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Solution (chemistry)
Solution (chemistry)
Solution (chemistry)
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Making a saline water solution by dissolving table salt (NaCl) in water. The salt is the solute and the water the solvent.

In chemistry, a solution is defined by IUPAC as "A liquid or solid phase containing more than one substance, when for convenience one (or more) substance, which is called the solvent, is treated differently from the other substances, which are called solutes. When, as is often but not necessarily the case, the sum of the mole fractions of solutes is small compared with unity, the solution is called a dilute solution. A superscript attached to the ∞ symbol for a property of a solution denotes the property in the limit of infinite dilution."[1] One parameter of a solution is the concentration, which is a measure of the amount of solute in a given amount of solution or solvent. The term "aqueous solution" is used when one of the solvents is water.[2]

Types

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Homogeneous means that the components of the mixture form a single phase. Heterogeneous means that the components of the mixture are of different phase. The properties of the mixture (such as concentration, temperature, and density) can be uniformly distributed through the volume but only in absence of diffusion phenomena or after their completion. Usually, the substance present in the greatest amount is considered the solvent. Solvents can be gases, liquids, or solids. One or more components present in the solution other than the solvent are called solutes. The solution has the same physical state as the solvent.

Gaseous mixtures

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If the solvent is a gas, only gases (non-condensable) or vapors (condensable) are dissolved under a given set of conditions. An example of a gaseous solution is air (oxygen and other gases dissolved in nitrogen). Since interactions between gaseous molecules play almost no role, non-condensable gases form rather trivial solutions. In the literature, they are not even classified as solutions, but simply addressed as homogeneous mixtures of gases. The Brownian motion and the permanent molecular agitation of gas molecules guarantee the homogeneity of the gaseous systems. Non-condensable gaseous mixtures (e.g., air/CO2, or air/xenon) do not spontaneously demix, nor sediment, as distinctly stratified and separate gas layers as a function of their relative density. Diffusion forces efficiently counteract gravitation forces under normal conditions prevailing on Earth. The case of condensable vapors is different: once the saturation vapor pressure at a given temperature is reached, vapor excess condenses into the liquid state.

Liquid solutions

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Liquids dissolve gases, other liquids, and solids. An example of a dissolved gas is oxygen in water, which allows fish to breathe under water. An examples of a dissolved liquid is ethanol in water, as found in alcoholic beverages. An example of a dissolved solid is sugar water, which contains dissolved sucrose.

Solid solutions

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If the solvent is a solid, then gases, liquids, and solids can be dissolved.

Solubility

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Main articles: Solubility and Solvation

The ability of one compound to dissolve in another compound is called solubility.[clarification needed] When a liquid can completely dissolve in another liquid the two liquids are miscible. Two substances that can never mix to form a solution are said to be immiscible.

All solutions have a positive entropy of mixing. The interactions between different molecules or ions may be energetically favored or not. If interactions are unfavorable, then the free energy decreases with increasing solute concentration. At some point, the energy loss outweighs the entropy gain, and no more solute particles[clarification needed] can be dissolved; the solution is said to be saturated. However, the point at which a solution can become saturated can change significantly with different environmental factors, such as temperature, pressure, and contamination. For some solute-solvent combinations, a supersaturated solution can be prepared by raising the solubility (for example by increasing the temperature) to dissolve more solute and then lowering it (for example by cooling).

Usually, the greater the temperature of the solvent, the more of a given solid solute it can dissolve. However, most gases and some compounds exhibit solubilities that decrease with increased temperature. Such behavior is a result of an exothermic enthalpy of solution. Some surfactants exhibit this behaviour. The solubility of liquids in liquids is generally less temperature-sensitive than that of solids or gases.

Properties

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The physical properties of compounds such as melting point and boiling point change when other compounds are added. Together they are called colligative properties. There are several ways to quantify the amount of one compound dissolved in the other compounds collectively called concentration. Examples include molarity, volume fraction, and mole fraction.

