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Questions and Answers
The value of $K_{sp}$ is ______ at constant temperature, regardless of the concentration of reactants.
The value of $K_{sp}$ is ______ at constant temperature, regardless of the concentration of reactants.
constant
When mixing solutions, precipitation occurs if the ionic product is ______ than the $K_{sp}$.
When mixing solutions, precipitation occurs if the ionic product is ______ than the $K_{sp}$.
greater
In precipitation problems, after mixing solutions, it is essential to calculate the ______ ion concentrations due to the change in volume.
In precipitation problems, after mixing solutions, it is essential to calculate the ______ ion concentrations due to the change in volume.
new
When solving precipitation problems, the ______ soluble salt is usually the one whose $K_{sp}$ is mentioned.
When solving precipitation problems, the ______ soluble salt is usually the one whose $K_{sp}$ is mentioned.
For the dissolution of $M(OH)3(s)$ into $M^{3+}(aq)$ and $3OH^-(aq)$,the ionic product expression is given by $[M^{3+}]_____^3$.
For the dissolution of $M(OH)3(s)$ into $M^{3+}(aq)$ and $3OH^-(aq)$,the ionic product expression is given by $[M^{3+}]_____^3$.
The reaction between halides and $AgNO_3$ solution forms the basis of qualitative analysis of halides such as $Cl^-$, $Br^-$, and ______.
The reaction between halides and $AgNO_3$ solution forms the basis of qualitative analysis of halides such as $Cl^-$, $Br^-$, and ______.
When aqueous ammonia is added to $AgCl(s)$, a soluble ______ $Ag(NH_3)_2]^+(aq)$ is formed, decreasing the concentration of $Ag^+(aq)$.
When aqueous ammonia is added to $AgCl(s)$, a soluble ______ $Ag(NH_3)_2]^+(aq)$ is formed, decreasing the concentration of $Ag^+(aq)$.
According to Le Chatelier's Principle, the equilibrium position shifts to increase the concentration of $Ag^+$ when a soluble complex $[Ag(NH_3)_2]^+(aq)$ is formed when dissolving $AgCl(s)$ in ______.
According to Le Chatelier's Principle, the equilibrium position shifts to increase the concentration of $Ag^+$ when a soluble complex $[Ag(NH_3)_2]^+(aq)$ is formed when dissolving $AgCl(s)$ in ______.
$AgBr$ and $AgI$ remain insoluble in dilute $NH_3(aq)$ because their ionic product of $[Ag^+][X^-]$ easily exceeds their ______ values.
$AgBr$ and $AgI$ remain insoluble in dilute $NH_3(aq)$ because their ionic product of $[Ag^+][X^-]$ easily exceeds their ______ values.
When concentrated ammonia is added to $AgBr(s)$, a soluble complex cation $[Ag(NH_3)_2]^+(aq)$ is formed, decreasing the concentration of ______.
When concentrated ammonia is added to $AgBr(s)$, a soluble complex cation $[Ag(NH_3)_2]^+(aq)$ is formed, decreasing the concentration of ______.
The addition of concentrated hydrochloric acid to solid lead(II) chloride, $PbCl_2$, causes it to dissolve because forming $Pb^{2+}$ and $Cl^-$ ions is ______.
The addition of concentrated hydrochloric acid to solid lead(II) chloride, $PbCl_2$, causes it to dissolve because forming $Pb^{2+}$ and $Cl^-$ ions is ______.
The solubility of $AgCl$ in aqueous $NH_3$ will be ______ than that in pure water due to the formation of the complex ion $[Ag(NH_3)_2]^+(aq)$.
The solubility of $AgCl$ in aqueous $NH_3$ will be ______ than that in pure water due to the formation of the complex ion $[Ag(NH_3)_2]^+(aq)$.
When concentrated ammonia is added to $Ag+(aq)$ and $Br^–(aq)$, the equilibrium position shifts as $[Ag+(aq)]$ decreases to a value below its low ______ value, hence $AgBr$ dissolves.
When concentrated ammonia is added to $Ag+(aq)$ and $Br^–(aq)$, the equilibrium position shifts as $[Ag+(aq)]$ decreases to a value below its low ______ value, hence $AgBr$ dissolves.
The ionic product of $M(OH)_3$ is calculated by the expression ______.
The ionic product of $M(OH)_3$ is calculated by the expression ______.
When mixing two solutions in equal volumes, the concentration of each ion is ______ due to the increased total volume.
When mixing two solutions in equal volumes, the concentration of each ion is ______ due to the increased total volume.
If the ionic product of a compound is less than its $K_{sp}$, ______ precipitation will occur.
If the ionic product of a compound is less than its $K_{sp}$, ______ precipitation will occur.
