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

Chlorek sodu jest uważany za rozpuszczalny, ponieważ duże jego ilości rozpuszczają się w wodzie, ale po dodaniu chlorku ołowiu do wody rozpuszcza się tylko niewielka ilość, podczas gdy reszta pozostaje nierozpuszczalna.
Nierozpuszczone ciało stałe współistnieje z jonami ołowiu i chlorku znajdującymi się w roztworze. Część stałego chlorku ołowiu nadal się rozpuszcza, podczas gdy niektóre jony w roztworze rekombinują, tworząc osad.
Gdy szybkość rozpuszczania jest równa szybkości wytrącania, ustala się równowagę rozpuszczalności.
Stałą równowagi można obliczyć na podstawie stężeń równowagowych jonów zgodnie z reakcją rozpuszczania - w której chlorek ołowiu dysocjuje na jeden jon ołowiu i dwa jony chlorkowe.
Tak więc stała równowagi jest określona przez stężenie molowe jonów ołowiu pomnożone przez kwadrat molowego stężenia jonów chlorkowych. Ponieważ stężenie stałego chlorku ołowiu pozostaje stałe, jest on wyłączony z obliczeń.
Ta stała równowagi nazywana jest iloczynem rozpuszczalności, oznaczanym przez Ksp. W temperaturze 25 °C Ksp chlorku ołowiu wynosi 1,17 × 10−5.
Wartość Ksp reprezentuje stopień, w jakim związek może się rozpuścić, tworząc nasycony roztwór wodny. W danej temperaturze Ksp związku jest stałe.
Rozpuszczalność związku w molach na litr, znana jako rozpuszczalność molowa, jest często używana do wyrażenia stężenia rozpuszczonego ciała stałego w nasyconym roztworze. Rozpuszczalność związku może się różnić w zależności od czynników, takich jak pH roztworu i obecność innych jonów.
Rozpuszczalność molową związku, x, można obliczyć na podstawie jego Ksp za pomocą tabeli ICE.
Początkowe stężenia jonów ołowiu i jonów chlorkowych w roztworze wynoszą zero.
W stanie równowagi stężenie molowe jonów ołowiu jest reprezentowane przez x, podczas gdy stężenie jonów chlorkowych wynosi 2x.
Podstawiając do wyrażenia równowagi, iloczyn rozpuszczalności chlorku ołowiu jest równy x razy 2x2, co równasię 4 x3.
Ponieważ Ksp dla chlorku ołowiu wynosi 1,17 × 10−5, x jest rozpuszczone jako 1,43 × 10−2 molowe.
W przypadku związków, które mają taką samą stechiometrię dysocjacji, takich jak chlorek ołowiu i fluorek wapnia, gdzie 1 mol każdego związku wytwarza 3 mole rozpuszczonych jonów, odpowiednie wartości Ksp można bezpośrednio wykorzystać do porównania ich względnych rozpuszczalności.

<ol>
	<li>Clifford, A. F. "Przewidywanie stałych produktu rozpuszczalności1." Journal of the American Chemical Society 79, nr 20 (1957): 5404-5407.</li>
	<li>Hill, Arthur E. "Niestałość produktu rozpuszczalności". Journal of the American Chemical Society 32, nr 10 (1910): 1186-1193.</li>
	<li>Reynolds, John P. "Eksperyment Ksp: Produkt rozpuszczalności wodorotlenku baru." Journal of Chemical Education 52, nr 8 (1975): 521.</li>
</ol>

równowagi rozpuszczalności

Równowagi rozpuszczalności powstają, gdy rozpuszczanie i wytrącanie substancji rozpuszczonej następuje z równą szybkością. Te równowagi leżą u podstaw wielu procesów naturalnych i technologicznych, począwszy od próchnicy zębów po oczyszczanie wody. Zrozumienie czynników wpływających na rozpuszczalność związku jest zatem niezbędne do skutecznego zarządzania tymi procesami. W tej sekcji zastosowano wcześniej wprowadzone koncepcje i narzędzia równowagi do układów obejmujących rozpuszczanie i wytrącanie.

Produkt rozpuszczalności

Przypomnijmy, że rozpuszczalność substancji może zmieniać się od zasadniczo zera (nierozpuszczalna lub trudno rozpuszczalna) do nieskończoności (mieszalna). Substancja rozpuszczona o skończonej rozpuszczalności może dać roztwór nasycony, gdy zostanie dodana do rozpuszczalnika w ilości przekraczającej jej rozpuszczalność, co daje niejednorodną mieszaninę nasyconego roztworu i nadmiaru nierozpuszczonej substancji rozpuszczonej. Na przykład nasycony roztwór chlorku srebra to taki, w którym została ustalona równowaga pokazana poniżej.

