Learning Objectives

Learning Objectives

  • Measure pH in acidic, basic and neutral solutions and collect data
  • Examine factors that impact the pH of a solution
  • Demonstrate how to adjust the pH of a solution
  • Demonstrate how to maintain the pH of a solution using a buffer solution
  • Discuss differences in the pH of different household solutions

The Importance of pH

You might not expect it but your life depends on the hydrogen content of every solution in your life whether it is your own blood or drinks such as lemonade or coffee. Another way to say this is: The hydrogen ion H+ concentration of a solution (or its pH) determines the suitability of that solution for various tasks. For example, household ammonia has a very high pH while vinegar has a low pH. Ammonia solutions are regarded as bases while vinegar is considered an acid.

Acids and Bases

Acids donate hydrogen ions, H+, during chemical reactions. We can represent an acid as a hydrogen containing compound, HA. In solution these compounds tend to dissociate into ions, HA → H+ + A−. The A− ion is known as the conjugate base of the acid. The dissociation of hydrochloric acid, HCl → H+ + Cl−, is a common example.

Bases accept hydrogen ions, H+, during chemical reactions. If we represent a base as B, then during a chemical reaction this compound tends to accept a hydrogen ion, H+ + B → HB+. The HB+ ion is known as the conjugate acid of the base. The reaction of ammonia in the presence of hydrogen ion, H+ + NH3 → NH4+, is an example of this behavior.

Self-ionization of Water

Interestingly, water can dissociate into ions. In pure water this reaction produces equal quantities of hydrogen ions, H+, and hydroxide ions, OH−.

H2O ⇄ H+ + OH−

The right and left arrows indicate the reaction can proceed in either direction. However, the reaction indicated by the left directed arrow is strongly favored and therefore only small quantities of hydrogen and hydroxide ions are found in water.

The pH Scale

Concentration is a measure of the amount of a substance contained per unit volume. One commonly used unit for the amount of substance is the mole (mol). When used along with a metric unit for volume, concentration may be expressed in mole per liter (mol/L).

The range of observed hydrogen ion concentrations in solutions is enormous. A typical battery may have a 1 mol/L concentration while the concentration in an oven-cleaning compound might be as low as 10−14 mol/L. Due to this large range of values, it is convenient to express hydrogen ion concentrations on a scale based on powers of ten, the pH scale.1

On the pH scale, every unit increase in pH is directly related to a factor of 10 decrease in hydrogen ion concentration. The approximate hydrogen ion concentration and pH for various solutions is shown in the table below (where [H+] is the hydrogen ion concentration in mol/L).

Solutionbattery acidvinegartomato juicemilkpure watersea waterantacidsoapy waterdrain cleaner[H+]10010010−210−210−410−410−610−610−710−710−810−810−1010−1010−1210−1210−1410−14pH024678101214

Using the pH scale, a solution can be categorized as follows below.

  • pH < 7.00 ⇒ The solution is acidic
  • pH = 7.00 ⇒ The solution is neutral
  • pH > 7.00 ⇒ The solution is basic

Measuring pH in Solutions


There are many ways to measure pH or the hydrogen ion concentration in a solution. You might have heard of pH paper, Litmus paper or pH meters. In this set of experiments, you will measure the pH of various household solutions, biological fluids, and a buffer solution. An example of a measurement made using a portable pH meter is shown on the right.

Buffer Solutions and pH

In this laboratory activity, you will get the opportunity to explore how to maintain pH within a specific range. It can be a matter of life and death to maintain biological fluids within a limited pH range. For example, if the pH of your blood was to deviate far outside of normal range (average pH of 7.3), cell structure and function would be disrupted to a life-threatening degree. Fortunately, in humans and other animals, there is a naturally occurring buffer called bicarbonate. In fact, you might have purchased a form of this buffer if you have ever bought an antacid at a pharmacy or grocery store. Can you already guess the action of this buffer?

