Sodium acetate
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| Names | |
|---|---|
| IUPAC name
Sodium acetate
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| Systematic IUPAC name
Sodium ethanoate | |
| Other names
Hot ice (Sodium acetate trihydrate)
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| Identifiers | |
3D model (JSmol)
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| ChEBI | |
| ChEMBL | |
| ChemSpider | |
| ECHA InfoCard | 100.004.386 |
| EC Number |
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| E number | E262 (preservatives) |
PubChem CID
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| RTECS number |
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| UNII | |
CompTox Dashboard (EPA)
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| Properties | |
| C2H3NaO2 | |
| Molar mass | 82.034 g·mol−1 |
| Appearance | White deliquescent powder |
| Odor | Vinegar (acetic acid) odor when heated to decomposition[1] |
| Density | 1.528 g/cm3 (20 °C, anhydrous) 1.45 g/cm3 (20 °C, trihydrate)[2] |
| Melting point | 324 °C (615 °F; 597 K) (anhydrous) 58 °C (136 °F; 331 K) (trihydrate) |
| Boiling point | 881.4 °C (1,618.5 °F; 1,154.5 K) (anhydrous) 122 °C (252 °F; 395 K) (trihydrate) decomposes |
| Trihydrate: 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) Anhydrous: 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)[3] | |
| Solubility | Soluble in alcohol, hydrazine, SO2[4] |
| Solubility in methanol | 16 g/100 g (15 °C) 16.55 g/100 g (67.7 °C)[4] |
| Solubility in ethanol | Trihydrate: 5.3 g/100 mL |
| Solubility in acetone | 0.5 g/kg (15 °C)[4] |
| Acidity (pKa) | 24 (20 °C)[4] 4.75 CH3COOH[5] |
| Basicity (pKb) | 9.25 |
| −37.6·10−6 cm3/mol | |
Refractive index (nD)
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1.464 |
| Structure | |
| Monoclinic | |
| Thermochemistry | |
Heat capacity (C)
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100.83 J/mol·K (anhydrous)[6] 229 J/mol·K (trihydrate)[7] |
Std molar
entropy (S⦵298) |
138.1 J/mol·K (anhydrous)[6] 262 J/mol·K (trihydrate)[2] |
Std enthalpy of
formation (ΔfH⦵298) |
−709.32 kJ/mol (anhydrous)[4] −1604 kJ/mol (trihydrate)[2] |
Gibbs free energy (ΔfG⦵)
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−607.7 kJ/mol (anhydrous)[4] |
| Pharmacology | |
| B05XA08 (WHO) | |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards
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Irritant |
| NFPA 704 (fire diamond) | |
| Flash point | >250 °C (482 °F; 523 K) [5] |
| 600 °C (1,112 °F; 873 K)[5] | |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
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3530 mg/kg (oral, rat) |
| Safety data sheet (SDS) | External MSDS |
| Related compounds | |
Other anions
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Sodium formate Sodium propionate |
Other cations
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Potassium acetate Calcium acetate |
Related compounds
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Sodium diacetate |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Sodium acetate, CH3COONa, also abbreviated NaOAc,[8] also known as sodium ethanoate, is the sodium salt of acetic acid. This colorless deliquescent salt has a wide range of uses.
Applications
Industrial
Sodium ethanoate 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.
Concrete longevity
Sodium ethanoate 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.[9]
Food
Sodium ethanoate 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,[10] given the E-number E262. It is often used to give potato chips a salt and vinegar flavor.
Buffer solution
As the conjugate base of acetic acid, a solution of sodium acetate 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).
Heating pad
Sodium acetate is also used in heating pads, hand warmers, and hot ice. Sodium acetate trihydrate crystals melt at 136.4 °F/58 °C[11] (to 137.12 °F/58.4 °C),[12] 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.[13][14] The latent heat of fusion is about 264–289 kJ/kg.[11] 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.[15]
Preparation
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.
- CH3COOH + NaHCO3 → CH3COONa + H2CO
3 - H2CO
3 → CO
2 + H
2O
Industrially, sodium acetate trihydrate is prepared by reacting acetic acid with sodium hydroxide using water as the solvent.
- CH3COOH + NaOH → CH3COONa + H2O
Reactions
Sodium acetate can be used to form an ester with an alkyl halide such as bromoethane:
- CH3COONa + BrCH2CH3 → CH3COOCH2CH3 + NaBr
Caesium salts catalyze this reaction.
References
- ^ "Sodium Acetate". International Chemical Safety Cards. National Institute of Occupational Safety and Health.
- ^ a b c "sodium acetate trihydrate". chemister.ru.
- ^ Seidell, Atherton; Linke, William F. (1952). Solubilities of Inorganic and Organic Compounds. Van Nostrand.
- ^ a b c d e f "sodium acetate". chemister.ru.
- ^ a b c Sigma-Aldrich Co., Sodium acetate. Retrieved on 2014-06-07.
- ^ a b Acetic acid, sodium salt in Linstrom, Peter J.; Mallard, William G. (eds.); NIST Chemistry WebBook, NIST Standard Reference Database Number 69, National Institute of Standards and Technology, Gaithersburg (MD) (retrieved 2014-05-25)
- ^ Acetic acid, sodium salt, hydrate (1:1:3) in Linstrom, Peter J.; Mallard, William G. (eds.); NIST Chemistry WebBook, NIST Standard Reference Database Number 69, National Institute of Standards and Technology, Gaithersburg (MD) (retrieved 2014-05-25)
- ^ Clayden, Jonathan; Greeves, Nick; Warren, Stuart; Wothers, Peter (2001). Organic Chemistry (1st ed.). Oxford University Press. ISBN 978-0-19-850346-0.
- ^ "Potato Chip Flavoring Boosts Longevity Of Concrete". Science Daily. 8 August 2007.
- ^ AG, Jungbunzlauer Suisse. "Sodium Diacetate - Jungbunzlauer". www.jungbunzlauer.com.
- ^ a b Ibrahim Dincer and Marc A. Rosen. Thermal Energy Storage: Systems and Applications, page 155
- ^ Courty JM, Kierlik E, Les chaufferettes chimiques, Pour la Science, décembre 2008, p 108-110
- ^ "Crystallization of Supersaturated Sodium Acetate". Journal of Chemical Education.
- ^ "Fake latent heat and supersaturation]". www.atmos.washington.edu.
- ^ "How do sodium acetate heat pads work?". HowStuffWorks. Retrieved 2007-09-03.
External links
