Acid-Base Titration
Acid-base titration is a technique used to determine the concentration of an acid or base in a solution. It involves slowly adding a solution of known concentration (titrant) to the solution of unknown concentration (analyte) until the reaction between the two is complete. The point at which the reaction is complete, known as the equivalence point, is used to calculate the concentration of the analyte.
8 Key excerpts on "Acid-Base Titration"
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
...An Acid-Base Titration involves the determination of the concentration of an acid or base by exactly neutralizing the acid or base with an acid or base of known concentration. Neutralization is a type of chemical reaction in which an acid and a strong base react completely with each other resulting in a solution that is neither acidic nor basic. The net ionic equation for the neutralization reaction between an acid and a strong base is just the reaction of hydronium ion with hydroxide ion to give water, which is the reverse of the autoionization of water. H 3 O + + OH – → 2H 2 O The complete reaction of the acid and a base allows for determination of the concentration of the unknown acid or basic solution when the concentration of the other is known. A typical Acid-Base Titration begins by dispensing an accurately known volume of the acid or base whose concentration is unknown into a flask with a small amount of indicator. The indicator is designed to change color at the exact point that the reaction is complete. The titrant is contained in a calibrated burette shown in Fig. 5.10, a device used for dispensing accurately measured amounts of a solution. The burette is suspended above the flask containing the analyte and small volumes of the titrant are added to the analyte and indicator until the indicator changes color, signaling that the reaction is complete. This point in the titration is called the end point or equivalence point. When the endpoint of the reaction is reached, the volume of reactant consumed is measured and used to calculate the concentration of analyte as; Fig. 5.10 A typical Acid-Base Titration. In this experiment, the titrant in the burette is a base and the analyte in the flask is an acid...
eBook - ePub- James F. Pankow(Author)
- 2019(Publication Date)
- CRC Press(Publisher)
...For alkalimetric and acidimetric titrations, the titration curve is usually plotted as pH vs. amount (e.g., volume) of strong base or strong acid added. Here, however, since we want to discuss things in a general context, for the x -axis rather than volume, we will mostly use the value of C B − C A that is produced in the solution (or the related parameter f, defined below). If the concentration in the titrant is sufficiently large relative to the concentration of what is being titrated (say 20×), then near-constancy in the value of A T (or B T, etc.) in the solution can be assumed over the course of the titration: the volume of titrant added does not significantly increase the volume of solution being titrated so that A T (or B T, etc.) is not significantly lowered during the titration. This assumption simplifies computing titration curves. While the mathematics of titrations might be viewed as, well, tedious and confusing, it is fundamentally important for understanding how aqueous solutions behave in response to changes in acid or base content. First, most of the salt in the oceans is the result of the grand back-and-forth chemical neutralization that has occurred over geologic time of metal oxide bases from terrestrial solids reacting with added HCl from volcanism and vice versa (Schilling et al., 1978). Second, natural waters undergo titration changes whenever there are spills of strong acid or base, and when acid rain falls on a lake watershed system. Third, laboratory titrations are routinely used in analytical determinations of “alkalinity” in samples of: (1) natural water; and (2) water flowing through waste and drinking water treatment plants, to provide input information for calculations regarding pH control...
eBook - ePubBiermann's Handbook of Pulp and Paper
Volume 2: Paper and Board Making
- Pratima Bajpai(Author)
- 2018(Publication Date)
- Elsevier(Publisher)
...The format of the x -axis is going to be modified here, and the modified format will be used throughout the remainder of the book, unless stated otherwise. Instead of the ratio of acid to base, we will assume all of the acid reacts with all of the base. The x -axis will be reported in terms of percent of acid and percent of base form. One can use this to see where endpoints occur, but more generally, one can quickly determine the pH of a solution of mixed composition. Mixtures of acids or bases can be titrated sequentially if they differ by a factor of about 10 6 in their K a or K b values. Sometimes a species may be removed from solution to prevent interference. Other times a species may react with an added compound to change its acid or base characteristics. Examples of such titrations are given in the section on analysis of pulping liquors. 21.5. Reduction–Oxidation Titrations Reduction–oxidation (redox) titrations are performed in a manner similar to those of acid–base titrations. The titration is monitored by the potential (voltage) with respect to a standard electrode or by the use of indicators that are oxidized or reduced after the analyte. In redox titrations, equivalents are based on moles of electrons transferred in the reaction. A compound with a certain equivalent weight for an acid–base reaction may have a different equivalent weight in a redox reaction; it may even have an equivalent weight that varies depending on the particular redox reaction. In order for an analyte to be titrated with a titrant there should be at least 0.2–0.3 V difference between the standard two electrode potentials. This does not say the reaction will occur quickly, only that thermodynamically the reaction is predicted to occur. When the two materials have reacted, i.e., they are in equilibrium, both of their actual half reaction potentials are equal...
