9.2: Lab - The pH of Household Products - Acid and Base Solutions
- Page ID
- 438413
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Prepare 12 sets of equipment and materials for 24 students per class section. Include a few more for backup if needed. Each set should include
Equipment
- Computers and internet access
- pH probe and pH indicator paper
- mortar and pestle
- 250 mL Erlenmeyer Flask
- Graduated Cylinder
Materials
- Bromo green indicator
- 6M \(\ce{HCl}\) in dropper bottles
- ID unknowns as acid, base or neutral
- A baking soda solution labelled Unknown A.
- A soda solution labelled Unknown B
- A salt solution made with distilled water to get as close to pH 7 labelled Unknown C
- Clinical Case Unknowns: We will use three different brands of antacid as shown in the images below:
Antacid 1
Antacid 2
Antacid 3- Please place the tablets from each antacid type in lab vials and label them as antacid 1, 2 and 3
- We do not want them to read labels but concentrate only on their data
- Each group should use 1-2 tablets for each antacid type for their experiment
Note: We have tested these samples, with 6M \(\ce{HCl}\), Bromo Green indicator to show that antacid 2 > antacid 1 > antacid 3. We appreciate if this same samples are kept throughout for consistency.
- To measure the pH of different household chemicals
- To explore the health or environmental impact of the pH of different solutions
- To identify the ions involved in determining pH of acids and bases
- To calculate the pH of a solution using the hydronium ion concentration
- To calculate the pH of a solution using the hydroxide ion concentration
- To learn to use a pH indicator paper, laboratory pH meter and virtual pH meter
Laboratory Skill
- Practice recording data from experiments and computer simulations
- Practice making inferences from pH data
- Practice using a calculator to solve pH problems
Equipment and Materials
- An electronic device (Computer, tablet or phone) with internet access
- Sodium bicarbonate also called baking soda (\(\ce{NaHCO3}\))
- Milk
- Soda
- Salt solution (with distilled water to get pH 7)
- Isopropyl alcohol
- pH meter or paper
Safety and Hazard Information
Personal Protective Equipment (PPE) required: Safety goggle, closed-toe shoes
Background Information
pH, acids and bases
The pH of an aqueous solution is a measure of whether the solution is acidic or basic. Many biological processes are impacted by pH changes and can lead to negative and severe consequences for patients if not corrected. In humans, the maximum amount of oxygen is carried in blood when the pH is maintained between 7.35 and 7.45. A blood pH of 6.8-7.30 is an indication of a medical condition called respiratory acidosis. In the environment, acidic rain can lead to the dissolution and destruction of edifices like the Statue of Liberty. Many agricultural processes rely on soil with optimum pH (5.5 and 7.0) to ensure that plant roots absorb the right nutrients for growth. Understanding how acids and bases affect the pH of the chemicals we encounter every day, knowledge of biological processes and reactions is essential.
Definition of acids, bases and pH
A compound that dissociates (ionizes) in aqueous solution to produces hydronium ions \((\ce{H3O^+})\) or hydrogen ions \((\ce{H^+})\) is called an acid. An acid donates its proton to water. Bases on the other hand ionize in solution to produce hydroxide ions \((\ce{OH^-})\). Acids are classified as strong or weak. Strong acids dissociate completely in solution producing only ions. Whereas weak acids dissociate partially in solution producing fewer ions. As a result, the concentration of hydronium ions, \(\ce{H3O^+}\) (or hydrogen ions, \(\ce{H^+}\)) and pH values are different for various acid solutions. Scheme 1 below describes the ionization or dissociation in water of a strong acid (\(\ce{HCl}\)), while scheme 2 describes the ionization or dissociation in water of a weak acid (\(\ce{HF}\)).
