Determination of Acetic Acid in Vinegar using a pH Electrode
What is vinegar? Vinegar principally consists of acetic acid and water. Because of its chemical properties, vinegar can be used in a variety of ways: as a cleanser, salad dressing, disinfectant, preservative, or cooking ingredient. As such, it is a very common household item. Vinegar is produced from fermentation of sugars. Fermentation processes have been used for food preservation since ancient times. Fermentation is a chemical change produced in organic substances by the action of enzymes. Fermentation can produce lactic acid, as in fermentation of milk to yogurt; or alcohol or acetic acid, as in fermentation of fruit. As Ben Selinger explains in Chemistry in the Marketplace, "Fermentation of food to produce acid is common to all cultures and cuisines, e.g., pickles, sauerkraut, coffee beans, kimchi, salami, cheese, sour dough bread, soy sauce." Bacteria, yeasts, and molds are used to produce lactic acid or acetic acid or both. Fermentation to alcohol produces a pleasant, but not stable, product unless the concentration of alcohol is fairly high. Thus, wine turns to vinegar over time. Fermentation also can happen in human intestines when certain bacteria are present that work on undigested carbohydrates. The troublesome products are gases such as carbon dioxide or hydrogen sulfide that may be formed in quantities sufficient to cause distention and pain. The starting materials to vinegar production can be any fruit or source of sugars. They are fermented with yeast to convert the sugar to ethanol (CH3CH2OH). The ethanol is then oxidized to form acetic acid (CH3COOH) by bacterial digestion of the ethanol. Check the following link: A strong acid is an acid that completely ionizes in aqueous solutions. For example, you might have
already used hydrochloric acid, a strong acid. The hydrochloric acid bottle was probably labeled
"HCl," but the chemical species in such a solution are actually ions, often symbolized as H+ and Cl-.
Because the ions are "free" to move about the solution, the HCl is said to dissociate into its constituent
ions. By contrast, in a weak acid, only a small fraction of the molecules exist in their dissociated form
at any given moment. Not surprisingly, weak acids are present in many household products. Vinegar
contains the weak acid known as acetic acid, HC2H3O2. Most wart removal preparations contain
salicylic acid, a weak acid that slowly attacks and destroys wart tissue. Salicylic acid is also released
when aspirin tablets dissolve in the stomach. This is why normal aspirin (acetylsalicylic acid) tablets
can be very hard on the stomach. Some doctors recommend that individuals who need to use aspirin
take a buffered aspirin. Buffered aspirins usually dissolve more readily and spend less time in contact
with the stomach lining.
Visit the following links about aspirin, its uses, and its chemistry: Always wear your safety goggles in the laboratory. Your TA will ask you to leave the laboratory unless you are wearing proper protective eyewear. Caution! Acids and bases are hazardous if splashed on clothing, exposed skin or in the eyes. Prolonged exposure of the skin to even dilute solutions of acid and base can cause serious burns. If acids and bases splash on skin or clothes, remove the affected clothing and flush the affected areas thoroughly with cold water. Caution! Although acetic acid is a weak acid, prolonged exposure to it, or solutions like vinegar that contain it, can cause chemical burns. Prolonged exposure to many pure weak acids and bases or solutions of weak acids and bases can cause serious burns. In case of contact with the skin, remove any remaining material and flush the exposed areas with running water. Any chemical that comes in contact with your eyes should be flushed out with running water for at least ten minutes. Locate the safety shower and eyewash station prior to beginning the experiment. Notice on disposal of chemical wastes: All of the solutions used in this laboratory can be discarded down the drain after being neutralized. Consult your TA on the proper method of disposal. Ingestion: Never taste anything in the laboratory, and do not bring food or drink into the laboratory. Never put objects, such as pens, into your mouth. It is necessary that you wash your hands thoroughly when you leave the laboratory to ensure that you do not ingest anything that might be on your hands. In case of a spill: For this experiment, use a damp sponge to collect the spilled solution and wash the area with water. Rinse the sponge thoroughly, and wash your hands thoroughly. Use the internet, or other resources such as the labels on your favorite foods and drinks, to answer the following questions for submission to your TA at the start of lab. DO NOT taste any material that is not classified as a food. Can you think of any foods, whether processed or unprocessed, that have a sour taste? Make a short list of these foods and suggest chemical compounds that might be responsible for their sour flavor. Procedure
