Experiment 1 Ph Measurement And Buffer Preparation 6c1y44

PREPARATION OF BUFFER AND ELECTROMETRIC AND COLORIMETRIC DETERMINATION OF PH Eunice Aurelle T. Basco, Ian Lindley C. Cabral, Aira Mina A. Cayago, Jardine Mariel L. Ching, Leomariss M. Chua and Filjosh R. Cucueco Group 2 2A-Medical Technology Biochemistry Laboratory ABSTRACT Preparation of buffer solution, determination of pH of the buffer and samples electrometrically and colorimetrically were done in the experiment. The standard reagent was prepared by diluting 60g of NaOH pellets. The primary phosphate buffer was prepared by using 1.69 mL of 14.8 M H3PO4 and 0.88g NaOH pellets. To obtain the desired pH of 3.0, the buffer was measured electrometrically using the pH meter and was manipulated by using either 6.0 M HCl or 6.0 M NaOH. The pH of the prepared buffer solution was measured colorimetrically using acid-base indicators such as thymol blue, bromophenol blue, bromocresol green, bromocresol purple, phenol red, methyl red, methyl orange, and phenolphthalein. The sample, distilled water with a pH of 5.7, was subjected to colorometric determination giving the result colors of yellow for thymol blue, violet for bromophenol blue, blue for bromocresol green, yellow for bromocresol purple, yellow for phenol red, yellow for methyl red, orange for methyl orange, and colorless for phenolphthalein.

INTRODUCTION pH is defined as a way of expressing low concentration of hydrogen ions. [1] The term was first described by Danish biochemist Søren Peter Lauritz Sørensen in 1909. [2] pH is an abbreviation for "power of hydrogen" where "p" is short for the German word for power (potenz) and H is the element symbol for hydrogen. The H is capitalized because it is standard to capitalize element symbols. [2] A buffer system/solution is a mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid. [1] Since its components neutralize the excess hydrogen (H+) and hydroxyl ion (OH-), it allows solutions to resist large changes in pH. Therefore, a buffer helps maintain a near constant pH upon the addition of small amounts of H+ or OH- to a solution. [3] For buffers, The Henderson-Hasselbach equation shows the relationship of the pH of a solution to the pK of an acid and the ratio of the concentrations of the acid and its conjugate base. [1] The equation provides a convenient way to think about buffers and pH:

A−¿ ¿ ¿ pH= pKa+ log ¿

The Henderson-Hasselbalch equation is used to determine if an aqueous solution of a conjugate acid/base pair is functioning as a buffer. [3] With these in mind, we prepared different buffer solutions, determined the pH of the buffers and samples colorimetrically (by using liquid indicators) and electrometrically (by using the pH meter), and calculated the buffer capacity of the prepared buffer solutions in this experiment.

EXPERIMENTAL A. Compounds, solvents, and solutions utilized and consumed For the preparation of reagents, the samples used, solvents or compounds tested in the experiment were distilled water (H2O), Sodium hydroxide pellets (NaOH) and 12.2 M concentrated Hydrochloric Acid (HCl). For the preparation of primary phosphate buffer, the samples used, solvents or compounds tested in the experiment were distilled water (H2O), Sodium hydroxide pellets (NaOH) and 14.8 M concentrated Phosphoric Acid (H3PO4). For the electrometric determination of pH, the samples used, solvents or compounds tested in the experiment were distilled water (H2O), Coca Cola soft drink (assigned sample), prepared 250 mL Primary phosphate buffer, prepared 250 mL 6.0 M Hydrochloric acid reagent, and prepared 250 mL 6.0 M Sodium hydroxide reagent. For the colorimetric determination of pH, the samples used, solvents or compounds tested in the experiment were the prepared 250 mL Primary phosphate buffer and the different acidbase indicators: Thymol blue (C27H30O5S), Bromophenol blue (C19H10Br4O5S), Bromocresol green (C21H14Br4O5S), Bromocresol purple (C21H16Br2O5S), Phenol red (C19H14O5S), Methyl red (C15H15N3O2), Methyl orange (C14H14N3NaO3S), and Phenolphthalein (C20H14O4).

B.iProcedure 1. Preparation of Reagents The group was assigned to prepare a 250 mL 6.0 M NaOH aqueous solution. First, the amounts needed to prepare the reagents were computed

using the dilution factor as well as the formula for getting the molar concentration. For the preparation of 250 mL 6.0M NaOH solution, 60 g of Sodium hydroxide (NaOH) pellets were measured using a triple beam balance and was then dissolved with distilled water (H2O) in a 250 mL beaker. The aqueous solution was transferred to a 250 mL volumetric flask were additional distilled water (H2O) was added to fill up to the 250 mL mark. Afterwards, the 250 mL 6.0M NaOH reagent was transferred to a reagent bottle and then labeled properly. Given: NaOH pellets (40 MM) 250 mL 6.0 M NaOH aqueous solution

moles Buffer Solution=( L solution)( M solution ) moles Buffer Solution=(0.250 L)(0.1 M ) moles Buffer Solution=0.0250 moles [conjugate base] pH= pKa+ log [weak acid]