The properties of ideal solutions can be calculated by the linear combination of the properties of its components. If both solute and solvent exist in equal quantities (such as in a 50% ethanol, 50% water solution), the concepts of "solute" and "solvent" become less relevant, but the substance that is more often used as a solvent is normally designated as the solvent (in this example, water).

Liquid solution characteristics

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See also: Solvent § Solvent classifications

In principle, all types of liquids can behave as solvents: liquid noble gases, molten metals, molten salts, molten covalent networks, and molecular liquids. In the practice of chemistry and biochemistry, most solvents are molecular liquids. They can be classified into polar and non-polar, according to whether their molecules possess a permanent electric dipole moment. Another distinction is whether their molecules can form hydrogen bonds (protic and aprotic solvents). Water, the most commonly used solvent, is both polar and sustains hydrogen bonds.

Water is a good solvent for some polar materials because water molecules are polar and capable of forming hydrogen bonds.

Salts dissolve in polar solvents, forming positive and negative ions that are attracted to the negative and positive ends of the solvent molecule, respectively. If the solvent is water, hydration occurs when the charged solute ions become surrounded by water molecules. A standard example is aqueous saltwater. Such solutions are called electrolytes. Whenever salt dissolves in water ion association has to be taken into account.

Polar solutes dissolve in polar solvents, forming polar bonds or hydrogen bonds. As an example, all alcoholic beverages are aqueous solutions of ethanol. On the other hand, non-polar solutes dissolve better in non-polar solvents. Examples are hydrocarbons such as oil and grease that easily mix, while being incompatible with water.

An example of the immiscibility of oil and water is a leak of petroleum from a damaged tanker, that does not dissolve in the ocean water but rather floats on the surface.

See also

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Look up solution or solute in Wiktionary, the free dictionary.

References

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Revisions and contributorsEdit on WikipediaRead on Wikipedia

Solution (chemistry)

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In chemistry, a solution is a homogeneous of two or more substances in which the molecules or atoms of the substances are completely dispersed and uniformly distributed throughout the mixture, such that no distinct particles of the original substances can be observed. The component present in the greater amount, which serves as the dissolving medium, is called the , while the substance that dissolves in it, typically in smaller quantity, is the solute. Solutions are fundamental to numerous chemical processes and everyday phenomena, including the dissolution of salts in or gases in liquids, and they exhibit properties distinct from those of pure substances. Solutions can be classified based on the physical states of the solute and solvent, resulting in various types such as gas in gas (e.g., air), liquid in liquid (e.g., ethanol in water), solid in liquid (e.g., sugar in water), and gas in liquid (e.g., carbon dioxide in soda). Another classification considers the amount of solute relative to solubility limits: an unsaturated solution holds less solute than the maximum possible at a given temperature; a saturated solution contains the maximum amount, in equilibrium with undissolved solute; and a supersaturated solution exceeds this limit, often achieved by careful temperature changes and unstable without seeding. Solutions may also be described as dilute (low solute concentration) or concentrated (high solute concentration), though these terms are relative and depend on context. The concentration of a solution quantifies the amount of solute per unit volume or mass of the solution and is expressed in units such as molarity (moles of solute per liter of solution), molality (moles of solute per kilogram of solvent), mass percent (grams of solute per 100 grams of solution), or mole fraction (ratio of moles of solute to total moles). Solubility, the maximum amount of solute that can dissolve, is influenced by factors including temperature (often increasing with heat for solids in liquids), pressure (significant for gases, as per Henry's law), and the chemical nature of solute-solvent interactions, encapsulated in the principle "like dissolves like." A key set of solution properties, known as colligative properties, depend solely on the number of solute particles rather than their identity and include vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure. These properties arise from the solute particles disrupting the solvent's behavior and are crucial in applications like antifreeze formulations, desalination, and biological systems such as cell membranes. Overall, solutions play a central role in chemistry, enabling reactions in homogeneous media and influencing fields from environmental science to pharmaceuticals.

