Name: solubility limits and selective precipitation
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 3 min 11 s
Channel: JoVE

K12 Standards

ohios-learning-standards-and-model-curriculum

Ohio’s Learning Standards and Model Curriculum

9-12

Chemistry

Interactions Of Matter

Chemical reactions

Precipitation of Ions

Wenn die Lösungen zweier ionischer Verbindungen, wie Natriumiodid und Blei(II)-nitrat, gemischt werden, können sich die Kationen aus der einen Lösung mit den Anionen aus der anderen verbinden.
Da eines der Kreuzprodukte – Natriumnitrat – wasserlöslich ist, bleiben Natrium- und Nitrationen weiterhin in Lösung, während Blei- und Jodidionen einen Bleijodidniederschlag bilden.
Der Reaktionsquotient Q der Auflösungsreaktion ist gleich dem Produkt aus den Konzentrationen von Blei- und Iodidionen.
Im Gegensatz zum Löslichkeitsprodukt Ksp handelt es sich beim Reaktionsquotienten um Ionenkonzentrationen in jeder Phase, nicht nur im Gleichgewicht.
Die Werte von Q und Ksp können verglichen werden, um vorherzusagen, ob eine Fällungsreaktion auftreten wird.
Erwägen Sie eine tropfenweise Zugabe von Natriumiodidlösung zur Bleinitratlösung.
Zu Beginn ist Q kleiner als Ksp, und sowohl Blei- als auch Iodidionen sind mit den Natrium- und Nitrationen in Lösung. Dabei handelt es sich um eine ungesättigte Lösung.
Je mehr Natriumiodid zugesetzt wird, desto mehr nimmt die Konzentration der Iodidionen zu. Die Reaktion hat das Gleichgewicht erreicht, wenn Q = Ksp. In diesem Stadium befindet sich eine kleine Menge festes Bleiiodid im dynamischen Gleichgewicht mit den Ionen und bildet eine gesättigte Lösung.
Durch die weitere Zugabe von Natriumiodid wird Q > Ksp, und die Reaktion verlagert sich in Richtung des Niederschlags. Dabei handelt es sich um eine übersättigte Lösung, bei der die Ausfällung so lange anhält, bis die Ionenkonzentrationen auf ihre Gleichgewichtswerte abgesenkt sind.
Nehmen wir zum Beispiel an, dass das Mischen von Natriumiodid und Blei(II)-nitrat-Lösungen zu einer Lösung führt, die 1,6 × 10-4 M Blei-Ionen und 4,0 × 10-4 M Iodid-Ionen enthält.
Hier ist Q = 2,6 × 10−11, während Ksp für Bleiiodid 1,4 × 10−8 beträgt. Da Q < Ksp ist, fällt Bleiiodid nicht aus.
Die Vorhersage von Fällungsreaktionen kann bei der Trennung von ionischen Verbindungen sehr nützlich sein.
Stellen Sie sich eine Lösung mit zwei Metallionen vor – Blei(II) und Kupfer(II). Wenn dieser Lösung Salzsäure zugesetzt wird, fällt Blei(II)-chlorid aus, da es ein kleines Ksp hat, während Kupfer in Lösung verbleibt, da Kupfer(II)-chlorid in Wasser gut löslich ist.
Diese Technik wird als selektive Niederschlagung bezeichnet.

<ol>
	<li>Firsching, Ferdinand Henry. "Selektive Ausfällung von Silberhalogeniden aus homogener Lösung. Trennung von Iodid, Bromid und Chlorid durch Verflüchtigung von Ammoniak." Analytische Chemie 32, Nr. 13 (1960): 1876-1878.</li>
	<li>Reynolds, John P. "Ksp-Experiment: Das Löslichkeitsprodukt für Bariumhydroxid." Zeitschrift für chemische Bildung 52, Nr. 8 (1975): 521.</li>
	<li>Hou, Miaolin und George L. Baughman. "Vorhersage der Ausfällung von sauren und direkten Farbstoffen in natürlichen Gewässern." Farbstoffe und Pigmente 18, Nr. 1 (1992): 35-46.</li>
</ol>

niederschlag von ionen

Niederschlag vorhersagen Die Gleichung, die das Gleichgewicht zwischen festem Calciumcarbonat und seinen solvatisierten Ionen beschreibt, lautet: 

