Name: predicting precipitation with ksp
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 3 min 7 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

Solubility Equilibria

Il cloruro di sodio è considerato solubile perché grandi quantità di esso si dissolvono in acqua, ma quando all'acqua viene aggiunto il cloruro di piombo, solo una piccola quantità si dissolve, mentre il resto rimane insolubile. Il solido non disciolto coesiste con gli ioni piombo e cloruro che sono in soluzione. Parte del cloruro di piombo solido continua a dissolversi, mentre alcuni degli ioni nella soluzione si ricombinano per formare un precipitato.

Quando la velocità di dissoluzione è uguale alla velocità di precipitazione, si stabilisce un equilibrio di solubilità. La costante di equilibrio può essere calcolata dalle concentrazioni di equilibrio degli ioni in base alla reazione di dissoluzione dove il cloruro di piombo si dissocia in uno ione piombo e due ioni cloruro. Quindi, la costante di equilibrio è data dalla concentrazione molare degli ioni piombo moltiplicata per il quadrato della concentrazione molare degli ioni cloruro.

Poiché la concentrazione del cloruro di piombo solido rimane costante, viene esclusa dal calcolo. Questa costante di equilibrio è detta prodotto di solubilità, indicata con Ksp. A 25 C, il Ksp del cloruro di piombo è 1, 17 10⁻⁵.

Il valore di Ksp rappresenta la misura in cui un composto può dissolversi per formare una soluzione acquosa satura. A una data temperatura, il Ksp di un composto è costante. La solubilità di un composto in moli per litro, nota come solubilità molare, viene spesso usata per esprimere la concentrazione del solido disciolto in una soluzione satura.

La solubilità di un composto può variare a seconda di fattori, come il pH della soluzione e se sono presenti altri ioni. La solubilità molare di un composto, x, può essere calcolata dal suo Ksp, utilizzando una tabella ICE. Le concentrazioni iniziali di ioni piombo e ioni cloruro nella soluzione sono zero.

In equilibrio, la concentrazione molare degli ioni piombo è rappresentata da x, mentre quella degli ioni cloruro è 2x. Sostituendo nell'espressione di equilibrio, il prodotto di solubilità per il cloruro di piombo è uguale a x volte 2x², che è uguale a 4x³. Poiché il Ksp per il cloruro di piombo è 1, 17 10⁻⁵, x è risolto in 1, 43 10⁻² molare.

Per i composti che hanno la stessa stechiometria di dissociazione, come il cloruro di piombo e il fluoruro di calcio, dove 1 mole di ciascun composto produce 3 moli di ioni disciolti, i rispettivi valori Ksp possono essere utilizzati direttamente per confrontare le loro solubilità relative.

<ol>
	<li>Clifford, A. F. "La previsione delle costanti del prodotto di solubilità1." Giornale della Società Chimica Americana 79, n. 20 (1957): 5404-5407.</li>
	<li>Hill, Arthur E. "L'incostanza del prodotto di solubilità". Giornale della Società Chimica Americana 32, n. 10 (1910): 1186-1193.</li>
	<li>Reynolds, John P. "Esperimento Ksp: il prodotto di solubilità per l'idrossido di bario." Journal of Chemical Education 52, n. 8 (1975): 521.</li>
</ol>

equilibri di solubilità

Gli equilibri di solubilità vengono stabiliti quando la dissoluzione e la precipitazione di una specie di soluto avvengono a velocità uguali. Questi equilibri sono alla base di molti processi naturali e tecnologici, che vanno dalla carie alla depurazione dell’acqua. La comprensione dei fattori che influenzano la solubilità dei composti è quindi essenziale per la gestione efficace di questi processi. Questa sezione applica i concetti e gli strumenti di equilibrio precedentemente introdotti ai sistemi che coinvolgono dissoluzione e precipitazione.

Il prodotto di solubilità

Ricordiamo che la solubilità di una sostanza può variare da essenzialmente zero (insolubile o scarsamente solubile) a infinito (miscibile). Un soluto con solubilità finita può produrre una soluzione satura quando viene aggiunto a un solvente in una quantità superiore alla sua solubilità, risultando in una miscela eterogenea della soluzione satura e del soluto non disciolto in eccesso. Ad esempio, una soluzione satura di cloruro d'argento è quella in cui è stato stabilito l'equilibrio mostrato di seguito.

<img alt="Eq1" src="/files/ftp_upload/11417/11417_Equation_1.png" style="text-align: center; height: 50px;" />

In questa soluzione, un eccesso di AgCl solido si dissolve e si dissocia per produrre ioni Ag^+ e Cl^– acquosi alla stessa velocità con cui quest’ultimi si combinano e precipitano per formare AgCl solido. Poiché il cloruro d'argento è un sale scarsamente solubile, la concentrazione di equilibrio dei suoi ioni disciolti nella soluzione è relativamente bassa.

