Titration of a Strong Acid by a Strong Base

Transcription

1 Strong_Acid_Strong_Base_v15b.docx Titration of a Strong Acid by a Strong Base An exploration of the reaction between a strong acid and a strong base, monitored by ph, conductivity, and the color change of an indicator molecule. 1.1 PREREQUISITE SKILLS AND KNOWLEDGE 1 OBJECTIVES Students should be familiar with electrical conductivity as a physical concept and how it pertains to aqueous solutions, Students should also have experience with RStudio, Excel, and LabVIEW. 1.2 RESEARCH SKILLS After this lab, students will have had practice in: following laboratory protocols using a laboratory notebook organizing data choosing an appropriate size of micropipette using micropipettes using a graduated cylinder using a volumetric flask diluting stock solutions calculating the concentration of a diluted solution using LabVIEW to control and collect data from a sensor measuring the ph of a solution measuring the conductivity of a solution using Excel to analyze experimental data using R to graphically analyze experimental data making publishable plots titration following hazardous waste guidelines using the proper personal protection equipment 1.3 LEARNING OBJECTIVES After this lab, students will be able to: Identify the acid and base in a given acid-base reaction Write the balanced equation for a given acid-base reaction Predict the ph of a given acid-base mixture Describe how the dissociation of water affects the ph of an aqueous acid-base reaction Predict the shape of the conductivity curve for a given acid-base titration

2 2 T itration of a Strong Acid 2.1 ACIDS AND BASES 2 PRE-EXPERIMENT ASSIGNMENT For this experiment, you will be working with a strong acid and a strong base in a neutralization reaction. In aqueous solution, an acid is any substance that increases the concentration of hydronium ions (H 3 O + ) in solution. A hydronium ion is a water molecule with an additional proton, and for that reason acids are often called proton donors. Remember that elemental hydrogen is just a proton and an electron (it is the only element that does not have a neutron). Therefore, the hydrogen ion H +, which is a hydrogen that has lost its electron, is just a proton. Strong acids include hydrochloric acid, phosphoric acid, and sulfuric acid. When hydrochloric acid HCl reacts with water, a hydrogen ion (proton) transfers from the HCl to H 2 O to form a hydronium ion, leaving behind its electron, forming a chloride ion. HCl + H 2 O H 3 O + + Cl - In aqueous solution, a base is any substance that increases the concentration of hydroxide ions, OH -, in solution, or equivalently, reduces the concentration of hydronium ions. For this reason, bases are also called proton acceptors. Strong bases include sodium hydroxide and potassium hydroxide. Sodium hydroxide, for example, dissolves in water to form hydroxide ions and sodium ions. NaOH Na + + OH - When hydrochloric acid and sodium hydroxide are combined in aqueous solution, the hydronium and hydroxide ions combine to form water, while the remaining sodium and chloride ions form a salt, sodium chloride. HCl + NaOH + H 2 O H 3 O + + Cl - + Na + + OH - H 2 O + NaCl ph and Equilibrium You will use a ph electrode to monitor the ph of an aqueous solution over the course of a neutralization reaction. Although, as noted above, you will often read and hear about hydrogen ions (H + ) in discussions of ph, in aqueous solution it is more accurate to refer to ph as related to hydronium ions (H 3 O + ). The term ph has its origin in the French phrase puissance de hydrogen (the power of hydrogen), referring to the power to which ten needs to be raised in order to express the concentration hydronium ions. An exponential scale was developed to facilitate the expression of the extremely low concentrations of hydronium ions in a solution. ph is the negative logarithm of the hydronium ion concentration. (The square brackets [A] mean the concentration of A.) ph = -log [H 3 O + ] In order to understand the ph scale, it helps to know how water dissociates. Pure water is a very weak electrolyte. It reacts with itself and dissociates to form a hydronium ion and a hydroxide ion (OH - ): 2 H 2 O (l) H 3 O + (aq) + OH - (aq) The double arrow in the above reaction indicates that the reaction can occur in both the forward and reverse directions. A sample of pure water at 25 C exists in a state of dynamic equilibrium, wherein the forward and reverse reaction rates are equal, and therefore the ratio of products to reactants in the system remains constant over time. The constant K w, called the ion-product constant for water, indicates the extent of dissociation of water. K w = [H 3 O + ] [OH - ] = 10-14

3 Titration of a Strong Acid 3 Regardless of the source of the H 3 O + and OH - ions in water, the product of the concentrations of hydronium and hydroxide ions at equilibrium at 25 C is always 1.0 x This is a very small amount of dissociation. When the concentration of hydronium ion is equal to the concentration of hydroxide ion, it is easy to calculate their concentrations. In deionized water, the concentration of hydronium ion is the same as the concentration of hydroxide ion: [H 3 O + ] = [OH - ] If the concentration of hydroxide ion (OH - ) in water is equal to 1.0 x 10-7 M, what is the concentration of hydronium ion (H 3 O + )?What is the ph of the water? Now, consider what happens when an acid is added to water: the hydronium ion concentration increases, but K w must remain constant. If K w remains constant at 10-14, then what happens to the hydroxide ion concentration when an acid is added to the water: does it go up or down? Consider what would happen if you added 1.0 ml of 1.0 M hydrochloric acid (HCl) to 999 ml deionized water. Recall that K w = [H 3 O + ] [OH - ] = When you add the acid, you will increase the concentration of hydronium ions. That is, you will be adding (1.0 x 10-3 L)(1.0 moles/l) = 1.0 x 10-3 moles hydronium ion to the water. The new concentration of hydronium ion is [H 3 O + ] = (1.0 x 10-3 mol H 3 O + )/(0.999 L L) = 1.0 x 10-3 M * What is the ph of the water to which the acid has been added? The ph scale allows for greater ease in working with the small concentrations of these ions by referring to the proportions of H 3 O + and OH - on a scale of 0 to 14. The product of the concentration of hydronium ion and the concentration of hydroxide ion is constant at Complete Table 1 below. What pattern do you notice? Table 1: ph of Dilute Solutions of Acids and Bases at 25 C [H 3 O + ] (M) [OH - ] (M) ph poh 1.0 x x * The concentration will really be (1 x 10-3 moles + 1 x 10-7 moles)/l = x 10-3 M. Why is the smaller value ( x 10-3 M) ignored?

4 4 T itration of a Strong Acid x TITRATION AND THE EQUIVALENCE POINT Titration is the process of repeatedly adding a small quantity of a known concentration and volume of one solution (titrant) to a known volume of another solution until the reaction reaches a desired end point. In this lab, you will observe the neutralization of hydrochloric acid (HCl) by titration with sodium hydroxide (NaOH). Neutralization of a strong acid with a strong base occurs when the moles of hydronium ion in solution are equal to the moles of hydroxide ion. The point in the titration when the acid has been exactly neutralized by the base is known as the equivalence point. Before starting the titration, it is useful to know the volume of titrant required to reach the equivalence point Sample equivalence point volume calculation To calculate the volume of titrant required to reach equivalence point V e, you need to know the balanced equation of the reaction, the concentration of both solutions, and the initial volume of the solution being titrated. In this example, 20 ml of M hydrobromic acid (HBr) is titrated with 0.10 M potassium hydroxide (KOH), HBr + KOH + H 2 O H 3 O + + Br - + K + + OH - H 2 O + KBr (aq) The goal is to calculate the volume V e of 0.10 M KOH required to reach the equivalence point. At the equivalence point, the number of moles of hydroxide ion exactly equals the number of moles of hydronium ion: moles OH - = moles H 3 O + In this example, the number of moles of hydronium ion is equal to the number of moles of hydrobromic acid in the initial solution, moles H 3 O + = moles HBr and the number of moles of hydroxide ion is equal to the number of moles of potassium hydroxide added to the solution: moles OH - = moles KOH added to reach equivalence Therefore, moles KOH = moles HBr

5 Titration of a Strong Acid 5 Stop, and satisfy yourself that the above equation is true at the equivalence point. Recall that the number of moles is equal to the volume (in L) multiplied by the concentration (in M = moles/l). Thus, the previous equation can be rewritten as (1) ml 0.10 M KOH 20 ml M HBr V e = 4.0 x 10-3 L = 4.0 ml Confirm this calculation for yourself. V e = (20 ml acid)(0.020 M acid)/(0.1 M base) = 10 ml base Once the full 10.0 ml of 0.1 M KOH have been added, the reaction will be complete. What will be the ph at the equivalence point? Three Simultaneous ph Observations During this titration you will simultaneously observe the ph (by measuring the voltage produced by a ph electrode), conductivity (using a conductivity sensor), and color of a ph indicator dye as the solution is fully neutralized and then titrated beyond the equivalence point. You will plot two titration curves: the first plot will have "moles NaOH added" as the independent variable and ph as the dependent variable; the second plot will have "moles NaOH added" as the independent variable and conductivity as the dependent variable Titration using a Conductivity Probe In addition to the ph meter, you will use a conductivity probe to follow the titration reaction. Use the table below to predict the conductivity of your starting solution, ending solution, and the equivalence point. How will you know from the conductivity measurement when you ve reached the equivalence point? Table 2: Equivalent ionic conductivity at infinite dilution (S cm/mol) at 25 C Cations Anions H OH K Br Mg Cl Na CH 3 CO ph Indicators Several dye molecules will change color depending on the ph of a solution. In this lab, you will use the indicator bromothymol blue. In acid solutions with a ph of 6 or lower the indicator is yellow in color, between ph 6 and ph 7.6 it is green, and above ph 7.6 it is blue. Predict the color of the solution at the starting point, ending point, and equivalence point PREPARE FOR THE EXPERIMENT Prepare an Excel workbook to collect the required data. You will need two columns for the independent variables (one for volume of base added and one for moles of base added include the calculation to make this easy) and one column for each of your dependent variables. Be sure to include units. yourself a copy. Read ahead in the Laboratory Manual to help you prepare your lab notebook.

6 6 T itration of a Strong Acid You will also need a VI to collect two data at each point in the titration: conductivity and ph. Set up the VI to collect these two data once every ten seconds or so, one time for each volume of base you add to the acid.

7 Titration of a Strong Acid MATERIALS CHECK OFF LIST Each small group of (2-3) students will have: 3 LABORATORY MANUAL laptop computer with Excel, R, and LabVIEW SensorDAQ interface ph electrode conductivity probe 10-mL serological pipette with hand dispenser set of micropipettes with associated tips box of Kim-Wipes DI water bottle 100 ml graduated cylinder volumetric flask, 100 ml tall beaker, 200 ml large magnetic stir bar small magnetic stir bar 600 ml beaker with DI water for probes 400 ml beaker for liquid waste bench top lab waste container Each large group of 2 small groups will share: 500 ml of stock solution of 0.01 M hydrochloric acid 250 ml of stock solution of 0.10 M sodium hydroxide Dropper bottle of bromothymol blue ph indicator Each class will have: ph paper in range 1-3 and SAFETY AND WASTE DISPOSAL PROTOCOLS Lab goggles, a lab coat (or a loose, but not floppy, long-sleeved shirt) and long pant legs, and closed-toe shoes are mandatory for participation in this experiment. Concentrated HCl and NaOH solutions are highly reactive and corrosive; the diluted solutions you will be working with are relatively safe. If any contact with skin occurs, wash the area IMMEDIATELY with copious amounts of water. If you have any questions, please ask your TA immediately. Make sure all nearby surfaces are dry when working with electrical equipment. Test the ph of your waste solutions before disposing of them. Solutions with a ph 5 or ph 12 must be poured into the appropriate labeled waste bottle by the sink. Solutions with a ph 6 and 11 may be poured into the sink. If you are not sure how to dispose of a specific solution, ask an instructor. Pipette tips and lab wipes can be disposed of in the bench top lab waste bags, or in the indicated Lab Waste box. Paper towels can go in the regular trash. Clean all glassware (rinsing with DI water and setting it to dry in the appropriate rack) and leave your lab station clean and tidy. Leave the conductivity probe and ph probe in DI water at your station.

8 8 T itration of a Strong Acid 3.3 EXPERIMENTAL PROCEDURE Prepare the VI Your VI will need to collect two data at each point in the titration: conductivity and ph. Set up the VI to collect one point every ten seconds or so, once for each volume of base you add to the acid Begin Titration For your first titration, use the stock concentrations of acid (HCl) and base (NaOH). If you have time for a second (and third) titration, you may want to dilute the base, to get more data near the equivalence point. 1. Decant 100 ml of HCl into the 200 ml beaker and add approximately 10 drops of the bromothymol blue ph indicator to the solution. Record the initial color of bromothymol blue in your notebook. Note any color change in the acidic solution upon addition of the indicator Remove the bottle attached to the ph meter, containing the ph buffer storage solution. Tape the ph and conductivity probes together and insert them into the HCl solution. 3. Open your VI and create a new experiment displaying the data readings of ph and conductivity. Set the sampling time at about 10 seconds; this interval will prevent inclusion of excess data. Try to remember to add only one aliquot of NaOH in each interval. 4. At the same time, open an Excel workbook. While one partner titrates, the other can record the ph and conductivity reading after each addition of NaOH. Record in your notebook how much NaOH you are adding at each interval. 5. Continue adding volumes of NaOH after equivalence is reached and monitor the changes in ph and conductivity. Ideally, there should be as many data points before the equivalence point as there are after Data Analysis Use RStudio together with Excel to make figure-quality graphs of ph vs. Moles of NaOH added and Conductivity vs. Moles of NaOH added. This can be done by converting the volume of NaOH added into moles, by way of the concentration. Make sure your graphs look professional! Open and save a fresh RStudio script (though you might also want to open Linear_Fit.R for reference). First, highlight and copy just the data from the ml of NaOH added column, shown to the right. Type in the following first line: volume.naoh <- as.numeric(readclipboard()) Go down to the Console and type in: volume.naoh and hit Enter. You should see the data from your first column. Enter the data from the remaining two columns in the same way: conductivity <- as.numeric(readclipboard()) ph <- as.numeric(readclipboard()) Use the Console to check that the data is as expected. Next you will convert your volume data to concentration by creating a new object. In the new script, type:

9 Titration of a Strong Acid 9 moles.naoh<-volume.naoh/1000*0.1 This text performs the entered operation dividing by 1000 and multiplying by 0.1 on each entry in volume.naoh. Use the Console to print moles.naoh to check that the transformed data look as they should. Q1. What is purpose of the above operation? Why is the volume divided by 1,000 and then multiplied by 0.1? You will create two different plots: one of ph versus moles of NaOH added, and one of conductivity versus moles of NaOH. Use the plot command for each plot as shown below. (Notice the green text following the # symbol. These are comments, not commands.) Q2. Snip a copy of each publishable plot and paste them here. Note that there are lines between the points in this plot type. These lines are simply there as an aid to visualizing the points. They do not indicate a fit or regression of any kind. Q3. Using your conductivity data, determine the best way to fit or relate the points to find the equivalence point. Q4. Do the same for the ph data. 2 Q5. Do the ph and conductivity data agree? Do they agree with that indicated by the color change of the solution? Why or why not? 3 Determine where each set of data comes to equivalence point and how each behaves before, during, and after. Q6. How would the graph of ph vs. Moles added differ if you titrated a beaker of NaOH with aliquots of HCl? 4 Q7. How would the graph of Conductivity vs. Moles added differ if you titrated a beaker of NaOH with aliquots of HCl? 5 Determine where each set of data comes to equivalence point and how each behaves before, during, and after. Q8. How would the graph of ph vs. Moles added differ if you titrated a beaker of NaOH with aliquots of HCl? Q9. How would the graph of Conductivity vs. Moles added differ if you titrated a beaker of NaOH with aliquots of HCl? Prepare NaOH Solution for Second (and Third) Titration If there is time, prepare a new NaOH solution for a second (and even third) titration.

10 10 T itration of a Strong Acid Keep in mind that each titration will require you to titrate in twice the amount of NaOH required for equivalence. For example, if 15 ml NaOH are required to reach equivalence, you will need to titrate in at least 30 ml NaOH. Prepare the required amount of NaOH solution in a volumetric flask using your choice of pipettes. Q10. What new NaOH concentration will you use? Q11. What is the volume required to complete the titration? After you complete the second (and third) titration, prepare your graphs as described above and paste them into your document. Be sure to label them appropriately. Q12. Did your equivalence point value change with improved resolution? 3.4 POST-LAB ASSIGNMENT Work with your extended group to submit a one-page abstract in class, describing the experiment you just completed. Include your figure in your abstract. For details, refer to the Abstract Writing Guidelines posted in the Student Resources folder on e-learning.