Solubility Product

5 Key excerpts on "Solubility Product"

• 

  eBook - ePub

  Environmental Engineering

  Principles and Practice

  • Richard O. Mines(Author)
  • 2014(Publication Date)
  • Wiley-Blackwell

    (Publisher)

  2.7 Solubility (Solubility Product) So far, we have dealt with aqueous solutions in which the chemical species are highly soluble. In this section, our focus will be on liquid-solid species that are partially soluble or insoluble. All solids, no matter how seemingly insoluble, are soluble to some degree. When a solid is placed in water, the ions at the surface of the solid will migrate into the water. This is called dissolution. Simultaneously, ions in the solution will be redeposited on the surface of the solid; this is known as precipitation. Equilibrium will be reached between the crystals of the compound in the solid state and its ions in solution. In general, the solubility of most compounds increases with increasing temperature. Snoeyink & Jenkins (1980, page 251) indicate that the solubilities of and do not increase as temperature increases. Equation (2.112) shows the general equation of a solid compound dissolving in pure water to form its constituent ions. 2.112 The equilibrium expression is written as follows: 2.113 As described by Sawyer & McCarty (1994, page 37), at equilibrium or saturation, the surface area of the solid is the only portion that is in equilibrium with the ions in solution. Therefore, the concentration of solid as represented by in the denominator of Equation (2.113) can be considered a constant in equilibrium solubility problems. Equation (2.114) is rewritten to show the development of the solubility-product constant, 2.114 2.115 When the solution is saturated or at equilibrium. When the solution is under-saturated and no solids species are present. When the solution is super-saturated and solid species are being formed. The solubility-product constants for several solids of significance in environmental engineering are presented in Table 2.16. Partially soluble salts have small values, while soluble salts have relatively large values

  Sign up to read

  Learn more about book
• 

  eBook - ePub

  Rapid Review of Chemistry for the Life Sciences and Engineering

  With Select Applications

  • Armen S. Casparian, Gergely Sirokman, Ann Omollo(Authors)
  • 2021(Publication Date)
  • CRC Press

    (Publisher)

  ppm).

  4.12 Definition of Solubility Product Constant K

  sp

  Consider, for example, the slightly soluble salt CaF2 (s) in equilibrium with water. Its limited dissociation and ionization in water can be expressed as follows:

  CaF 2

  ( s )

  ⇌

  Ca

  2 +

  ( aq )

  + 2

  F −

  ( aq )

  (4.9)

  And its Solubility Product constant,K

  sp

  , as

  K

  s p

  =

  [

  Ca

  2 +

  ]

  [

  F −

  ]

  2

  (4.10)

  Note that [CaF2 (s)] does not appear in the Ksp expression since it represents the undissolved portion, and it is pure solid. It has no concentration.

  Also note that [Ca2+ ] represents the concentration of Ca2+ ion dissolved in water, while [F− ] represents the F− ion concentration dissolved in water.

  Note also that when a given amount of CaF2 (s) is dissolved in a known amount of water:

  [

  Ca

  2 +

  ( aq )

  ]

  = 2

  [

  F −

  ( aq )

  ]

  =

  [

  CaF 2

  ( aq )

  ]

  where CaF2 (aq) is the portion that is dissolved in water. In other words, whatever the proportion of CaF2 (s) that dissolves in water, the same molar concentration is present as Ca2+ (aq) ion, but twice that molar concentration is present as F− (aq) ion.

  The smaller the Ksp value, the lower the solubility of the salt. It follows that the Ksp value of a salt can be calculated by determining the solubility of each ion and then using Equation 4.10.

  4.13 Calculating the Molar Solubility from K

  sp

  The molar solubility of any slightly soluble salt, along with the concentration of any of its ions, can be calculated from its Ksp value. The formula of the salt must be known, however, so that its dissociation and ionization can be written.

  It is important to note that simply comparing the Ksp

  Sign up to read

  Learn more about book
• 

  No longer available |Learn more

  AP Chemistry Premium, 2024: 6 Practice Tests + Comprehensive Review + Online Practice

  • Neil D. Jespersen, Pamela Kerrigan(Authors)
  • 2023(Publication Date)
  • Barrons Educational Services

    (Publisher)

  M solution). Saturated solutions of insoluble salts have concentrations that are less than 0.1 molar. However, most insoluble salts do dissolve to a small extent.

  The solution process may be written in a form similar to a chemical reaction. For solid Fe2 (OH)3 we have

  Fe(OH)3 (s) ⇌ Fe3+ (aq) + 3OH− (aq)

  The equilibrium expression for this is written as

  K

  c

  = [Fe3+ ][OH− ]3

  In the special case of the solubility of slightly soluble compounds, the equilibrium expression always represents the product of the ions produced when the compound dissolves. The equilibrium constant is called the Solubility Product constant and is given the symbol Ksp :

  K

  sp

  = [Fe3+ ][OH− ]3

  Fe(OH)3 (s) does not appear in the equilibrium expression because it is a solid. However, in the equilibrium table, we should note in the column for the solid that initially we start with some solid. Then some small amount of solid, −x, will dissolve. At equilibrium, some solid must be left. This is shown in the next equilibrium table.