<img alt="Eq1" src="/files/ftp_upload/11417/11417_Equation_1.png" style="text-align: center; height: 50px;" />

W tym roztworze nadmiar stałego AgCl rozpuszcza się i dysocjuje, tworząc wodne jony Ag+ i Cl– z taką samą szybkością, z jaką te wodne jony łączą się i wytrącają, tworząc stały AgCl. Ponieważ chlorek srebra jest solą trudno rozpuszczalną, stężenie równowagowe jego rozpuszczonych jonów w roztworze jest stosunkowo niskie.

Stała równowagi dla równowag rozpuszczalności, takich jak ta, nazywana jest w tym przypadku stałą produktu rozpuszczalności, Ksp.

<img alt="Eq2" src="/files/ftp_upload/11417/11417_Equation_2.png" style="text-align: center; height: 50px;" />

Przypomnijmy, że w wyrażeniach na stałe równowagi reprezentowane są tylko gazy i substancje rozpuszczone, zatem Ksp nie obejmuje określenia nierozpuszczonego AgCl.

Ksp i rozpuszczalność

Ksp słabo rozpuszczalnego związku jonowego można po prostu powiązać z jego zmierzoną rozpuszczalnością, pod warunkiem, że proces rozpuszczania obejmuje jedynie dysocjację i solwatację, na przykład:

<img alt="Eq3" src="/files/ftp_upload/11417/11417_Equation_3.png" style="text-align: center; height: 50px;" />

W takich przypadkach można wyprowadzić wartości Ksp z podanych rozpuszczalności i odwrotnie. Obliczenia tego rodzaju najwygodniej przeprowadza się, stosując rozpuszczalność molową związku, mierzoną jako mol rozpuszczonej substancji rozpuszczonej na litr nasyconego roztworu.

Przewidywanie opadów

Równanie opisujące równowagę pomiędzy stałym węglanem wapnia i jego solwatowanymi jonami to:

<img alt="Eq4" src="/files/ftp_upload/11417/11417_Equation_4.png" style="text-align: center; height: 50px;" />

Należy pamiętać, że równowaga ta zachodzi w każdym roztworze wodnym zawierającym jony Ca2+ i CO32–, a nie tylko w roztworze powstałym w wyniku nasycenia wody węglanem wapnia. Rozważmy na przykład mieszanie wodnych roztworów rozpuszczalnych związków węglanu sodu i azotanu wapnia. Jeżeli stężenia jonów wapnia i węglanowych w mieszaninie nie dają współczynnika reakcji Q przekraczającego iloczyn rozpuszczalności Ksp, wówczas nie nastąpi wytrącanie. Jeżeli stężenia jonów dają iloraz reakcji większy niż produkt rozpuszczalności, wówczas nastąpi wytrącanie, obniżając te stężenia aż do ustalenia równowagi (Q = Ksp). Porównanie Q z Ksp w celu przewidywania opadów jest przykładem ogólnego podejścia do przewidywania kierunku reakcji, wprowadzonego po raz pierwszy na lekcjach równowagi. Dla konkretnego przypadku równowagi rozpuszczalności:

Q < Ksp: reakcja przebiega w kierunku do przodu (roztwór nie jest nasycony, nie obserwuje się wytrącania)

Q > Ksp: reakcja przebiega w odwrotnym kierunku (roztwór jest przesycony, nastąpi wytrącenie)

Równowagi rozpuszczalności

Równowaga rozpuszczalności ustala się, gdy rozpuszczanie i wytrącanie się substancji rozpuszczonej zachodzi w równym tempie. Te równowagi leżą u podstaw wielu ...

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Acid-base and Solubility Equilibria

Równowagi kwasowo-zasadowe i rozpuszczalności

KEY TERMS AND CONCEPTS

Kluczowe terminy i pojęcia

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+...

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titration-calculations-strong-acid-strong-base

pH at Key Points in Acid-Base Titration

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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...

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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

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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

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In 1923, G. N. Lewis proposed a generalized definition of acid-base behavior in which acids and bases are identified by their ability to accept or to ...

Common Ion Effect

Compared with pure water, the solubility of an ionic compound is less in aqueous solutions containing a common ion (one also produced by dissolution of ...

Chemical reactions
• Types of reactions
• Kinetics
• Energy
• Equilibrium
• Acids/bases

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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