Name: common ion effects on solubility
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Ohio’s Learning Standards and Model Curriculum

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Chemistry

Interactions Of Matter

Chemical reactions

Factors Affecting Solubility

A solubilidade de um composto iónico pode variar consoante sobre outros solutos presente na solução. O soluto pode ser um íon comum para o composto, ou um ácido ou base. Assim, os dois fatores principais que afetam a solubilidade são o efeito iónico comum e o pH de uma solução.Se cloreto de chumbo for adicionado a uma solução de cloreto de sódio, ambos os sais ionizarão na água, produzindo sódio e catíons de chumbo e o aníon comum, cloreto. Porque os íons do cloreto na solução vêm tanto do completo ionização do cloreto de sódio e a ionização parcial de cloreto de chumbo a sua concentração seria mais alta do que a de qualquer um dos íons de sódio ou os íons de chumbo. Para compensar, o equilíbrio entre o cloreto de chumbo sólido e os seus íons na solução desloca-se para o sal não dissociado, causando mais cloreto de chumbo que fica por dissolver.Assim, a presença de um íon comum diminui a solubilidade de uma substância moderadamente solúvel. Por exemplo, a solubilidade molar, x, do cloreto de chumbo em 0, 100 molar de uma solução de cloreto de sódio pode ser calculada a partir da tabela ICE. A concentração inicial dos íons de chumbo na solução é zero, enquanto o cloreto é 0, 100 molar.Cada molécula de cloreto de chumbo irá dissociar-se num só chumbo e dois íons de cloreto. Assim, a alteração na concentração de íons de chumbo será x e do cloreto será 2x. A concentração de equilíbrio dos íons de chumbo será x, mas para os íons de cloreto, será a soma de 2x e 0, 100.A 25 C, o valor de Ksp para o cloreto de chumbo é 1, 17 10-⁵, e a expressão de equilíbrio é o produto das concentrações de equilíbrio dos íons de chumbo e de cloreto, que é x vezes o quadrado de Porque a Ksp é pequena, 2x é muito menor que 0, 100 molar e a sua soma pode ser aproximado a 0, 100 molar. A substituição no rendimento de expressão 1, 17 10-⁵ igual a x vezes 0, 100 molar ao quadrado. Solucionando para x, a solubilidade molar do cloreto de chumbo em 0, 100 molar cloreto de sódio a solução é 1, 17 10-³ molar.Em comparação com a solubilidade molar do cloreto de chumbo na água, que é 1, 43 10-² molar, o íon comum diminui a solubilidade do cloreto de chumbo 12 vezes. A solubilidade de um pode também ser afetado pelo pH da solução. Consideremos a dissociação parcial do hidróxido de cálcio em íons de cálcio e hidróxido de cálcio.Se o pH for aumentado, em adição ao hidróxido de potássio por exemplo, o íon hidróxido comum irá baixar a solubilidade do hidróxido de cálcio. Inversamente, se o pH é diminuído, acrescentando, por exemplo ácido clorídrico, os prótons combinarão com íons de hidróxidos, baixando a concentração de íons hidróxidos na solução. O equilíbrio da solubilidade do hidróxido de cálcio mudaria então em relação aos produtos, levando a um aumento da solubilidade numa solução ácida.

<ol>
	<li>Koubek, E. "Demonstração do Efeito do Íon Comum". Jornal de educação química 70, nº 2 (1993): 155.</li>
	<li>Amaral, L. F., I. R. Oliveira, R. Salomão, E. Frollini e V. C. Pandolfelli. "Temperatura e efeito de íons comuns na hidratação do óxido de magnésio (MgO)." Ceramics International 36, nº 3 (2010): 1047-1054.</li>
	<li>Cassens, Jan, Anke Prudic, Feelly Ruether e Gabriele Sadowski. "Solubilidade de produtos farmacêuticos e seus sais em função do pH." Pesquisa em Química Industrial e de Engenharia 52, nº 7 (2013): 2721-2731.</li>
</ol>