  Questions involving Ksp are solved in the same way as other equilibrium problems. For example, Ksp has a value of 1.6 × 10− 39 for Fe(OH)3 . The molar solubility is calculated by setting up an equilibrium table:

  Entering the information from the EQUILIBRIUM line into the Ksp equilibrium expression gives

  The value of x is used to calculate the molar concentrations of Fe3 + and OH− in the solution. We enter these data in the table:

  The solubility of Fe(OH)3 may be deduced from this table also. The CHANGE line indicates that −x of the compound dissolves, and therefore x represents the solubility of Fe(OH)3 , or 8.8 × 10− 11 mol L− 1 .

  If Fe(OH)3 is dissolved in a solution that already contains the Fe3 + ion, the common ion effect is observed. The common ion effect is a decrease in the solubility of a compound when it is dissolved in a solution that already contains an ion in common with the salt being dissolved. As an example, we can calculate the solubility of Fe(OH)3 in a solution that already has a 6.5 × 10− 5 M concentration of Fe3

  Sign up to read

  Learn more about book
• 

  eBook - ePub

  Dosage Form Design Considerations

  Volume I

  • (Author)
  • 2018(Publication Date)
  • Academic Press

    (Publisher)

  olubility may be defined as the concentration of the solute in concentrated solution at a certain temperature. When it comes to qualitative measurement, solubility is assumed to be a spontaneous interaction of two or more components to produce a homogenous molecular dispersion in a given solvent.

  15.1.1.1 Importance of Solubility and Solubilization in Product Development

  The efficient delivery of drugs is the issue of prime importance to the makers of pharmaceuticals. Approximately 40% of marketed drugs have low solubility and almost 80%–90% drug candidates in the R&D product development pipeline fail due to the solubility concerns. As discussed above, the therapeutic efficacy of a drug depends on the bioavailability, and ultimately upon the solubility of drug molecules. The solubility is one of the imperative parameters to accomplish the desired drug concentration in systemic circulation for achieve the required pharmacological.

  For the effective absorption of a drug, it must be present in the form of an aqueous solution at the site of absorption. Therefore, water is the solvent of choice for liquid pharmaceutical formulations. Most drugs are weakly acidic or weakly basic with poor aqueous solubility. For this reason, several techniques are used to improve the solubility of poorly water-soluble drugs.

  15.1.1.2 Process of Solubilization

  The method of dissolving solute involves the breakage of intermolecular or interionic bonds in the solute molecule, the split-up of the solvent component to provide space in the solvent for the solute, and interaction between the solvent and the solute molecule or ion. The change in enthalpy of the solution denotes the total quantity of heat released/absorbed during solubilization.

  15.1.1.3 Solvent–Solute Interactions

  During the product development, choice of the appropriate solvent depends on the principle of “like dissolves like,” which means solute preferentially dissolves in the solvent with alike physicochemical characteristics. In other words, two substances with similar intermolecular forces are soluble in each other. For example, polar solutes like common salt and sugar dissolve in polar solvents like water. Similarly, the nonpolar solutes dissolve in nonpolar solvents, for example, naphthalene solubilized in benzene.

  Sign up to read

  Learn more about book
• 

  eBook - ePub

  Essentials of Pharmaceutical Preformulation

  • Simon Gaisford, Mark Saunders(Authors)
  • 2012(Publication Date)
  • Wiley-Blackwell

    (Publisher)

  4.5 ) the reason for the term equilibrium solubility noted earlier.

  It appears from Equation (4.2 ) that the crystal lattice energy might affect solubility. It also seems from Equation (4.1 ) that there should be an effect of temperature on solubility, since the position of equilibrium will change. Both of these effects can be explored further through the concept of ideal solubility .

  Summary box 4.1

  • Solubility is the maximum concentration of a given solute that can be attained in a given solvent.
  • Solids transition to solution by dissolution.
  • Thermodynamic solubility is a position of equilibrium.
  • Dissolution governs the rate at which solubility is achieved.
  • As a general rule, solubility below 1 mg mL−1 is likely to hinder development while solubility above 10 mg mL−1 is acceptable.

  4.2.1 Ideal solubility

  In the special case where the enthalpy of any solute–solvent interaction is equal to the enthalpy of any solvent–solvent interaction then solvation of the solute may occur with no change in enthalpy (i.e. Δmix H = 0) and dissolution is said to be ideal . Formation of an ideal solution also occurs with the following change in entropy ( ):

  (4.6)

  where R is the universal gas constant (8.314 J K−1 mol−1 ). Ideal dissolution (although unlikely, because the solute and solvent molecules would need to possess identical properties, such as size, shape and chemical nature) leads to ideal solubility and is an interesting theoretical position because it can be described in thermodynamic terms, which allows calculation of the dependence of solubility on temperature.

  From Equation (4.2 ) if Δmix H = 0 then Δf H is equal to Δsol H (note that since Δf H must be positive, i.e. endothermic, Δsol H must also be positive for ideal dissolution). For a process to occur spontaneously the Gibbs free energy (ΔG ) must be negative. The familiar thermodynamic relationship for dissolution is

  (4.7)

  where T is absolute temperature. Δsol G is most likely to be negative when Δsol H is negative but, as noted above, Δsol H is frequently positive for dissolution and must be so when dissolution is ideal. This means that for dissolution to occur spontaneously the driving force can only be a significant increase in entropy. Since the mole fractions of both solvent and solute must be less than 1, the logarithmic terms in Equation (4.6

  Sign up to read

  Learn more about book