\[\ce{HCl} (aq) + \ce{H2O} (l) \rightarrow \ce{H3O^+} (aq) + \ce{Cl^-} (aq) \nonumber\]
Scheme 1: Ionization reaction of a strong acid
\[\ce{HF} (aq) + \ce{H2O} (l) \rightleftharpoons \ce{H3O^+} (aq) + \ce{F^-} (aq) \nonumber\]
Scheme 2: Ionization reaction of a weak acid
The pH scale (Figure \(\PageIndex{1}\)) is a logarithmic scale (each number represents a tenfold change) defined as the negative logarithm of the molar concentration of hydronium ions present in solution
\[\text{pH} = -\log[\ce{H3O^+}] \label{1}\]
The pH values of chemicals range from 0 to 14, with pH 7 being neutral. A pH value below 7 is acidic, and above 7 is basic (alkaline). Common body fluids (e.g. human tears) have pH 7.4. If the pH of the liquid near your eyes is changed by a small amount you will feel irritation. This is why you experience eye irritation in a swimming pool or from a shampoo. If the pH of the solution near the eye is changed by a large amount, such as dropping below pH = 4 or going above pH = 10, one will not only experience irritation, but permanent damage to the eye may also occur due to breakdown of tissue on the surface of the eye. The typical home in the United States contains several commercial products that are acidic or basic enough to be eye hazards. To be safe, eye protection should always be worn when working with these products. Understanding how chemicals in our food create pH imbalances can help healthcare professionals think critically about their patients and the overall impact that pH fluctuations have on their wellbeing.
We will use a paper strip that changes its color when a small amount of solution is added to it. This pH indicator paper provides a visual way to estimate the pH of solution by comparing color change to the indicator scale (Figure \(\PageIndex{1}\)). These pH indicators can be used to test water samples from swimming pools.
In addition to pH strips, accurate and quantitative measurements of the pH of different solutions can be achieved using a pH meter.
In this lab, common household chemicals that represent eye hazards will be classified as either acidic or basic. Students will also learn how the ions present in solution (\(\ce{H^+}\), \(\ce{H3O^+}\), and \(\ce{OH^-}\)) can be used to classify a solution as acidic or basic, calculate the pH. Incidentally, this is critical in the application of tums to fight acid reflux of heartburns.
How to calculate pH?
In order to calculate the pH of an aqueous solution, you need to determine concentration of the hydronium ion in units of moles per liter (Molarity). For example, if the \([\ce{H3O^+}]\) in an aqueous solution is determined to be \(1.0 \times 10^{-3}\) M:, the pH can be calculated using Equation \ref{1} as follows:
\[ \begin{align*} & \text{pH} = -\log [\ce{H3O^+}] \\ & \text{pH} = -\log (1.0 \times 10^{-3}) = -(-3.00) = 3.00 \end{align*} \]
Note here the number of decimal places in the pH value is the number of significant digits in the \([\ce{H3O^+}]\). The number to the left of the decimal point represents a power of 10.
Incidentally we can calculate the \([\ce{H3O^+}]\) if the pH value is known or measured, by reversing Equation \ref{1}. This is shown in Equation \ref{2} below
\[[\ce{H3O^+}] = 10^{-\text{pH}} \text{ or } [\ce{H3O^+}] = \text{antilog} (-\text{pH}) \label{2}\]
In water, there is an equilibrium between water molecules, where one water molecule serves as an acid, the other a base. This acid/base ionization reaction produces \(\ce{H3O^+}\) and \(\ce{OH^-}\) (Scheme 3).
\[\ce{H2O} (l) + \ce{H2O} (l) \rightleftharpoons \ce{H3O^+} (aq) + \ce{OH^-} (aq) \nonumber\]
Scheme 3: Autoionization reaction of a water and the ion product
This equilibrium can also be described by an equilibrium constant, \(K_w\), given in Equation \ref{3}
\[K_w = [\ce{H3O^+}][ \ce{OH^-}] = 1 \times 10^{-14} \label{3}\]
\(K_w\) is also called the ion product of water and has been experimentally determined. It shows that in every aqueous solution, at 25 \(^{\circ}\)C, the product of the concentrations of \(\ce{H3O^+}\) and \(\ce{OH^-}\) is equal to \(1.00 \times 10^{-14}\). This expression also depicts how the concentrations of each of the ions can be used to determine whether a solution is acidic \(([\ce{H3O^+}] > [ \ce{OH^-}])\) or basic \(([\ce{H3O^+}] < [ \ce{OH^-}])\) or neutral \(([\ce{H3O^+}] = [ \ce{OH^-}])\).
