Chemistry
Acid-Base Indicators
Acid-base indicators are substances that change color depending on the pH of a solution. They are commonly used in chemistry to visually determine whether a solution is acidic, basic, or neutral. Indicators work by undergoing a reversible chemical change that is sensitive to the hydrogen ion concentration in the solution.
Written by Perlego with AI-assistance
Related key terms
1 of 5
11 Key excerpts on "Acid-Base Indicators"
eBook - PDF- Mel Schwartz(Author)
- 2008(Publication Date)
- CRC Press(Publisher)
The indicator should not be more expensive than the product it is protecting. Application : Easy to attach to a variety of containers or pack-ages. Once installed, the device must remain intact and readable during the service life of the packages. Response : The response mechanism must have fast kinetics on the order of seconds to a few hours and be reproducible. There should be no time delay in response between reactions involving solid, liquid, or gas. Most applications require the response in the indicator to be irreversible to preserve the needed indication record. Sensitivity : The indicator must be highly sensitive, accurate, and easily activated. A user-friendly indicator that can provide useful information when needed will make both the product manufacturer and the consumer happy. Shelf-life : The indicator must have a shelf-life equal to or longer than that of the product it is monitoring. It should be technically diffi cult to duplicate or counterfeit the indicator’s response. In this case, the indicator is acting as a “smart” locking mechanism. 30.2.4 pH Indicators The pH indicator is probably the oldest and simplest smart chemical indicating device known. The chemistry of acid–base indicators is well documented [7] and involves proton or elec-tron exchange reactions. It will be mentioned only briefly here. By Brönsted and Ostwald’s definition, indicators are weak acids or bases. Upon dissociation, they exhibit a structural and color change that is different from the undissociated form. Indicators commonly used in acid–base and redox titrations are conjugated organic dyes containing one or more light-absorbing groups called chromophores. The electronic structure of these dyes can be changed by redox or proton exchange reactions. This also changes the absorption energy in the visible region of the elec-tromagnetic spectrum, and consequently, the color. The color change therefore provides a visual indication of an endpoint of a titration [1,7].
eBook - PDF- Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
- 2013(Publication Date)
- Wiley(Publisher)
8.3 Detection of the End Point: Indicators Carrying out the titration process is of little value unless we can tell exactly when the The goal is for the end point to coincide with the equivalence point. acid has completely neutralized the base, i.e., when the equivalence point has been reached. Therefore, we wish to determine accurately when the equivalence point is reached. The point at which the reaction is observed to be complete is called the end point. A measurement is chosen such that the end point coincides with or is very close to the equivalence point. The difference between the equivalence point and the end point is referred to as the titration error; as with any measurement, we want to minimize error. The most obvious way of determining the end point is to measure the pH at different points of the titration and make a plot of this versus milliliters of titrant. This is done with a pH meter, which is discussed in Chapter 13. It is often more convenient to add an indicator to the solution and visually detect a color change. An indicator for an acid–base titration is a weak acid or weak base that is highly colored. The color of the ionized form is markedly different from that of the nonionized form. One form may be colorless, but at least one form must be colored. These substances are usually composed of highly conjugated organic constituents that give rise to the color (see Chapter 16). Assume the indicator is a weak acid, designated HIn, and assume that the nonionized form is red while the ionized form is blue: Hln (red) H + + In − (blue) (8.10) We can write a Henderson – Hasselbalch equation for this, just as for other weak acids: pH = pK In + log [In − ] [HIn] (8.11) The indicator changes color over a pH range. The transition range depends on the Your eyes can generally discern only one color if it is 10 times as intense as the other. ability of the observer to detect small color changes.