Fundamentals

Definition

In chemistry, a solution is a homogeneous consisting of a solute dissolved in a , in which the solute particles are uniformly dispersed at the molecular or ionic level, resulting in a single phase with no visible boundaries between components. The solute is typically the substance present in smaller quantity that dissolves, while the is the medium, usually in greater amount, that facilitates the dispersion. The term "solution" derives from the Latin solutio, meaning a loosening or dissolving. The chemical meaning of the term 'solution,' referring to a containing a dissolved substance, was first recorded in the late ; it was used by in his discussions of dissolution processes in (1661). Its modern conceptualization was formalized in the late 19th century through the emergence of , with key contributions from , who advanced solution theory via studies on dilution and conductivity in electrolyte solutions. Unlike heterogeneous mixtures such as suspensions, where larger particles settle out over time and create visible separation, solutions form stable, uniform dispersions spontaneously upon mixing compatible substances or with minimal energy input, as the solute integrates fully into the without phase boundaries. Classic examples include saltwater, an ionic solution of (NaCl) dispersed in (H₂O), and ethanol in , a molecular solution where the alcohol molecules mix evenly with molecules.

Components

In a solution, the two primary components are the solute and the solvent. The solute is the substance present in the smaller amount that dissolves into the solvent to form the homogeneous mixture. Solutes can exist in various states of , including , liquids, or gases, depending on the type of solution being formed. The solvent, in contrast, is the medium present in the larger amount that dissolves the solute and determines the physical state of the resulting solution. While solvents are most commonly liquids, such as or , they can also be gases or in certain solution types. The process by which a solute dissolves in a solvent is known as , during which solvent molecules surround and interact with the solute particles to stabilize them within the solution. This stabilization occurs primarily through intermolecular forces, such as ion-dipole interactions when ionic solutes are involved, where the partial charges on polar solvent molecules attract oppositely charged ions. For molecular solutes, dipole-dipole forces between the polar solute and solvent molecules play a key role in facilitating . If the solvent is , this process is specifically termed hydration. A classic example of solute-solvent interaction is seen in saltwater, where (NaCl), a solid ionic compound, acts as the solute and serves as the ; the polar molecules surround the Na⁺ and Cl⁻ ions via ion-dipole forces to form the solution. The of a solute in a given is largely governed by the principle that "like dissolves like," meaning polar solutes tend to dissolve well in polar like , while nonpolar solutes are more soluble in nonpolar such as . This compatibility arises from the similar intermolecular forces between the solute and , ensuring effective .

Types

Gaseous Solutions

Gaseous solutions are homogeneous mixtures in which a gas acts as the solute dissolved in another gas serving as the , resulting in a uniform distribution of components at the molecular level. These solutions form a single gaseous phase and are prevalent in both natural and industrial environments. A prominent example is Earth's atmosphere, where constitutes approximately 78% of the and functions as the primary , while oxygen (about 21%), (0.93%), (about 0.043% as of 2025), and trace gases serve as solutes; this composition enables the mixture to approximate behavior under low-pressure conditions typical of the . The formation of gaseous solutions arises from the spontaneous intermingling of gas molecules, a process favored by a substantial increase in as the gases expand to occupy the available volume uniformly. Unlike solutions involving condensed phases, gaseous solutes exhibit complete with gaseous solvents in all proportions, with no upper limit on under ideal conditions, due to the negligible intermolecular forces at typical temperatures and pressures. This mixing is quantitatively described by of partial pressures, which states that the total pressure exerted by the mixture equals the sum of the partial pressures of the individual gases:
Ptotal=P1+P2+⋯+PnP_{\text{total}} = P_1 + P_2 + \cdots + P_n
where each partial pressure PiP_i is proportional to the of the component gas, Pi=xiPtotalP_i = x_i P_{\text{total}}. This law assumes non-reacting gases and ideal behavior, allowing prediction of mixture properties from individual gas contributions. Distinct properties of gaseous solutions include their exceptionally high diffusivity, stemming from the rapid, random motion of gas molecules as described by kinetic molecular theory, which enables quick homogenization even over large volumes. These solutions lack distinct phases or boundaries between components, maintaining uniformity throughout. In contrast to gas-liquid systems governed by for limited , gas-gas interactions impose no such restrictions, permitting arbitrary compositions while the overall mixture adheres to the at low densities. These characteristics make gaseous solutions highly relevant for modeling atmospheric dynamics and optimizing handling. Representative examples illustrate the practical significance of gaseous solutions. In the atmosphere, functions as a dilute solute in air, reaching peak concentrations of about 10 parts per million in the , where it absorbs harmful ultraviolet radiation. Industrially, —primarily a mixture of (solvent in some contexts) and , with minor and —exemplifies a engineered gaseous solution produced via of carbonaceous feedstocks for use in Fischer-Tropsch synthesis and production.

