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

K12 Standards

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Ohio’s Learning Standards and Model Curriculum

9-12

Chemistry

Interactions Of Matter

Chemical reactions

Precipitation of Ions

Als de oplossingen van twee ionische verbindingen worden gemengd, zoals natriumjodide en lood(II)nitraat, kunnen de kationen van de ene oplossing worden gecombineerd met de anionen van de andere. Omdat een van de kruisproducten natriumnitraat in water oplosbaar is, blijven natrium-en nitraationen in oplossing, terwijl lood-en jodide-ionen een loodjodideneerslag vormen. Het reactiequotiënt, Q, van de oplossingsreactie is gelijk aan het product van de concentraties lood-en jodide-ionen.In tegenstelling tot het oplosbaarheidsproduct Ksp, omvat het reactiequotiënt ionconcentraties in elk stadium, niet alleen bij evenwicht. De waarden van Q en Ksp kunnen worden vergeleken om te voorspellen of er een neerslagreactie zal optreden. Overweeg een druppelsgewijze toevoeging van natriumjodide-oplossing aan de loodnitraatoplossing.In het begin is Q kleiner dan Ksp, en zowel lood-als jodide-ionen zijn in oplossing met de natrium-en nitraationen. Dit is een onverzadigde oplossing. Naarmate er meer natriumjodide wordt toegevoegd, blijft de concentratie van jodide-ionen toenemen.De reactie heeft een evenwicht bereikt wanneer Q gelijk is aan Ksp. In dit stadium is een kleine hoeveelheid vast loodjodide in dynamisch evenwicht met de ionen, waardoor een verzadigde oplossing wordt gevormd. Verdere toevoeging van natriumjodide maakt Q groter dan Ksp, en de reactie verschuift naar de neerslag.Dit is een oververzadigde oplossing, waarbij het neerslaan doorgaat totdat de ionenconcentraties zijn verlaagd tot hun evenwichtswaarden. Stel dat het mengen van natriumjodide en lood(II)nitraatoplossingen resulteert in een oplossing die 1, 6 10⁻⁴ molaire loodionen en 4, 0 10⁻⁴ molaire jodide-ionen bevat. Hier is Q gelijk aan 2, 6 10⁻¹¹, terwijl Ksp voor loodjodide 1, 4 10⁻⁸ is.Omdat Q kleiner is dan Ksp, zal loodjodide niet neerslaan. Het voorspellen van neerslagreacties kan erg nuttig zijn tijdens de scheiding van ionaire verbindingen. Overweeg een oplossing met twee metaalionen:lood(II)en koper(II)Als aan deze oplossing zoutzuur wordt toegevoegd, slaat lood(II)chloride neer omdat het een kleine Ksp heeft, terwijl koper in oplossing blijft omdat koper(II)chloride zeer goed oplosbaar is in water.Deze techniek wordt selectieve neerslag genoemd.

