Analysis of Soda Ash and Carbonate-Bicarbonate Mixture
Analysis of Soda Ash and Carbonate-Bicarbonate Mixture
Submitted: February 27, 2013
Department of Chemical Engineering, Faculty of Engineering
University of Santo Tomas
España, Manila
Abstract
A standard acid solution like HCl can be used as titrant for the analysis of both soda ash and a carbonate-bicarbonate mixture. In the analysis of soda ash, the volume needed to neutralize the soda ash is used to compute for its alkalinity, in this experiment we obtained a 17.6 % alkalinity with an error of 15.14% In the analysis of a carbonate-bicarbonate mixture two indicators (phenolphthalein and methyl orange) were used. The first endpoint determines the half-neutralization of the carbonate and the second determines that of the bicarbonate. The experiment results to an analysis of 4.92% carbonate and 5.07% bicarbonate content in the unknown sample.
Introduction Soda ash or sodium carbonate is a sodium salt of carbonic acid. It is commonly known for its everyday use as water softener.[1] It can be naturally extracted from plants of synthetically produced from large amounts of sodium. It may contain small to moderate amounts of chlorides and hydroxides as impurities. The hydroxide present in sodium carbonate reacts with an acid titrant like HCl and its total alkaline strength is increased. Titration of soda ash with a standard acid solution with methyl orange as indicator, neutralizes its carbonate ions. The usual endpoint of this titration is at pH 4.
A standard HCl can also be used to neutralize a carbonate-bicarbonate mixture through the use of the double indicator method. A bicarbonate is an immediate form of deprotonation of carbonic acid while carbonates as mentioned above is a salt from carbonic acid.[2] In the analysis of a carbonate- mixture, there are two endpoints, first is the phenolphthalein endpoint which is the volume needed to half-neutralize the carbonate content. The second endpoint uses another indicator, the methyl orange, when the solution
Submitted: February 27, 2013
Department of Chemical Engineering, Faculty of Engineering
University of Santo Tomas
España, Manila
Abstract
A standard acid solution like HCl can be used as titrant for the analysis of both soda ash and a carbonate-bicarbonate mixture. In the analysis of soda ash, the volume needed to neutralize the soda ash is used to compute for its alkalinity, in this experiment we obtained a 17.6 % alkalinity with an error of 15.14% In the analysis of a carbonate-bicarbonate mixture two indicators (phenolphthalein and methyl orange) were used. The first endpoint determines the half-neutralization of the carbonate and the second determines that of the bicarbonate. The experiment results to an analysis of 4.92% carbonate and 5.07% bicarbonate content in the unknown sample.
Introduction Soda ash or sodium carbonate is a sodium salt of carbonic acid. It is commonly known for its everyday use as water softener.[1] It can be naturally extracted from plants of synthetically produced from large amounts of sodium. It may contain small to moderate amounts of chlorides and hydroxides as impurities. The hydroxide present in sodium carbonate reacts with an acid titrant like HCl and its total alkaline strength is increased. Titration of soda ash with a standard acid solution with methyl orange as indicator, neutralizes its carbonate ions. The usual endpoint of this titration is at pH 4.
A standard HCl can also be used to neutralize a carbonate-bicarbonate mixture through the use of the double indicator method. A bicarbonate is an immediate form of deprotonation of carbonic acid while carbonates as mentioned above is a salt from carbonic acid.[2] In the analysis of a carbonate- mixture, there are two endpoints, first is the phenolphthalein endpoint which is the volume needed to half-neutralize the carbonate content. The second endpoint uses another indicator, the methyl orange, when the solution
References: Zumdahl, Steven S. (2009). Chemical Principles 6th Ed.. Houghton Mifflin Company. p. A23 “Clinical correlates of pH levels: bicarbonate as a buffer". Biology.arizona.edu. October 2006.