factores que afetam a solubilidade

Em comparação com a água pura, a solubilidade de um composto iónico é menor em soluções aquosas contendo um ião comum (um também produzido por dissolução do composto iónico). Este é um exemplo de um fenómeno conhecido como efeito de ião comum, que é uma consequência da lei de ação das massas que pode ser explicada utilizando o princípio de Le Chȃtelier. Considere a dissolução do iodeto de prata:
<img alt="Eq1" src="/files/ftp_upload/11418/11418_Equation_1.png" style="text-align: center; height: 50px;" />
Este equilíbrio de solubilidade pode ser deslocado pela adição de iões de prata ou iodeto, resultando na precipitação de AgI e na diminuição das concentrações de Ag+ e I– dissolvidos. Em soluções que já contêm qualquer destes iões, menos AgI pode ser dissolvido do que em soluções sem estes iões.
Este efeito pode também ser explicado em termos de ação das massas, tal como representado na expressão do produto de solubilidade:
<img alt="Eq2" src="/files/ftp_upload/11418/11418_Equation_2.png" style="text-align: center; height: 50px;" />
O produto matemático das molaridades dos iões de prata e iodeto é constante em uma mistura de equilíbrio, independentemente da fonte dos iões, e assim um aumento na concentração de um ião deve ser equilibrado por uma diminuição proporcional no outro.
O Papel da Precipitação no Tratamento de Águas Residuais
Os equilíbrios de solubilidade são ferramentas úteis no tratamento de águas residuais realizadas em instalações que podem tratar a água municipal em uma cidade ou município. Especificamente, a precipitação selectiva é utilizada para remover contaminantes das águas residuais antes de serem libertadas para corpos naturais de água. Por exemplo, os iões fosfato (PO43−) estão frequentemente presentes na água descarregada das instalações de fabrico. Uma abundância de fosfato faz com que algas cresçam em excesso, o que afeta a quantidade de oxigénio disponível para a vida marinha, assim como torna a água inadequada para consumo humano.
Uma maneira comum de remover fosfatos da água é pela adição de hidróxido de cálcio, ou cal, Ca(OH)2. À medida que a água se torna mais básica, os iões de cálcio reagem com iões fosfato para produzir hidroxiapatita, Ca5(PO4)3·OH, que depois precipita da solução:
<img alt="Eq3" src="/files/ftp_upload/11418/11418_Equation_3.png" style="text-align: center; height: 50px;" />
Como a quantidade de iões de cálcio adicionada não excede os produtos de solubilidade para outros sais de cálcio, os aniões desses sais permanecem para trás nas águas residuais. O precipitado é então removido por filtração, e a água é trazida de volta para um pH neutro pela adição de CO2 em um processo de recarbonização. Outros produtos químicos também podem ser utilizados para a remoção de fosfatos por precipitação, incluindo cloreto de ferro(III) e sulfato de alumínio.
 Este texto é adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/15-1-precipitation-and-dissolution">Openstax, Chemistry 2e, Section 15.1: Precipitation and Dissolution.</a>

Factores que Afetam a Solubilidade

Comparado com a água pura, a solubilidade de um composto iônico é menor em soluções aquosas contendo um íon comum (um também produzido pela dissolução de ...

Education

JoVE Core

Química

Acid-base and Solubility Equilibria

Equilíbrios Ácido-Base e de Solubilidade

KEY TERMS AND CONCEPTS

TERMOS-CHAVE E CONCEITOS

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

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

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

chemistry

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Physical Science

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actors Affecting Dissolution Rates

Chemical Equations

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Reaction Stoichiometry

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Limiting Reactant

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Chemical Reactions in Aqueous Solutions

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Oxidation&ndash;Reduction Reactions

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Oxidation Numbers

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The Integrated Rate Law: The Dependence of Concentration on Time

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Internal Energy

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Dynamic Equilibrium

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The Equilibrium Constant

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Homogeneous Equilibria for Gaseous Reactions

Homogeneous Equilibria for Gaseous Reactions
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Calculating the Equilibrium Constant

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Reaction Quotient

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Relative Strengths of Conjugate Acid-Base Pairs

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Calculating Equilibrium Concentrations

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Strong Acid and Base Solutions

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The Small x Assumption

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Weak Acid Solutions

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Le Chatelier's Principle: Changing Concentration

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Weak Base Solutions

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Ions as Acids and Bases

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Bronsted-Lowry Acids and Bases

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Determining the pH of Salt Solutions

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Acid/Base Strengths and Dissociation Constants

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Polyprotic Acids

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Water: A Bronsted-Lowry Acid and Base

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Acid Strength and Molecular Structure

Binary Acids and Bases
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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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