eBook - PDFChemistry
An Atoms First Approach
- Steven Zumdahl, Susan Zumdahl, Donald J. DeCoste, , Steven Zumdahl, Steven Zumdahl, Susan Zumdahl, Donald J. DeCoste(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
9. What is an acid–base indicator? Define the equivalence (stoichiometric) point and the end point of a titration. Why should you choose an indicator so that the two points coincide? Do the pH values of the two points have to be within 60.01 pH unit of each other? Explain. 10. Why does an indicator change from its acid color to its base color over a range of pH values? In general, when do color changes start to occur for indicators? Can the indicator thymol blue contain only a single OCO 2 H group and no other acidic or basic functional group? Explain. A third method that can be used to solve for the pH of a buffer solution is the Henderson–Hasselbalch equa- tion. What is the Henderson–Hasselbalch equation? What assumptions are made when using this equation? 3. One of the most challenging parts of solving acid–base problems is writing out the correct equation. When a strong acid or a strong base is added to solutions, they are great at what they do, and we always react them first. If a strong acid is added to a buffer, what reacts with the H 1 from the strong acid and what are the prod- ucts? If a strong base is added to a buffer, what reacts with the OH 2 from the strong base and what are the products? Problems involving the reaction of a strong acid or strong base are assumed to be stoichiometry problems and not equilibrium problems. What is assumed when a strong acid or strong base reacts to make it a stoichiometry problem? 4. A good buffer generally contains relatively equal con- centrations of weak acid and conjugate base. If you wanted to buffer a solution at pH 5 4.00 or pH 5 10.00, how would you decide which weak acid– conjugate base or weak base–conjugate acid pair to use? The second characteristic of a good buffer is good buffering capacity.
- Premasis Sukul(Author)
- 2023(Publication Date)
Indicators 49 3.3. Fluorescent indicator Fluorescence refers to a phenomenon where a molecule absorbs light to form a short-lived excited state and re-emit light at a longer wavelength. The fluorescent indicators whose fluorescent properties in solution are influenced by a change in hydrogen-ion concentration, oxidation potential, or metal-ion concentration can be used for endpoint determination. The main principle of fluorescence indicators in acid-base titrations involves giving rise of fluorescence in either its acidic or basic form or vice versa. However, basic fluorescence is most common. For instance, eosin (pH range, 0.0-3.0; green fluorescence), dichlorofluorescein (pH range, 4.0-6.6; blue fluorescence), coumaric acid (pH range, 7.2-9.0; green fluorescence), salicylic acid (pH range, 2.0-4.0; violet fluorescence), etc are all base fluorescence which means that their acid form does not generate any fluorescence, but their basic form normally produces fluorescence. The fluorescence helps to detect the endpoint of any titration reaction, where conventional Acid-Base Indicators cannot function properly due to solution turbidity or highly coloured solution. BCECF [2′,7′-bis-(2-carboxyethyl)-5-(and-6)-carboxyfluorescein] is a commonly used fluorescent pH sensor to detect the changes in cytosolic pH, possessing pKa of 6.97 which is close to physiological pH. It has been evidenced that the indicator is poorly fluorescent at low pH but the fluorescence becomes strong with increasing pH. The excitation spectrum of the BCECF faces a little shift with pH change, while the wavelength of the emission maximum remains the same. The pH is determined ratiometrically by the relative fluorescent intensities at 535 nm when the dye is excited at 439 nm and 505 nm. 3.4. Precipitation indicator In this case, after completion of the titration, a coloured precipitation is obtained to identify the end point.
eBook - PDFChemistry
Structure and Dynamics
- James N. Spencer, George M. Bodner, Lyman H. Rickard(Authors)
- 2011(Publication Date)
- Wiley(Publisher)
A common type of titration is an acid–base titration in which an acid of unknown concentra- tion is titrated with a base that has an accurately known concentration. Titrations can therefore be used to determine the concentration of an acid or base solution. Indicators are weak acids or weak bases whose conjugate acid–base pairs have different colors in aqueous solution. A commonly used indicator is phe- nolphthalein, which is colorless in its acid form and pink in its base form. The two forms of an indicator can be represented by HIn for the acid form and In for the base form. The ionization of the indicator can therefore be represented by the following equation, HIn(aq) + H 2 O(aq) uv H 3 O + (aq) + In - (aq) K rxn = 5.1 * 10 - 4 * 4.8 * 10 - 4 * 1.0 * 10 14 = 2.5 * 10 7 HA(aq) + B(aq) uv BH + (aq) + A - (aq) K rxn = K a * K b * 1.0 * 10 14 H 3 O + (aq) + OH - (aq) uv 2 H 2 O(l) K = 1.0 * 10 14 B(aq) + H 2 O(l) uv BH + (aq) + OH - (aq) K b = ? HA(aq) + H 2 O(l) uv H 3 O + (aq) + A - (aq) K a = ? CH 3 NH 2 (aq) + HNO 2 (aq) ¡ CH 3 NH 3 + (aq) + NO 2 - (aq) 11.19 pH TITRATION CURVES 513 for which K a could be written as follows: This equation can be rearranged to give: The color of the indicator solution depends on the concentration of the H 3 O ion. When the concentration of the H 3 O ion is large, the value of [HIn] is larger than [In ], and the solution has the characteristic color of the acid form of the indi- cator. If phenolphthalein is the indicator, the solution is colorless in acid. When the concentration of the H 3 O ion is small, the value of [In ] is larger than [HIn], and the color of the base form of the indicator predominates. For phenolphthalein, the color of the base form is pink. For most indicators color changes occur when the ratio [In ]/[HIn] changes from 1/10 to about 10. Some common indicators are listed in Table 11.7 along with the pH interval over which the indicator changes color.
