Skip to main content
Chemistry LibreTexts

Labs 2/3: Standardization of HCl & Potentiometry: Determination of an Unknown Soda Ash

  • Page ID
  • \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    \( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)

    ( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\id}{\mathrm{id}}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\kernel}{\mathrm{null}\,}\)

    \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\)

    \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\)

    \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)

    \( \newcommand{\vectorA}[1]{\vec{#1}}      % arrow\)

    \( \newcommand{\vectorAt}[1]{\vec{\text{#1}}}      % arrow\)

    \( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vectorC}[1]{\textbf{#1}} \)

    \( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)

    \( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)

    \( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)

    \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    In this experiment a pH meter equipped with an ion selective electrode (selective toward protons) will be used to construct the titration curve. The unknown carbonate/bicarbonate mixture which contains sodium bicarbonate and potassium carbonate in an unknown ratio is to be titrated with standard HCl using a potentiometric (pH) end point measured with a pH meter employing a pH glass electrode. The end point breaks are compared with indicator color changes.


    The important reactions for carbonate and bicarbonate are:

    \[ H^+ + CO_3^{2-} \rightleftharpoons HCO_3^- \]

    (phenolphthalein end point)

    \[ H^+ + HCO_3^- \rightleftharpoons H_2CO_3 + H_2O + CO_2 \]

    (bromocresol green end point)

    Note that between the first and second end points a buffer system will occur. This will give a poor visual end point, unless the buffer is destroyed. In practice, the visual titration used for standardization is continued until the bromocresol green end point is reached, at which time the solution is gently boiled to remove the \(CO_2 \), leaving only the remaining \(HCO_3^-\), which is then titrated to completion as was done in an earlier experiment.

    Part I. Calibration of 25 ml Pipet

    In this section of the experiment you will calibrate your 25 ml pipet. This experiment illustrates the importance of careful measurements for accurate determinations. Procedure

    1. Draw about 500 ml of deionized water into a clean beaker and let it stand 30 minutes to reach room temperature. During this time you should clean your burets and begin to dry and weigh samples for the next experiment. READ AHEAD! (See step 4 below and part 2.) Measure and record the temperature of the water.
    2. If the pipet is dirty, wash it with 6M \(H_2SO_4\). If you cannot clean it, replace it at the storeroom. You may confirm it is clean by filling the pipet to about 1 inch above the mark with deionized water and then letting it drain. There should be no water drops left on the walls. A similar criterion applies to burets and to all glassware.
    3. Calibrate a 25 ml pipet as follows:
      1. Place a 125 ml Erlenmeyer flask on the balance (your TA may explain this further). Tare the balance to give a reading of 0 grams. Since the absolute accuracy of the balances can vary slightly from one to the other, it is important that you use the same balance throughout an experiment. In this way any systematic error in the balance is removed.
      2. Deliver 25.00 ml of the equilibrated water into the 125 ml Erlenmeyer flask. Be careful to insure that the pipet is held vertically when the meniscus is lowered to the mark, and that the tip of the pipet is free from pendant drops or "creep". The pipet should be allowed to drain slowly so that no splashing occurs. After the liquid level has reached the bottom of the pipet wait 30 seconds and then touch the tip to the side wall of the flask to remove pendant drops. Never blow out a volumetric pipet – they are calibrated "to deliver" (TD) not "to contain" (TC). Weigh the flask.
      3. Each member of the group should do triplicate calibrations of the same 25 ml pipet. The average of the multiple calibrations should be used in all future quantitative experiments. Make sure that you can reproducibly pipet 25mls. There should be less than 1% variation between your samples. All your results for Chem 105, 115 and 125 depend on you being able to pipet accurately. Now is the time to practice.
    4. If you haven’t done so, now it would be an excellent time to clean your most valuable pieces of volumetric glassware, your burets. Secure the correct buret from the rack appropriate to your lab. Clean it by using the soap found near the sinks. The clean burets should always be stored in the racks, filled only with deionized water, and with the stopcocks tightly closed. Use an inverted vial for a cover during storage. Label the vial, not the buret. This is your buret, treat it carefully!

    Helpful hintS

    It is important that you always clean your buret and fill it with deionized water before storage. It is also important that when you use a buret that you follow the recommended procedure for changing a solution. First empty the buret out the top and half-fill it with deionized water. Open the stopcock and drain ~5 ml out the tip. Empty the buret out the top and repeat the water washing, this time also opening the stopcock when the buret is inverted so as to drain most of the water from the tip. Wait ~30 seconds for drainage and close the stopcock. Add ~10 ml of the new solution to the buret. Drain some of it out the tip and close the stopcock. At the sink, cradle the top of the buret between the thumb and index finger of one hand and turn the buret horizontal. Then twirl the buret and slowly empty it through the top, being careful to wet the entire interior wall with the new solution. Repeat this operation two more times. Finally, fill the buret above 0.00 ml mark and drain the excess out the tip until a reasonable starting mark is reached. Follow this procedure whenever you change the solution in a buret.

