| Terephthalic acid | |
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Other names
Benzene-1,4-dicarboxylic acid |
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| Identifiers | |
| CAS number | 100-21-0 |
| ChemSpider | 7208 |
| ChEBI | CHEBI:15702 |
| RTECS number | WZ0875000 |
| Jmol-3D images | Image 1 |
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| Properties | |
| Molecular formula | C8H6O4 |
| Molar mass | 166.13 g/mol |
| Appearance | white crystals or powder |
| Density | 1.522 g/cm³ |
| Melting point |
300°C in a sealed tube (sublimes at 402°C (675 K) in air) |
| Boiling point |
sublimes |
| Solubility in water | 0.0017 g/100 mL at 25°C |
| Solubility | polar organic solvents aqueous base |
| Acidity (pKa) | 3.51, 4.82[1] |
| Structure | |
| Dipole moment | zero |
| Hazards | |
| MSDS | External MSDS |
| EU classification | not listed |
| Related compounds | |
| Related carboxylic acids | Phthalic acid Isophthalic acid Benzoic acid p-Toluic acid |
| Related compounds | p-Xylene Polyethylene terephthalate Dimethyl terephthalate |
| Supplementary data page | |
| Structure and properties |
n, εr, etc. |
| Thermodynamic data |
Phase behaviour Solid, liquid, gas |
| Spectral data | UV, IR, NMR, MS |
| (verify) (what is: /?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) |
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| Infobox references | |
Terephthalic acid is the organic compound with formula C6H4(COOH)2. This colourless solid is a commodity chemical, used principally as a precursor to the polyester PET, used to make clothing and plastic bottles. Several billion kilograms are produced annually. It is one of three isomeric phthalic acids.
Contents |
Terephthalic acid is poorly soluble in water and alcohols, consequently up until around 1970 most crude terephthalic acid was converted to the dimethyl ester for purification. It sublimes when heated.
Terephthalic acid is produced by oxidation of p-xylene by oxygen in air:
This reaction proceeds through a p-toluic acid intermediate which is then oxidized to terephthalic acid. In p-toluic acid, deactivation of the methyl by the electron withdrawing carboxylic acid group makes the methyl one tenth as reactive as xylene itself, making the second oxidation significantly more difficult. [2] The commercial process utilizes acetic acid as solvent and a catalyst composed of cobalt and manganese salts, with a bromide promoter. The yield is nearly quantitative. The most problematic impurity is 4-formylbenzoic acid, which is removed by hydrogenation of a hot aqueous solution. This solution is then cooled in a stepwise manner to crystallize highly pure terephthalic acid.
Despite optimized yields greater than 95% with excellent purity, the synthesis has proven to have shortcomings. Due to high reaction temperature, approximately 5% of the acetic acid solvent is lost by decomposition or ‘burning.’ Solvent burning is a significant economic factor in the oxidation process. In addition, product loss by decarboxylation to benzoic acid is common. The high temperature diminishes oxygen solubility in an already oxygen starved system. Pure oxygen cannot be used in the traditional system due to hazards of flammable organic-O2 mixtures. Atmospheric air can be used in its place, but once reacted needs to be purified of toxins and ozone depleters such as methylbromide before being released. Additionally, the corrosive nature of bromides at high temperatures, requires the reaction be run in expensive titanium reactors.[3] [4]
Alternatively, but not commercially significant, is the so-called "Henkel process" or "Raecke process," named after the company and patent holder, respectively. This process involves the rearrangement of phthalic acid to terephthalic acid via the corresponding potassium salts.[5][6] Terephthalic acid can be prepared in the laboratory by oxidizing various para-disubstituted derivatives of benzene, including Caraway Oil or a mixture of cymene and cuminol with chromic acid.
The use of CO2 overcomes many of the problems with the original industrial process. Because CO2 is a better flame inhibitor than nitrogen gas, pure oxygen is used directly instead of air with reduced flammability hazards. The solubility of molecular oxygen in solution is also enhanced in the CO2 environment. Because more oxygen is available to the system, supercritical carbon dioxide (Tc = 31 ⁰C) has more complete oxidation with fewer byproducts, lower CO production, less decarboxylation and higher purity than the commercial process. [7] [8]
When reaction run is in supercritical water can be effectively catalyzed by MnBr2 with pure O2 in a medium-high temperature. Use of supercritical water instead of acetic acid as a solvent diminishes environmental impact and offers a cost advantage. However, the scope of such reaction systems is limited by the even harsher conditions than the industrial process (T = 300−400 °C, P > 200 bar).[9]
Ketones have been found to act as promoters for formation of the active Co(III) catalyst. In particular, ketones with a-methylene groups oxidize to hydroperoxides that are known to oxidize Co(II). Viable ketones were butanone, triacetylmethane (TAM), 2,3-pentanedione (2,3-PD), and acetylacetone; all of which can stabilize radical formation through resonance.[10]
Reactions run at temperatures as low as 100 ⁰C are possible by using zirconium salts as a cocatalyst in place of bromide and manganese acetate. It is thought that the Zr(IV) acts to oxidize Co(II) to the active Co(III). This alone shortens the induction period, and has been shown to have a synergistic effect with ketones. However, a greater amount of cobalt acetate is required than the common industrial process and is ineffective over 160 ⁰C.3
The addition of a small portion of metalloporphyrin, in particular T(p-Cl)PPMnCl, has a cocatalytic effect with the traditional Co(OAc)2 catalyst. This requires less acetic acid and does not require bromides. The catalytic effect has been attributed to the ease of peroxide formation over the metalloporphyrin.[11][12]
The autoxidation of p-xylene is known to proceed through a free radical process. The Mn(III) and Co(III) metals alone are not strong enough oxidizers to start the radical chain reaction, but instead initiate it by forming bromine radicals from the ions in solution. These bromine radicals then decompose hydroperoxides that are ligated to the metals as well as abstract hydrogens from the methyl groups on p-xylene to form free radicals and propagate the reaction. The following are the proposed initiation, propagation and terminations steps for the first of four oxidations involved in the autoxidation:
The radical chain reaction proceeds through a series of intermediates, starting with the oxidation of p-xylene to p-tolualdehyde (TALD), then p-toluic acid (PT), 4-carboxybenzaldehyde (4-CBA), and finally to the terephthalic acid (TA) product.
The kinetics of the oxidation hare exceedingly complex, but a general understanding of the mechanism has been established.[13]
Virtually the entire world's supply of terephthalic acid and dimethyl terephthalate are consumed as precursors to polyethylene terephthalate (PET). World production in 1970 was around 1.75 million tonnes.[14] By 2006, global purified terephthalic acid (TPA) demand had exceeded 30 million tonnes.
There is a smaller, but nevertheless significant, demand for terephthalic acid in the production of polybutylene terephthalate and several other engineering polymers.[15]
In the research laboratory, terephthalic acid has been popularized as a component for the synthesis of metal-organic frameworks.
The anxiolytic, analgesic, and antidepressant oxycodone occasionally comes as a terephthalate salt; however, the more usual salt of oxycodone is the hydrochloride. Pharmacologically, one milligram of terephthalas oxycodonae is equivalent to 1.13 mg of hydrochloridum oxycodonae.