 <img alt="Eq1" src="/files/ftp_upload/11420/11420_Equation_1.png" style="text-align: center; height: 50px;"> 

 Es ist wichtig zu wissen, dass dieses Gleichgewicht in jeder wässrigen Lösung hergestellt wird, die Ca2+ und CO32–-Ionen enthält, und nicht nur in einer Lösung, die durch Sättigung von Wasser mit Calciumcarbonat entsteht. Denken Sie beispielsweise an das Mischen wässriger Lösungen der löslichen Verbindungen Natriumcarbonat und Calciumnitrat. Wenn die Konzentrationen von Calcium- und Carbonationen in der Mischung keinen Reaktionsquotienten Q ergeben, der das Löslichkeitsprodukt Ksp übersteigt, kommt es zu keiner Ausfällung. Wenn die Ionenkonzentrationen einen Reaktionsquotienten ergeben, der größer als das Löslichkeitsprodukt ist, kommt es zu einer Ausfällung, wodurch diese Konzentrationen sinken, bis sich ein Gleichgewicht einstellt (Q = Ksp). Der Vergleich von Q mit Ksp zur Vorhersage von Niederschlägen ist ein Beispiel für den allgemeinen Ansatz zur Vorhersage der Richtung einer Reaktion. Für den speziellen Fall von Löslichkeitsgleichgewichten: 

 Q < Ksp: Die Reaktion verläuft in Vorwärtsrichtung (Lösung ist nicht gesättigt; kein Niederschlag beobachtet) Q > Ksp: Die Reaktion verläuft in umgekehrter Richtung (Lösung ist übersättigt; es kommt zur Ausfällung) 

 In Lösungen, die zwei oder mehr Ionen enthalten, die mit demselben Gegenion unlösliche Verbindungen bilden können, kann eine experimentelle Strategie namens selektive Fällung verwendet werden, um einzelne Ionen aus der Lösung zu entfernen. Durch eine kontrollierte Erhöhung der Gegenionenkonzentration können in Lösung befindliche Ionen einzeln ausgefällt werden, sofern die Löslichkeiten ihrer Verbindungen ausreichend unterschiedlich sind. In Lösungen mit gleichen Konzentrationen an Zielionen fällt das Ion, das die am wenigsten lösliche Verbindung bildet, zuerst aus (bei der niedrigsten Konzentration an Gegenionen). Die anderen Ionen fallen anschließend aus, wenn die Löslichkeiten ihrer Verbindung erreicht sind.

 Fällung von Silberhalogeniden 

 Eine Lösung enthält 0,00010 Mol KBr und 0,10 Mol KCl pro Liter. Dieser Lösung wird nach und nach AgNO3 zugesetzt. Was bildet sich zuerst, festes AgBr oder festes AgCl? 

 Die beiden beteiligten Gleichgewichte sind: 

 <img alt="Eq2" src="/files/ftp_upload/11420/11420_Equation_2.png" style="text-align: center; height: 90px;"> 

Würde die Lösung etwa gleiche Konzentrationen an Cl– und Br– enthalten, würde zuerst das Silbersalz mit dem kleineren Ksp (AgBr) ausfallen. Da die Konzentrationen jedoch nicht gleich sind, müssen der [Ag+]-Wert, bei dem AgCl auszufallen beginnt, und der [Ag+] -Wert, bei dem AgBr auszufallen beginnt, berechnet werden. Das Salz, das sich am unteren [Ag+] bildet, fällt zuerst aus. 

AgBr fällt aus, wenn Q = Ksp für AgBr <img alt="Eq3" src="/files/ftp_upload/11420/11420_Equation_3.png" style="text-align: center; height: 112px;"> AgBr beginnt auszufallen, wenn [Ag+] 5,0 × 10−9 M beträgt. 

Für AgCl: 

AgCl fällt aus, wenn Q gleich Ksp für AgCl ist (1,6  × 10-10). Wenn [Cl–] = 0,10 M: 

<img alt="Eq4" src="/files/ftp_upload/11420/11420_Equation_4.png" style="text-align: center; height: 112px;"> 

AgCl beginnt auszufallen, wenn [Ag+] 1,6  × 10-9 M beträgt. 

