course chemistry laboratory

0H measurement of the reaction hydrochloric acid HCI + distilled water + sodium hydroxide NAOH.... The objectives of the laboratory The lab objectives include differentiating between wea

EXPERIMENT 2:

pH and BUFFERS

By Group 1 For

Dr Doan Hoai Linh

Course: Chemistry Laboratory

Group 1

Bui Thi Nhat Linh BEBEIU22246

Vũ Thị Ngân Hà IELSIU22242 Lương Quốc Duy BTFTIU22188

Phạm Tran Quang Minh ITITIU22103

Huynh Van Quang Chién ITCSIU22022

Experiment performed: 19/03/2024 Submitted: 26/03/2024

1 affirm that I have carefully proofread each report section, and that each satisfies all applicable

criteria listed on the report checklist

TABLE OF CONTENTS

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ConcÏUSIOH c0 21 21111111121 1111211 1211111111111 11111151 11111 T1 HH kh kh HH HH HH Hy 17

List of Table

Table 1 0H measurement of the reaction hydrochloric acid (HCI) + distilled water + sodium hydroxide (NAOH) 0 cece cee ete eee cee nh nhe hà Ha ty na Ha sex se se ssseseexsee TỔ Table 2 pH in experiment and theory c.c cọ LH nàn He nh HH khe hs se se các 12 Table 3 pH 1n expertment and theory about 3 types of salts solution Ì3 Table 4 The acid, base, calculated and measured pH table of volume (mL) 14 Table 5 The changed pH after add HCL table of buffer A, B TỔ Table 6 The changed pH after add NaOH table of buffer A, B 17

List of Figure

Figure 1 The portable pH meter 2.0 cence cee ences cet eee teers ceneaneeeeerenseeeee D Figure 2 pH concentration over different từne perIods 8 Figure 3 pH measured by pH meter at (1) the initial time, (2) the second time, (3) the third time, (4) the fourth time, (5) the fifth trne, (6) the sixth từme - - - 9 Figure 4 pH of experiment 2 measured by 2 pH meters at (1) the initial time, (2) the second time, (3) the third time, (4) the fourth time cece ct eee crete treet tettteterteeeee LL Figure 5 The pH of mixed 10mL 0.1M CH3COOH and 40mL 0.1M CH3:COONa experiment measured at (1) the initial time, PO the second time, a the third time, a the fourth time, ©) me fifth time, (6) the sixth time

Figure 6 pH recorded on pH meter of experiment 3 „ 14 Figure 7 pH recorded on pH meter of experiment 4 cc eee ee TS Figure 8 The pH bar chart of mixed 10mL 0.1M CH3COOH and 40mL 0.1M CH3COONa following each periods 6c cece ee cee cence aáaA eee LS Figure 9 The pH of experiment mixed 40mL 0.1M CH3COOH and 10mL 0.1M CH3COONa measured by pH meter at (1) the initial time, Nt the second time, " the third time, a the foun time, (5) the fifth time, (6) the sixth time

Figure 10 The pH bar chart of mixed 40mL 0.1M CH3COOH and 10mL 0.1M CH3COONa following each periods 0 cece cee ce cee cence cents cen nee ttettaee tances eae teens tesaee esse LT

Assignments

1 Bui Thi Nhat Linh BEBEIU22246 100%

2 Vũ Thị Ngân Hà IELSIU22242 100%

3 Lương Quốc Duy BTFTIU22188 100%

4 Pham Tran Quang Minh ITITIU22103 100%

5 Huynh Van Quang Chién ITCSIU22022 100%

I Introduction

1.1 The objectives of the laboratory

The lab objectives include differentiating between weak and strong acids, learning how to prepare buffer solutions, comprehending the capacity of buffers to withstand significant changes

in pH, and becoming proficient with a pH meter Students will distinguish between mild and strong acids by performing acid-base titrations and pH analyses In addition, they will acquire practical experience in preparing buffer solutions and observe how buffers maintain a constant pH In addition, students will learn how to effectively use a pH meter, including calibration and accurate

pH measurements The purpose of these objectives is to provide students with a solid foundation

in acid classification, buffer preparation, and pH meter use

1.2 The description of strong acids and strong bases

In 1887, the Swedish physicist Svante Arrhenius devised the Arrhenius hypothesis, which states that acids dissolve in water into electrically charged atoms or molecules known as ions, one

of which is a hydrogen ion (H"), whereas bases ionize in water into hydroxide ions (OH’)

It is now understood that the hydrogen ion cannot exist alone in water; rather, it exists as the hydronium ion (H3;0°) in combination with a water molecule In common parlance, the

hydronium ion is still known as the hydrogen ion LH,

In chemistry, a base is a substance that donates electrons, accepts protons, or releases hydroxide ions (OH’) when dissolved in water Bases possess properties such as a fluid consistency (e.g., detergent), an acerbic taste, the ability to react with acids to form salts, and the ability to catalyze chemical reactions There are a variety of bases, including Arrhenius, Bronsted-Lowry, and Lewis bases Soap and alkali metal hydroxides and alkaline earth metal hydroxides are

examples of bases 71,

Strong acids and strong bases: In aqueous solutions, strong acids and bases undergo

complete dissociation into hydronium and hydroxide ions, respectively "1,

Weak acids and weak bases: A weak acid or base undergoes partial dissociation in an aqueous medium, resulting in the formation of a relatively low concentration of H30* or OH

HA (aq) + H20 = H30*(aq) + A (aq)

Equation 1 The dissociation of acid

A’ (aq) + H20 = HA (aq) + OH (aq)

Equation 2 The dissociation of base

[H;0* ]|A~]

Equation 3 The acid dissociation constant equation

_ [HA][OH™]

PA]

Equation 4 The base dissociation constant equation 1.3 Buffers

A buffer solution is composed of a weak acid or weak base and its conjugate species, which

is salt By absorbing or donating protons to resist changes in pH, the weak acid or weak base provides buffering capacity Through the common ion effect, the salt helps maintain the ionic balance in the solution and prevents significant pH changes As they can absorb or release protons

to maintain equilibrium between the acid and its conjugate base, or the base and its conjugate acid, buffer solutions are used to stabilize the pH of a system To attain a desired pH range and capacity,

it is possible to modify the buffer solution's constituents Changes in pH of buffer solutions can be determined using the Henderson-Hasselbach equation:

_ [A ]

pH = pK, + log [HA]

Equation 5 The Henderson-Hasselbach equation 1.4 pH meter

By comparing the electrical potential between a pH-sensitive electrode and a stable reference electrode, a pH meter measures the acidity or basicity of a solution It is utilized in numerous disciplines, including chemistry, biology, agriculture, and industry A pH-meter is a piece of laboratory apparatus used to measure hydrogen ion concentrations (i.e acidity and alkalinity) in a water-based solution In addition, the pH scale ranges from 1 to 14, where pH = 7 indicates a neutral solution, pH < 7 indicates an acidic solution, and pH > 7 indicates a basic solution

II Experimental methods

1.1 Equipment

Portable pH meter was used to measure the pH index in solutions or water to determine acidity or basicity The solution is basic when the index is in the range of 7 < pH < 14, and acidic

if the meter displays an index of 0 < pH < 7 The chemical solution required for measurement should be ready before using The pH meter should then be turned on and allowed to warm up 30 minutes in advance Next, remove the electrode from the storage solution, rinse it with distilled water, and pat dry with paper that is free of dust Insert the electrode tip into the solution sample

that needs to be measured after that The pH index of the solution will be shown by the device after two to three minutes Clean the electrode by rinsing it with distilled water and drying it after using the equipment

Figure 1 The portable pH mete 1.2 Chemicals

1.2.1 pH of deionized water

50mL distilled water was poured into a cylindrical beaker Next, a stirring rod was used to stir well for 20 seconds, then record the pH and repeat this process two more times Continue performing the same procedure until there was no significant change in the pH value Record the data written to the data table

1.2.2 pH of strong acid

This experiment required two sections In the first section, the step to prepare the chemical solution to be measured included nearly 10mL of 0.1M hydrogen chloride (HCI) into beaker 1, 20mL of 0.1M sodium hydroxide (NaOH) into beaker 2, and 100mL of 0.01M NaOH solution (using the dilution process) In the second section, take 10mL of 0.1M HCl into the beaker and record the pH Then, add 90mL of distilled water and record the pH Next, add 10mL of 0.1M NaOH into the beaker and record the pH Then, add 90mL of 0.01M NaOH and record the pH Wait about 2 minutes and record the results, respectively

1.2.3 pH of weak acid

This experiment required two sections In the first section, the step to prepare the chemical solution to be measured included 0.1M acetic acid (CH3COOH), 0.01M CH3sCOOH

(dilute solution A 10 times) and 0.001M CHsCOOH (dilute solution A 100 times or dilute solution

B 10 times) To dilute solution B, use 5mL solution Added with 45mL deionized water that created 50mL 0.01M CH3COOH, also apply this method to solution C In the second section, take 20mL

of each solution and record the pH and Ka

1.2.4 pH of salts

This experiment was to measure the pH concentration of each salt solution including 20mL 0.1M sodium chloride (NaCl), 20mL 0.2M sodium acetate (CHsCOONa) and 20mL 0.1M ammonium chloride (NH4Cl) After preparing, use the pH meter to measure the pH index

