Name: common ion effects on solubility
Uploaded: 2020-09-24T15:06:51.000Z
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Ohio’s Learning Standards and Model Curriculum

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Chemistry

Interactions Of Matter

Chemical reactions

Factors Affecting Solubility

La solubilità di un composto ionico può variare a seconda degli altri soluti presenti nella soluzione. Il soluto può essere uno ione comune al composto, oppure un acido o una base. Pertanto, due fattori principali che influenzano la solubilità sono l'effetto ionico comune e il pH di una soluzione.

Se il cloruro di piombo è aggiunto ad una soluzione di cloruro di sodio, ambo i sali si ionizzano nell'acqua, producendo sodio, cationi di piombo e l'anione comune, il cloruro. Poiché gli ioni cloruro in soluzione provengono sia dalla completa ionizzazione del cloruro di sodio sia dalla ionizzazione parziale del cloruro di piombo la loro concentrazione sarebbe maggiore a quella degli ioni sodio o degli ioni piombo. Per compensare, l'equilibrio fra il cloruro di piombo solido e i suoi ioni in soluzione si sposta verso il sale non dissociato, facendo sì che più cloruro di piombo rimanga indisciolto.

Pertanto, la presenza di uno ione comune riduce la solubilità di una sostanza scarsamente solubile. Per esempio, la solubilità molare, x, del cloruro di piombo in una soluzione di cloruro di sodio 0, 100 molare può essere calcolata dalla tabella ICE. La concentrazione iniziale di ioni piombo nella soluzione è zero, mentre il cloruro è 0, 100 molare.

Ogni molecola di cloruro di piombo si dissocerà in uno ione piombo e due ioni cloruro. Quindi, il cambiamento nella concentrazione degli ioni piombo sarà x e quello del cloruro sarà 2x. La concentrazione di equilibrio degli ioni piombo sarà x, ma per gli ioni cloruro sarà la somma di 2x e 0.100.

A 25 gradi Celsius, il valore di Ksp per il cloruro di piombo è 1, 17 10⁻⁵ e l'espressione di equilibrio è il prodotto delle concentrazioni di equilibrio di ioni di piombo e cloruro, che è x volte il quadrato di Poiché il Ksp è piccolo, 2x è molto più piccolo di 0, 100 molare e la loro somma può essere approssimata a 0, 100 molare. La sostituzione nell'espressione produce 1, 17 10⁻⁵ è uguale a x volte 0, 100 molare al quadrato. Risolvendo per x, la solubilità molare del cloruro di piombo in una soluzione di cloruro di sodio molare 0, 100 è 1, 17 10⁻³ molare.

Rispetto alla solubilità molare del cloruro di piombo in acqua, che è 1, 43 10 ² molare, lo ione comune riduce la solubilità del cloruro di piombo di 12 volte. La solubilità di un composto può anche essere influenzata dal pH della soluzione. Considerate la parziale dissociazione dell'idrossido di calcio in ioni di calcio e idrossido.

Se il pH viene aumentato, aggiungendo per esempio idrossido di potassio, il comune ione idrossido abbasserà la solubilità dell'idrossido di calcio. Viceversa, se il pH è diminuito, per esempio aggiungendo acido cloridrico, i protoni si combineranno con ioni idrossido, abbassando la concentrazione di ioni idrossido nella soluzione. L'equilibrio di solubilità dell'idrossido di calcio si sposterebbe quindi verso i prodotti, portando a una maggiore solubilità in una soluzione acida.

<ol>
	<li>Koubek, E. "Dimostrazione dell'effetto ionico comune". Giornale di educazione chimica 70, n. 2 (1993): 155.</li>
	<li>Amaral, LF, IR Oliveira, R. Salomão, E. Frollini e VC Pandolfelli. "Temperatura ed effetto degli ioni comuni sull'idratazione dell'ossido di magnesio (MgO)". Ceramics International 36, n. 3 (2010): 1047-1054.</li>
	<li>Cassens, Jan, Anke Prudic, Feelly Ruether e Gabriele Sadowski. "Solubilità dei prodotti farmaceutici e dei loro sali in funzione del pH". Industrial & Engineering Chemistry Research 52, n. 7 (2013): 2721-2731.</li>
</ol>

fattori che influenzano la solubilità

Rispetto all'acqua pura, la solubilità di un composto ionico è inferiore nelle soluzioni acquose contenenti uno ione comune (prodotto anche dalla dissoluzione del composto ionico). Questo è un esempio di un fenomeno noto come effetto dello ione comune, che è una conseguenza della legge dell’azione di massa che può essere spiegata utilizzando il principio di Le Chȃtelier. Consideriamo la dissoluzione dello ioduro d'argento:

<img alt="Eq1" src="/files/ftp_upload/11418/11418_Equation_1.png" style="text-align: center; height: 50px;" />

Questo equilibrio di solubilità può essere spostato a sinistra mediante l'aggiunta di ioni argento o ioduro, con conseguente precipitazione di AgI e riduzione delle concentrazioni di Ag^+ e I^– disciolti. Nelle soluzioni che già contengono uno di questi ioni, può essere disciolta meno AgI rispetto alle soluzioni senza questi ioni.