Two common reactions involving acids and bases form the basis for the action of antacids (TUMS, Maalox, Mylanta) to relieve heartburn and indigestion. In these reactions commonly referred to as neutralization reactions, bases \((\ce{Al(OH)3}, \ce{Mg(OH)2})\) and carbonates of calcium and sodium react to diminish the impact of gastric juice \((\ce{HCl})\) as describe in Schemes 4 and 5 below.
\[2\ce{HCl} (aq) + \ce{Mg(OH)2} (s) \rightarrow 2\ce{H2O} (l) + \ce{MgCl2} (aq) \nonumber\]
Scheme 4: Neutralization of \(\ce{HCl}\) with magnesium hydroxide base in Mylanta
\[2\ce{HCl} (aq) + \ce{CaCO3} (s) \rightarrow \ce{H2O} (l) + \ce{CaCl2} (aq) + \ce{CO2} (g) \nonumber\]
Scheme 5: Neutralization of \(\ce{HCl}\) with calcium carbonate base in TUMS
Special Instructions (if any)
Personal Protective Equipment (PPE) required: Safety goggle, closed-toe shoes
Procedure
\(\PageIndex{A}\): Measure the pH of household samples using virtual pH meter
1. Open the PhET lab simulation link here or at https://phet.colorado.edu/sims/html/...-scale_en.html
2. Open the Macro section to explore and get used to the controls.
3. Select a sample from the list provided on the page.
4. Move the probe over to the solution to measure and record the pH.
5. Report your data in Table \(\PageIndex{1}\) below.
\(\PageIndex{B}\): Identify the type and ratio of the predominant ions in household products
2. Open the Micro section to explore and get used to the controls.
3. Select an acid, a base or a neutral sample.
4. Click to check the box for \(\ce{H3O^+}/\ce{OH^-}\) ratio.
5. For each sample, record the pH, concentration of \(\ce{H3O^+} ([\ce{H3O^+}])\) and concentration of \(\ce{OH^-} ([\ce{OH^-}])\).
6. Record your data in Table \(\PageIndex{2}\).
\(\PageIndex{C}\): Identify an unknown solution as acidic or basic using pH strip or pH meter
1. Obtain a liquid sample from the stock table and record the sample number (Sample ID).
2. Use a pH indicator strip and pH meter to measure and record the pH of the sample.
3. For the pH indicator strip:
a. Dip a glass rod into the tested solution.
b. Deliver a drop of the unknown solution to a small piece of pH test strip.
c. After a couple of seconds while strip is still wet, compare the color of the pH strip to the color chart provided with the pH paper kit.
d. Record your pH value.
4. If using a pH probe:
a. Dip the pH probe into clean distilled water and then a buffer solution of pH 7.
b. Place your probe in the unknow sample and record the pH after it stabilizes.
c. Clean probe with distilled water and replace it in buffered container.
5. Calculate the \([\ce{H3O^+}]\), \([\ce{OH^-}]\), and pOH for the sample.
\(\PageIndex{D}\): Clinical Antacid Case study to explore the strength of different antacids in the market
In this section you will perform an experiment to estimate the dosage or the amount in grams of an over the counter antacid that will be needed to neutralize a certain quantity of gastric acid (\(\ce{HCl}\)). You will then analyze your data to determine the antacid with higher neutralization strength
1. Obtain 1-tablet of antacid provided and record its ID in Table \(\PageIndex{3}\).
2. Using a mortar and pestle crush the tablets provided to fine powder.
3. Transfer crushed tablet to a weigh boat.
4. Measure and record the total mass of the crushed tablets and weigh boat (Initial Mass).
5. Place 100 mL of water in a 250 mL Erlenmeyer flask. Add 3-5 drops of bromo green indicator, then stir the mixture thoroughly using a stirring rod and record its initial color and pH.
6. To the solution in the flask add 2-5 drops of 6M \(\ce{HCl}\) acid solution until the mixture changes color. Record and measure the color and pH of the mixture at this point.
7. Slowly add about a tip of a spatula full of the crushed tablets to the colored solution in the flask while continuously mixing the components thoroughly until the solution reverts to the initial color. This is the end point of the neutralization reaction.
8. You may need to add more of the crushed antacid to reach the neutralization point (original solution color) and record the pH.
9. Measure and record the mass of the weigh boat with left-over crushed tablets. This is the Final Mass of weighing boat + Crushed tablets.
10. Subtract the final mass from the initial mass to determine the amount of antacid used to neutralize the acid.
11. Repeat steps 1-9 above, using antacids sample 2 and 3 one at a time.
12. Use the data obtained from the three experiments (trials) to identify the antacid that reached the neutralization point fastest, and explain why?