eBook - ePubChemistry
Concepts and Problems, A Self-Teaching Guide
- Richard Post, Chad Snyder, Clifford C. Houk(Authors)
- 2020(Publication Date)
- Jossey-Bass(Publisher)
...In a neutralization reaction taking place between an Arrhenius acid and an Arrhenius base, one product will always be ____________. Answer: H 2 O If the concentration of an acid solution is unknown, a basic solution of known concentration can be added slowly in measured amounts until the solution is neutralized. The concentration of the acid can then be determined. It is also possible to do the reverse, adding an acid solution to a base. This process of adding a basic solution to an acid or vice versa for neutralization is called titration. An indicator is used in the solution being titrated. The indicator shows one color at a specific level of hydrogen ion concentration and another color at another level of hydrogen ion concentration. The concentration of a basic solution can be determined by a process called ___________. Answer: titration The hydrogen ions of an acid react with the hydroxide ions of a base during neutralization to produce water. The ions other than H + and OH − that make up an acid or a base are spectator ions that do not enter the reaction. Complete and balance the following neutralization reaction. The complete ionic equation for this reaction is __________. The spectator ions are __________ and __________. Answer: CaCl 2 + 2H 2 O 2H + + 2Cl − + Ca 2+ + 2OH − → Ca 2+ + 2Cl − + 2H 2 O Ca 2+ ; 2Cl − The net ionic equation for a neutralization reaction is simply: Answer: H 2 O In a neutralization reaction, one H + ion reacts. with one OH − ion to form one molecule of H 2 O. (The ratio of H + to OH − is 1 to 1.) Molarity is a measure of concentration. Molarity equals moles of solute per liter of solution. Remember that M indicates moles/liter (this was introduced in Chapter 11). A 1 liter solution of 1 M H + ions could be expected to completely neutralize 1 liter of 1 M OH − ions...
eBook - ePubHandbook of Environmental Analysis
Chemical Pollutants in Air, Water, Soil, and Solid Wastes, Third Edition
- Pradyot Patnaik(Author)
- 2017(Publication Date)
- CRC Press(Publisher)
...17 Alkalinity Alkalinity of water is a measure of its acid-neutralizing ability. The titrable bases that contribute to the total alkalinity of a sample are generally the hydroxides, carbonates, and bicarbonates. However, other bases such as phosphates, borates, and silicates can also contribute to the total alkalinity. The alkalinity value depends on the pH end point designated in the titration. The two end points commonly fixed in the determination of alkalinity are the pH 8.3 and pH 4.5 (or between 4.3 and 4.9, depending on the test conditions). When the alkalinity is determined to pH 8.3, it is termed phenolphthalein alkalinity. In such alkalinity titration, phenolphthalein or metacresol purple may be used as an indicator. On the other hand, the total alkalinity is measured by titrating the sample to pH 4.5 using bromocresol green as the indicator. Alkalinity may also be determined by potentiometric titration to the preselected pH. An acid standard solution, usually 0.02 N H 2 SO 4 or HCl, is used in all titrations. The procedure for potentiometric titration is presented in Chapter 6. In this titration, a standard acid titrant is added to a measured volume of sample aliquot in small increments of 0.5 mL or less, that would cause a change in pH of 0.2 unit or less per increment. The solution is stirred after each addition and the pH is recorded when a constant reading is obtained. A titration curve is constructed, plotting pH versus cumulative volume titrant added. The volume of titrant required to produce the specific pH is read from the titration curve. Calculation Alkalinity, mg CaCO 3 /L = V × N × 50,000 mL sample where V is mL standard acid titrant used N is normality of the standard acid Since the equivalent weight of CaCO 3 is 50, the milligram equivalent is 50,000...
eBook - ePub- Peter Kam, Ian Power(Authors)
- 2015(Publication Date)
- CRC Press(Publisher)
...The ability of a substance to donate or accept a proton (i.e., to act as an acid or a base) depends on the concentration of H + ions in solution (pH of the solution) and the degree of dissociation (p K) of the substance. P H SYSTEM H + ion concentration may be measured in two ways: directly as concentrations in nanomoles per litre or indirectly as pH. pH is defined as the negative logarithm (to the base 10) of the concentration of hydrogen ions. The pH is related to the concentration of H + as follows: pH = log 10 1 [ H + ] pH = log 10 [ H + ] H + = 10 − pH pH = p K + log base/acid Table 8.1 Relationship between pH and hydrogen ion concentration pH Hydrogen ion concentration (nmol/L) 7.7 20 7.4 40 7.3 50 7.1 80 It is important to note that pH and hydrogen ion concentration [H + ] are inversely related such that an increase in pH describes a decrease in [H + ] (Table 8.1). However, the logarithmic scale is nonlinear and, therefore, a change of one pH unit reflects a 10-fold change in [H + ] and equal changes in pH are not correlated with equal changes in [H + ]. For example, a change of pH from 7.4 to 7.0 (40 nmol/L [H + ] to 100 nmol/L [H + ]) represents a change of 60 nmol/L [H + ], although the same pH change of 0.4, but from 7.4 to 7.8 (40 nmol/L [H + ] to 16 nmol/L [H + ]), represents a change of only 24 nmol/L [H + ]. BUFFERS A buffer is a solution consisting of a weak acid and its conjugate base, which resists a change in pH when a stronger acid or base is added, thereby minimizing a change in pH. The most important buffer pair in extracellular fluid (ECF) is carbonic acid (H 2 CO 3) and bicarbonate (HCO 3 −). The interaction between this buffer pair forms the basis of the measurement of acid–base balance. HYDROGEN ION BALANCE Cellular hydrogen ion turnover can be described in terms of processes that produce or consume H + ions in the body (Table 8.2). The total daily H + ion turnover in a normal adult is approximately 150 moles...
