Trimer (chemistry)

From Wikipedia, the free encyclopedia
  (Redirected from Trimerization)
Jump to: navigation, search

In chemistry, a trimer (/ˈtrmər/) (tri-, "three" + -mer, "parts") is an oligomer derived from three identical precursors. An example is the procedure of production of polymers. At first, a monomer is made. By combining two monomers, a dimer is produced. With further additions, a trimer and eventually a polymer is made. Often, trimerization competes with polymerization; for example, dimethylsilanediol polymerizes to polydimethylsiloxane, even though a trimer is made:

Me2Si(OH)2 + (HO)2SiMe2 → Me2(OH)Si-O-Si(OH)Me2dimer
Me2(OH)Si-O-Si(OH)Me2 + (HO)2SiMe2 → Me2(OH)Si-O-SiMe2-O-Si(OH)Me2trimer
Me2(OH)Si-O-SiMe2-O-Si(OH)Me2 + (HO)2SiMe2 → Me2(OH)Si-O-SiMe2-O-SiMe2-O-Si(OH)Me2tetramer
et cetera, until Me2(OH)Si-O-[SiMe2-O-]nSi(OH)Me2 (n>100) – polymer

Trimers are typically cyclic. Chemical compounds that often trimerise are aliphatic isocyanates and cyanic acids.

Trimerisation via breaking triple bond[edit]

Alkyne trimerisation[edit]

Further information: Alkyne trimerisation
The trimerization cyclization reaction can be understood with this scheme.

In 1866, Marcellin Berthelot reported the first example of cyclotrimerization leading to aromatic products, the cyclization of acetylene to benzene.[1] In the Reppe synthesis, the trimerization of acetylene gives benzene:

Reppe-chemistry-benzene.png

Breaking carbon-hetero triple bonds forms symmetrical unsaturated 1,3,5-heterocycles[edit]

Symmetrical 1,3,5-triazines are prepared by trimerization of certain nitriles such as cyanogen chloride or cyanimide.

Cyanuric chloride is prepared in two steps from hydrogen cyanide via the intermediacy of cyanogen chloride, which is trimerized at elevated temperatures over a carbon catalyst:

HCN + Cl2 → ClCN + HCl
Cyanurchloride Synthesis V.1.svg

In 2005, approximately 200,000 tonnes were produced.[2]

Cyanuric bromide is analogously prepared from cyclotrimerization of cyanogen bromide:

3 BrCN → (BrCN)3

An industrial route to cyanuric acid entails the thermal decomposition of urea, with release of ammonia. The conversion commences at approximately 175 °C:[2]

3 H2N-CO-NH2 → [C(O)NH]3 + 3 NH3

Synthesis of melamine[edit]

The endothermic synthesis of melamine can be understood in two steps. First, urea decomposes into cyanic acid and ammonia in an endothermic reaction:

(NH2)2CO → HOCN + NH3

Then in the second step, cyanic acid polymerizes to form cyanuric acid which condenses with the liberated ammonia from the first step to release melamine and water.

3 HOCN → [C(O)NH]3
[C(O)NH]3 + 3 NH3 → C3H6N6 + 3 H2O

This water then reacts with cyanic acid present, which helps drive the trimerization reaction, generating carbon dioxide and ammonia.

3 HOCN + 3 H2O → 3 CO2 + 3NH3

In total, the second step is exothermic:

6 HCNO + 3 NH3 → C3H6N6 + 3 CO2 + 3NH3

but the overall process is endothermic.

-Ene trimerisation[edit]

Diene trimerisation[edit]

The 1,5,9-trans-trans-cis isomer of cyclododecatriene C12H18 has some industrial importance [3] and is obtained by cyclotrimerization of butadiene with titanium tetrachloride and an organoaluminium co-catalyst:[4]

Cyclododeca-1,5,9-triene

Catalyzing and dehydrating by sulfuric acid, trimerization of acetone via aldol condensation affords mesitylene.

Breaking carbon-hetero double bonds forms symmetrical saturated 1,3,5-heterocycles[edit]

Cyclotrimerization of formaldehyde affords 1,3,5-Trioxane:

Trioxane Synthesis V.1.svg

1,3,5-Trithiane is the cyclic trimer of the otherwise unstable species thioformaldehyde. This heterocycle consists of a six-membered ring with alternating methylene bridges and thioether groups. It is prepared by treatment of formaldehyde with hydrogen sulfide.[5]

Three molecules of acetaldehyde condense[how?] to form “paraldehyde,” a cyclic trimer containing C-O single bonds.

References[edit]

  1. ^ Schetter, M. C. R. (1866). Hebd. Seances Acad. Sci. 62: 905. 
  2. ^ a b Klaus Huthmacher, Dieter Most "Cyanuric Acid and Cyanuric Chloride" in Ullmann's Encyclopedia of Industrial Chemistry, 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a08_191.
  3. ^ 1995 figures for global annual capacity 8000 ton
  4. ^ Industrial Organic Chemistry, Klaus Weissermel, Hans-Jurgen Arpe John Wiley & Sons; 3rd 1997 ISBN 3-527-28838-4
  5. ^ Bost, R. W.; Constable, E. W. "sym-Trithiane" Organic Syntheses, Collected Volume 2, p.610 (1943). http://www.orgsyn.org/orgsyn/pdfs/CV2P0610.pdf