In the first part of this investigation you will use a computer interface to gather data on the acidity of solutions. An interface is an electronic device that takes an electronic signal from a device (usually called a probe) and translates it into a signal that the computer reads. The computer requires a software package called LoggerPro to import the electrical current (signal) produced by the probe and translate it into numbers that can displayed in a spreadsheet or graph on the monitor. There are several solution-phase household items available in the lab. You will use thand H of these items. 1. Make sure that the pH electrode and temperature probe are connected to the LabPro box, and that it is, in turn, connected to the LabPro box via a port, either CH1, CH2, CH3, or CH4. 2. Start LoggerPro by clicking its icon on the desktop. Go to the Experiment pull-down menu and select "Set Up Sensors" followed by "Show All Interfaces." The resulting Dialog Box should show a picture of the pH electrode in the box corresponding to its port. If not, right click on the box corresponding to the port you're using for your pH probe. Select "Choose Sensor" followed by "pH/mV/ORP Amplifiers," and finally, "pH." 3. Unscrew the cap holding the pH electrode in the bottle, rinse with a small amount of distilled water, gently shake dry, and touch with a Kimwipe to dry. Place the pH probe successively into each of the 4. Obtain a small amount (~3 mL) of one of the supplied household substances in a 10 mL-size graduated cylinder. 5. Place the pH probe directly into the solution in the graduated cylinder, making certain that the glass bulb at the end of the probe is submerged and not surrounded by large pockets of air bubbles. Record the displayed pH value. 6. Rinse the probe with distilled water. Use a beaker to catch the rinse water. Gently shake the probe to remove any excess water, and if you wish, touch it with a Kimwipe to dry. 7. Repeat steps 4-6 for the other household items. What trends do you see for the measured pHs? For example, are there similarities among substances with similar pHs? Part 2 – Determining the Concentration of Acetic Acid in Vinegar Use your current understanding of acids, bases, and concentration, along with references available within this document and on the web, to develop a plan for determining the concentration of acetic acid in vinegar. Discuss the following with your group members as you consider how the titration experiment described below will allow you to accomplish this task: What key things do you need to know before starting the experiment? What chemical reaction will be taking place during the titration? How will you calculate the acetic acid concentration from your titration data? How will you know whether your results are accurate? Do you wish to compare your results with those of other lab groups who may be studying different brands or different types of vinegar? In performing a titration, the relative concentrations of the acid and base need to be considered. In this experiment the NaOH concentration is about 0.15 M. As you perform the experiment, consider what problems might arise if, for example, 1 M or 0.01 M NaOH was employed instead. In this experiment only 5.00 mL of vinegar will be titrated. It can be difficult to see the color of the indicator in this small volume. Thus, deionized water will be added to increase the volume, rendering the endpoint to visually easier to observe. 1. Obtain 12-15 mL of vinegar in a small beaker. Using a pipet, transfer 5.00 mL of vinegar into a 100-mL beaker. The carefully measured sample of vinegar is a liquid aliquot. Add approximately 20 mL deionized water. Why is it unnecessary to add a precise volume of H2O? 2. Add a magnetic stir bar to the beaker and place the beaker atop a magnetic stir plate. (Continuous stirring during the titration will homogenize the sample as drops are added from the buret. To prevent breakage of the pH electrode, be careful that the stir bar does not touch the electrode during stirring.) 3, Add approximately 3 drops of phenolphthalein pH indicator to the beaker. (While titrating with the base, the phenolphthalein will exhibit a faint pink color that disappears as the solution is mixed.) 4. Obtain about 50 mL of NaOH solution for titrating your vinegar sample. Note that the specific molarity (M) is shown on the reagent bottle. Should you use this molarity value, or the nominal value of 0.15 M, in your calculations? 5. To prepare the buret for use, rinse the buret with distilled water and then with a small quantity of the NaOH solution. Allow the some of the solution to run out the tip; then slowly rotate the buret while tipping it to a horizontal position to rinse the walls of the buret. Dispose of this wash solution. This process removes leftover rinse water from the buret, so that the NaOH is not diluted when the buret is filled. Fill the buret with the NaOH solution until the level of the solution is around the 1-mL mark. 6. Using a piece of paper with a blackened area behind and below the meniscus (to reflect off the meniscus), read the initial buret volume. How many decimal places should this reading contain? 7. Calibrate the pH Probe according to the directions (if necessary, see 8. To prepare your computer to plot the titration data as pH vs. volume of titrant using LoggerPro, open the file called "Exp 24a Acid-Base Titration" under the File pull-down menu. (This file resides in a folder under Program Files\Vernier Software\Logger Pro3\Experiments\_Chemistry with Computers\Exp 24a Acid-Base Titration.) 