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

염화나트륨은 다량의 염화물이 물에 용해되기 때문에 용해 가능한 것으로 간주되지만 물에 염화 납을 첨가하면 소량만 용해되고 나머지는 용해되지 않습니다. 용해되지 않은 고체는 용액 속에 있는 납 및 염화물 이온과 공존합니다. 고체상태의 염화납의 일부는 계속 용해되는 반면 용액 속의 일부 이온은 재결합하여 침전물을 형성합니다.

해리 속도가 침전 속도와 같을 경우 용해의 평형이 이루어집니다. 평형 상수는 해리 반응에 따른 이온의 평형 농도로부터 계산할 수 있습니다. 여기서 염화납은 1개의 납과 2개의 염화 이온으로 해리됩니다.

따라서 평형 상수는 납 이온의 몰 농도에 염화 이온의 몰 농도의 제곱을 곱하여 계산합니다. 고체 상태의 염화납의 농도는 일정하기 때문에 계산에서 제외됩니다. 이 평형 상수를 용해도적이라고 하며 Ksp로 나타냅니다.

25°C에서 염화 납의 Ksp는 1.17 10⁻⁵입니다. Ksp 값은 화합물이 해리되어 포화 수용액을 형성할 수 있는 범위를 나타냅니다. 주어진 온도에서 화합물의 Ksp는 일정합니다.

몰 용해도로 알려진 리터당 몰로 구성된 화합물의 용해도는 종종 포화 용액에 용해된 고체의 농도를 표현하기 위해 사용됩니다. 화합물의 용해도는 용액의 pH와 같은 요인 및 다른 이온의 존재 여부에 따라 달라질 수 있습니다. 화합물의 몰 용해도 x는 ICE 표를 사용하여 Ksp로부터 계산할 수 있습니다.

용액 내 납 이온과 염화 이온의 초기 농도는 0입니다. 평형 상태에서 납 이온의 몰 농도가 x로 표시된다면 염화 이온의 몰 농도는 2x입니다. 평형식으로 대체하면 염화연의 용해도적은 x 2x²이며 값은 4x³입니다.

염화납의 Ksp는 1.17 10⁻⁵이므로 x는 1.43 10⁻² 몰이 됩니다. 염화납 및 플루오르화 칼슘과 같이 동일한 해리 화학양론을 가지며 각 화합물의 1 몰이 3몰의 해리 이온을 생성하는 화합물의 경우 각 Ksp 값을 사용하여 상대 용해도를 직접 비교할 수 있습니다.

<ol>
	<li>Clifford, A. F. "용해도 생성물 상수의 예측1." 미국 화학 학회 저널 79, no. 20 (1957) : 5404-5407.</li>
	<li>Hill, Arthur E. "용해도 생성물의 불변성." 미국 화학 학회 저널 32, no. 10 (1910) : 1186-1193.</li>
	<li>Reynolds, John P. "Ksp 실험 : 수산화 바륨의 용해도 생성물." 화학 교육 저널 52, no. 8 (1975) : 521.</li>
</ol>

용해도 평형

용해도 평형은 용질 종의 용해와 침전이 동일한 속도로 발생할 때 확립됩니다. 이러한 평형은 충치부터 수질 정화에 이르기까지 많은 자연적, 기술적 과정의 기초가 됩니다. 따라서 화합물 용해도에 영향을 미치는 요인에 대한 이해는 이러한 공정을 효과적으로 관리하는 데 필수적입니다. 이 섹션에서는 이전에 소개된 평형 개념과 도구를 용해 및 침전과 관련된 시스템에 적용합니다.

용해도 제품

물질의 용해도는 본질적으로 0(불용성 또는 거의 용해되지 않음)부터 무한대(혼화성)까지 다양할 수 있다는 점을 기억하십시오. 유한한 용해도를 갖는 용질은 용해도를 초과하는 양으로 용매에 첨가되면 포화 용액을 생성할 수 있으며, 그 결과 포화 용액과 과량의 용해되지 않은 용질의 이질적인 혼합물이 생성됩니다. 예를 들어, 염화은 포화 용액은 아래 그림과 같은 평형이 이루어진 용액입니다.