    In Part 3, weigh your carbonate unknown to ensure the sample is greater than 1.0 gram. Only one unknown sample will be issued to each group. It is the student's responsibility to make sure that four replicates can be run from the one single unknown sample!


    Calculate the average volume, relative deviation, and the standard deviation of your calibration measurements. This can easily be done on a computer spreadsheet.

    Part 2: The Preparation and Standardization of HCl

    Introduction For quantitative determination, solutions of HCl must be standardized to obtain accurate concentrations. One of the best methods to standardize HCl solutions is to use sodium carbonate. However, the use of this compound is complicated by the fact that carbonate is a weak acid and that a final product of the titration is a gas. The two acid/base reactions are:

    \[ H^+ + CO_3^{2-} \rightarrow HCO_3^- \]

    \[ H^+ + HCO_3^- \rightarrow H_2CO_3 + H_2O + CO_2(g) \]

    These two reactions combine to from a solution that is called a buffer. We will boil the solution near the end of the titration to drive out the carbon dioxide and thus destroy this buffer and get a much sharper end point. Safety Precautions: Wear your goggles throughout the entire procedure. Be especially careful when using the strong acid or strong base solutions as they can cause severe burns. Since concentrated HCl has a pungent odor, you must dispense this solution in the hood.

    Preliminary Steps

    1. Obtain approximately one gram of pure sodium carbonate, which is in the laboratory. Place the solid in a clean, dry weighing bottle. Using a folded paper strip or crucible tongs put the weighing bottle into a 50 ml beaker that has your initials on it. With tongs lay the stopper crosswise on the mouth of the weighing bottle and then put the beaker and its contents into a 110°C oven for at least one hour.
    2. Thoroughly wash a 1-liter glass bottle and its cap. Rinse everything well with deionized water. Dry the outside of each bottle but not the inside. Label the bottle as 0.1 M HCl. Preparation of solutions of approximately 0.1 M HCl solution could be prepared by dilution of concentrated HCl, which is about 12 M. A safer alternative is to dilute the 2.5 M HCl. Either way, you must compute the volume of this solution that is required to prepare 1000 ml of 0.1 M HCl using the equation M1V1 = M2V2. If you are unsure of your answer confirm it with your TA. Obtain approximately 1000 ml of degassed, deionized water as described above. If using concentrated HCl, pour about 250 ml of this water into the bottle labeled HCl. Pour concentrated HCl into a clean, dry 50 ml beaker. Use a graduated cylinder to transfer the correct amount of the concentrated HCl into the 1/4 full bottle. Always remember to add acid to water and never the reverse. Neutralize the excess concentrated acid with sodium or potassium carbonate by pouring the concentrated acid into a large beaker that contains the carbonate dissolved in a minimum amount of water. Once the bubbling subsides, ensure the pH is neutral using litmus paper. Pour it in the waste jug or wash it down the sink with a rapid water flow. Cap your HCl bottle tightly and shake it vigorously. In about three more stages add the remainder of the deionized water needed, each time shaking the bottle vigorously to mix its contents thoroughly. The final solution must be homogeneous. Set the bottle aside to allow the solution to return to room temperature. Record all volumes in your notebook.

    Standardization of HCl Solution

    1. After the sodium carbonate (Na2CO3) has dried for at least 1 hour, remove the beaker from the oven using tongs, and set it on a pad of paper on the bench top. Allow it to cool there until you can put the inside of your forearm against the beaker without discomfort. Then put the beaker into a desiccator and allow it to cool for at least another half-hour. The weighing bottle should remain uncapped throughout this operation.
    2. By difference, weigh (to the nearest 0.1 mg) triplicate 0.2000-0.2500 gram samples of primary standard (Na2CO3) and put into three separate 250 or 300 ml Erlenmeyer flasks. Since the absolute accuracy of the balances can vary slightly, it is important that you use the same balance throughout an experiment. In this way any systematic error in the balance is removed.
    3. To do such a weighing, remove the beaker from the desiccator and, using tongs, seat the top of the weighing bottle. Immediately replace the desiccator lid to avoid breakage. With a folded paper strip transfer the stoppered bottle to the balance pan and weigh. Record the weight of the full bottle. Remove the bottle with the paper strip, remove the stopper with tongs, tap out some of the contents into the Erlenmeyer flask, replace the stopper, and reweigh the bottle. Continue this operation until the desired sample size has been obtained. Weighing the second and third samples will be faster because you can match the pile size of the first sample.
    4. Dissolve the first sample of sodium carbonate in 50 ml of deionized water and add 10 drops of bromocresol green indicator. Titrate the solution with HCl until the solution just changes from blue to green. Boil the solution for 2 to 3 minutes, during this time the color should revert back to blue. If it does not, then an accurately measured amount of base of known concentration must be prepared and added to change the color to blue and this volume of added base must be included in the calculations. Cool to room temperature, and complete the titration to the final green endpoint. Repeat the procedure with the other two samples. Do not dissolve the sample until you are ready to do the titration.
    5. Compute the HCl molarity. Report an average, relative deviation, and a standard deviation.
    6. Carefully store and cap the HCl solution for use in later experiments.