Liquid Solutions

Liquid solutions, where a liquid acts as the solvent, are among the most prevalent types of solutions in chemistry and everyday life, as they facilitate the dissolution of diverse solutes in a fluid medium. , often termed the universal solvent, exemplifies this due to its polar nature and ability to form bonds, which enable it to dissolve a broad array of polar, ionic, and some nonpolar substances more effectively than other liquids. This polarity arises from water's bent molecular structure, with the oxygen atom bearing a partial negative charge and the hydrogen atoms partial positive charges, promoting interactions with solutes like salts and sugars. Preparation of liquid solutions typically involves dissolving the solute in the through physical or chemical means. Common methods include mechanical stirring to enhance mixing and contact between solute particles and molecules, or gentle heating to increase the and of the solute, as seen in the creation of aqueous salt solutions like in . Solutions can also form via chemical reactions, such as the neutralization of acids and bases to produce ionic species in . In these solutions, the constitutes the major component, while the solute is the minor one, whether solid, liquid, or gas. The phase behavior of liquid solutions differs from that of pure solvents, maintaining liquidity over a range influenced by solute presence, which causes modest shifts in and freezing points. Notable examples include , a solution of various salts in essential for marine ecosystems; carbonated sodas, where gas dissolves in under pressure for ; and ethanol-water mixtures, which form homogeneous liquids used in beverages and fuels.

Solid Solutions

Solid solutions form when solute atoms or ions are incorporated into the crystal lattice of a solid solvent, creating a homogeneous mixture without phase separation. In this structure, the solute substitutes for or occupies positions within the solvent's lattice, maintaining the overall crystalline order while altering properties such as strength or conductivity. There are two primary types of solid solutions: substitutional and . In substitutional solid solutions, solute atoms replace atoms at lattice sites, requiring the solute to have an within about 15% of the 's to minimize lattice strain. A classic example is , where atoms substitute into the lattice, enhancing malleability and corrosion resistance. In contrast, solid solutions involve smaller solute atoms occupying voids between the atoms, which does not require size matching but limits due to available space. exemplifies this type, with carbon atoms fitting into the iron lattice to increase . The formation of solid solutions typically occurs through cooling from a molten state, where atoms arrange into a lattice during solidification; high-temperature , allowing solute migration into the lattice; or , where powdered metals are heated below their to promote atomic bonding and homogenization. For substitutional solutions to form extensively, the must be satisfied: the solute and solvent should share the same , have similar electronegativities to avoid compound formation, exhibit comparable valences (with lower-valence solutes often more soluble in higher-valence solvents), and maintain the aforementioned atomic compatibility. These empirical guidelines, derived from studies of binary metal systems, predict the extent of and stability of the resulting . Beyond metallic alloys, solid solutions are crucial in semiconductors, where controlled doping introduces impurities into the host lattice to modify electrical properties. For instance, atoms substitute into the lattice in n-type semiconductors, donating excess electrons due to phosphorus's higher valence, with solubility limits up to approximately 3 at%. Similarly, doping creates p-type by substituting atoms that accept electrons, forming carriers, and is widely used in integrated circuits. Another example is , a substitutional of silver with about 7.5% , which improves durability for jewelry and coinage.

