|Preferred IUPAC name
Hot ice (sodium acetate trihydrate)
3D model (JSmol)
|E number||E262 (preservatives)|
CompTox Dashboard (EPA)
|Appearance||White deliquescent powder|
|Odor||Vinegar (acetic acid) odor when heated to decomposition|
|Density||1.528 g/cm3 (20 °C, anhydrous)|
1.45 g/cm3 (20 °C, trihydrate)
|Melting point|| 324 °C (615 °F; 597 K) |
58 °C (136 °F; 331 K)
|Boiling point|| 881.4 °C (1,618.5 °F; 1,154.5 K) |
122 °C (252 °F; 395 K)
119 g/100 mL (0 °C)
123.3 g/100 mL (20 °C)
125.5 g/100 mL (30 °C)
137.2 g/100 mL (60 °C)
162.9 g/100 mL (100 °C)
32.9 g/100 mL (-10 °C)
36.2 g/100 mL (0 °C)
46.4 g/100 mL (20 °C)
82 g/100 mL (50 °C)
|Solubility||Soluble in alcohol, hydrazine, SO2|
|Solubility in methanol||16 g/100 g (15 °C)|
16.55 g/100 g (67.7 °C)
|Solubility in ethanol||Trihydrate:|
5.3 g/100 mL
|Solubility in acetone||0.5 g/kg (15 °C)|
|Acidity (pKa)||24 (20 °C)|
4.75 (when mixed with CH3COOH as a buffer)
Refractive index (nD)
Heat capacity (C)
|100.83 J/mol·K (anhydrous)|
229 J/mol·K (trihydrate)
|138.1 J/mol·K (anhydrous)|
262 J/mol·K (trihydrate)
Std enthalpy of
|-709.32 kJ/mol (anhydrous)|
-1604 kJ/mol (trihydrate)
Gibbs free energy (ΔfG?)
|-607.7 kJ/mol (anhydrous)|
|Safety data sheet||External MSDS|
|NFPA 704 (fire diamond)|
|Flash point||>250 °C (482 °F; 523 K) |
|600 °C (1,112 °F; 873 K)|
|Lethal dose or concentration (LD, LC):|
LD50 (median dose)
|3530 mg/kg (oral, rat)|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Sodium acetate is used in the textile industry to neutralize sulfuric acid waste streams and also as a photoresist while using aniline dyes. It is also a pickling agent in chrome tanning and helps to impede vulcanization of chloroprene in synthetic rubber production. In processing cotton for disposable cotton pads, sodium acetate is used to eliminate the buildup of static electricity.
Sodium acetate is used to mitigate water damage to concrete by acting as a concrete sealant, while also being environmentally benign and cheaper than the commonly used epoxy alternative for sealing concrete against water permeation.
Sodium acetate may be added to food as a seasoning, sometimes in the form of sodium diacetate, a one-to-one complex of sodium acetate and acetic acid, given the E-number E262. It is often used to give potato chips a salt and vinegar flavor. Sodium acetate (anhydrous) is widely used as a shelf-life extending agent, pH control agent It is safe to eat at low concentration.
A solution of sodium acetate (a basic salt of acetic acid) and acetic acid can act as a buffer to keep a relatively constant pH level. This is useful especially in biochemical applications where reactions are pH-dependent in a mildly acidic range (pH 4-6).
Sodium acetate is also used in heating pads, hand warmers, and hot ice. Sodium acetate trihydrate crystals melt at 136.4 °F/58 °C (to 137.12 °F/58.4 °C), dissolving in their water of crystallization. When they are heated past the melting point and subsequently allowed to cool, the aqueous solution becomes supersaturated. This solution is capable of cooling to room temperature without forming crystals. By pressing on a metal disc within the heating pad, a nucleation center is formed, causing the solution to crystallize back into solid sodium acetate trihydrate. The bond-forming process of crystallization is exothermic. The latent heat of fusion is about 264-289 kJ/kg. Unlike some types of heat packs, such as those dependent upon irreversible chemical reactions, a sodium acetate heat pack can be easily reused by immersing the pack in boiling water for a few minutes, until the crystals are completely dissolved, and allowing the pack to slowly cool to room temperature.
For laboratory use, sodium acetate is inexpensive and usually purchased instead of being synthesized. It is sometimes produced in a laboratory experiment by the reaction of acetic acid, commonly in the 5-8% solution known as vinegar, with sodium carbonate ("washing soda"), sodium bicarbonate ("baking soda"), or sodium hydroxide ("lye", or "caustic soda"). Any of these reactions produce sodium acetate and water. When a sodium and carbonate ion-containing compound is used as the reactant, the carbonate anion from sodium bicarbonate or carbonate, reacts with hydrogen from the carboxyl group (-COOH) in acetic acid, forming carbonic acid. Carbonic acid readily decomposes under normal conditions into gaseous carbon dioxide and water. This is the reaction taking place in the well-known "volcano" that occurs when the household products, baking soda and vinegar, are combined.
The crystal structure of anhydrous sodium acetate has been described as alternating sodium-carboxylate and methyl group layers. Sodium acetate trihydrate's structure consists of distorted octahedral coordination at sodium. Adjacent octahedra share edges to form one-dimensional chains. Hydrogen bonding in two dimensions between acetate ions and water of hydration links the chains into a three-dimensional network.
|Degree of hydration||Na coordination||Strongly bonded aggregation||Weakly bonded aggregation|
sheets stacked with
hydrophobic surfaces in contact
chains linked by hydrogen bonds
(one chain highlighted in light blue)
Sodium acetate undergoes decarboxylation to form methane (CH4) under forcing conditions (pyrolysis in the presence of sodium hydroxide):