neerslag van ionen

Neerslag voorspellen 
De vergelijking die het evenwicht beschrijft tussen vast calciumcarbonaat en zijn gesolvateerde ionen is: 
<img alt="Eq1" src="/files/ftp_upload/11420/11420_Equation_1.png" style="text-align: center; height: 50px;"> 
Het is belangrijk om te beseffen dat dit evenwicht tot stand komt in elke waterige oplossing die Ca2+- en CO32–-ionen bevat, en niet alleen in een oplossing die wordt gevormd door water te verzadigen met calciumcarbonaat. Denk bijvoorbeeld aan het mengen van waterige oplossingen van de oplosbare verbindingen natriumcarbonaat en calciumnitraat. Als de concentraties calcium- en carbonaationen in het mengsel geen reactiequotiënt Q opleveren dat het oplosbaarheidsproduct Ksp overschrijdt, zal er geen neerslag optreden. Als de ionenconcentraties een reactiequotiënt opleveren dat groter is dan het oplosbaarheidsproduct, zal er neerslag optreden, waardoor die concentraties worden verlaagd totdat een evenwicht is bereikt (Q = Ksp). De vergelijking van Q met Ksp om neerslag te voorspellen is een voorbeeld van de algemene benadering voor het voorspellen van de richting van een reactie. Voor het specifieke geval van oplosbaarheidsevenwichten: 
Q < Ksp: de reactie verloopt in voorwaartse richting (oplossing is niet verzadigd; geen neerslag waargenomen) Q > Ksp: de reactie verloopt in de omgekeerde richting (oplossing is oververzadigd; er zal neerslag optreden) In oplossingen die twee of meer ionen bevatten die onoplosbare verbindingen kunnen vormen met hetzelfde tegenion, kan een experimentele strategie, selectieve precipitatie genaamd, worden gebruikt om individuele ionen uit de oplossing te verwijderen. Door de tegenionenconcentratie op een gecontroleerde manier te verhogen, kunnen ionen in oplossing individueel worden neergeslagen, ervan uitgaande dat hun oplosbaarheid van verbindingen voldoende verschillend is. In oplossingen met gelijke concentraties doelionen zal het ion dat de minst oplosbare verbinding vormt als eerste neerslaan (bij de laagste concentratie tegenion). De andere ionen slaan vervolgens neer als de oplosbaarheid van hun verbinding wordt bereikt. Neerslag van zilverhalogeniden Een oplossing bevat 0,00010 mol KBr en 0,10 mol KCl per liter. Aan deze oplossing wordt geleidelijk AgNO3 toegevoegd. Wat vormt het eerst, vast AgBr of vast AgCl? De twee betrokken evenwichten zijn: <img alt="Eq2" src="/files/ftp_upload/11420/11420_Equation_2.png" style="text-align: center; height: 90px;"> Als de oplossing ongeveer gelijke concentraties Cl– en Br– zou bevatten, zou het zilverzout met de kleinere Ksp (AgBr) als eerste neerslaan. De concentraties zijn echter niet gelijk, dus moeten de [Ag+] waarbij AgCl begint te precipiteren en de [Ag+] waarbij AgBr begint te precipiteren worden berekend. Het zout dat zich aan de lagere [Ag+] vormt, slaat als eerste neer. AgBr slaat neer wanneer Q = Ksp voor AgBr <img alt="Eq3" src="/files/ftp_upload/11420/11420_Equation_3.png" style="text-align: center; height: 112px;"> AgBr begint te precipiteren wanneer [Ag+] 5,0 × 10−9 M bedraagt. Voor AgCl: AgCl slaat neer wanneer Q gelijk is aan Ksp voor AgCl (1,6 × 10-10). Wanneer [Cl–] = 0,10 M: <img alt="Eq4" src="/files/ftp_upload/11420/11420_Equation_4.png" style="text-align: center; height: 112px;"> AgCl begint te precipiteren wanneer [Ag+] 1,6 × 10-9 M is. AgCl begint te precipiteren bij een lagere [Ag+] dan AgBr, dus AgCl begint eerst te precipiteren. Merk op dat de chloride-ionenconcentratie van het aanvankelijke mengsel aanzienlijk groter was dan de bromide-ionenconcentratie, en dus sloeg zilverchloride als eerste neer, ondanks dat het een Ksp had die groter was dan die van zilverbromide. Deze tekst is aangepast van <a href="https://openstax.org/books/chemistry-2e/pages/15-1-precipitation-and-dissolution">Openstax, Chemistry 2e, Hoofdstuk 15.1: Neerslag en oplossende kracht</a>.

Neerslag van ionen

Predicting Precipitation
The equation that describes the equilibrium between solid calcium carbonate and its solvated ions is:

It is important to realize ...

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Acid-base and Solubility Equilibria

Zuur-base- en oplosbaarheidsevenwichten

KEY TERMS AND CONCEPTS

BELANGRIJKE TERMEN EN CONCEPTEN

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...

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...

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,...

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

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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. 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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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