AgCl beginnt bei einem niedrigeren [Ag+] als AgBr auszufallen, daher beginnt AgCl zuerst auszufallen. Beachten Sie, dass die Chloridionenkonzentration der anfänglichen Mischung deutlich höher war als die Bromidionenkonzentration, sodass Silberchlorid zuerst ausgefällt wurde, obwohl es einen Ksp hatte, der größer als der von Silberbromid war.

Dieser Text wurde angepasst von <a href="https://openstax.org/books/chemistry-2e/pages/15-1-precipitation-and-dissolution">Openstax, Chemistry 2e, Chapter 15.1: Precipitation and Dissolution</a>.

Niederschlag von Ionen

Vorhersage von Niederschlag
Die Gleichung, die das Gleichgewicht zwischen festem Calciumcarbonat und seinen solvatisierten Ionen beschreibt, lautet:

Es ist wichtig, sich darüber im Klaren zu sein, dass ...

Education

Lehre

JoVE Core

Chemie

Eins

Acid-base and Solubility Equilibria

Säure-Base- und Löslichkeitsgleichgewichte

KEY TERMS AND CONCEPTS

KERNBEGRIFFE UND KONZEPTE

solubility limits and selective precipitation

Solubility limits and selective precipitation explain when ions form a solid in water. The key idea is the balance between an insoluble compound and its dissolved ions. In aqueous solutions that contain calcium ions and carbonate ions, this equilibrium can exist even if the solution was not made by saturating water with calcium carbonate.

When two soluble salts are mixed, such as sodium carbonate and calcium nitrate, the ion concentrations determine whether a precipitate forms. If the reaction quotient, Q, does not exceed the solubility product, Ksp, no solid appears. If Q is greater than Ksp, precipitation occurs until the solution returns to equilibrium, where Q = Ksp.

For solubility equilibria, the comparison of Q and Ksp predicts the reaction direction. When Q Ksp, the solution is supersaturated and a precipitate will form.

Selective precipitation uses this same idea to remove one ion from a mixture that contains more than one ion that can form insoluble compounds with the same counter ion. By adding the counter ion in a controlled way, ions can be precipitated one at a time if their solubilities are different enough. If the target ions start at equal concentrations, the ion that forms the least soluble compound precipitates first.

A silver halide example shows how concentration affects the order of precipitation. A solution contains 0.00010 mol of KBr and 0.10 mol of KCl per liter, and AgNO3 is added gradually. Although AgBr has the smaller Ksp, AgCl begins to precipitate first because it reaches its precipitation point at a lower Ag+ concentration. AgBr begins to precipitate when [Ag+] is 5.0 × 10−9 M, while AgCl begins when [Ag+] is 1.6 × 10−9 M. The chloride ion concentration is much higher than the bromide ion concentration, so silver chloride forms first.

Solubility Limits and Selective Precipitation

solubility-limits-and-selective-precipitation

common ion effect

Common Ion Effect on Solubility

The common ion effect reduces the solubility of an ionic compound in a solution that already contains one of its ions. In aqueous solutions, this means less of the solid dissolves than it would in pure water. The effect follows the law of mass action and can be explained with Le Châtelier’s principle. Silver iodide shows this equilibrium clearly. When silver(I) ions or iodide ions are added, the dissolution equilibrium shifts left. More AgI precipitates, and the concentrations of dissolved Ag+...

buffers

Buffer Solutions and pH Stability

Buffer solutions are mixtures that help keep pH steady. A buffer contains an appreciable amount of a weak conjugate acid-base pair. It resists a change in pH when small amounts of a strong acid or a strong base are added. One common buffer uses a weak acid and its salt. Acetic acid and sodium acetate form this type of system, written as CH3COOH(aq) + CH3COONa(aq). Another buffer uses a weak base and its salt. A solution of ammonia and ammonium chloride, NH3(aq) + NH4Cl(aq), is an example. The...

calculating ph changes in a buffer solution

calculating-ph-changes-in-a-buffer-solution

Buffer pH After Acid or Base Addition

A buffer solution can resist sudden pH changes when a strong acid or strong base is added. It does not stop the pH from changing completely. The shift is usually small, as long as the addition stays within the buffer’s capacity. To find the new pH, first determine how the added acid or base changes the amounts of the buffer components. A buffer contains a weak acid and its conjugate base, or the reverse pair. These amounts are found with stoichiometric calculations, which track how much of...

buffer effectiveness

Buffer Capacity and pH Stability

Buffer solutions help keep pH fairly constant, but they do not work forever. Their resistance to pH change depends on having enough of a weak acid and its conjugate base, or a weak base and its conjugate acid, in the mixture. When too much strong acid or strong base is added, one part of the buffer pair is used up and the buffer action becomes weaker. Buffer capacity describes how much acid or base can be added to a fixed volume of buffer before the pH changes noticeably, usually by about one...

titration calculations: strong acid - strong base

titration-calculations-strong-acid-strong-base

pH at Key Points in Acid-Base Titration

pH at key points in a strong acid and strong base titration can be calculated from the amounts of acid and base present. The example uses 25.00 mL of 0.100 M HCl, a strong acid, and 0.100 M NaOH, a strong base. The pH changes as base is added, and each stage follows a different calculation. At 0 mL of titrant, the pH comes from the ionization of HCl. Because HCl is a strong acid, its ionization is complete. The hydronium ion molarity is 0.100 M, and the pH is found from that value. At 12.50...

titration calculations: weak acid - strong base

titration-calculations-weak-acid-strong-base

pH at Key Points in a Weak Acid Titration

pH at key points in a weak acid titration can be calculated for an acetic acid solution as sodium hydroxide is added. The example uses 25.00 mL of 0.100 M CH3CO2H titrated with 0.100 M NaOH. The reaction is shown as the weak acid reacts with the strong base. The initial pH is found from the acetic acid solution using the usual ICE approach. ICE means initial, change, and equilibrium, and it is a common way to set up equilibrium problems. The solution pH is calculated before any titrant is...

indicators

Acid-Base Color Changes

Acid-base indicators show color changes that help identify the pH of a solution. These are often weak organic acids or weak organic bases, and they respond to changes in hydronium ion concentration in dilute aqueous solution. Phenolphthalein is one example. It is colorless when the hydronium ion concentration is greater than 5.0 × 10^-9 M, which means the pH is below 8.3. In more basic solutions, where the hydronium ion concentration is lower, it turns red or pink. Methyl orange shows how an...

titration of a polyprotic acid

Polyprotic Acid Equivalence Points

Polyprotic acids have more than one ionizable hydrogen, so they lose protons in steps during titration. When the acid dissociation constants for those protons are far enough apart, the titration curve shows a separate equivalence point for each hydrogen that can ionize. Diprotic acids have two ionizable hydrogens and produce two equivalence points. Triprotic acids have three ionizable hydrogens and produce three equivalence points. Carbonic acid, H2CO3, is a weak diprotic acid. Carbonic acid...

solubility equilibria

Predicting Precipitation with Ksp

Solubility equilibria explain how a dissolved ionic compound and its solid form can exist in balance. In a saturated solution, the solute dissolves and the ions precipitate at equal rates. These ideas matter in real settings such as tooth decay and water purification, where compound solubility affects the outcome. A substance with finite solubility can form a saturated solution when more solute is added than the solvent can dissolve. The result is a heterogeneous mixture of the saturated...

factors affecting solubility

Common Ion Effects on Solubility

The common ion effect lowers the solubility of an ionic compound in water when the solution already contains one of the ions made by that compound. It follows the law of mass action and can be explained with Le Châtelier’s principle. When a common ion is added, the dissolution equilibrium shifts toward the solid, so less of the salt stays dissolved. Silver iodide shows this effect clearly. If silver ions or iodide ions are added to a solution of AgI, the equilibrium shifts left and more AgI...

formation of complex ions

Ligands and Complex Ion Stability

Complex ions form when a central atom, often a transition metal cation, binds to ligands. Ligands are ions or molecules around the metal ion. They can be neutral molecules such as H2O or NH3, or ions such as CN− or OH−. In many cases, the ligands act as Lewis bases by donating a pair of electrons to the central atom. These reactions are part of coordination chemistry, which studies coordination complexes. When a metal ion reacts with one or more ligands to form a complex, the equilibrium...

