Dissolution (chemistry)

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

Gold, formerly dissolved in crystal of pyrite, is left behind after the cubic crystal of pyrite dissolved away. Note a corner of the former cube seen in center of rock.

The dissolution of gases, liquids, or solids into a liquid or other solvent is a process by which these original states become solutes (dissolved components), forming a solution of the gas, liquid, or solid in the original solvent. Solid solutions are the result of dissolution of one solid into another, and occur, e.g., in metal alloys, where their formation is governed and described by the relevant phase diagram. In the case of a crystalline solid dissolving in a liquid, the crystalline structure must be disintegrated such that the separate atoms, ions, or molecules are released. For liquids and gases, the molecules must be able to form non-covalent intermolecular interactions with those of the solvent for a solution to form.

The free energy of the overall, isolated process of dissolution must be negative for it to occur, where the component free energies contributing include those describing the disintegration of the associations holding the original solute components together, the original associations of the bulk solvent, and the old and new associations between the undissolved and dissolved materials.

Dissolution is of fundamental importance in all chemical processes, natural and unnatural, from the decomposition of a dying organism and return of its chemical constituents into the biosphere, to the laboratory testing of new, man-made soluble drugs, catalysts, etc. Dissolution testing is widely used in industry, including in the pharmaceutical industry to prepare and formulate chemical agents of consistent quality that will dissolve, optimally, in their target millieus as they were designed.

Some distinctions can be made between solvation, dissolution, and solubility.

Dissolution by class of compound

Gases

Gaseous elements and compounds will dissolve in liquids dependent on the interaction of their bonds with the liquid solvent.

Liquids

Gaseous elements and compounds may also dissolve in another liquid depending on the compatibility of the chemical and physical bonds in the substance with those of the solvent. Hydrogen bonds play an important role in aqueous dissolution.

Ionic compounds

For ionic compounds, dissolution takes place when the ionic lattice breaks up and the separate ions are then solvated. This most commonly occurs in polar solvents, such as water or ammonia:

NaCl_(s) → Na⁺_(aq) + Cl⁻_(aq)

In a colloidal dispersed system, small dispersed particles of the ionic lattice exist in equilibrium with the saturated solution of the ions, i.e.

NaCl_(aq) $\rightleftharpoons$ Na⁺_(aq) + Cl⁻_(aq)

The solubility of ionic salts in water is generally determined by the degree of solvation of the ions by water molecules. Such coordination complexes occur by water donating spare electrons on the oxygen atom to the ion. The behavior of this system is characterised by the activity coefficients of the components and the solubility product, defined as:

${\displaystyle {\ce {{\mathit {a}}_{Na+}.{\mathit {a}}_{Cl^{-}}={\mathit {K_{sp}}}}}}$

The ability of an ion to preferentially dissolve (as a result of unequal activities) is classified as the Potential Determining Ion. This in turn results in the remaining particle possessing either a net positive/negative surface charge.

Oxides

The dissolution of oxide minerals such as silicates occurs by several mechanisms which depend on the composition of the mineral and the chemistry of the solution.^[1] Dissolution rates partially depend on solution pH. Adsorbed protons or hydroxides polarize the mineral surface and weaken cation-oxygen bonds, accelerating dissolution. Silicate minerals containing metal cations undergo incongruent dissolution as the cations leach out of the mineral faster than the silica lattice degrades. Incongruent dissolution results in a surface layer with different composition than the bulk, called an alteration layer.

The reductive dissolution of a transition metal oxide can occur when a redox event in solution reduces a cation. Dissolution occurs when the reduced cation is unstable in the solid material.^[2] In minerals such as ferric oxides, reduction may be caused by electron transfer from organic molecules^[3] or bacteria^[4] in anoxic waters or soils. Charge carriers responsible for reductive dissolution may also be introduced by photoexcitation or by electrochemical poising at negative potentials. Reductive dissolution is integral to natural geochemical phenomena such as the iron cycle.^[5]

Using ferric oxide (${\displaystyle {\ce {Fe2O3}}}$) as an example, the fundamental formula of reductive dissolution is:

${\displaystyle {\ce {{Fe^{3}+_{(}s)}+{e^{-}}->{Fe^{2}+_{(}aq)}}}}$

Here, an ${\displaystyle {\ce {Fe^{3}+}}}$ cation at the oxide surface captures an electron (${\displaystyle {\ce {e^{-}}}}$), converting the cation to ${\displaystyle {\ce {Fe^{2}+}}}$. However, ${\displaystyle {\ce {Fe^{2}+}}}$ is unstable in the oxide lattice relative to the solution and is subsequently solvated.