Experimental Report
\(\PageIndex{A}\): Measure the pH of household samples using virtual pH meter
Table \(\PageIndex{1}\): pH data collected from virtual pH meter
| Item | pH | Type of Solution | Eye Hazard | |
|---|---|---|---|---|
| 1 | Water | |||
| 2 | Drain Cleaner | |||
| 3 | Hand Soap | |||
| 4 | Blood | |||
| 5 | Spit | |||
| 6 | Milk | |||
| 7 | Chicken Soup | |||
| 8 | Coffee | |||
| 9 | Orange Juice | |||
| 10 | Soda Pop | |||
| 11 | Vomit | |||
| 12 | Battery acid |
Follow-up Questions:
Put A next to the substances from Table \(\PageIndex{1}\) which are acidic.
Put B next to the substances from Table \(\PageIndex{1}\) which are basic.
Put N next to the substances from Table \(\PageIndex{1}\) which are neutral.
Put a star next to the pH of those substances in Table \(\PageIndex{1}\) which are eye hazards (cause severe damage).
\(\PageIndex{B}\): Identify the type and ratio of the predominant ions in household products
Table \(\PageIndex{2}\): The ions dominant in acidic and basic solutions
| Solution | \( [\ce{H3O^+}]\) (M) | \([\ce{OH^-}]\) (M) | pH | Product |
|---|---|---|---|---|
| Acid | ||||
| Base | ||||
| Neutral |
Follow-up Questions:
Find the product of the \([\ce{H3O^+}]\) and \([\ce{OH^-}]\) for all your solutions and record value on Table \(\PageIndex{2}\).
Show all work here:
Reflect on your data and come up with a relationship (or trend) between pH and \([\ce{H3O^+}]\) or \([\ce{OH^-}]\).
Suppose you are allowed to mix drain cleaner and battery acid, what type of solution would you expect?
\(\PageIndex{C}\): Identify the unknown solution as acidic or basic using pH strip or pH meter
| Sample ID:_________________________ | Sample pH:_________________________ |
| A | \([\ce{H3O^+}] =\) |
| B | \([\ce{OH^-}] =\) |
| C | pOH = |
Show work here:
Reflect on your data and explain whether your sample is acidic, basic or neutral using multiple pieces of data.
\(\PageIndex{D}\): Clinical Antacid Case study to explore the strength of different antacids in the market
| Data | Antacid -1 | Antacid -2 | Antacid -3 |
|---|---|---|---|
| Initial color of the water-indicator mixture | |||
| Initial pH of the water-indicator mixture | |||
| pH of the acidic mixture | |||
| Initial Mass of weighing boat + Crushed tablets | |||
| Final Mass of weighing boat + Crushed tablets | |||
| Final Color of the mixture | |||
| Final pH of the mixture | |||
| Mass of tablets used in neutralization |
Which brand has stronger antacid effect, explain your answer choice?
Follow up questions:
Identify each solution as acidic, basic, or neutral.
a. ___ Stomach acid with a pH = 1.5
b. ___ pancreatic fluid, \([\ce{H3O^+}] = 1 \times 10^{-9}\) M
c. ___ A soft drink, pH = 3.0
d. ___ pH = 7.0
e. ___ \([\ce{OH^-}] = 3 \times 10^{-11}\) M
f. ___ \([\ce{H3O^+} ] = 5 \times 10^{-13}\) M
Post Lab Clinical Application Case:
In patients with Chronic Kidney Disorder (CKD), the kidneys can’t remove enough acid, which can lead to a disorder known as metabolic acidosis. Metabolic acidosis results in a decrease of plasma bicarbonate concentration below the normal range (22-29 mEq/L) and in the lowering of pH. Bicarbonate therapy is recommended for treatment of metabolic acidosis. Use the data from the table below to identify the patient(s) that will need a sodium bicarbonate 5% IV solution, and explain your choices?
| Laboratory Data | Patient 1 | Patient 2 | Patient 3 |
| Bicarbonate(mEq/L) | 19 | 18 | 23 |
| pH | 6.2 | 6.0 | 7.0 |
Calculate the concentration of ions in liquid products:
Calculate the \([\ce{H3O^+}]\) or \([\ce{OH^-}]\) and pH of the solutions shown below using one of the equations below:
\[\ce{H3O^+} = 10^{-\text{pH}}, \ 1 \times 10^{-14} = [\ce{H3O^+}][\ce{OH^-}], \ \text{pH} = -\log [\ce{H3O^+}], \ \text{pOH} = -\log [\ce{OH^-}] \nonumber\]
\([\ce{OH^-}] = 1.1 \times 10^{-9}\) M
\([\ce{H3O^+}] = 5.0 \times 10^{-7}\) MShow your work here: