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Solubility
The definition of solubility is the maximum quantity of solute that can dissolve in a certain quantity of solvent or quantity of solution at a specified temperature or pressure (in the case of gaseous solutes). In CHM1045 we discussed solubility as a yes or no quality. But the reality is that almost every solute is somewhat soluble in every solvent to some measurable degree.
As stated in the definition, temperature and pressure play an important role in determining the degree to which a solute is soluble.
Let's start with temperature:
For Gases, solubility decreases as temperature increases (duh...you have seen water boil, right?) The physical reason for this is that when most gases dissolve in solution, the process is exothermic. This means that heat is released as the gas dissolves. This is very similar to the reason that vapor pressure increases with temperature. Increased temperature causes an increase in kinetic energy. The higher kinetic energy causes more motion in the gas molecules which break intermolecular bonds and escape from solution. Check out the graph below:
As the temperature increases, the solubility of a gas decreases as shown by the downward trend in the graph.
For solid or liquid solutes:
CASE I: Decrease in solubility with temperature:
If the heat given off in the dissolving process is greater than the heat required to break apart the solid, the net dissolving reaction is exothermic (See the solution process). The addition of more heat (increases temperature) inhibits the dissolving reaction since excess heat is already being produced by the reaction. This situation is not very common where an increase in temperature produces a decrease in solubility. But is the case for sodium sulfate and calcium hydroxide.
CASE II: Increase in solubility with temperature:
If the heat given off in the dissolving reaction is less than the heat required to break apart the solid, the net dissolving reaction is endothermic. The addition of more heat facilitates the dissolving reaction by providing energy to break bonds in the solid. This is the most common situation where an increase in temperature produces an increase in solubility for solids.
The use of first-aid instant cold packs is an application of this solubility principle. A salt such as ammonium nitrate is dissolved in water after a sharp blow breaks the containers for each. The dissolving reaction is endothermic - requires heat. Therefore the heat is drawn from the surroundings, the pack feels cold.
The effect of temperature on solubility can be explained on the basis of Le Chatelier's Principle. Le Chatelier's Principle states that if a stress (for example, heat, pressure, concentration of one reactant) is applied to an equilibrium, the system will adjust, if possible, to minimize the effect of the stress. This principle is of value in predicting how much a system will respond to a change in external conditions. Consider the case where the solubility process is endothermic (heat added). An increase in temperature puts a stress on the equilibrium condition and causes it to shift to the right. The stress is relieved because the dissolving process consumes some of the heat. Therefore, the solubility (concentration) increases with an increase in temperature. If the process is exothermic (heat given off). A temperature rise will decrease the solubility by shifting the equilibrium to the left.
Now let's look at pressure:
Solids and liquids show almost no change in solubility with changes in pressure. But gases are very dependent on the pressure of the system. Gases dissolve in liquids to form solutions. This dissolution is an equilibrium process for which an equilibrium constant can be written. For example, the equilibrium between oxygen gas and dissolved oxygen in water is O2(aq) <=> O2(g). The equilibrium constant for this equilibrium is K = p(O2)/c(O2). The form of the equilibrium constant shows that the concentration of a solute gas in a solution is directly proportional to the partial pressure of that gas above the solution. This statement, known as Henry's law, was first proposed in 1800 by J.W. Henry as an empirical law well before the development of our modern ideas of chemical equilibrium.
Henry's Law:
Sg stands for the gas solubility, kH is the Henry's Law constant and Pg is the partial pressure of the gaseous solute.
Table: Molar Henry's Law Constants for Aqueous Solutions at 25oC
Gas | Constant | Constant |
(Pa/(mol/dm3)) | (atm/(mol/dm3)) | |
He | 282.7e6 | 2865 |
O2 | 74.68e6 | 756.7 |
N2 | 155.0e6 | 1600 |
H2 | 121.2e6 | 1228 |
CO2 | 2.937e6 | 29.76 |
NH3 | 5.69e6 | 56.9 |
The inverse of the Henry's law constant, multiplied by the partial pressure of the gas above the solution, is the molar solubility of the gas. Thus oxygen at one atmosphere would have a molar solubility of (1/756.7)mol/dm3 or 1.32 mmol/dm3. Values in this table are calculated from tables of molar thermodynamic properties of pure substances and aqueous solutes
Summary of Factors Affecting Solubility
Normally, solutes become more soluble in a given solvent at higher temperatures. One way to predict that trend is to use Le Chatelier's principle. Because DHsoln is positive for most solutions, the solution formation reaction is usually endothermic. Therefore, when the temperature is increased, the solubility of the solute should also increase. However, there are solutes that do not follow the normal trend of increasing solubility with increasing temperature. One class of solutes that becomes less soluble with increasing temperature is the gasses. Nearly every gas becomes less soluble with increasing temperature.