Reductants causing reductive dissolution include natural electron donors such as ascorbic acid and ${\displaystyle {\ce {Fe^{2}+_{(}aq)}}}$. Chelating species such as oxalate accelerate the process by detaching surface-bound ${\displaystyle {\ce {Fe^{2}+}}}$, opening surface sites for further attack by reductants. Reductive dissolution is also promoted by light.^[3]

Reductive dissolution does not necessarily occur at the site of reductant adsorption, particularly for conductive specimens. Excess electrons injected into a hematite particle during a redox event can travel through the particle, causing reductive dissolution elsewhere on the particle.^[6] The transport of charge across a hematite particle is driven by differences in the surface potential of different crystal terminations.^[7]

Semiconductors

Photocorrosion is the light-induced degradation and dissolution of semiconductor materials used as electrodes in photoelectrochemical cells. This can occur when photoexcited charge carriers change the oxidation state of surface atoms or ions, destabilizing the material. Materials with smaller bandgaps which can absorb larger regions of the solar spectrum are more susceptible to photocorrosion.^[8] In photocatalytic water splitting using a cadmium sulfide photoelectrode, for example, it is desired that holes (${\displaystyle {\ce {h^{+}}}}$) generated in ${\displaystyle {\ce {CdS}}}$ by absorption of photons will oxidize hydroxyl species in solution:

${\displaystyle {\ce {2{h^{+}}+{OH^{-}}->{1/2O_{2}}+{H^{+}}}}}$

However, in a competing pathway, holes may instead degrade ${\displaystyle {\ce {CdS}}}$:^[9]

${\displaystyle {\ce {2{h^{+}}+{CdS}->{Cd^{2}+}+{S}}}}$

The photocorrosion of some photo-absorbing electrodes can be mitigated by using protective thin film coatings.^[10]

Polar compounds

Polar solid compounds can be amorphous or crystalline. Crystalline solids dissolve with breakdown of their crystal lattice, and due to their polarity, or non-polarity, mix with the solvent.

Polymers

The solubility of polymers depends on the chemical bonds present in the backbone chain and their compatibility with those of the solvent. The Hildebrand solubility parameter is commonly used to evaluate polymer solubility. The closer the value of the parameters, the more likely dissolution will occur.

Rate of dissolution

The rate of dissolution quantifies the speed of the dissolution process. It depends on the chemical natures of the solvent and solute, the temperature (and possibly to a small degree, the pressure), the degree of undersaturation, the presence of a means of mixing during the dissolution, the interfacial surface area, and the presence of "inhibitors" (e.g., substances adsorbed on the surface).

The rate can be often expressed by the Noyes–Whitney equation or the Nernst and Brunner equation^[11] of the form:

${\displaystyle {\frac {dm}{dt}}=A{\frac {D}{d}}(C_{\mathrm {s} }-C_{\mathrm {b} })}$

where:

m = mass of dissolved material

t = time

A = surface area of the interface between the dissolving substance and the solvent

D = diffusion coefficient

d = thickness of the boundary layer of the solvent at the surface of the dissolving substance

C_s = mass concentration of the substance on the surface

C_b = mass concentration of the substance in the bulk of the solvent

For dissolution limited by diffusion, C_s is equal to the solubility of the substance. When the dissolution rate of a pure substance is normalized to the surface area of the solid (which usually changes with time during the dissolution process), then it is expressed in kg/m²s and referred to as "intrinsic dissolution rate". The intrinsic dissolution rate is defined by the United States Pharmacopeia.

