# 1.7: pH and Buffers


##### Learning Objectives

Goals:

• Accurately measure the pH of solutions using pH indicator strips and a pH meter.
• Create buffer solutions and test the effects of adding acid and base to each.

Student Learning Outcomes:

Upon completion of this lab, students will be able to:

• Describe the pH scale.
• Correctly use pH indicator strips and a pH meter.
• Explain the function and composition of a buffer.

## Introduction

The pH of solutions is an important characteristic. Cells must maintain a constant pH so that the enzymes and processes taking place inside the cells will continue as needed. Chemical and enzymatic reactions are typically dependent on a specific pH range. Thus, it is important to understand pH and be able to determine the pH of various solutions.

The pH scale is a familiar concept for students who study science. The pH value of a solution reflects the relative concentration of hydrogen ions (H+) or protons to the concentration of hydroxide ions (OH-) in a solution. Solutions with a pH value less than 7 are acidic and those with a value greater than 7 are basic, or alkaline. The value 7 is neutral meaning the amount of H+ in a solution is equal to the amount of OH- in a solution. Pure water H2O, which can dissociate naturally into H+ and OH- ions, would have a value of 7.

##### Equation 1

$\ce{H2O <=> H^{+} + OH^{-}} \nonumber$

Table 1. The pH Scale

[H+] in mol/L

pH

[OH-] in mol/L

pH Classification

1.0

0

10-14

Acidic

0.1

1

10-13

Acidic

0.01

2

10-12

Acidic

0.001

3

10-11

Acidic

10-4

4

10-10

Acidic

10-5

5

10-9

Acidic

10-6

6

10-8

Acidic

10-7

7

10-7

Neutral

10-8

8

10-6

Basic

10-9

9

10-5

Basic

10-10

10

10-4

Basic

10-11

11

0.001

Basic

10-12

12

0.01

Basic

10-13

13

0.1

Basic

10-14

14

1.0

Basic

#### Chemistry Review

In a chemical equation, variables that are surrounded by brackets “[“ and “]” are expressions of concentration, or the specific amount of a molecule in a given volume of solution. For example, if you see “[H+]” in an equation, this is read as “the concentration of hydrogen ion”.

The concentration of a solution is often expressed in units of moles per liter (mol/L). Just as one “dozen” represents a quantity of 12 items, one “mole” represents a quantity of approximately 6.022 X 1023 items.

one dozen molecules = 12 molecules

one mole of molecules = 602,200,000,000,000,000,000,000 molecules!

Note: “n” is used in equations to indicate a quantity measured in moles. For example if you see “nAcid” in an equation, this is read as “moles of acid”.

The term “Molarity” indicates that a solution’s concentration is in units of moles per liter. A one molar solution (1 M) contains one mole of solute within each liter of that solution. Reagents used in the laboratory will often be labeled with their concentrations expressed in terms of molarity.

The relative concentration of H+ or OH- may change very dramatically in solutions, so a logarithmic scale (called pH) instead of a linear scale is used to express concentration. Equations 2 and 3 can be used to calculate the pH based on hydrogen ion concentration or vice versa.

##### Equation 2

To calculate pH based on hydrogen ion concentration [H+]:

pH = -log [H+]

##### Equation 3

To calculate hydrogen ion concentration [H+] based on pH:

[H+] = 10-pH

#### Buffers

A buffer is a mixture of a weak acid (HA) and its salt (e.g., NaA), and is sometimes referred to as a conjugate acid-base pair. As mentioned above, buffers have a major role in stabilizing the pH of living systems. Vertebrate organisms maintain the pH of blood using a buffer composed of a mixture of carbonic acid (H2CO3) and sodium bicarbonate (Na+HCO3-). The weak acid in this buffer is carbonic acid and the salt is sodium bicarbonate. When dissolved in water, sodium bicarbonate disassociates completely into sodium ions (Na+) and bicarbonate ions (HCO3-). The H2CO3 is the conjugate acid of HCO3- and the HCO3- is the conjugate base of H2CO3 . Together, this conjugate acid-base pair functions as the bicarbonate buffer system.

Buffer systems are also of particular importance to experimental cell biology.