- James N. Jensen(Author)
- 2022(Publication Date)
- Wiley(Publisher)
A pH indicator is an acid–base pair where the acid is one color or colorless in solution and the conjugate base is a different color in solution. Only a small amount of the indicator is added to the sample to avoid changing the sample pH. If the pK a of the acid is near the pH of the endpoint, then a color change is observed at pH values near the endpoint. Characteristics of several acid–base indicators are listed in Table 13.2. In the older literature, the endpoint approximating the f = 0 equivalence point is called the methyl orange endpoint because the pH indicator methyl orange was used to visualize the endpoint. The preferred pH indicator currently is bromcresol green (or the bromcresol 4 5 6 7 8 9 10 11 0.0 0.5 1.0 1.5 2.0 pH f C T = 1×10 –3 M C T = 5×10 –4 M C T = 1×10 –4 M FIGURE 13.4 Carbonic Acid Titration Curves for Different C T Values 4 5 6 7 8 9 10 11 12 0.0001 0.01 pH 0.001 C T (M, log scale) f = 0 (g = 2) g = 0 (f = 2) f = 1 (g = 1) FIGURE 13.5 pH Values at Integral Equivalence Points as a Function of C T Common Endpoint Indicators Indicator Acid Color Base Color pK a Transition pH Range Near f = 0 (g = 2): Methyl orange Red (HB + ) Orange (B) 3.8 4.5–3.2 Bromcresol green Yellow (HA) Blue (A – ) 4.8 5.4–3.8 Near f = 2 (g = 0): Phenolphthalein Clear (H 2 A) Pink (A – ) 9.7 8.2–10.0 Metacresol purple Yellow (HA) Purple (A – ) 8.3 7.6–9.2 Table 13.2 methyl orange endpoint: a pH endpoint of about pH 4.5 to 3.2, previously used to estimate the alka- linity equivalence point (f = 0 for carbonic acid) 296 Chapter 13 Alkalinity and Acidity green-methyl red mixed indicator). Phenolphthalein is used for endpoints near neu- tral pH (e.g., the titration of a strong acid with a strong base; see also Section 13.4.3).
eBook - PDF- Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2015(Publication Date)
- Openstax(Publisher)
We base our choice of indicator on a calculated pH, the pH at the equivalence point. At the equivalence point, equimolar amounts of acid and base have been mixed, and the calculation becomes that of the pH of a solution of the salt resulting from the titration. 822 Chapter 14 | Acid-Base Equilibria This OpenStax book is available for free at http://cnx.org/content/col11760/1.9 acid ionization acid ionization constant (K a ) acid-base indicator acidic amphiprotic amphoteric autoionization base ionization base ionization constant (K b ) basic Brønsted-Lowry acid Brønsted-Lowry base buffer buffer capacity color-change interval conjugate acid conjugate base diprotic acid diprotic base Henderson-Hasselbalch equation ion-product constant for water (K w ) leveling effect of water monoprotic acid neutral oxyacid Key Terms reaction involving the transfer of a proton from an acid to water, yielding hydronium ions and the conjugate base of the acid equilibrium constant for the ionization of a weak acid organic acid or base whose color changes depending on the pH of the solution it is in describes a solution in which [H 3 O + ] > [OH − ] species that may either gain or lose a proton in a reaction species that can act as either an acid or a base reaction between identical species yielding ionic products; for water, this reaction involves transfer of protons to yield hydronium and hydroxide ions reaction involving the transfer of a proton from water to a base, yielding hydroxide ions and the conjugate acid of the base equilibrium constant for the ionization of a weak base describes a solution in which [H 3 O + ] < [OH − ] proton donor proton acceptor mixture of a weak acid or a weak base and the salt of its conjugate; the pH of a buffer resists change when small amounts of acid or base are added amount of an acid or base that can be added to a volume of a buffer solution before its pH changes significantly (usually by one pH unit) range in pH over which the color change of an indicator takes place substance formed when a base gains a proton substance formed when an acid loses a proton acid containing two ionizable hydrogen atoms per molecule.