Dissolution rates vary by orders of magnitude between different systems. Typically, very low dissolution rates parallel low solubilities, and substances with high solubilities exhibit high dissolution rates, as suggested by the Noyes-Whitney equation. However, this is not a rule.

See also

• Solvation
• Hansen solubility parameter
• Rate of solution
• Solubilization

References

1. Brantley, Susan L.; Kubicki, James D.; White, Art F. (2008). Kinetics of Water-Rock Interaction. pp. 151–210. doi:10.1007/978-0-387-73563-4.
2. Cornell, R. M.; Schwertmann, U. (2003). The Iron Oxides: Structure, Properties, Reactions, Occurences and Uses, Second Edition. p. 306. doi:10.1002/3527602097.
3. Sulzberger, Barbara; Suter, Daniel; Siffert, Christophe; Banwart, Steven; Stumm, Werner (1989). "Dissolution of fe(iii)(hydr)oxides in natural waters; laboratory assessment on the kinetics controlled by surface coordination". Marine Chemistry. 28 (1-3): 127–144. doi:10.1016/0304-4203(89)90191-6. ISSN 0304-4203.
4. Roden, Eric E. (2008). "Microbiological Controls on Geochemical Kinetics 1: Fundamentals and Case Study on Microbial Fe(III) Oxide Reduction": 335–415. doi:10.1007/978-0-387-73563-4_8.
5. Cornell, R. M.; Schwertmann, U. (2003). The Iron Oxides: Structure, Properties, Reactions, Occurences and Uses, Second Edition. p. 323. doi:10.1002/3527602097.
6. Yanina, S. V.; Rosso, K. M. (2008). "Linked Reactivity at Mineral-Water Interfaces Through Bulk Crystal Conduction". Science. 320 (5873): 218–222. doi:10.1126/science.1154833. ISSN 0036-8075.
7. Chatman, S.; Zarzycki, P.; Rosso, K. M. (2015). "Spontaneous Water Oxidation at Hematite (α-Fe2O3) Crystal Faces". ACS Applied Materials & Interfaces. 7 (3): 1550–1559. doi:10.1021/am5067783. ISSN 1944-8244.
8. Li, Jiangtian; Wu, Nianqiang (2015). "Semiconductor-based photocatalysts and photoelectrochemical cells for solar fuel generation: a review". Catal. Sci. Technol. 5 (3): 1360–1384. doi:10.1039/C4CY00974F. ISSN 2044-4753.
9. Ashokkumar, M (1998). "An overview on semiconductor particulate systems for photoproduction of hydrogen". International Journal of Hydrogen Energy. 23 (6): 427–438. doi:10.1016/S0360-3199(97)00103-1. ISSN 0360-3199.
10. Hu, Shu; Lewis, Nathan S.; Ager, Joel W.; Yang, Jinhui; McKone, James R.; Strandwitz, Nicholas C. (2015). "Thin-Film Materials for the Protection of Semiconducting Photoelectrodes in Solar-Fuel Generators". The Journal of Physical Chemistry C. 119 (43): 24201–24228. doi:10.1021/acs.jpcc.5b05976. ISSN 1932-7447.
11. Dokoumetzidis, Aristides; Macheras, Panos (2006). "A century of dissolution research: From Noyes and Whitney to the Biopharmaceutics Classification System". Int. J. Pharm. 321 (1–2): 1–11. doi:10.1016/j.ijpharm.2006.07.011.

Further reading

• Brady, Patrick V.; House, William A. (1996). "4. Surface-controlled dissolution and growth of minerals". In Brady, Patrick V. Physics and chemistry of mineral surfaces. Boca Raton, Fla.: CRC Press. pp. 226–298. ISBN 9780849383519.
• Kumar, Vijai; Tewari, Divya (2006). "Dissolution". In Troy, David B.; Beringer, Paul. Remington: The Science and Practice of Pharmacy. Lippincott Williams & Wilkins. pp. 672–690. ISBN 9780781746731.
• Whitten, Kenneth W. (2014). Chemistry (10th ed.). Belmont, CA: Brooks/Cole Cengage Learning. pp. 506–514. ISBN 9781133610663.