The $K_{sp}$ of calcium sulfate is given as $2.0 \times 10^{-5} \text{ mol}^2 \text{ dm}^{-6}$. If the ionic product is greater than this value, a precipitate of calcium sulfate will be ______.
The $K_{sp}$ of calcium sulfate is given as $2.0 \times 10^{-5} \text{ mol}^2 \text{ dm}^{-6}$. If the ionic product is greater than this value, a precipitate of calcium sulfate will be ______.
To determine if a precipitate will form, one must compare the ionic product ($Q$) with the ______.
To determine if a precipitate will form, one must compare the ionic product ($Q$) with the ______.
In the calculation of the new concentration of $M^{3+}$ after mixing solutions, if the stoichiometry of $M_2(SO_4)_3$ to $M^{3+}$ is 1:2, the initial concentration of $M_2(SO_4)_3$ must be multiplied by ______.
In the calculation of the new concentration of $M^{3+}$ after mixing solutions, if the stoichiometry of $M_2(SO_4)_3$ to $M^{3+}$ is 1:2, the initial concentration of $M_2(SO_4)_3$ must be multiplied by ______.
To calculate the solubility of $Mn(OH)2$, given its $K{sp}$ value, one sets up an equilibrium expression where the concentrations of $Mn^{2+}$ and $OH^-$ are related by the stoichiometry of the dissolution, leading to the equation $K_{sp} = s \times (2s)^2$, where s represents the ______ solubility.
To calculate the solubility of $Mn(OH)2$, given its $K{sp}$ value, one sets up an equilibrium expression where the concentrations of $Mn^{2+}$ and $OH^-$ are related by the stoichiometry of the dissolution, leading to the equation $K_{sp} = s \times (2s)^2$, where s represents the ______ solubility.
To calculate the minimum pH required to precipitate $Mn(OH)2$, one first determines the $[OH^-]$ concentration needed to reach the $K{sp}$ when $[Mn^{2+}]$ is at its minimum allowed concentration, and then uses the relationship $pOH + pH = ______$ to find the corresponding pH.
To calculate the minimum pH required to precipitate $Mn(OH)2$, one first determines the $[OH^-]$ concentration needed to reach the $K{sp}$ when $[Mn^{2+}]$ is at its minimum allowed concentration, and then uses the relationship $pOH + pH = ______$ to find the corresponding pH.
Flashcards
Ksp (Solubility Product)
Ksp (Solubility Product)
Equilibrium constant for solid dissolving into ions. Constant at a given temperature; reflects solubility.
Ionic Product (IP)
Ionic Product (IP)
A calculation using current, non-equilibrium concentrations, predicting if precipitation will occur.
Precipitation Condition
Precipitation Condition
Precipitation occurs when the Ionic Product (IP) exceeds the Solubility Product (Ksp).
Solving Precipitation Problems
Solving Precipitation Problems
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Concentration Change on Mixing
Concentration Change on Mixing
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AgCl Solubility in NH3
AgCl Solubility in NH3
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Why AgBr/AgI insoluble in dilute NH3?
Why AgBr/AgI insoluble in dilute NH3?
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AgBr Solubility in conc. NH3
AgBr Solubility in conc. NH3
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PbCl2 Dissolving in conc. HCl
PbCl2 Dissolving in conc. HCl
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Complex Ion Formation
Complex Ion Formation
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Solubility Product (Ksp)
Solubility Product (Ksp)
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Le Chatelier's Principle
Le Chatelier's Principle
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Ionic Product
Ionic Product
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IP < Ksp
IP < Ksp
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Mixing Solutions Effect
Mixing Solutions Effect
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CaSO4 Dissolution
CaSO4 Dissolution
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Mn(OH)2 Solubility
Mn(OH)2 Solubility
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Precipitation
Precipitation
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g dm−3 to mol dm−3
g dm−3 to mol dm−3
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Study Notes
- Solubility Equilibria is Part II of the Chemistry of Aqueous Solution
- 2024 JC1 H2 CHEMISTRY 9729
- Candidates should understand/apply the concept of solubility product, Ksp through calculation and discussion
Solubility Equilibria Topics
- Solubility of a sparingly soluble salt
- Solubility Product (Ksp)
- Common ion effect
- Calculation of Solubility of Salt in the Presence of a Common Ion