Solubility

Measurement and Expression

Solubility is defined as the maximum quantity of solute that can dissolve in a certain quantity of or solution at a specified and pressure. This measure quantifies the equilibrium state where the rate of dissolution equals the rate of , establishing the limit beyond which additional solute remains undissolved. Solubility is commonly expressed using units such as grams of solute per 100 grams of (g/100 g), which provides a mass-based suitable for many practical applications. Other units include molarity, defined as moles of solute per liter of solution (M), which is widely used in settings for its relation to . Additionally, (X=nsolutentotalX = \frac{n_{\text{solute}}}{n_{\text{total}}}), a representing the of moles of solute to total moles in the solution, facilitates comparisons across different systems. Solubility curves offer a graphical representation of how the solubility of solids in liquids varies with temperature, typically plotting solubility on the y-axis against temperature on the x-axis. These curves illustrate trends, such as increasing solubility for most ionic solids like potassium nitrate with rising temperature, aiding in the prediction of dissolution behavior under varying conditions. For example, the solubility of sodium chloride (NaCl) in water is approximately 35.7 g per 100 mL at 20°C, often approximated as 36 g/100 g solvent given water's density near 1 g/mL. In contrast, the solubility of gases in liquids exhibits an inverse temperature dependence, decreasing as temperature increases due to the exothermic nature of gas dissolution.

Factors Affecting Solubility

The solubility of a solute in a is influenced by several key factors, including the chemical nature of the solute and , , , and the solution's or presence of common ions. These variables determine the extent to which intermolecular interactions allow the solute particles to disperse uniformly in the , affecting the equilibrium between dissolved and undissolved states. A fundamental principle governing solubility is the "like dissolves like" rule, which states that polar solutes tend to dissolve in polar solvents, while nonpolar solutes dissolve in nonpolar solvents, due to favorable intermolecular forces such as dipole-dipole interactions or dispersion forces matching between solute and solvent. For instance, ionic compounds like exhibit high solubility in polar solvents such as because the polar molecules can effectively solvate the charged ions through ion-dipole interactions, whereas nonpolar substances like are more soluble in nonpolar solvents like . This rule arises from the thermodynamic favorability of solute-solvent interactions over solute-solute or solvent-solvent interactions, as quantified in empirical models of liquid . Temperature significantly affects solubility, with the direction of change depending on whether the dissolution process is endothermic or exothermic. For many solid solutes where dissolution is endothermic (absorbing heat), solubility increases with rising temperature; (KNO₃), for example, shows a marked increase in solubility from about 13 g/100 mL at 0°C to over 240 g/100 mL at 100°C in . In contrast, for exothermic dissolution processes, solubility decreases with increasing temperature; (Ca(OH)₂) solubility drops from approximately 0.173 g/100 mL at 10°C to 0.065 g/100 mL at 100°C. For gases dissolved in liquids, solubility generally decreases with increasing temperature because higher disrupts the weak van der Waals forces holding gas molecules in solution; oxygen (O₂) in , for instance, has a solubility of approximately 0.070 g/L at 0°C but only 0.040 g/L at 25°C at 1 partial pressure. Pressure has a pronounced effect on the solubility of gases but minimal impact on and . According to , the solubility (S) of a gas in a is directly proportional to the partial pressure (P_gas) of the gas above the at constant temperature, expressed as: S=kH⋅PgasS = k_H \cdot P_{\text{gas}} where kHk_H is the Henry's law constant specific to the gas-solvent pair. For example, increasing the pressure of over enhances its solubility, which is the basis for carbonated beverages under pressure. This linear relationship holds for dilute solutions and low pressures, breaking down at high pressures due to non-ideal behavior. The of the solution and the also influence , particularly for ionic compounds with limited solubility. The occurs when an from an added soluble salt shifts the dissolution equilibrium toward the undissolved form via , reducing . For sparingly soluble (AgCl), which has a solubility product Ksp=1.8×10−10K_{sp} = 1.8 \times 10^{-10}, adding chloride ions from a soluble salt like NaCl decreases AgCl ; in pure , AgCl is about 1.34×10−51.34 \times 10^{-5} M, but in 0.1 M NaCl, it drops to approximately 1.8×10−91.8 \times 10^{-9} M. Similarly, affects the of weak acids or bases by altering their ionization state; acidic drugs like aspirin are more soluble in basic solutions where they form ionized, more polar species. These effects are crucial in controlling precipitation in and pharmaceutical formulations.

