Acid salt

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Acid salts are a class of salts that produce an acidic solution after being dissolved in a solvent. Its formation as a substance has a greater electrical conductivity than that of the pure solvent.[1] An acidic solution formed by acid salt is made during partial neutralization of diprotic or polyprotic acids. A half-neutralization occurs due to the remaining of replaceable hydrogen atoms from the partial dissociation of weak acids that have not been reacted with hydroxide ions (OH) to create water molecules. Acid salt is an ionic compound consisted of an anion, contributed from a weak parent acid, and a cation, contributed from a strong parent base.

Acidic solution and examples of acid salts[edit]

Structure of ammonium chloride

Acid-base property of the resulting solution from a neutralization reaction depends on the remaining salt products. A salt containing reactive cations undergo hydrolysis by which they react with water molecules, causing deprotonation of the conjugate acids.
For example, the acid salt ammonium chloride is the main species formed upon the half neutralization of ammonia in hydrochloric acid solution:[2]

NH3 (aq) + HCl (aq) → NH4Cl (aq)


Example of acid salts[3]
Identity: Sodium bicarbonate Sodium bisulfate Monosodium phosphate Disodium phosphate
Structural formula
SodiumBicarbonate.png
Sodium bisulfate.png
Monosodium phosphate.png
Disodium hydrogen phosphate.png
Chemical formula NaHCO
3
NaHSO
4
NaH
2
PO
4
Na
2
HPO
4
IUPAC Name Sodium hydrogen carbonate Sodium hydrogen sulfate Sodium dihydrogen phosphate Disodium hydrogen phosphate
Other name
  • Baking Soda
  • Carbonic Acid Monosodium Salt
  • Sodium acid sulfate
  • Bisulfate of soda
  • Monobasic sodium phosphate
  • Sodium acid phosphate
  • Sodium biphosphate
  • Disodium hydrogen orthophosphate
  • Sodium phosphate dibasic
  • disodium phosphate
Molecular Weight 84.006 g/mol 120.054 g/mol 119.976 g/mol 141.957 g/mol
Formal Charge zero zero Zero Zero
Odour Odourless[4] Odourless Odourless Odourless[5]
Appearance White crystalline powder or lumps[6] White crystals or granules White crystalline powder[7] White, hygroscopic powder[8]
Structure Monoclinic[9]
  • triclinic (anhydrous)
  • monoclinic (monohydrate)
Monoclinic crystals[10] Monoclinic crystals (anhydrous)[11]
Solubility
  • Soluble in water
  • Insoluble in ammonia
  • Soluble in water
  • Insoluble in ethanol or ether
  • Soluble in water
  • Insoluble in ethanol
Density 2.1 g/cm3
  • 2.742 g/cm3 (anhydrous)
  • 1.8 g/cm3 (monohydrate)
0.5-1.2 g/cm3 1.7 g/cm3
Decomposition

(through heating)

Emits acrid smoke, fumes, and carbon dioxide[12] Forms sodium carbonate, water, and carbon dioxide Emits toxic fumes of phosphoxides and sodium oxide[13] Emits toxic fumes of phosphorus- and sodium oxides[13]
Uses
  • Drug indication; treatment of metabolic acidosis
  • Automotive Care Products
  • Food additives
  • Bleaching agents
  • Plating agents and surface treating agents
  • Cleaning and Furnishing Care Products
  • Treat constipation
  • Clean the bowel before a colonoscopy
  • Bleaching agents
  • A source of phosphorus
  • Visicol tablets are indicated for cleansing of the colon
  • Corrosion inhibitors and anti-scaling agents

Use in food[edit]

Some acid salts such as sodium bicarbonate, NaHCO3 are used in baking. They are found in baking soda, bread soda or cooking soda and are typically divided into low-temperature (or single-acting) and high-temperature (or double-acting) acid salts. Common low-temperature acid salts react at room temperature to produce a leavening effect. They include cream of tartar, calcium phosphate, and citrates. High-temperature acid salts produce a leavening effect during baking and are usually aluminium salts such as calcium aluminium phosphate. Some acid salts may also be found in non-dairy coffee creamers. And also disodium phosphate, Na2HPO4 is used in foods and monosodium phosphate, NaH2PO4 is used in animal feed, toothpaste and evaporated milk.