3.00=2.12+ log

[ primary phosphate] [ phosphoric acid]

[ primary phosphate] 7.5858 M = [ phosporic a cid ] 1M

[Buffer solution ]=[ primary phosphate ] +[ phosphoric acid ] [Buffer solution]=7.5858 M+ 1 M moles NaOH =(Molarity NaOH )(Volume NaOH )moles NaOH =( 6.0solution]=8.5858 M ) (0.250 L) [Buffer M moles NaOH =1.5 moles NaOH

Computation:

Moles of Primary Phosphate (H2PO4-1)

mass NaOH=(moles NaOH )(molar mass NaOH ) mass NaOH =(1.5 moles)(40 g /moles) mass NaOH=60 g NaOH 2. Preparation of Buffer Solution The group was assigned to prepare a 250 mL primary phosphate buffer solution (pK=2.12) with a desired pH of 3.00. After identifying the weak acid and conjugate base components of the buffer solution, the amounts needed to prepare the buffer were computed using the HendersonHasselbalch equation, dilution equation, and as well as the formula for getting the molar concentration. For the preparation of 250 mL primary phosphate buffer solution, 0.88 g of Sodium hydroxide (NaOH) pellets was reacted with 1.69 mL of 14.8 M concentrated Phosphoric Acid (H3PO4). It was then diluted with distilled water (H2O) in a 250 mL beaker. Then, the aqueous solution was transferred to a 250 mL volumetric flask were additional distilled water (H2O) was added to fill up to the 250 mL mark. Afterwards, the 250 mL primary phosphate buffer solution was transferred to a reagent bottle and then labeled properly.

[Primary phosphate] moles primary phosphate = [buffer solution] moles buffer solution 7.5858 M moles primary phosphate = 8.5858 M 0.0250 moles Moles Primary phosphate=0.022 moles Moles of Phosphoric Acid

[Phosphoric acid ] moles phosphoric acid = [ buffer solution ] moles buffer solution 1M moles primary phosphate = 8.5858 M 0.0250 moles Moles Phosphoric acid=0.0029moles Volume of Acid Used

Volume Phosphoric acid=

0.0250 moles 14.8 M

Volume Phosphoric acid=0.00169 L∨1.69mL Mass of Base Needed

Given: 14.8 M Concentrated H3PO4 (pK = 2.12) NaOH pellets (40 MM) 250 mL 0.10 M buffer solution Desired pH = 3.00 Weak acid = Phosphoric acid (H3PO4) Conjugate base = Primary Phospate (H2PO4-1)

Mass Sodium hydroxide=(0.022moles)(40 g/moles) Mass Sodium hydroxide=0.88 g 3. Electrometric Determination of pH The pH meter was calibrated first at pH 4, 7, and 10. First, the pH and temperature of the distilled water (H2O) was measured and noted.

The same procedure was done for the Coca Cola soft drink (assigned sample) and prepared Primary phosphate buffer. For each of the pH measured, the [H+] were calculated using the formula 10-pH. Using a pH meter, the pH of the prepared buffer solution was measured. The buffer solution was then manipulated to pH 3.0 (which is the desired pH) using the 6.0M NaOH solution (to make it more basic, increase pH) and 6.0M HCl solution (to make it more acidic, decrease pH). 4. Colorimetric Determination of pH Table 1. Organic Dyes Used for Colorimetric Determination of pH Indicator

pH Range Acidic: 1.2-2.8

Thymol Blue Alkaline: 8.0-9.6 Bromophenol Blue Bromocresol Green Bromocresol Purple

3.0-4.6 3.8-5.4 5.2-6.8

Color Change Acidic: Red to yellow Alkaline: Yellow to blue Yellow to Blueviolet

The preparation of the buffer was done by altering the amount of weak acid and its conjugate base in the solution. The pH of the buffer was obtained by the use of a pH meter. This is a device utilized to measure the expression of the degree of activity of an acid or base according to hydrogen ion [H+] activity or the acidity or alkalinity of a solution - also known as pH [5]. pH is the unit of measure that describes the degree of acidity or alkalinity of the given solution and is measured on a scale of 0 to 14 [5]. In the experiment, if the actual pH was less than the desired pH, the buffer would be treated with an appropriate amount of the standard Sodium hydroxide (NaOH). However, if the actual pH was greater than the desired pH, the buffer would be treated with a small amount of the standard Hydrochloric acid (HCl).