precipitation of ions

Solubility limits and selective precipitation explain when ions form a solid in water. The key idea is the balance between an insoluble compound and its dissolved ions. In aqueous solutions that contain calcium ions and carbonate ions, this equilibrium can exist even if the solution was not made by saturating water with calcium carbonate. When two soluble salts are mixed, such as sodium carbonate and calcium nitrate, the ion concentrations determine whether a precipitate forms. If the reaction...

qualitative analysis

Cation Identification by Precipitation Tests

Cation identification by precipitation tests uses selective reactions to find metal ions in a mixture. In qualitative analysis, different reagents are added step by step. Each reaction can produce a characteristic precipitate for a specific group of cations. Other steps, such as changing pH, heating the mixture to dissolve a precipitate again, or adding reagents that form complex ions, help separate ions within the same group. Group IV cations such as Ba2+, Ca2+, and Mg2+ are identified...

henderson-hasselbalch equation

Buffer pH with the Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation helps relate buffer pH to the weak acid in the solution. It comes from the ionization-constant expression for a weak acid. By rearranging that expression to solve for hydronium ion concentration, [H3O+], and then taking the negative logarithm of both sides, the equation can be written in terms of pH. In this form, pKa is defined as the negative logarithm of the weak acid ionization constant, or pKa = −log Ka. The equation shows the connection between pH, pKa,...

acid-base titration curves

Tracking pH During Acid-Base Titration

Acid-base titration uses pH changes to track how a solution responds as titrant is added. A titration curve is a graph of a solution property versus the amount of added titrant. In acid-base titrations, pH is the main property to follow because it changes in a predictable way with the solution makeup and helps show the endpoint. This pattern can be seen with a strong acid and strong base, a strong acid and weak base, or a strong base and weak acid. For example, 25.00 mL of 0.100 M HCl titrated...

key terms and concepts

Learn how Q and Ksp predict precipitation and how selective precipitation separates ions. Includes a silver chloride and silver bromide example.

Solubility Limits and Selective Precipitation | JoVE K12

chemistry

JoVE Core Chemie erweckt allgemeine Konzepte der Chemie durch kompakte und leicht verständliche animierte Videolektionen zum Leben. Außerdem zeigen wir in den Videos Wissenschaftler in Aktion bei klassischen und neuartigen Forschungsexperimenten, wie sie heutzutage in Laboren weltweit durchgeführt wurden.