A buffer solution can be made by adding a weak acid (an acid that contributes only a small percentage of its available hydrogen ions to the solution) and its conjugate base to water. Two examples of weak acids and their conjugate bases, and the associated buffer systems, are given in the table below. We can use a convenient shorthand notation to represent our weak acid + conjugate base buffer solution: HA = weak acid, and A− = conjugate base.

Weak Acid
HAConjugate Base
A−Buffer Solution
HA + A−acetic acid
HC2H3O2acetate ion
C2H3O2−acetic acid + acetate ion
HC2H3O2 + C2H3O2−carbonic acid
H2CO3bicarbonate ion
HCO3−carbonic acid + bicarbonate ion
H2CO3 + HCO3−

The concentrations of weak acid and its conjugate base don’t have to be the same in a buffer solution. In fact, by choosing different relative concentrations, buffers with different pHs can be constructed from the same weak acid conjugate base pair.

How does a buffer solution help maintain pH? Consider the situation shown in the figure below. On the far left we have a buffer solution that consists of a weak acid (HA) and its conjugate base (A−). This solution will have a specific pH value that depends on the particular acid-base pair used for the buffer and their relative concentrations.

A set of three diagrams. The first shows a buffer solution with equal amounts of weak acid and conjugate base. The second shows a small amount of acid being added to the buffer solution. The third shows the net result, the buffer responds to addition of acid by reacting with some of the conjugate base in the solution thereby negating most of the impact of adding acid to the solution.

In the center diagram a small amount of acid (indicated by H+ ions) is added to the buffer solution. How will the buffer respond? The net result is shown in the diagram on the far right. Notice that most of the added acid has reacted with the conjugate base (A−) in the solution to form more weak acid (HA). In this illustration, only one of three added hydrogen ions remains free in the solution and contributes to a change in solution pH. It is in this manner that a buffer system responds to additions of acid and maintains pH. Note that a buffer system can be overwhelmed by the addition of too much acid. Basically, a buffer system will lose its ability to maintain pH once all of the conjugate base has been consumed by reacting with added acid.

The buffer system above was shown responding to additions of acid to the solution. Can it also cope with additions of base? Yes, added base reacts with weak acid to form water and releases the conjugate base.

Finally, it should be noted that buffer systems can also be constructed from a weak base and its conjugate acid. The principle and actions are similar to what is shown above.

Orientation to the Acid and Base Lab Activities

Procedure I and II Overview

You will perform pH testing of household and biological solutions and use that information to classify solutions as acids or bases.

Procedure III and IV Overview

You will compare pH changes of a buffer solution and water when acid or base are added to each solution.

Summary of Formulas and Concepts Needed for Calculations

Percent Change of pH

Adding acid or base to a solution causes changes to the pH of the solution. Calculating the percent change of pH allows us to understand how large (or small) a specific pH change is.

Percent Change of pH is calculated using the formula below

Percent Change of pH=100%×final pH value – initial pH valueinitial pH valuePercent Change of pH=100%×final pH value – initial pH valueinitial pH value

Sample Calculation: Determine the percent change of pH of a solution given the data below:

  • initial pH of solution = 5.25
  • final pH of solution = 3.75

Percent Change of pH=100%×final pH value – initial pH valueinitial pH valuePercent Change of pH=100%×final pH value – initial pH valueinitial pH value

Percent Change of pH=100%×3.75−5.255.25Percent Change of pH=100%×3.75−5.255.25

Percent Change of pH=100%×−1.505.25Percent Change of pH=100%×−1.505.25

Percent Change of pH=−28.57%Percent Change of pH=−28.57%

Percent Change can be positive or negative. A positive result indicates an increase and a negative result indicates a decrease.

1. The pH of a solution is calculated as follows: pH = −log[H+], where [H+] is the hydrogen ion concentration in mol/L, and log is the common logarithm function (base 10 logarithm).

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