The pH of a buffer solution may be calculated as follows:

##### Equation 4

The pH of a buffer solution may be calculated as follows:

$pH=pK_a + log \frac{n_A}{n_{HA}}\nonumber$

Where pKa = dissociation constant of the acid, nA = initial number of moles of salt in the buffer, and nHA = initial number of moles of acid in the buffer.

If you know these values, it is possible to accurately calculate the pH of a buffer system before you create it!

The pKa of acetic acid (used in today’s experiment) is 4.75

##### Equation 5

To find the volume of the conjugate base or conjugate acid:

nA = volume of conjugate base (mL) $$\times \dfrac{1\: L}{100\: mL} \times$$ concentration of conjugate base (mol/L)

nHA = volume of conjugate acid (mL) $$\times \dfrac{1\: L}{100\: mL} \times$$ concentration of conjugate acid (mol/L)

#### Use of pH Indicator Strips

The pH of a solution can be roughly approximated using strips of paper treated with color changing indicator reagents. The strips are dipped into the solution to be tested for several seconds and then removed. The color of the indicator strip is then compared to a reference chart, often printed on the side of the strip’s container. The reference color on the chart that most closely matches the color of the reacted strip will have a pH value printed below it and that will be the approximate pH. One advantage to using pH indicator strips is that they are relatively inexpensive, easy to use, and are adequate for determining pH where an error of +/- 1 pH unit is acceptable. A more accurate method of determining pH is to use a calibrated pH meter, which can determine the exact pH to one or more decimal places depending on the quality of the device.

#### Use of a pH Meter

The pH meter measures the acidity of a solution. It is a scientific instrument that uses electrodes to measure the hydrogen ion (proton) concentration of water-based solutions. Essentially, the pH meter is a voltmeter that will measure the difference between two electrodes. The probe you place into the solution contains a reference electrode and a detector electrode. The reference electrode is not affected by the solution being measured and is in contact with a solution of potassium chloride. The detector electrode comes in contact with the test solution. The hydrogen ions in the test solution interact with the electrode and the difference in electrical potential between the two electrodes is detected and reported as millivolts or converted to a pH value.

For accurate measurements, it is important to calibrate your pH meter before use with buffer solutions of known values. It is best to calibrate your meter with buffer solutions that are near the anticipated or desired pH of your test solution. You should also blot the probe with laboratory wipes in between solutions to avoid contamination but avoid rubbing. Rubbing the probe may cause a static electricity charge to build up on the electrode which will cause inaccurate readings to occur. Accidentally letting the probe dry out will also cause it to stop working so always keep the end of the probe immersed in a holding solution when not taking measurements. Remember to return it to the storage solution as soon you are finished with the experiment.

Calibrate the pH meter for pH 4, 7, and 10 before taking measurements. If calibrated properly, your pH meter should produce measurements with an accuracy of +/- 0.06 pH units. Always test your meter after calibration using the standard buffers and recalibrate the meter if necessary before proceeding.

Your instructor will demonstrate the proper calibration, care, and use of the meter. Be sure to take good notes!

## Activity 1: Measuring pH

### Materials

Per group of 4:

• 1 Set of 4 unknown solutions (in 30 mL tubes with screw top lids)
• 1 container of pH indicator strips and color reference chart
• 1 pair of forceps
• 1 pH meter (calibrated – See instructor for directions)

### Procedure

1. Obtain a set of unknown solutions from instructor.
2. Measure the pH of each solution using the pH indicator strips first. Hold the strips with the forceps. Use a new strip for each solution!
3. Record your data in Table 1.
4. Measure the pH of each solution using the pH meter. Be sure to rinse the tip of the probe with DI water before putting the probe into each sample! (Ask the instructor for instructions if you are not sure how to properly calibrate and use the pH meter).
5. Record your data in Table 2.
Table 2. Measured pH values of Known Test Solutions

Unknown Solution

pH value measured
using indicator strips

pH value measured
using pH meter

Expected pH value

A

B

C

D

### Data Analysis

• How do your pH indicator strip values compare to your pH meter values?
• Check your measured pH values with those of the other teams. Are your values similar?
• Check with your instructor to see what the actual pH values should be. How accurate were you?