9. Each x-y data point of the titration curve will consist of a titration vessel pH and its corresponding buret volume reading. To begin, click on the “Collect” button, monitor the pH reading for 5-10 seconds or until it is stable, and then click on the “Keep” button. An "Events with Entry" dialogue box should appear; type in the buret reading for this first data point. Now click the OK button. 10. For the second and subsequent data points, you will first add the desired volume of titrant, let the pH reading stabilize, and click Keep. The volume that you enter for each point should be the buret reading, not the individual increment that you just allowed to drip into the beaker! 11. Continue in this manner to construct the titration curve point by point; as the color of the indicator takes longer and longer to fade, add smaller incremental volumes of NaOH. The equivalence point, at which the amount of base added is equal to the amount of acid present, occurs when the faint pink color persists. Make special note of the buret volume and pH at this visually determined equivalence point. 12. Continue titrating beyond the equivalence point to assess how the pH changes beyond the equivalence point. Once the pH is no longer changing rapidly upon addition of NaOH, the titration experiment can be considered complete. Click the Stop button. 13. You'll notice that there are columns for the first and second derivatives of your data. To look at these graphs, simply click the vertical axis label of your graph, select the Axes Options tab, and check the box for first or second derivative, labeled d1 or d2. With your knowledge of calculus, can you use one or both of these derivative plots to rapidly identify the titration equivalence point? 14. Use Word and Excel, if directed by your TA, to save electronic copies of your numerical data and graphical titration curve. 15. Perform at least two additional titration trials. 16. When you are finished with the pH probe, rinse the tip with distilled water, place it back into the storage solution vial, and screw the cap back on. NOTE: The storage solution is a special buffer solution that preserves the probe. If the solution is spilled, visit the stockroom, 2019 Malott, to have the bottle refilled. Appendix
Acids are proton (i.e., H+) donors. Accordingly, chemists often use the symbol H+ to represent the ion produced by dissociation of an acid molecule. However, in aqueous solutions, "bare" H+ ions do not exist. Instead, they reside on water molecules--that is, the H+ is "donated" to the water molecule by the acid. The symbol H3O+, often called the hydronium ion, thus provides a better representation of reality than H+. Nonetheless, for describing aqueous solutions, H+ and H3O+ are used interchangeably by scientists. The stronger the acid, the greater the number of H+ (H3O+) ions found in a solution of a given concentration. Bases are proton acceptors and will accept protons from acids. When combined, aqueous acids and bases react to neutralize each other, forming water and a salt. A solution resulting from a mixture of a simple strong acid and strong base will be neutral if exactly the same amounts (a stoichiometric ratio) of each are mixed. If an excess of one or the other is present, the solution will exhibit either acidic or basic character. The pH of a solution can be viewed as a measure of the solution's acidity (or basicity). The pH is related to the concentration of H+ (i.e., H3O+) which is the ion responsible for the acid’s reactions. A ion concentration by measuring the voltage created in the probe. The software automatically converts the voltage to a pH value. The pH is determined based on the hydronium ion concentration using the following equation: (1) pH = - log [H3O+]
"Log" is short for logarithm and is a mathematical calculation useful in working with exponential numbers. Note: If you are unfamiliar with logarithms it might be a good idea to review the properties of logarithms in a math book or in your text. The log (base 10) of any number is the exponent of 10 raised to that power. The subscript 10 indicates a base ten logarithm, because this is a common base for logarithms; in practice, the 10 is usually not explicitly written. Another common logarithm is the natural logarithm, base e, symbolized as ln. Both logarithms are found on scientific calculators. Most calculators also have a key, or sequence of keys, for the inverse logarithm (usually shown as log-1 or 10x for base 10 logarithms) to find the value of the number. Using the definitions of pH and logs we can rearrange the formula to solve for the hydronium ion concentration: Thus, if you know the pH of a solution you can determine the hydrogen ion concentration by raising 10
to the power of -pH. For example a pH of 3.5 would give (plugging 3.5 into eqn. 2):