Questo effetto può anche essere spiegato in termini di azione di massa come rappresentato nell'espressione del prodotto di solubilità:

<img alt="Eq2" src="/files/ftp_upload/11418/11418_Equation_2.png" style="text-align: center; height: 50px;" />

Il prodotto matematico delle molarità degli ioni argento e ioduro è costante in una miscela di equilibrio indipendentemente dalla fonte degli ioni, quindi un aumento della concentrazione di uno ione deve essere bilanciato da una diminuzione proporzionale dell'altro.

Il ruolo delle precipitazioni nel trattamento delle acque reflue

Gli equilibri di solubilità sono strumenti utili nel trattamento delle acque reflue effettuato in impianti che possono trattare l'acqua comunale in una città o paese. Nello specifico, la precipitazione selettiva viene utilizzata per rimuovere i contaminanti dalle acque reflue prima che vengano rilasciate nuovamente nei corpi idrici naturali. Ad esempio, gli ioni fosfato (PO_4^3−) sono spesso presenti nell'acqua scaricata dagli impianti di produzione. L’abbondanza di fosfati provoca la crescita di alghe in eccesso, il che influisce sulla quantità di ossigeno disponibile per la vita marina e rende l’acqua inadatta al consumo umano.

Un modo comune per rimuovere i fosfati dall'acqua è mediante l'aggiunta di idrossido di calcio, o calce, Ca(OH)_2. Quando l'acqua diventa più basica, gli ioni calcio reagiscono con gli ioni fosfato per produrre idrossiapatite, Ca_5(PO_4)_3·OH, che poi precipita dalla soluzione:

<img alt="Eq3" src="/files/ftp_upload/11418/11418_Equation_3.png" style="text-align: center; height: 50px;" />

Poiché la quantità di ioni calcio aggiunti non comporta un superamento dei prodotti di solubilità per altri sali di calcio, gli anioni di tali sali rimangono nelle acque reflue. Il precipitato viene quindi rimosso mediante filtrazione e l'acqua viene riportata a pH neutro mediante l'aggiunta di CO_2 in un processo di ricarbonatazione. Per la rimozione dei fosfati mediante precipitazione possono essere utilizzati anche altri prodotti chimici, tra cui il cloruro di ferro (III) e il solfato di alluminio.

 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>

Fattori che influenzano la solubilità

Rispetto all'acqua pura, la solubilità di un composto ionico è minore nelle soluzioni acquose contenenti uno ione comune (prodotto anche dalla dissoluzione di ...

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Acid-base and Solubility Equilibria

Equilibri acido-base e di solubilità

KEY TERMS AND CONCEPTS

TERMINI E CONCETTI CHIAVE

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 precipitates out. The result is lower concentrations of dissolved Ag+ and I− than in pure water.

The same idea appears in the solubility product expression, or Ksp. In an equilibrium mixture, the product of the silver ion and iodide ion molarities stays constant no matter where the ions came from. If one ion’s concentration rises, the other must fall to keep the equilibrium balance.

Solubility equilibria also help in wastewater treatment. Facilities often use selective precipitation to remove contaminants before water is released into rivers or lakes. One important target is phosphate ion, PO43−, because too much phosphate can cause excess algae growth, reduce oxygen for marine life, and make water unsafe for people.

A common treatment step is adding calcium hydroxide, or lime, Ca(OH)2. This makes the water more basic and allows calcium ions to react with phosphate ions to form hydroxylapatite, Ca5(PO4)3·OH, which precipitates out. The calcium added is kept low enough that other calcium salts do not exceed their solubility products, so their anions stay in the wastewater.

After filtration removes the solid, the water is brought back to a neutral pH by adding CO2 in a recarbonation step. Other chemicals can also remove phosphate by precipitation. Iron(III) chloride and aluminum sulfate are two examples.

Common Ion Effects on Solubility

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

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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

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 the common ion effect changes solubility, why AgI precipitates, and how selective precipitation removes phosphate in wastewater treatment.

Common Ion Effects on Solubility | JoVE K12

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actors Affecting Dissolution Rates

Chemical Equations

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Homogeneous Equilibria for Gaseous Reactions

Homogeneous Equilibria for Gaseous Reactions
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Relative Strengths of Conjugate Acid-Base Pairs

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Strong Acid and Base Solutions

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Determining the pH of Salt Solutions

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pH Scale

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Chemical reactions
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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.

Common Ion Effects on Solubility

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