Introduction to Buffer Solutions

 

Name of Candidate

mayurathan

Introduction

Buffer solution is a solution consisting a mixture of weak acid and its conjugate base . Furthermore they declare that buffer solutions are very essential to keep the pH value nearly constant in variety of chemical application. ( Crowe and Bradshaw 2010)

Many life forms thrives only in a relatively small pH range so they utilize a buffer solution to maintain a constant pH, one common example of a buffer solution found in nature is blood. buffer solution is very vital t keep the pH at a constant value in variety of enzymes in many organisms to work this is due to different enzymes work at different précised pH. On the same way they state that if the pH range moves above or below the range, the enzymatic action either stops or can denature where there are many cases denaturation can permanently disable their catalytic activity. ( Kotz, Treichel and Townsend ,2009)

Buffer of carbonic acid and bicarbonate is present in blood plasma maintain pH between 7.35 to 7.45.Over more they clarify that there are two main types of buffer system,

• Acid buffer system: the solution containing large amount of weak acid and its salt with strong base is termed as acidic buffer solution.

pH = pKa + log [salt] / [acid]

• Basic buffer system: the solution containing large amount of weak base and its salt with strong acid is termed as basic buffer solution.

pOH = pKb + log [salt] / [base]

(Moore, Stanitski and Jurs (2009)

In living organisms including human the important buffer solution to maintain the ph of the blood is bicarbonate buffering system. On the other hand they state that this bicarbonate buffer system tends to maintain relatively constant plasma pH, where carbon dioxide combines with water to form carbonic acids which in turn rapidly dissociate to form hydrogen ions and bicarbonate ions.( Lanham et al ,2011)

CO₂ + H₂O H₂CO₃ HCO₃^– + H⁺

Blood is dumped by excess hydrogen ions, some of those hydrogen ions associate with bicarbonate forming carbonic acid results in increase in acidity of the blood. As well as they elaborate that due to this incident the buffering system becomes powerful regulator of acidity by accompanying with respiratory compensation in which breathing is altered to modify the carbon dioxide in circulation which results in increase in ventilation therefore increase the loss of carbon dioxide into the atmosphere. (Rhoades and Bell (2012),

pH meter is a electronic device that is used to measure the ph of the solution, where a typical pH meter consists of a special measuring probe connected to an electronic meter that measures and display the reading. (Kenkel,2013)

Objectives

• To know how to prepare buffer solution.
• To practice again how to prepare standard solution.
• To know how to measure the pH using pH meter.

Materials Used

• Electronic balance
• Pipette
• Measuring cylinder
• Watch glass
• Beaker
• pH meter
• Volumetric flask (100ml)
• NaHPO₄ powder (0.7 g)
• Na₂HPO₄ powder (1.56 g)
• Sodium hydroxide pellets
• Ammonium chloride powder
• Glass rod
• Spatula

Methodology

At pH 7

NAHPO₄, 0.70g and Na₂HPO₄, 1.56g was taken and measured using the electronic balance. Then both NAHPO₄ and Na₂HPO₄ were mixed together into the volumetric flask. First half of the volumetric flask was filled with distilled water and dissolved by inverting. Later on the flask was filled with distilled water till it reaches the meniscus level. Finally the pH was measured.

At pH 10

• NH₄OH NH₄⁺ + OH; pKb = 4.74
• pOH = pKb + log [NH₄⁺] / [NH₃]
• pKw = pH + pOH

14 = 10 + pOH

pOH = 4

• pOH = pKb + log [NH₄⁺] / [NH₃]

4 = 4.74 + log [NH₄⁺] / [NH₃]

• [NH₄⁺] / [NH₃] = 10 -0.74 = 0.181

NH4CL, 0.1m standard solution was made. Afterwards NaOH, 0.34g was measured using the electronic balance. NaOH of 0.34g was obtained by,

• NH₄⁺ + OH- NH3 + H2O
  1. x

0.1 – x 0 x

• [NH₄⁺] / [NH₃] = 0.181
• X = 0.085M of OH-
• Number of mole of NaOH = 0.085 x (100 / 1000)

= 0.0085 mols

• Mass of NaOH = 0.0085 x 40

= 0.34g

Then afterwards 0.34g of NaOH was dissolved in 0.1 M of NH₄Cl to form a buffer solution.