eBook - PDF- William R. Robinson, Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley(Authors)
- 2016(Publication Date)
- Openstax(Publisher)
We base our choice of indicator on a calculated pH, the pH at the equivalence point. At the equivalence point, equimolar amounts of acid and base have been mixed, and the calculation becomes that of the pH of a solution of the salt resulting from the titration. 788 Chapter 14 | Acid-Base Equilibria This OpenStax book is available for free at http://cnx.org/content/col12012/1.7 acid ionization acid ionization constant (K a ) acid-base indicator acidic amphiprotic amphoteric autoionization base ionization base ionization constant (K b ) basic Brønsted-Lowry acid Brønsted-Lowry base buffer buffer capacity color-change interval conjugate acid conjugate base diprotic acid diprotic base Henderson-Hasselbalch equation ion-product constant for water (K w ) leveling effect of water monoprotic acid neutral oxyacid Key Terms reaction involving the transfer of a proton from an acid to water, yielding hydronium ions and the conjugate base of the acid equilibrium constant for the ionization of a weak acid organic acid or base whose color changes depending on the pH of the solution it is in describes a solution in which [H 3 O + ] > [OH − ] species that may either gain or lose a proton in a reaction species that can act as either an acid or a base reaction between identical species yielding ionic products; for water, this reaction involves transfer of protons to yield hydronium and hydroxide ions reaction involving the transfer of a proton from water to a base, yielding hydroxide ions and the conjugate acid of the base equilibrium constant for the ionization of a weak base describes a solution in which [H 3 O + ] < [OH − ] proton donor proton acceptor mixture of a weak acid or a weak base and the salt of its conjugate; the pH of a buffer resists change when small amounts of acid or base are added amount of an acid or base that can be added to a volume of a buffer solution before its pH changes significantly (usually by one pH unit) range in pH over which the color change of an indicator takes place substance formed when a base gains a proton substance formed when an acid loses a proton acid containing two ionizable hydrogen atoms per molecule.
- Subrayal M Reddy(Author)
- 2014(Publication Date)
- Royal Society of Chemistry(Publisher)
The total colour difference of the mixed pH dye-based indicator correlated well with CO 2 levels of intermediate moisture levels in the dessert. Spoilage could be monitored at various constant temperatures and even correlated with temperature fluctuation, potentially allowing the label to also act as a temperature indicator. Simple methods to indicate levels of CO 2 can therefore be used to deter-mine the quality and degree of deterioration of the packed food during the various stages of production and transport to outlets, as well as for the consumer to make an informed choice prior to and post-purchase. Food traceability throughout the food supply chain is increasingly becoming a necessary task, mandatory in the European Union (EU) since 2005. In light of the above, there is an increasing need to develop low-cost, accurate, rapid, robust, non-invasive and non-destructive indicators. The indicators also need to be non-toxic, and insoluble in water. 31 They should also exhibit an irreversible response towards the analyte and be tamper proof. While the latter is a little more difficult to police, the former is important in order to confirm traceability and history of the handling of the product. Table 5.2 Some commonly used pH indicators. Indicator Acid Colour Base Colour pH Transition Range Thymol blue Red Yellow 1.2–2.8 Methyl orange Red Yellow 3.1–4.4 Bromocresol green Yellow Blue 3.8–5.4 Bromocresol purple Yellow Purple 5.2–6.8 p -Nitrophenol Colourless Yellow 5.6–7.6 Cresol red Yellow Red 7.2–8.8 Phenolphthalein Colourless Pink 8.0–9.6 Alizarin yellow Yellow Orange-red 10.1–12.0 Smart Indicator Technologies for Chemical and Biochemical Detection 151 5.3 Oxygen Indicators Often taken for granted, packaging is important for much of the quality we expect from food products. Traditionally, packaging fulfils four basic roles: containment, protection, communication and convenience.