<img alt="Eq1" src="/files/ftp_upload/11417/11417_Equation_1.png" style="text-align: center; height: 50px;" />

이 용액에서는 과량의 고체 AgCl이 용해 및 해리되어 수용성 Ag^+ 및 Cl^- 이온을 생성합니다. 이러한 수성 이온이 결합하고 침전되어 고체 AgCl을 형성하는 것과 동일한 속도입니다. 염화은은 난용성 염이기 때문에 용액에 용해된 이온의 평형 농도는 상대적으로 낮습니다.

이와 같은 용해도 평형에 대한 평형 상수를 용해도 곱 상수, K_sp라고 합니다. 이 경우

<img alt="Eq2" src="/files/ftp_upload/11417/11417_Equation_2.png" style="text-align: center; height: 50px;" />

평형 상수 표현에는 기체와 용질만 표시되므로 K_sp에는 용해되지 않은 AgCl에 대한 항이 포함되지 않습니다.

K_sp 및 용해도

용해 과정이 해리 및 용매화만 포함하는 경우 용해도가 낮은 이온 화합물의 K_sp는 측정된 용해도와 간단하게 관련될 수 있습니다. 예를 들면 다음과 같습니다.

<img alt="Eq3" src="/files/ftp_upload/11417/11417_Equation_3.png" style="text-align: center; height: 50px;" />

이와 같은 경우 제공된 용해도로부터 K_sp 값을 도출할 수 있으며 그 반대의 경우도 마찬가지입니다. 이러한 종류의 계산은 포화 용액 1리터당 용해된 용질의 몰수로 측정되는 화합물의 몰 용해도를 사용하여 가장 편리하게 수행됩니다.

침전 예측

고체 탄산칼슘과 그 용매화 이온 사이의 평형을 설명하는 방정식은 다음과 같습니다.

<img alt="Eq4" src="/files/ftp_upload/11417/11417_Equation_4.png" style="text-align: center; height: 50px;" />

이 평형은 물을 탄산칼슘으로 포화시켜 형성된 용액뿐만 아니라 Ca^2+ 및 CO_3^2- 이온을 함유한 모든 수용액에서 확립된다는 점을 인식하는 것이 중요합니다. 예를 들어, 가용성 화합물인 탄산나트륨과 질산칼슘의 수용액을 혼합하는 것을 고려해 보십시오. 혼합물의 칼슘 및 탄산 이온 농도가 용해도 곱 K_sp를 초과하는 반응 지수 Q를 생성하지 않으면 침전이 발생하지 않습니다. 이온 농도가 용해도 곱보다 큰 반응 지수를 생성하는 경우 침전이 발생하여 평형이 확립될 때까지 해당 농도가 낮아집니다(Q = K_sp). 침전을 예측하기 위해 Q와 K_sp를 비교하는 것은 평형 수업에서 처음 소개된 반응 방향을 예측하는 일반적인 접근법의 예입니다. 용해도 평형의 특정 경우:

Q < K_sp: 반응이 정방향으로 진행됩니다(용액이 포화되지 않고 침전이 관찰되지 않음).

Q > K_sp: 반응이 반대 방향으로 진행됩니다(용액이 과포화되어 침전이 발생함).

이 문서는 에서 발췌되었습니다 <a href="https://openstax.org/books/chemistry-2e/pages/15-1-precipitation-and-dissolution">Openstax, Chemistry 2e, Section 15.1: Precipitation and Dissolution.</a>

용해도 평형은 용질 종의 용해와 침전이 동일한 비율로 발생할 때 성립됩니다. 이러한 평형은 많은 기초가 됩니다 ...

Education

교육

JoVE Core

화학

하나

Acid-base and Solubility Equilibria

산-염기 및 용해도 평형

KEY TERMS AND CONCEPTS

주요 용어 및 개념

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

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Predicting Precipitation with Ksp

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