NaOH, 0.1m standard solution was made. Afterwards NaOH, 0.34g was measured using the electronic balance. NH₄Cl of 0.63g was obtained by,

• NH₄⁺ + OH- NH3 + H2O

x 0.1

x – 0.1 0 0.1

• [NH₄⁺] / [NH₃] = 0.181
• X = 0.1181M of NH₄⁺
• Number of mole of NH₄Cl = 0.1181x (100 / 1000)

= 0.01181 mols

• Mass of NH₄Cl = 0.01181 x 35

= 0.63g

Then afterwards 0.63g of NH4Cl was dissolved in 0.1 M of NaOH to form a buffer solution. Later on using the pH meter the final ph of each buffer solution was measured approximately.

Results

Buffer solution	pH
HPO42- / H2po4–	7.35
NH4Cl 0.1M / NaOH	10.54
NaOH 0.1M / NH4Cl	10.76

Discussion / Conclusion

Every precise work the pH meter should be calibrated before each measurement, at every experiment the calibration is done because the glass rode does not give reproducible electro motive force over longer periods of time. Additionally further he states that pH meter calibration should be performed with at least two standard buffer solutions that span the range of pH values to be measured. (Hauser 2001)

Single measurement the probe should be rinsed with distilled water to remove any traces of solution being measured and then it should be blotted using the scientific wipe to absorb any remaining water which could dilute the sample which alters the reading. (Prichard 2003)

Reference

Crowe, J and Bradshaw, T (2010). Chemistry of biosciences: the essential concepts. Google Books [Online].Availableat:http://books.google.lk/books?id=onacAQAAQBAJ&pg=PA578&dq=buffer+solution&hl=en&sa=X&ei=UkOuU7OtJJO78gXkvIDYDQ&redir_esc=y#v=onepage&q=buffer%20solution&f=false(Accessed: 28 Aug 2014)

Kotz, J., Treichel, P and Townsend, J (2009). Chemistry and chemistry reactivity, enhanced edition. Google Books [Online].

Available at: http://books.google.lk/books?id=IBESYmQcb0sC&pg=PA851&dq=buffer+solution&hl=en&sa=X&ei=u0CuU4PYBYS78gXVpoGIAw&redir_esc=y#v=onepage&q=buffer%20solution&f=false

(Accessed: 28 aug 2014)

Moore, J., Stanitski, C and Jurs, P (2009). Principles of chemistry: the molecular science. Googke Books [Online].

Available at: http://books.google.lk/books?id=ZOm8L9oCwLMC&pg=PA575&dq=buffer+solution&hl=en&sa=X&ei=u0CuU4PYBYS78gXVpoGIAw&redir_esc=y#v=onepage&q=buffer%20solution&f=false

(Accessed: 27 aug 2014)

Lanham, S.A., Stear,S., Shirreffs,S and Colins, A (2011). Sports and exercise nutrition. Google Books [online].

Available at: http://books.google.lk/books?id=YePJM98Np5MC&pg=RA1PT115&dq=blood+acting+as+buffer&hl=en&sa=X&ei=E0KuU8ypKYb_8QWLh4G4Aw&redir_esc=y#v=onepage&q=blood%20acting%20as%20buffer&f=false

(Accessed: 28 aug 2014)

Rhoades, R.A and Bell, D.R (2012). Medical physiology: principles of clinical medicine. Google Books [Online].

Available at: http://books.google.lk/books?id=1kGcFOKCUzkC&pg=PA457&dq=blood+acting+as+buffer&hl=en&sa=X&ei=E0KuU8ypKYb_8QWLh4G4Aw&redir_esc=y#v=onepage&q=blood%20acting%20as%20buffer&f=false

(Accessed: 28 June 2014)

Kenkel, J (2013). Analytical chemistry for technicians fourth edition. Google books [Online].

Available at: http://books.google.lk/books?id=JZAAAAAAQBAJ&pg=PA400&dq=ph+meter&hl=en&sa=X&ei=0LOuU8T0Os7r8AWY0oHwAg&redir_esc=y#v=onepage&q=ph%20meter&f=false

(Accessed: 27aug 2014)

Hauser, B (2001). Drinking water chemistry: laboratory manual. Google Books [Online].

Available at:

http://books.google.lk/books?id=SVxcRu68YGwC&pg=PA28&dq=calibration+of+ph+meter&hl=en&sa=X&ei=lQ2wU_GhI8yB8gWx9oDAAw&redir_esc=y#v=onepage&q=calibration%20of%20ph%20meter&f=false

(Accessed: 26 aug 2014)

Prichard, E (2003). Measurement of pH. Google books [Online].

Available at: http://books.google.lk/books?id=HNJy5rtJLjAC&pg=PA2&dq=ph+meter&hl=en&sa=X&ei=0LOuU8T0Os7r8AWY0oHwAg&redir_esc=y#v=onepage&q=ph%20meter&f=false

(Accessed: 28 aug 2014)