[H3O+] = 10-3.5 = 10 0.5 x 10-4 = 3.16 x 10-4 A titration is the controlled addition of one reagent to another using the stoichiometric relationship between the two reagents. Precisely calibrated glassware is used to measure accurately the volumes involved. The fixed volume of one reagent may be measured with a pipet or automatic pipetter (auto-pipetter). A solution may be added using a buret, which allows for the precise measurement of the volume of solution. The endpoint of the titration can be determined by observing a change in some property of the solution. Two different methods are employed this semester: one uses a pH indicator; the other requires potentiometric measurement of pH. A suitable indicator should react with H3O+ at the pH of the endpoint of the titration. To qualify as an indicator, the species produced at the endpoint must be distinguishable from the species in the original solution--e.g., it must exhibit a different color. To learn more about titrations, examine the following Websites. Designing and interpreting a titration requires an understanding of acid strength. Students are often surprised to learn that an acid’s strength refers to its chemical reactivity rather than its concentration. A strong acid is an acid that completely ionizes in aqueous solutions. The titration of a monoprotic (i.e., containing only one H+) strong acid with a strong base results in neutralization as shown in the following chemical equation: (3) H3O+ + OH- ֖ 2 H2O
In the titration of a strong monoprotic acid with a strong base, the pH of the solution is solely based on the concentration of H3O+ in the solution. The titration curve can be divided into three different regions that differ accordingly: 1. Before the equivalence point, the pH of the solution is determined by the concentration of H3O+ from the unreacted acid; 2. At the equivalence point the pH is determined by a phenomenon called the autoprotolysis of H2O (2H2O ֖ H3O+ + OH- ), which affects a small subset of water molecules at any given moment; and 3. After the equivalence point, when an excess of base has been added, the pH is determined by the concentration of the excess OH-. In the titration of a weak acid, the determination of pH is a bit more complicated because the limited
dissociation of a weak acid must be considered:
(4) HA + H2O ֖ H3O+ + A-
When considering the pH of a weak acid solution, the [H3O+] in the solution at the beginning (and
throughout the titration) is dependent not only on the concentration of the acid, but also on the strength
of the acid as indicated by a constant specific to each acid, known as the acid dissociation constant, or
Ka. During the titration the hydroxide ion reacts (eqn. 3) with the H3O+ in solution formed from the
dissociation of the acid.
To determine the pH of the weak acid solution before the equivalence point, an equilibrium expression may be written for the reaction: (5) Ka = ( [H3O+] [A-] ) / [HA]
Which can be written in logarithmic form and rearranged to give: (6) pH = pKa + log ( [A-] / [HA] )
This equation shows the relationship between pH of the solution and the strength of the acid, Ka, before the acid is completely neutralized. At the equivalence point, the number of moles of added base is exactly equal to the number of moles of HA originally present. Past this point, the pH of a titration solution changes rapidly, as the added base no longer reacts with the acid but increases the OH- concentration. An acid-base indicator is a compound that changes color over a specific pH range. A pH indicator is a weak acid, and thus has an available proton. Each indicator has a different chemical structure which may be complex. For simplicity, assume the symbol for the protonated indicator is HIn. The HIn can lose its proton; shown in the dissociation equation: (7) HIn + H2O ֖ H3O+ + In-
The equilibrium constant for this dissociation is written: (8) Ka = ( [H3O+] [In-] ) / [HIn]
This expression can be rewritten in the form similar to eqn. 6,
(9) pH = pKa + log ( [In-] / [HIn] )
This equation shows that the relative abundances of the acid and base forms of the indicator change as the solution pH is changed. HIn and In- are different colors. As one form changes into the other (by the loss of a hydrogen ion), a change in the color of the solution will result. In the experiment described here, phenolphthalein is used as the indicator. The HIn form of phenolphthalein is colorless, while the In- form is red. When the indicator is in a solution of low pH (acidic), the major form is HIn.
Thus the human eye sees the color of the HIn. As the pH is increased (becoming more basic), more In-
is present in solution, as illustrated by LeChâtelier's principle (eqn. 7). As In- becomes the dominant
species, the eye sees the color of the In- in the solution (red or pink). What the human eye detects
depends upon the relative amounts of the two forms, or the [In-] / [HIn] ratio.
The indicator used for a given acid-base titration experiment must be chosen carefully. The indicator must change close to the endpoint. Thus, when enough base has been added to neutralize all of the acid, the pH change resulting from further base addition will cause the indicator to change. Indicator error will be introduced if the indicator does not change at the precise equivalence point. If an indicator is chosen correctly this error will be minimized. For the titration of acetic acid with sodium hydroxide, the equivalence point is approximately 8. Phenolphthalein changes from colorless to red in a pH range of 8.0-9.2, suggesting minimal indicator error.


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