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Content elaboration: In the Physical Science course, coefficients were used to balance simple equations. Other representations, including Lewis structures and three-dimensional models, were also used and manipulated to demonstrate the conservation of matter in chemical reactions. In this course, more complex reactions will be studied, classified and represented with balanced chemical equations and three-dimensional models. Classifying reactions into types can be a helpful organizational tool for recognizing patterns of what may happen when two substances are mixed. Teachers should be aware that the common reaction classifications that are often used in high school chemistry courses may lead to misconceptions because they are not based on the actual chemistry, but on surface features that can be similar from one system to another (e.g., exchanging partners), even though the underlying chemistry is not the same. However, these classifications may be useful in making predictions about what happens when two substances are mixed. Some general types of chemical reactions are oxidation/reduction, synthesis, decomposition, single replacement, double replacement (including precipitation reactions and some acid-base neutralizations) and combustion reactions. Some reactions can fit into more than one category. For example, a single replacement reaction can also be classified as an oxidation/reduction reaction. Identification of reactions involving oxidation and reduction as well as indicating what substance is being oxidized and what is being reduced are appropriate in this course. However, balancing complex oxidation/reduction reactions is reserved for more advanced study. Organic molecules release energy when undergoing combustion reactions and are used to meet the energy needs of society (e.g., oil, gasoline, natural gas) and to provide the energy needs of biological organisms (e.g., cellular respiration). When a reaction between two ionic compounds in aqueous solution results in the formation of a precipitate or molecular compound, the reaction often occurs because the new ionic or covalent bonds are stronger than the original ion-dipole interactions of the ions in solution. Laboratory experiences (3-D or virtual) with different types of chemical reactions should be provided. Reactions occur when reacting particles collide in an appropriate orientation and with sufficient energy. The rate of a chemical reaction is the change in the amount of the reactants or products in a specific period of time. Increasing the probability or effectiveness of the collisions between the particles increases the rate of the reaction. Therefore, changing the concentration of the reactants, changing the temperature or the pressure of gaseous reactants, or using a catalyst, can change the reaction rate. Likewise, the collision theory can be applied to dissolving solids in a liquid solvent and can be used to explain why reactions are more likely to occur between reactants in the aqueous or gaseous state than between solids. The rate at which a substance dissolves should not be confused with the amount of solute that can dissolve in a given amount of solvent (solubility). Mathematical treatment of reaction rates is reserved for more advanced study. Computer simulations can help visualize reactions from the perspective of the kinetic-molecular theory. In middle school, the differences between potential and kinetic energy and the particle nature of thermal energy were introduced. For chemical systems, potential energy is in the form of chemical energy and kinetic energy is in the form of thermal energy. The total amount of chemical energy and/or thermal energy in a system is impossible to measure. However, the energy change of a system can be calculated from measurements (mass and change in temperature) from calorimetry experiments in the laboratory. Conservation of energy is an important component of calorimetry equations. Thermal energy is the energy of a system due to the movement of its particles. The thermal energy of an object depends upon the temperature and chemical composition. Some materials require little energy to change their temperature and other materials require a great deal to change their temperature by the same amount. Specific heat is a measure of how much energy is needed to change the temperature of a specific mass of material a specific amount. Specific heat values can be used to calculate the thermal energy change, the temperature (initial, final or change in) or mass of a material in calorimetry. Water has a particularly high specific heat capacity, which is important in regulating Earth's temperature. As studied in middle school, chemical energy is the potential energy associated with chemical systems. Chemical reactions involve valence electrons forming bonds to yield more stable products with lower energies. Energy is required to break interactions and bonds between the reactant atoms and energy is released when an interaction or bond is formed between the atoms in the products. Molecules with weak bonds (e.g., ATP) are less stable and tend to react to produce more stable products, releasing energy in the process. Generally, energy is transferred out of the system (exothermic) when the products have stronger bonds than the reactants and is transferred into the system (endothermic) when the reactants have stronger bonds than the products. Predictions of the energy requirements (endothermic or exothermic) of a reaction can be made given a table of bond energies. Graphic representations can be drawn and interpreted to represent the energy changes during a reaction. The role of energy in determining the spontaneity of chemical reactions is dealt with conceptually in this course. Entropy and its influence on the spontaneity of reactions are reserved for more advanced study. All reactions are reversible to a degree and many reactions do not proceed completely toward products but appear to stop progressing before the reactants are all used up. At this point, the amounts of the reactants and the products appear to be constant and the reaction can be said to have reached dynamic equilibrium. Dynamic equilibrium means the rate of the reverse reaction is equal to the rate of the forward reaction so there is no apparent change in the reaction. If a chemical system at equilibrium is disturbed by a change in the conditions of the system (e.g., increase or decrease in the temperature, pressure on gaseous equilibrium systems, concentration of a reactant or product), then the equilibrium system will respond by shifting to a new equilibrium state, reducing the effect of the change (Le Chatelier's Principle). If products are removed as they are formed during a reaction, then the equilibrium position of the system is forced to shift to favor the products. In this way, an otherwise unfavorable reaction can be made to occur. Mathematical treatment of equilibrium reactions is reserved for advanced study. Computer simulations can help visualize the progression of a reaction to dynamic equilibrium and the continuation of both the forward and reverse reactions after equilibrium has been attained. Properties of acids and bases and the ranges of the pH scale were introduced in Physical Science. In this course, the structural features of molecules are explored to further understand acids and bases. Acids often result when hydrogen is covalently bonded to an electronegative element and is easily dissociated from the rest of the molecule to bind with water to form a hydronium ion (H3O+). The acidity of an aqueous solution can be expressed as pH, where pH can be calculated from the concentration of the hydronium ion. Bases are likely to dissociate in water to form a hydroxide ion. Acids can react with bases to form a salt and water. Such neutralization reactions can be studied quantitatively by performing titration experiments.

Solubility Limits and Selective Precipitation

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