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Salicylic Acid & Aspirin

Views: 0     Author: Site Editor     Publish Time: 2023-08-14      Origin: Site

First marketed under the trademark Aspirin in 1899, acetylsalicylic acid quickly gained a worldwide reputation as a non-prescription pain relief drug. Even today, it remains one of the most widely used drugs for treating headaches, body and muscle pains, arthritic pain, and more.

Salicylic acid is the major hydrolysis product of acetylsalicylic acid, as shown in Equation 1 below. The detection and quantification of salicylic acid in samples containing acetylsalicylic acid can be important for several reasons. Firstly, the degradation of the active ingredient during the shelf life of the drug is a major concern in the pharmaceutical industry. Secondly, the production of counterfeit drugs is a growing concern, and inadequate or unregulated synthesis of acetylsalicylic acid can lead to contamination with significant amounts of unreacted salicylic acid, which can have adverse health effects.

A simple yet effective method for detecting free salicylic acid is based on the differing coordination abilities of salicylic compounds with the trivalent iron cation Fe+3. Being a transition metal, Fe+3 can form coordination complexes with various salicylic compounds, resulting in highly colored solutions for qualitative analysis. Additionally, calibration standards can be prepared for quantitative spectrophotometric analysis. The reaction of phenols with ferric chloride is a well-known test in this regard. Phenol groups form a purple complex with Fe+3, with the color intensity being related to the coordination capacity of the phenol group. Both acetylsalicylic acid and salicylic acid contain phenol groups, but in acetylsalicylic acid, the phenol group is bonded to an acetyl group, thereby reducing its coordination capacity. As a result, while both compounds will form a colored complex with Fe+3, acetylsalicylic acid forms a slightly yellow-orange colored complex, whereas salicylic acid forms a highly colored deep purple complex.

Most transition metal ions act as Lewis acids bySalicylic Acidforming coordinate covalent bonds with ligands. When transition metal ions are bonded to ligands, they form complex ions, which can be cationic, anionic, or even neutral. These complex ions usually exhibit distinct colors, depending on the specific transition metal and ligands involved.

Substances exhibit color when they absorb specific wavelengths of light within the visible region of the electromagnetic spectrum (approximately 400-700 nm) while reflecting or transmitting other wavelengths. This absorption occurs due to the quantized energy levels within molecules, allowing only the absorption of light energies corresponding to the difference between electronic levels.

In the case of Fe+3, octahedral complexes are formed in aqueous solution. Crystal Field Theory considers ligands as point charges, leading to the assumption of an octahedral field and the calculation of its effect on the d-orbitals.

Uncomplexed Fe+3 has all d-orbitals at the same energy level. However, the presence of ligands alters the energy levels of certain d-orbitals in octahedral complexes. The ligands approach the central metal atom along the x-, y-, and z-axes, causing the dxy, dxz, and dyz orbitals to fill due to covalent bonding. Consequently, the dx2-y2 and dz2 orbitals have higher energy due to repulsion from the filled orbitals. The energy difference between these two groups of d-orbitals is known as the Crystal Field Splitting, D. In the case of Fe+3 complexed with phenolic compounds, this splitting corresponds to wavelengths in the visible light region. The complex formed between salicylic acid ligands and Fe+3 exhibits crystal field splitting with an energy difference corresponding to deep purple light, while the complex between acetylsalicylic acid ligands and Fe+3 results in crystal field splitting with an energy difference corresponding to yellowish light.

The intensity of light absorption depends on the concentration of the complex in solution, following the Beer-Lambert Law. Standard spectrophotometric tests for salicylic acid Fe+3 complexes involve measuring absorbance at a wavelength around 540 nm. For quick qualitative detection of acetylsalicylic acid and salicylic acid, a simple visual test is sufficient. By adding an excess of Fe+3 in a suitable form, such as ferric chloride, to a test solution, complexes of Fe+3 will form. 

If the solution turns deep purple, salicylic acid is present. If it turns yellow-orange, no salicylic acid is present, suggesting that only acetylsalicylic acid is present in the test solution. This visual test allows for quick assessment of aspirin samples. Commercial pharmaceutical preparations of aspirin should not contain significant quantities of salicylic acid, while acetylsalicylic acid preparations made in a university lab may contain unreacted salicylic acid.

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