Yellow to Blue Yellow to Purple

Phenol Red

6.4-8.2

Methyl Red Methyl Orange Phenolphthalei n

4.8-6.0 3.2-4.4

Yellow to Redviolet Red to yellow Red to yellow

8.2-10.0

Colorless to pink

a. Preparation of Color Standards Using the Buffer Solutions A certain number of vials (8) was prepared and properly labeled with their respective acid-base indicators. Using a serological pipette and an aspirator, each vial was filled with 5mL of the prepared primary phosphate buffer solution. A certain amount (2 drops) of acid-base indicator was dropped in the corresponding labeled vials. The vials were shaken and the color was noted down. b. Determination of the pH of samples The sample used in this experiment was distilled water (H2O). A certain number of vials (8) was prepared and properly labeled with their respective acid-base indicators. Using a serological pipette and an aspirator, each vial was filled with 5mL of distilled water (H2O). A certain amount (2 drops) of acid-base indicator was dropped in the corresponding labeledvials. The vials were shaken and the color was noted down.

RESULTS AND DISCUSSION 1. Electrometric Determination of pH

Figure 1. A pH meter connected combination pH electrode.

to

a

2. Colorimetric Determination of pH Colorimetry is a technique that is commonly used in biochemistry. This process makes use of a device called colorimeter that measures the absorbance of specific wavelengths of light by a specific solution and thus, involves the quantitative estimation of colors [4] [6]. The produced difference in color results in the variation in the absorption of light, which is made use of here in this technique. Moreover, this method is usually applied to determine the concentration of a known solute in a given solution by the adaptation of the Beer-Lambert law [4]. The colorimeter is based on Beer-Lambert's law, according to which the absorption of light transmitted through the medium is directly proportional to the medium concentration. The results depicted in Table 1 shows that at pH=8, the acid-base indicators, when mixed with the solution, yielded their corresponding colors and those colors are due to the concentration of the

colored complex that was formed. The color or the amount of light absorbed is also related to the concentration of the absorbing compound, therefore it is one of the most useful techniques in obtaining the pH. According to Beer’s law when monochromatic light es through the colored solution, the amount of light transmitted decreases exponentially with increase in concentration of the colored substance [6]. In addition, the Lambert’s law states that the amount of light transmitted decreases exponentially with increase in thickness of the colored solution [6]. Colorimetric determination of pH exhibits the varying color changes an acid-base indicator undergoes when added to a solution of a certain pH. This property is used to distinguish different substances by adjusting their pH ranges.

Figure 1. Indicators

Distilled water with Acid-Base

Figure 2. Buffer solution with AcidBase Indicators Table 2. Results of Colorometric

Colorimetric Analysis use the acquired various Determination of pH (R = Red, O= colors as a means of determining the pH since Orange, G= Green, Y = Yellow, b = Blue, the intensity of the color of a solution changes P=Pink, V= Violet, C= Colorless, YO= with its concentration or pH. The color may be Yellow Orange, OY= Orange Yellow, due to inherent property of a substance in the solution or due to the formation of a product as a GB= Green Blue, BV= Blue Violet, VB= result of the addition of a suitable reagent or Violet Blue) (Tb = Thymol Blue, Bb= acid-base indicator. The Acid pH Color Standard Bromophenol pH of a solution can be Base 2 3 5 7 7.5 8 12 D. blue, Bg= determined by Indic. 2.2 3 4.6 6.9 7.4 7.5 7.9 12.1 H2O Bromocresol comparing the color green, Bp = Tb O YO Y Y Y Y Y B Y intensities of the Bromocresol Bb G Y B VB V V V V V solution with unknown purple, Pr = Bg Y Y GB B B B B B B pH with the intensities Phenol Red, Mr Bp Y Y Y BV V V V V Y of the solutions with Pr Y Y Y YO R R P P Y known pH. = Methyl Red, Mr P P R Y Y Y Y Y Y Mo= Methyl Table 1. List of AcidMo R R OY YO O O O O O Orange, Pp = Base Indicators with Pp C C C C C C C P C their corresponding Phenolphthalein) colors Acid - Base Indicator

pH 8.0

Thymol blue

Yellow

Bromophenol blue

Purple

Bromocresol green

Blue

Bromocresol purple

Purple

Phenol red

Red

Methyl red

Yellow

Methyl orange

Orange

Phenolphthalein

Colorless

REFERENCES [1] Crisostomo, A. D., Daya, M. L., de Guia, R. M., Farrow, F. L., Gabona, M. G.,… Ysrael, M.C. (2009). Laboratory Manual in General Biochemistry. Ph Measurement and buffer preparation, 1, 1. [2] Anne Marie Helmenstine, Ph. D. (2015, November 13). What Does pH stand for? Retrieved from http://chemistry.about.com/od/ph/f/What-DoesPh-Stand-For.htm [3] Chemistry Department-Xavier University of St. Louisiana (2014, December 3). Preparation of buffers at a desired pH. Retrieved from

www.xula.edu_Chemistry_documents_biolab_Car oll...(14-12-03pdf).pdf [4] Chhabra, N. (September, 2015). ColorimetrySimplified. Retrieved from http://www.namrata.co/colorimetry-simplified/ [5] No author (2003). pH Meter. Retrieved from http://www.omega.co.uk/prodinfo/ph-meter.html [6] No author (June, 2011). Principles of Colorimetry. Retrieved from

https://ecoplants.wordpress.com/2011/06/23/pri nciples-of-colorimetry/