La costante di equilibrio per equilibri di solubilità come questo è chiamata costante del prodotto di solubilità, K_sp, in questo caso:

<img alt="Eq2" src="/files/ftp_upload/11417/11417_Equation_2.png" style="text-align: center; height: 50px;" />

Ricordiamo che solo i gas e i soluti sono rappresentati nelle espressioni della costante di equilibrio, quindi il K_sp non include un termine per l'AgCl non disciolto.

K_sp e solubilità

Il K_sp di un composto ionico leggermente solubile può essere semplicemente correlato alla sua solubilità misurata a condizione che il processo di dissoluzione implichi solo dissociazione e solvatazione, ad esempio:

<img alt="Eq3" src="/files/ftp_upload/11417/11417_Equation_3.png" style="text-align: center; height: 50px;" />

Per casi come questi, è possibile derivare i valori K_sp dalle solubilità fornite o viceversa. Calcoli di questo tipo vengono eseguiti più convenientemente utilizzando la solubilità molare di un composto, misurata come moli di soluto disciolto per litro di soluzione satura.

Prevedere le precipitazioni

L’equazione che descrive l’equilibrio tra il carbonato di calcio solido e i suoi ioni solvatati è:

<img alt="Eq4" src="/files/ftp_upload/11417/11417_Equation_4.png" style="text-align: center; height: 50px;" />

È importante rendersi conto che questo equilibrio viene stabilito in qualsiasi soluzione acquosa contenente ioni Ca^2+ e CO^3^2–, non solo in una soluzione formata saturando l'acqua con carbonato di calcio. Consideriamo, ad esempio, la miscelazione di soluzioni acquose di composti solubili come il carbonato di sodio e il nitrato di calcio. Se le concentrazioni di ioni calcio e carbonato nella miscela non danno un quoziente di reazione, Q, che supera il prodotto di solubilità, K_sp, allora non si verificherà alcuna precipitazione. Se le concentrazioni di ioni producono un quoziente di reazione maggiore del prodotto di solubilità, si verificherà la precipitazione, abbassando tali concentrazioni fino a stabilire l'equilibrio (Q = K_sp). Il confronto tra Q e K_sp per prevedere le precipitazioni è un esempio dell'approccio generale alla previsione della direzione di una reazione introdotto per la prima volta nelle lezioni sull'equilibrio. Per il caso specifico degli equilibri di solubilità:

Q < K_sp: la reazione procede in avanti (la soluzione non è satura; non si osserva alcuna precipitazione)

Q > K_sp: la reazione procede in senso inverso (la soluzione è sovrasatura; si avrà precipitazione)

Questo testo è adattato da <a href="https://openstax.org/books/chemistry-2e/pages/15-1-precipitation-and-dissolution">Openstax, Chemistry 2e, Section 15.1: Precipitation and Dissolution.</a>

Equilibri di solubilità

Gli equilibri di solubilità si stabiliscono quando la dissoluzione e la precipitazione di una specie di soluto avvengono a velocità uguali. Questi equilibri sono alla base di molti ...

Education

Didattica

JoVE Core

Chimica

Acid-base and Solubility Equilibria

Equilibri acido-base e di solubilità

KEY TERMS AND CONCEPTS

TERMINI E CONCETTI CHIAVE

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 solution and excess solid. Silver chloride is a classic example. In that equilibrium, solid AgCl dissolves into Ag+ and Cl− ions, and those ions also combine to form solid AgCl. For sparingly soluble salts like this, the dissolved ion concentration stays low.

The equilibrium constant for a solubility equilibrium is called the solubility product constant, or Ksp. Only gases and dissolved solutes appear in equilibrium expressions, so the undissolved solid is not included. For slightly soluble ionic compounds, Ksp can be related to measured solubility when dissolution involves only dissociation and solvation. In those cases, molar solubility is useful. It is the number of moles of dissolved solute in one liter of saturated solution.

Ksp also helps predict whether a precipitate will form. The equilibrium between solid calcium carbonate and its dissolved ions can exist in any aqueous solution that contains Ca2+ and CO32− ions. That is true even if the solution did not start with calcium carbonate as a solid. For example, when aqueous sodium carbonate and calcium nitrate are mixed, the ion concentrations determine the outcome.

The reaction quotient, Q, is compared with Ksp to test for precipitation. If Q does not exceed Ksp, no precipitation occurs. If Q is greater than Ksp, precipitation occurs and lowers the ion concentrations until equilibrium is reached again, when Q equals Ksp. In this way, Q and Ksp show whether a solution is unsaturated, saturated, or supersaturated.

Predicting Precipitation with Ksp

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

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

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 Ksp, molar solubility, and the reaction quotient Q help predict when a saturated solution forms or when a precipitate will appear.

Predicting Precipitation with Ksp | JoVE K12

chemistry

JoVE Core: Chimica illustra concetti chiave di chimica attraverso video animazioni concise e di immediata comprensione. Inoltre, i video della sezione "scienziati in azione" mostrano procedure metodologiche realizzate da ricercatori provenienti dai migliori laboratori internazionali.

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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.

Predicting Precipitation with Ksp

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