Another property of gaseous solutes in summarized by Henry's law which predicts that gasses become more soluble when their pressures above a liquid solution are increased. That property of gaseous solutes can be rationalized by using Le Chatelier's principle. Imagine that you have a glass of water inside of a sealed container filler with nitrogen gas. If the size of that container were suddenly halved, the pressure of nitrogen would suddenly double. To decrease the pressure of nitrogen above the solution (as is required by Le Chatelier's principle), more nitrogen gas becomes dissolved in the glass of water.
Dissolution [1][2][3]
Dissolution is the process where a solute in gaseous, liquid, or solid phase dissolves in a solvent to form a solution.
Solubility
Solubility is the maximum concentration of a solute that can dissolve in a solvent at a given temperature. At the maximum concentration of solute, the solution is said to be saturated. The units of solubility can be provided in mol/L or g/L.
Factors that affect solubility include:
The concentration of the solute
The temperature of the system
Pressure (for gases in solution)
The polarity of the solute and the solvent
Dissolution
The rate of dissolution is represented by the Noyes-Whitney equation: dm/dt = D*A*(Cs - C)/h
Where:
dm/dt represents the rate of dissolution
D represents the diffusion coefficient for the compound
A represents the surface area available for dissolution
Cs represents the solubility of the compound
C represents the solute concentration in bulk solution at time t
h represents the thickness of the dissolution layer
Solubility
Temperature
Effect of temperature on liquid and solid solutes
As temperature increases, the solubility of a solid or liquid can fluctuate depending on whether the dissolution reaction is exothermic or endothermic.
Increasing solubility with increasing temperature
In endothermic dissolution reactions, the net energy from breaking and forming bonds results in heat energy being absorbed into the system as the solute dissolves. When the temperature of the system increases, additional head energy is introduced into the system.
So according to Le Chatelier’s Principle, the system will adjust to this increase in the heat by promoting the dissolution reaction to absorb the added heat energy. Increasing the temperature will therefore increase the solubility of the solute.
An example of a solute whose solubility increases with greater temperature is ammonium nitrate, which can be used in first-aid cold packs. Ammonium nitrate dissolving in solution is an endothermic reaction. As the ammonium nitrate dissolves, heat energy is absorbed from the environment causing the surrounding environment to feel cold.
Decreasing solubility with increasing temperature
In exothermic reactions, heat energy is released when the solute dissolves in a solution. Increasing temperature introduces more heat into the system. Following Le Chatelier’s Principle, the system will adjust to this excess heat energy by inhibiting the dissolution reaction. Increasing temperature, therefore, decreases the solubility of the solute.
An example of a solute that decreases in solubility with increasing temperature is calcium hydroxide, which can be used to treat chemical burns and as an antacid.
Effect of temperature on gas solutes
In general, heat energy is released as gas dissolves in solution, meaning the dissolution reaction is exothermic. As such, a gas becomes less soluble as temperate increases.
Increasing temperature results in increased kinetic energy. Gas molecules with greater kinetic energy move more rapidly resulting in the intermolecular bonds between the gas solute and solvent breaking.
Pressure: Henry’s law
The solubility of gas is affected by changes in pressure on the system. A gas dissolves in liquids to form solutions. This results in equilibrium in the system where a proportion of gas molecules is dissolved in liquid while the rest remains in gaseous phase above the liquid.
Henry’s law states that: “At constant temperature, the amount of gas that dissolves in a volume of liquid is proportional to the partial pressure of the gas in equilibrium with the liquid.”
Henry's law results in the following equation: C = kP
Where:
C represents the solubility of the gas at a certain temperature in a specific solvent
K represents Henry’s law constant
P represents the partial pressure of the gas i.e. the pressure the gas exerts on the system at a given volume and temperature
Hence as the pressure of the gas above the liquid in the system increases, the gas molecules become more soluble in the solvent. Likewise, if the pressure of the gas in the system decreases, gas becomes less soluble in the solvent.
Limitations of Henry’s Law on gas solubility:
Only applies if the gas molecules are in equilibrium
Does not apply if there is a chemical reaction between the solvent and the solute
Does not apply to gas at high pressures
Dissolution
Methods to enhance dissolution to improve oral bioavailability include: [4][5][6]
Micronization to increase surface area for dissolution
Salt formation of the active ingredient
Use of co-solvents and micelle solutions to aid solubilization
Complexation through the use of cyclodextrins
Use of lipidic systems (for lipophilic drugs)
Solubility [7][8][9]
Le Chatelier’s principle:
If stressors like pressure and heat are applied to the equilibrium, the system will respond by adjusting to minimize the effects of the stress.