- Effect of complex ion formation on the solubility of salt
- Precipitation of Sparingly Soluble Salts
Solubility of a Sparingly Soluble Salt
- Sparingly soluble salts dissociate partially into their constituent ions
- Solubility of a salt refers to the concentration, specifically the maximum amount of solute (in g or mol) that can dissolve in 1 dm³ of solvent
- This creates a saturated solution at a provided temperature
- A2B3 (s) dissociates into 2A3+ (aq) + 3B2- (aq)
Solubility Product (Ksp)
- Equilibrium constant, representing the product of molar concentrations of dissolved dissociated ions, each raised to its appropriate power, in a saturated solution at a given temperature
- For a sparingly soluble binary salt, MX (s) dissociates into M+ (aq) + X- (aq) and the Solubility product: Ksp = [M+][X-]
- Ksp remains constant at a constant temperature, reflecting its nature as an equilibrium constant
- For Ag2CrO4 (s), the Ksp expression is: Ksp = [Ag+]2 [CrO42-] with units of mol³ dm–9
Ksp Values of Sparingly Soluble Salts
- AgCl: [[Ag+][Cl-] Ksp expression; s x s = s² in terms of solubility; Ksp value is 1.77 x 10-10 mol2 dm-6
- PbI2: [Pb2+][I-]2 is the Ksp expression; s x (2s)² = 4s3 in terms of solubility; Ksp value is 9.80 x 10-9 mol³ dm-9
- Ga(OH)3: [Ga³+][OH-]³ is the Ksp expression; s x (3s)³ = 27s4 in terms of solubility; Ksp value is 7.28 x 10-36 mol4 dm-12
- Ca3(PO4)2: [Ca2+]3[PO43-]2 is the Ksp expression; (3s)³ x (2s)² = 108s5 in terms of solubility; Ksp value is 2.07 x 10-33 mol5 dm-15
- "s" denotes the solubility, in mol dm¯³, of the sparingly soluble salt at 25°C
Common Ion Effect
- Le Chatelier's Principle dictates that the equilibrium position shifts to reduce the concentration of the common ion
- Common Ion Effect: The solubility of a sparingly soluble salt decreases in the presence of a common ion from an external source
- Equilibrium shifts left to reduce the concentration of the common ion
Effect of Complex Ion Formation on Solubility
- The presence of a reagent that reacts with the cation or anion to form complex ions increases the solubility of a sparingly soluble salt
- Equilibrium shifts to increase the concentration of the cation or anion to help dissolve more of the sparingly soluble salt
Comparing Solubility of Silver Halides (AgX) in Aqueous NH3 and Concentrated NH3
- Ksp of AgX/mol² dm-6: AgCl is 1.77 x 10-10; AgBr is 5.35 x 10-13; AgI is 8.52 x 10-17
- X-(aq) with AgNO3(aq) followed by NH3(aq): AgCl is a white ppt, soluble in excess NH3(aq); AgBr is a cream ppt, partially soluble in NH3(aq), soluble in conc. NH3; AgI is a yellow ppt, insoluble in NH3(aq)
- When aqueous ammonia is added, forms a complex [Ag(NH3)2]+(aq), decreasing the [Ag+(aq)]
Why AgBr and AgI Insoluble in Excess Dilute / Aqueous Ammonia?
- Ksp of AgBr and AgI are lower than AgCl.
- AgBr and AgI remain insoluble in dilute NH3(aq) as their ionic product of [Ag+][X¯] exceeds their Ksp values
Why is AgBr Soluble in Concentrated Ammonia?
- When concentrated ammonia is added, a soluble complex cation [Ag(NH3)2]+(aq) is formed, decreasing [Ag+(aq)]
- Equilibrium position shifts left to increase [Ag+]
- [Ag+(aq)] still decreases in eqm (1) and ionic product of [Ag+][Br¯] decreases to a value below its low Ksp value
Precipitation of Sparingly Soluble Salts
- Precipitation depends on the ionic product of dissolved ionic species compared to the Ksp of MX
- Ionic product = Ksp: Saturation point, where no precipitation happens
- Ionic product < Ksp: Lower than saturation point, where no precipitation occurs
- Ionic product > Ksp: Beyond saturation point, which means precipitation does occur
- Ionic product expression is similar to Ksp; the difference to predict precipitation occurring
- Ksp is an equilibrium constant as a product of the molar concentrations of dissolved dissociated ions in a saturated solution at a given temperature, regardless of reactant concentrations
- Ionic Product is a product of the molar concentrations of the dissolved dissociated ions in the solution at a given temperature and depends on the concentration of reactants or situation
- Expressions: Ksp = [M+]eqm[X-]eqm and Ionic product = [M+][X-]
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Description
This lesson covers key principles related to the solubility product constant ($K_{sp}$) and precipitation reactions. It explores how $K_{sp}$ remains constant at a given temperature and how to predict precipitation based on the ionic product. The lesson also discusses the importance of considering volume changes when mixing solutions and introduces the application of these principles in qualitative analysis.