Intensity of acid[edit]

An acid with higher Ka value dominates the chemical reaction. It serves as a better contributor of proton (H+). A comparison between the Ka and Kb indicates the acid-base property of the resulting solution by which:

  1. The solution is acidic if Ka > Kb. It contains a greater concentration of H+ ions than concentration of OH ions due more extensive of cation hydrolysis compared to that of anion hydrolysis.
  2. The solution is alkali if Ka < Kb. Anions hydrolyze more than cations, causing an exceeding concentration of OH ions.
  3. The solution is expected to be neutral only if Ka and Kb are identical.[14]

Other possible factors that could vary pH level of a solution are the relevant equilibrium constants and the additional amounts of any base or acid.

For example, in ammonium chloride solution, NH4+ is the main influence for acidic solution. It has greater Ka value compared to that of water molecules; Ka of NH4+ is 5.6 x 10−10 and Kw of H2O is 1.0 x 10−14. This ensures its deprotonation when reacting with water, and is responsible for the ph below 7 at room temperature. Cl will have no affinity for H+ nor tendency to hydrolyze, as its Kb value is very low (Kb of Cl is 7.7 x 10−21).[15]
Hydrolysis of ammonium at room temperature produces:
NH4+ (aq) + H2O (aq) ⇌ NH3 (aq) + H3O+ (aq)
= 5.6 x 10−10

See also[edit]

References[edit]

  1. ^ Cady, H. P.; Elsey, H. M. (1928). "A general definition of acids, bases, and salts". Journal of Chemical Education. 5 (11): 1425. Bibcode:1928JChEd...5.1425C. doi:10.1021/ed005p1425. Retrieved 27 February 2018.
  2. ^ Dekock, Roger L.; Gray, Harry B. (1989). Chemical bonding and structure (Second ed.). Sausalito, California: University Science Book. pp. 97–98. ISBN 0-935702-61-X. Retrieved 8 February 2018.
  3. ^ "Sodium Bicarbonate". PubChem Compound Database. National Center for Biotechnology Information. Retrieved 2 March 2018.
  4. ^ Osol, A. and J.E. Hoover, et al. (eds.). Remington's Pharmaceutical Sciences. 15th ed. Easton, Pennsylvania: Mack Publishing Co., 1975., p. 736
  5. ^ U.S. Coast Guard, Department of Transportation. CHRIS - Hazardous Chemical Data. Volume II. Washington, D.C.: U.S. Government Printing Office, 1984-5.
  6. ^ O'Neil, M.J. (ed.). The Merck Index - An Encyclopedia of Chemicals, Drugs, and Biologicals. 13th Edition, Whitehouse Station, NJ: Merck and Co., Inc., 2001., p. 1536
  7. ^ Lewis, R.J. Sr.; Hawley's Condensed Chemical Dictionary 15th Edition. John Wiley & Sons, Inc. New York, NY 2007., p. 1153
  8. ^ Lide, D.R. CRC Handbook of Chemistry and Physics 88TH Edition 2007-2008. CRC Press, Taylor & Francis, Boca Raton, Florida 2007, p. 4-90
  9. ^ a b Raton, Boca (2005). CRC Handbook of Chemistry and Physics (86th Edition Edited by David R. Lide ed.). CRC Press. pp. 4–87. ISBN 0-8493-0486-5.
  10. ^ Haynes, W.M. (ed.). CRC Handbook of Chemistry and Physics. 95th Edition. CRC Press LLC, Boca Raton: FL 2014-2015, p. 4-89
  11. ^ Somov, N.V.; Chausov, F.F.; Russ, J. (2017). "High-symmetry polymorph of anhydrous disodium hydrogen phosphate". Russian Journal of Inorganic Chemistry. 62 (2): 172–174. doi:10.1134/S0036023617020176. Retrieved 2 March 2018.
  12. ^ Sax, N.I. Dangerous Properties of Industrial Materials. 6th ed. New York, NY: Van Nostrand Reinhold, 1984., p. 2413
  13. ^ a b Wiley, John; Hoboken, NJ (2004). Sax's Dangerous Properties of Industrial Materials (11th Edition By Richard J. Lewis ed.). Wiley-Interscience. p. 3274. ISBN 0-471-47662-5. Retrieved 4 March 2018.
  14. ^ Raymond, Chang (2010). Chemistry (tenth ed.). Americas, New York: McGraw-Hill. pp. 725–727. ISBN 0077274318. Retrieved 9 February 2018.
  15. ^ Lower, S.K., (1999). Introduction to acid-base chemistry. Chem1 Genneral Chemistry Text. Retrieved from http://www.chem1.com/acad/pdf/c1xacid1.pdf