Properties

Concentration Units

In chemistry, the concentration of a solution quantifies the amount of solute present relative to the solvent or the total solution, providing a standardized way to describe solution composition for experimental and theoretical purposes. Various units are employed depending on the context, such as the nature of the solute and solvent, precision requirements, and temperature sensitivity. These units range from simple percentage-based measures to molar-based expressions that account for molecular quantities. Mass percentage, also known as weight/weight percent (w/w), expresses the concentration as the mass of solute divided by the total mass of the solution, multiplied by 100%. The formula is msolutemsolution×100%\frac{m_{\text{solute}}}{m_{\text{solution}}} \times 100\%, where mm denotes mass. This unit is straightforward and commonly used for solid solutes or non-volatile liquids, as it is independent of volume changes and easy to calculate from weighed components. Volume percentage, or volume/volume percent (v/v), is defined as the volume of solute divided by the total volume of the solution, multiplied by 100%. It is given by VsoluteVsolution×100%\frac{V_{\text{solute}}}{V_{\text{solution}}} \times 100\%. This measure is particularly applicable to liquid-liquid solutions, such as alcoholic beverages, where volumes are directly measurable, though it can be affected by mixing contractions or expansions. Molarity (M) represents the number of moles of solute per liter of solution and is calculated as M=nsoluteVsolutionM = \frac{n_{\text{solute}}}{V_{\text{solution}}}, where nn is moles and VV is volume in liters. It is widely used in stoichiometric calculations and preparations due to its relation to Avogadro's number, but it is temperature-dependent because solution volume varies with . For example, a 1 M NaCl solution contains 58.44 g of NaCl per liter, based on the of NaCl (58.44 g/mol). Molality (m) is the number of moles of solute per of , expressed as m=nsolutemsolventm = \frac{n_{\text{solute}}}{m_{\text{solvent}}}, with in kilograms. Unlike molarity, molality is temperature-independent since it relies on rather than , making it suitable for studies involving varying temperatures. Mole fraction (X) denotes the ratio of moles of solute to the total moles in the solution, Xsolute=nsolutentotalX_{\text{solute}} = \frac{n_{\text{solute}}}{n_{\text{total}}}. It is dimensionless and particularly useful for gaseous solutions or ideal mixtures, as it aligns with derivations and facilitates comparisons across different systems. For highly dilute solutions, such as environmental pollutants in , parts per million (ppm) is often used, defined as milligrams of solute per of solution (or equivalently, μ\mug/mL for aqueous solutions assuming ≈ 1 g/mL). This unit simplifies expression of trace concentrations, where 1 ppm corresponds to 1 mg/L in typical samples. can be viewed as the maximum concentration achievable before saturation occurs.