eBook - PDF- L. Pataki, E. Zapp, R. Belcher, D Betteridge, L Meites(Authors)
- 2013(Publication Date)
- Pergamon(Publisher)
5-1. The screened indicator consisting of methyl red and méthylène blue is violet in acidic media, the colour being a mixture of red and blue. In alkaline media the indicator turns green, i.e. a mixture of yellow and blue. As the colours are almost complementary, a grey tint appears at the equivalence point. This indicator allows the equivalence point to be determined with an accuracy of 0-2 pH.. The indicator exponent is 5-4. QUANTITATIVE CHEMICAL ANALYSIS 223 Acid-base titrations in aqueous solutions, and their main fields of application In acid-base titrations, strong and weak acids, or the salts of a weak base and a strong acid, are titrated with standard alkali solutions, whereas strong and weak bases, or the salts of a weak acid and a strong base are titrated with standard acids. First the basic conditions for such measurements will be discussed. The feasibility can be decided from the titration curves and the transi-tion intervals of the indicators. In most instances it is sufficient to know the ρΉ. of the solution at the equivalence point to select an appropriate indicator. As recommended by Bjerrum, the hydrogen exponent at the equivalence point of a titration is termed the titration exponent, and is denoted by ρΎ. When strong acids are titrated, the equivalence point is at ρΉ. 7-0, where, according to figure 28, the titration curve is steep. The end point of the titration can be detected by any indicator with a transition in the pH range 4-10 for 0 1 N solutions, or 5-9 for 0-01 N solutions. An indicator error will occur when the titration and indicator exponents are different. The error is, by definition, the difference between the volumes of standard solution required to reach the end point and the equivalence point, respectively.
eBook - PDF- Douglas P. Heller, Carl H. Snyder(Authors)
- 2015(Publication Date)
- Wiley(Publisher)
Some change color with an increase in pH as a re- sult of structural changes brought on as the molecule releases a proton, H + . To what general class of compounds do such indicators belong? 26. Acid deposits around a car’s battery terminals can be cleaned with a solution of sodium bicarbonate. Assuming that the acid is sulfuric acid, complete and balance the equation for the neutralization reaction: NaHCO 3 + H 2 SO 4 → ? 27. Malonic acid, with a molecular formula C 3 H 4 O 4 , is a dicar- boxylic acid containing three carbons. Draw its molecular structure. 28. Blood is buffered by the presence of the acid H 2 CO 3 and its conjugate base HCO 3 − . Another combination of an acid and a base—H 2 PO 4 − and HPO 4 2− —helps maintain the fluid within our cells (intracellular fluid) within a narrow pH range. Rec- ognizing that H 2 PO 4 − is the more acidic of these two compo- nents, and that HPO 4 2− is the more basic component, write a chemical reaction that shows how this system changes as the intracellular fluid becomes (a) more acidic and (b) more basic. 29. Can a substance act as an acid in the absence of a base? Explain your answer. 30. Baking powder, one of several types of leavening agents used in preparing bread, cakes and similar baked goods, causes dough to rise during the baking process by releasing tiny bubbles of gas. This gas is generated through an acid- base neutralization between two components of the baking powder: sodium bicarbonate (NaHCO 3 ), and an acid, such as calcium acid phosphate (CaHPO 4 ). The gas is released as these two chemicals react with each other within the wet dough. (a) What is the name and the chemical formula of the gas released during this reaction? (b) What role does the sodium bicarbonate serve in this process? 5JN 0 8BMLFS Exercises 263 Calculate 36. For a solution at 25°C, what is the pH when (a) [H 3 O + ] is 0.0001 M? (b) [H 3 O + ] is 0.0000000001 M? (c) equal amounts of hydronium and hydroxide ions are present? 37.
Index pages curate the most relevant extracts from our library of academic textbooks. They’ve been created using an in-house natural language model (NLM), each adding context and meaning to key research topics.