1.2.5 pH of buffers

This experiment required three sections In the first section, the step to prepare the chemical solution to be measured included nearly 50mL 0.1M CH3COQQOH, nearly 50mL 0.1M CH3COONa, approximately 40mL 0.1M HCI and 40mL 0.1M NaOH

In the second section, adding 10mL 0.1M CHsCOOH and 40mL 0.1M CH3sCOONa in

a beaker which obtained 50mL buffer A Then, record pH (2 times) of buffer A Dividing buffer

A into two separated breakers 1 and 2, then labeling buffer A1 and A2, respectively Record pH

of each buffer A1 and A2 Adding 10 drops 0.1M HCI into buffer A1 and 10 drops 0.1M into buffer A2 and then recorded the pH of both buffers Adding more drops 0.1M HCI into buffer A1 and 0.1M into buffer A2 until the pH changed by one unit from the start, recording Vuci and Vnaou

in mL

In the third section, adding 40mL 0.1M CHsCOOH and 10mL0.1M CH3COONa ina beaker which obtained 50mL buffer B Then, record pH (2 times) of buffer B Dividing buffer B into two separated breakers 1 and 2, then labeling buffer B1 and B2, respectively Record pH of each buffer B1 and B2 Adding 10 drops 0.1M HCI into buffer B1 and 10 drops 0.1M into uffer B2 and then recorded the pH of both buffers Adding more drops 0.1M HCI into buffer B1 and 0.1M into buffer B2 until the pH changed by one unit from the start, recording Vuci and Vnaon in

mL

III Result and discussion

3 Result

3.1 pH of deionized water

Figure 2 pH concentration over different time periods

As shown in Table 1, Observed pH can be considered as a function of Time (second) There is a gradually decreasing tendency that could be seen in the data Observed pH over the period This is because of the formation of carbonic acid (H2CO3) When carbon dioxide (CO2) dissolves in water, it reacts with water molecules to form H2CO3 through the following equation:

CO2 fait) + H2Ogiquiay H2COsqiquay PT Carbonic acid is a weak acid, meaning it partially dissociates in water to release hydrogen ions (H*) This increase in hydrogen ion concentration leads to a decrease in pH, making the solution more acidic Therefore, continuous stirring of deionized water in the air facilitates the dissolution of COz and the subsequent formation of carbonic acid, resulting in a decrease in pH

(1) (2) (3) (4) (5) (6)

Figure 3 pH measured by pH meter at (1) the initial time, (2) the second time, (3)

the third time, (4) the fourth time, (5) the fifth time, (6) the sixth time

3.2 pH of strong acid

Equation using in calculation below:

pH = -log[H' ] eq1El

[HT]: the molar hydrogen ion concentration

Mi: initial concentration

V1: initial volume in liters

Mo: final concentration

V2: final volume in liters

[7]

pH + pOH = 14 eqs

Table 1 pH measurement of the reaction hydrochloric acid (HCI) + distilled water + sodium hydroxide (NaOH)

acid (HCI)

hydroxide (NaOH)

As eq, with concentration of [H*] = [HCI] = 0.1M, the theoretical pH of initial solution:

pH = -log[H*] = -log [0.1] = 1 However, the pH of the first solution was in large error compared

to the theoretical pH because when measuring, using the broken machine led to a huge difference between the practical data as can be seen in table 2 and the theoretical one

After adding 90mL distilled water into the first solution, the solution was increased from

10 mL to100 mL The concentration of H* ions decreases while the concentration of OH’ ions remains relatively constant This change in ion concentrations affects the pH of the solution As eq2, concentration of H* could be calculated:

CiVi= C2V2 0.1 x 0.01 =C2 x 0.1 %C2= 0.01

So, the concentration of H* equaled 0.1: [H"]=0.01 Based on eq1, the theorical of the

second solution: PH=-log[H*] = -log[0.01]=2 By the error of machine range +0.5, the practical data as can be seen in table 2 and the theoretical data was approximately matched, this trial matched with the theory

Upon the addition of 0.1 M NaOH to a solution containing 0.1 M HCl, both solutions

being 10 mL each and with a total volume of 110 mL, the resulting solution exhibited neutrality

This neutrality indicates an equal concentration of H* and OH’ ions, commonly observed in water

This balance is a direct consequence of the complete reaction between HCI] and NaOH, resulting

in the formation of water and sodium chloride (NaCl), a neutral salt The neutral pH value of 7

signifies the equilibrium between acidic and basic species in the solution, demonstrating an equal

strength of both acidic and basic components