## Activity 2: Preparation of an Acetate Buffer

### Materials

Per Class:

• 1 bottle stock solution of 0.1 M acetic acid (CH3COOH)
• 1 bottle stock solution of 0.1 M sodium acetate (Na+CH3COO-)

Per Group of 4:

• 6 clean 30 mL plastic tubes
• 2 clean 5 mL serological pipettes
• 2 pipette pumps (10 mL capacity)
• 1 Sharpie Marker

### Procedure

1. Using a sharpie marker, label the two 30 mL tubes - one as “Acetic Acid” and the other “Sodium Acetate”. Fill each tube up with the correct stock solution.
2. Using a sharpie marker, label each of the two 5 ml pipettes - one as “AA” and the other as “SA”. To avoid contamination, DO NOT dip pipettes into stock solution bottles and ONLY use the designated pipette to transfer either acetic acid or sodium acetate from your group’s labeled tubes.
3. Using a sharpie marker, label a clean 30 mL tube as “Buffer 1”, another as “Buffer 2”, the third as “Buffer 3”, and the fourth as “H2O”. Each student in your group will take one tube. If there are only 3 students, one of you can also take the “H2O” tube. Write your names into the first column of table 2 next to the tube(s) you will be working with.
4. Create the acetate buffers using your marked serological pipettes and the specified volumes of acetic acid and sodium acetate in Table 2.
• Be sure to accurately pipet the volumes indicated to get good results! Review proper pipetting technique with your instructor if necessary.
• For the “H2O” tube, simply fill the tube about a third full with pure deionized water
5. Close the lids and gently shake each tube for about 20 seconds or more to mix the contents.
6. Measure the pH of each solution with the pH meter using proper technique and enter your measurements in table 2.

## Activity 3: Effects of Adding Acid and Base to Acetate Buffer

### Materials

Per Group of 4:

• Everything From Activity (2 above)
• 30 mL dropper bottle of 0.1 M HCl (Hydrochloric Acid)
• 30 mL dropper bottle of 0.1 M NaOH (Sodium Hydroxide)

### Procedure

1. Add a single drop of HCl to each of your team’s 4 tubes. Close the lids and gently shake the tubes to thoroughly mix the contents.
2. Measure the pH of each solution and enter the pH values in table 4.
3. Continue adding drops of HCl according to the table, measuring pH, and recording values.
4. When you have completed Table 3, you will now start adding drops of NaOH (base) to your tubes according to table 5.
5. Be sure to shake the tubes to mix the contents thoroughly before measuring pH and entering the values in Table 5.
6. Look at your results and compare the pH changes in your 4 tubes. What do you notice about the pH changes when you compare them?
7. Compare your pH values to those of the other teams. Ask your instructor for the expected values.
Table 3. Experimental Acetate Buffers Mixing Chart

Student

Tube

Volume of Acetic Acid (mL)

Volume of Sodium Acetate (mL)

Measured pH

Expected pH

Buffer 1

5.0

5.0

Buffer 2

7.0

3.0

Buffer 3

3.0

7.0

H2O

None

None

Table 4. Effect of Adding 0.1 M HCl (acid) to Acetate Buffers and Water
Tube pH after 1 drop HCl added pH after 2 drops HCl added pH after 3 drops HCl added

Buffer 1

Buffer 2

Buffer 3

H2O

Table 4. Effect of Adding 0.1 M NaOH (base) to Acetate Buffers and Water
Tube

pH after 1 drop

pH after 2 drops
pH after 3 drops

Buffer 1

Buffer 2

Buffer 3

H2O

## Study Questions

1. What range of pH values indicates that a solution is acidic? Basic?
2. In general, how does the relative concentration of hydrogen ions [H+] compare to that of hydroxide ions [OH-] in a neutral, acidic, and basic solution?
3. Based on your observations, how would you describe what a buffer does?
4. What factors determine the accuracy of a reading with a pH meter?
5. What is the pH of a solution that has a hydrogen ion concentration of 2.46 X 10-5 M?
6. What is the expected pH of a buffer made from 25.7 mL of 2.0 M Acetic acid and 0.0492 L of 0.90-M Sodium acetate?

This page titled 1.7: pH and Buffers is shared under a CC BY 4.0 license and was authored, remixed, and/or curated by Orange County Biotechnology Education Collaborative (ASCCC Open Educational Resources Initiative) .