Colligative Properties

Colligative properties of solutions are physical properties that depend solely on the number of solute particles present, rather than on the chemical identity or nature of those particles. These properties arise from the interactions between solute particles and the , particularly in dilute solutions, and are proportional to the of the solute, which measures concentration in moles of solute per of . The primary include vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure. Vapor pressure lowering occurs when a nonvolatile solute is added to a solvent, reducing the solvent's tendency to evaporate. This effect is quantitatively described by , formulated by François-Marie Raoult in the late , which states that the partial vapor pressure of the solvent in the solution (PsolutionP_{\text{solution}}) equals the of the (XsolventX_{\text{solvent}}) multiplied by the of the pure (Psolvent∘P^\circ_{\text{solvent}}): Psolution=Xsolvent⋅Psolvent∘P_{\text{solution}} = X_{\text{solvent}} \cdot P^\circ_{\text{solvent}} The resulting decrease in vapor pressure (ΔP=Psolvent∘−Psolution\Delta P = P^\circ_{\text{solvent}} - P_{\text{solution}}) is directly proportional to the mole fraction of the solute. Boiling point elevation and freezing point depression are temperature-related colligative effects influenced by the solute's impact on the solvent's phase transitions. The elevation in boiling point (ΔTb\Delta T_b) is given by the equation ΔTb=Kb⋅m⋅i\Delta T_b = K_b \cdot m \cdot i, where KbK_b is the molal boiling point elevation constant (specific to the ), mm is the molality of the solute, and ii is the van't Hoff factor accounting for the number of particles produced by dissociation or association of the solute. Similarly, the depression in freezing point (ΔTf\Delta T_f) follows ΔTf=Kf⋅m⋅i\Delta T_f = K_f \cdot m \cdot i, with KfK_f as the molal freezing point depression constant. The van't Hoff factor ii, introduced by , equals the ratio of the actual number of particles in solution to the number expected without dissociation; for example, i=2i = 2 for NaCl in ideal complete dissociation into Na+^+ and Cl−^-. Osmotic pressure (π\pi) represents the pressure required to prevent the net flow of across a separating a solution from pure , driven by the concentration difference. For dilute solutions, it is expressed as π=i⋅M⋅R⋅T\pi = i \cdot M \cdot R \cdot T, where MM is the molarity of the solute, RR is the (0.08206 L·atm·mol−1^{-1}·K−1^{-1}), and TT is the absolute in . This formula, also derived by van 't Hoff, treats the solute particles as exerting an effective pressure analogous to an . Practical applications of abound in everyday and industrial contexts. For instance, is added to in automotive to lower the freezing point via depression, preventing radiator fluid from solidifying in cold weather; a 50% - mixture can depress the freezing point to about -37°C. In biological and medical settings, dialysis exploits to remove waste solutes from blood across a , allowing and small molecules to pass while retaining larger proteins and cells.

Liquid Solution Characteristics

Saturation States

In liquid solutions, saturation states describe the relative amount of solute dissolved compared to the maximum possible at a given , determining whether additional solute can be accommodated. These states are fundamental to understanding solution behavior and are defined in relation to the solubility limit, which varies with —typically increasing for most solutes in solvents as rises. An unsaturated solution contains less solute than the solubility limit, allowing further dissolution without occurring. In this state, the solution remains stable, and adding more solute results in complete dissolution until the saturation point is approached. A saturated solution holds exactly the maximum amount of solute dictated by its at that , with undissolved solute potentially present in equilibrium with the dissolved form. This condition represents a dynamic equilibrium where the rate of solute dissolution equals the rate of , maintaining a constant solute concentration over time. Supersaturated solutions exceed the solubility limit, containing more dissolved solute than would be stable at equilibrium, rendering them metastable and prone to rapid upon disturbance. Such states are typically achieved by preparing a saturated solution at elevated —where is higher—and then cooling it slowly to avoid . A classic example is the supersaturated sodium acetate solution in reusable hand warmers, which crystallizes exothermically when triggered by flexing a metal disc to initiate , releasing . in supersaturated solutions can also be induced by seeding with a or mechanical agitation, converting the excess solute to solid form.

