Eriochrome Black T for the titration of calcium and magnesium ions, and the chelating agent EDTA used to titrate metal ions in solution. Zeta potential titrations are titrations in which the completion is monitored by the zeta potential, rather than by an indicator, in order to characterize heterogeneous systems, such as colloids. One of the uses is to determine the iso-electric point when surface charge becomes zero, achieved by changing the pH or adding surfactant. Another use is to determine the optimum dose for flocculation or stabilization. An assay is a type of biological titration used to determine the concentration of a virus or bacterium. Serial dilutions are performed on a sample in a fixed ratio (such as 1:1, 1:2, 1:4, 1:8, etc.) until the last dilution does not give a positive test for the presence of the virus. The positive or negative value may be determined by inspecting the infected cells visually under a microscope or by an immunoenzymetric method such as enzyme-linked immunosorbent assay (ELISA). This value is known as the titer. Different methods to determine the endpoint include: Indicator: A substance that changes color in response to a chemical change. An acid–base indicator (e.g., phenolphthalein) changes color depending on the pH. Redox indicators are also used. A drop of indicator solution is added to the titration at the beginning; the endpoint has been reached when the color changes. Potentiometer: An instrument that measures the electrode potential of the solution. These are used for redox titrations; the potential of the working electrode will suddenly change as the endpoint is reached. pH meter: A potentiometer with an electrode whose potential depends on the amount of H+ ion present in the solution. (This is an example of an ion-selective electrode.) The pH of the solution is measured throughout the titration, more accurately than with an indicator; at the endpoint there will be a sudden change in the measured pH. Conductivity: A measurement of ions in a solution. Ion concentration can change significantly in a titration, which changes the conductivity. (For instance, during an acid–base titration, the H+ and OH− ions react to form neutral H2O.) As total conductance depends on all ions present in the solution and not all ions contribute equally (due to mobility and ionic strength), predicting the change in conductivity is more difficult than measuring it. Color change: In some reactions, the solution changes color without any added indicator

Answer 2

To calculate the error of titration, the formula is Error = pH - pKw - pKa. First, calculate the pH using the formula pH = pKSP + log([Cu2+]/[SCN-]). Substituting the given values gives you pH = 13.4 + log(0.01/0.01) = 13.4. Next, substitute this pH value into the error formula: Error = 13.4 - 14 - 9 = -9.6. Therefore, the error of titration in this case is -9.6.