eBook - ePub- John Kenkel(Author)
- 2020(Publication Date)
- CRC Press(Publisher)
...500 mL of the 100 ppm solution would be measured into a 500 mL flask and diluted to volume. 3.7 BUFFER SOLUTIONS Buffer solutions, or solutions which resist changes in pH even when a strong acid or base is added, are almost always composed of a weak acid or weak base and the salt of this weak acid or base. The reason for the resistance to pH change is that the weak acid or weak base ionization equilibrium shifts in these solutions such that the H + or OH − added are consumed, thus resulting in no net pH change. Although commercially prepared buffer solutions are available, these are most often utilized solely for pH meter calibration and not for adjusting or maintaining a chemical reaction system at a given pH. It is not surprising, therefore, that the analyst often needs to prepare his/her own solutions for this purpose. It then becomes a question of what proportions of the acid, or base, and its salt should be mixed to give the desired pH. The answer is in the expression for the ionization constant, K a or K b, where the ratio of the salt concentration to the acid concentration is found. In the case of a weak acid, HA ⇌ H + A − (3.33) K a = [ H + ] [ A − ] [ HA ] (3.34) and in the case of a weak base, B+H 2 O ⇌ BH + + OH − (3.35) K b = [ BH + ] [ OH − ] [ B ] (3.36) Knowing the value of K a or K b, for a given weak acid or base and knowing the desired pH value, one can calculate the ratio of salt concentration to acid (or base) concentration that will produce the given pH. Rearranging Equation 3.34, for example, would show the method for calculating this ratio in the case of a weak acid and its salt. K a [ H + ] = [ A − ] [ HA ] (3.37) This is one form of the Henderson-Hasselbalch equation for dealing with buffer solutions. It says that one would simply divide the K a by the [H + ] to obtain the required ratio...
eBook - ePub- Arnost Kotyk, Jan Slavik(Authors)
- 2020(Publication Date)
- CRC Press(Publisher)
...Its pH follows from the dissociation reaction (Equation 11) governed by K A and can be expressed (for low degrees of dissociation) as pH = P K A + log (c A − / c HA) (44) (cf. Equation 12). On adding a strong base to the solution, to a final concentration C B, a salt is formed at practically the same concentration so that C B′ ≅ c A−. This necessarily leads to a decrease of the concentration of undissociated acid, such that c HA ≅ C T − C B'. This then turns into the so-called Henderson-Hasselbalch equation pH = p K A + log [ c B′ / (c T − c Bʹ) ] (45) Analogously, starting with a weak base (c T = c B + c BH +) we may write pH = p K BH + log(c B / c BH +) (46) Addition of a strong acid to a final concentration c A' leads then to salt formation such that c A ' ≅ c BH+ and, necessarily c B ≅ c T − c A'. The pH is then defined by pH = p K BH + log[(c T − c Aʹ) / c Aʹ ] (47) It is easy to show that the “titration” of, say, a weak acid with a strong base, proceeds through a “buffer zone” where pH changes very little in spite of adding large amounts of the base (Figure 2). The figure indicates that the buffer zone centers about a point where pH = P K Aʹ, this being also the point where c Bʹ = 1/2 c A. FIGURE 2. Titration of a weak acid A with a strong base B′ The pH observed at the different base-to-acid ratios is defined by the expressions on the right. The zone of high buffering power lies here at about 0.2 to 0.8. The efficiency of a buffer solution is usually expressed by its buffering power β which is the expression for a change of pH brought about by adding a certain amount of strong acid or base. Thus, β = d c Bʹ / dpH = 2.3 c Bʹ (1 − c Bʹ / c A) for an acid buffer (48a) β = d c Aʹ / dpH = 2.3 c Aʹ (c Aʹ / c B − 1) for an alkaline buffer (48b) The buffering power thus depends on c Bʹ, or c Aʹ and its maximum is reached (when the first derivative of Equation 48a or 48b is equal to zero) at c Bʹ = 1/2 c A or c Aʹ = 1/2 c B (Figure 3). FIGURE 3...