For example, if pressure is applied to a system, the dissolution reaction will respond to minimize this stress by reducing the pressure in the system.
Heat of solution
Solids and liquids form as a result of individual particles being held together by inter-particulate bonds. To form a solution, energy is required to break the bonds between the particles within the solid or liquid. Heat energy is also required to break the bonds in a solvent to insert one of the molecules into the solution. Both of these processes are endothermic. Heat energy is released when the solute molecules form bonds with the solvent molecules i.e. this process is exothermic.
Depending on whether more energy is used to break the bonds within the solute and solvent or is released when new bonds are formed between the solute and solvent, the reaction overall can be exothermic or endothermic.
If more energy is required to break the bonds within the solute and solvent than is released when new bonds are formed between the solute and solvent, the reaction is considered endothermic.
If more energy is released when new bonds are formed between the solute and solvent than is required to break the bonds within the solute and solvent, the reaction is considered exothermic.
The total amount of heat energy released from or absorbed by the system = sum of heat energy absorbed when bonds are broken – the sum of heat energy released when bonds are formed
If the total amount of heat energy released/absorbed from the system is greater than zero, the reaction is endothermic.
If the total amount of heat energy released/absorbed from the system is less than zero, the reaction is exothermic.
Application of Henry’s Law: Decompression Sickness
Henry’s Law explains the phenomena of decompression sickness. When scuba divers submerge themselves in deep water, the pressure of the water increases the pressure in their bodies. Nitrogen, a gas in our blood, dissolves under the increased pressure. Nitrogen is physiologically inert, so it is not used in tissue metabolism. If the scuba diver ascends to the surface too quickly, the rapid drop in pressure decreases the solubility of nitrogen, causing nitrogen bubbles to come out of solution. The nitrogen bubbles can form painful and potentially fatal gas embolisms.
Dissolution
Dissolution is important for health practitioners because, for drugs to be absorbed and have a physiological effect in the human body, they must be in solution. For solid preparations, such as tablets and suppositories, the rate of dissolution affects how fast a drug is absorbed in the body.
Solubility
Aqueous solubility is often considered when formulating drugs. Poorly soluble formulations provide difficulties in the development of pharmaceuticals. Chloramphenicol, phenytoin, and digoxin are some examples. Drugs, particularly those for oral administration, may have poor aqueous solubility. This may result in low bioavailability leading to insufficient exposure and physiologic effect in the body.
Review Questions
1.
Joshi K, Chandra A, Jain K, Talegaonkar S. Nanocrystalization: An Emerging Technology to Enhance the Bioavailability of Poorly Soluble Drugs. Pharm Nanotechnol. 2019;7(4):259-278. [PMC free article: PMC6967137] [PubMed: 30961518]
2.Itai S. [Development of Novel Functional Formulations Based on Pharmaceutical Technologies]. Yakugaku Zasshi. 2019;139(3):419-435. [PubMed: 30828022]
3.Karaźniewicz-Łada M, Bąba K, Dolatowski F, Dobrowolska A, Rakicka M. The polymorphism of statins and its effect on their physicochemical properties. Polim Med. 2018 Jul-Dec;48(2):77-82. [PubMed: 30916495]
4.Sujka M, Pankiewicz U, Kowalski R, Nowosad K, Noszczyk-Nowak A. Porous starch and its application in drug delivery systems. Polim Med. 2018 Jan-Jun;48(1):25-29. [PubMed: 30657655]
5.Modica de Mohac L, Keating AV, de Fátima Pina M, Raimi-Abraham BT. Engineering of Nanofibrous Amorphous and Crystalline Solid Dispersions for Oral Drug Delivery. Pharmaceutics. 2018 Dec 24;11(1) [PMC free article: PMC6359107] [PubMed: 30586871]
6.Couillaud BM, Espeau P, Mignet N, Corvis Y. State of the Art of Pharmaceutical Solid Forms: from Crystal Property Issues to Nanocrystals Formulation. ChemMedChem. 2019 Jan 08;14(1):8-23. [PubMed: 30457705]
7.Ribeiro ACF, Esteso MA. Transport Properties for Pharmaceutical Controlled-Release Systems: A Brief Review of the Importance of Their Study in Biological Systems. Biomolecules. 2018 Dec 17;8(4) [PMC free article: PMC6315691] [PubMed: 30563024]
8.Radivojev S, Zellnitz S, Paudel A, Fröhlich E. Searching for physiologically relevant in vitro dissolution techniques for orally inhaled drugs. Int J Pharm. 2019 Feb 10;556:45-56. [PubMed: 30529665]
9.Kadokawa JI. Dissolution, derivatization, and functionalization of chitin in ionic liquid. Int J Biol Macromol. 2019 Feb 15;123:732-737. [PubMed: 30465832]