Thermodynamic Aspects

The formation and stability of liquid solutions are governed by thermodynamic principles, particularly the change associated with the mixing process. The of mixing, ΔGmix\Delta G_{\text{mix}}, determines whether a solution forms spontaneously and is given by ΔGmix=ΔHmix−TΔSmix\Delta G_{\text{mix}} = \Delta H_{\text{mix}} - T \Delta S_{\text{mix}}, where ΔHmix\Delta H_{\text{mix}} is the , TT is the absolute , and ΔSmix\Delta S_{\text{mix}} is the . For the process to be spontaneous, ΔGmix<0\Delta G_{\text{mix}} < 0, which typically occurs when the favorable entropy increase from dispersing solute molecules among solvent molecules outweighs any unfavorable enthalpy changes due to intermolecular interactions. In ideal solutions, the interactions between unlike molecules are similar in strength to those between like molecules, leading to ΔHmix=0\Delta H_{\text{mix}} = 0 and adherence to over the entire composition range. states that the partial vapor pressure of each component is proportional to its mole fraction, Pi=xiPi∘P_i = x_i P_i^\circ, where Pi∘P_i^\circ is the vapor pressure of the pure component. A classic example is the benzene-toluene system, where the similar molecular structures result in nearly ideal behavior, with minimal deviations in vapor-liquid equilibrium data. Non-ideal solutions, in contrast, exhibit deviations from Raoult's law due to differing intermolecular forces, resulting in either positive or negative deviations in vapor pressure. Positive deviations occur when solute-solvent interactions are weaker than average solute-solute and solvent-solvent interactions, increasing volatility and often forming minimum-boiling azeotropes; the ethanol-water mixture is a prominent example, forming an azeotrope at approximately 95.6% ethanol by weight that boils at 78.2°C, preventing complete separation by simple distillation. Negative deviations arise from stronger solute-solvent attractions, leading to lower volatility and maximum-boiling azeotropes. To account for non-ideality in thermodynamic calculations, the concept of activity is introduced, where the effective concentration is the activity ai=γixia_i = \gamma_i x_i, with γi\gamma_i as the activity coefficient that corrects for deviations from ideality. For volatile components, fugacity fif_i serves a similar role, defined as fi=γixifi∘f_i = \gamma_i x_i f_i^\circ, where fi∘f_i^\circ is the fugacity of the pure component, providing a measure of the escaping tendency adjusted for non-ideal behavior. In the HCl-water system, strong ion-dipole interactions between hydrated chloride and hydronium ions lead to significant negative deviations, with activity coefficients varying markedly with concentration due to these electrostatic forces. The enthalpy of solution, ΔHsoln\Delta H_{\text{soln}}, reflects the net heat absorbed or released during dissolution and contributes to ΔHmix\Delta H_{\text{mix}}. Endothermic processes (ΔHsoln>0\Delta H_{\text{soln}} > 0) occur when breaking solute-solute and solvent-solvent bonds requires more energy than is released by forming solute-solvent bonds, as seen in dissolving in , which cools the solution. Exothermic processes (ΔHsoln<0\Delta H_{\text{soln}} < 0) result from stronger solute-solvent attractions, such as in the dissolution of HCl in , where the heat of hydration drives an overall energy release. These enthalpic contributions, combined with entropic effects, dictate the overall ΔGmix\Delta G_{\text{mix}} and solution stability.

References

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  2. https://www.chem.purdue.edu/gchelp/solutions/whatis.html
  3. https://pressbooks.online.ucf.edu/chemistryfundamentals/chapter/solubility-2/
  4. https://www.chem.purdue.edu/gchelp/howtosolveit/Solutions/concentrations.html
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  7. https://chem.libretexts.org/Courses/College_of_Marin/CHEM_114%253A_Introductory_Chemistry/13%253A_Solutions/13.02%253A_Solutions-_Homogeneous_Mixtures
  8. https://www.etymonline.com/word/solution
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  10. https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch3/solution.html
  11. https://genchem1.chem.okstate.edu/BDA/Lecture/Chapter13/Lec20507.html
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