    Part 3: Potentiometric determination of pH with proton selective electrode (pH meter)

    Safety Precautions: All waste solutions may be disposed of by rinsing them down the drain. Treat the pH meter with great care as it is an expensive instrument. It is critical that you are extremely careful with the very expensive and fragile pH electrodes. These must always be stored in a solution to remain moist. You should also NEVER wipe a pH electrode with a paper towel, always use a "blotting" action. WEAR YOUR GOGGLES.


    1. Obtain an unknown from the stockroom 3462. Make sure to record the total weight of your unknown to ensure it is greater than 1.0 gram. You will need to get four replicate samples from this unknown. Each group will only receive one unknown. Requiring an additional unknown comes with a point penalty. The unknowns are a mixture of sodium bicarbonate and potassium carbonate. You will determine the weight percent of sodium bicarbonate and potassium carbonate in the unknown. Do not dry the unknown in the oven. Heat enables the conversion of HCO3 to CO3.
    2. Calibrate the pH meter as described below, using the pH_standard buffer. This will consist essentially of adjusting the meter to read the proper pH with the electrodes immersed in the buffer solution. You will need to check out the pH electrode and magnetic stir bar from the stockroom.
    3. Set up the titration assembly, consisting of a pH meter, pH electrode, a magnetic stirrer, and a buret filled with your standard 0.1 M HCl that you prepared in an earlier experiment.
    4. Just before you are ready to prepare a titration curve you should standardize the pH meter using buffer solutions at pH 4.00 and 10.00.
    5. Trial titration. The purpose of this titration is to locate quickly and approximately the two end points.
    6. Accurately weigh by difference a dried sample of unknown carbonate (i.e. 25% of the unknown's initial mass) and add it to a 400 ml beaker containing a magnetic stirring bar. Add approximately 100 ml deionized water and a few drops of phenolphthalein indicator. The indicators are for the purpose of making a comparison between the potentiometric end points and the indicator color changes.
    7. Wash the standardized electrode with your wash bottle, blot it dry, and quickly position it in the titration beaker. Turn on the magnetic stirrer so it is rotating slowly. Read the pH meter and the buret, and record both readings in your notebook. The solution must be well mixed, but do not splash or draw a whirlpool. Titrate with standard HCl, taking readings about every 2 ml. It is critical that you wait the same time between adding the acid and measuring the pH, a time span of about 5 seconds works well. Continue this process by using 2 ml increments, until the phenolphthalein color disappears. Record this volume. Record the pH range for this indicator. Then add a few drops bromocresol green indicator and titrate at 2 ml increments until the second end point is reached. Record this volume. Record the pH range for this indicator. Add a few increments beyond the end point. Plot a curve of pH (on the ordinate) versus volume of HCl (on the abscissa) and locate the approximate end points.
    8. Remove the electrode, wash it thoroughly, and store it in a beaker of deionized water until you are ready to use it again. Remove the stirring bar from the beaker with a spatula. Clean and dry the bar.
    9. Final titration. Now that you have a rough idea of where the endpoints are located you are now ready to do an accurate determination. Accurately weigh another sample of the dry unknown and titrate as before, but make pH readings every 2 ml to within 1 ml of each end point (both sides of end point). Then, make readings at 0.2 ml intervals near the end point. Be sure that you are consistent and add exactly 0.2 ml each time. Near the second end point, take readings as quickly as possible because the pH will tend to drift as CO2 escapes from the solution. Some students find that adding a same set number of drops and then coverting to a volume is easier than trying to read the buret. Use the method that you feel will be most accurate for you. Note and record the points at which the indicators change color. Plot a curve of pH (on the ordinate) versus HCl volume (on the abscissa) and indicate on this curve the range in which the indicators change color. Determine the end points from the inflection points of the curve. Be sure to rinse the electrodes between titrations.
    10. Repeat the titration 2 more times.
    11. To accurately determine the endpoints of the titration from the inflection points of the curve it is very useful to make plot of ΔpH/ΔV vs. average volume of added HCl at points close to each endpoint. This curve will reach a maximum at the endpoint. An even better method is to make a plot of Δ2pH/ΔV2 vs. average volume of added HCl at points close to each endpoint. This curve will pass through the x-axis at the endpoint. Make these first and second derivative curves for each titration to determine the accurate endpoints for each titration. This can easily be done on a computer using spreadsheet software.
    12. Calculate and report the percent K2CO3 by mass and the percent NaHCO3 by mass in your unknown for each portion analyzed. They may not sum to 100% due to water in the sample. Report the mean percent value, standard deviation, 95% confidence limit, and relative deviation using a spreadsheet. Hand in the plots of the titration curves, the first and second derivatives along with your report.


    Wash all glassware with deionized water. Carefully rinse and return the pH electrodes. Carefully clean the buret, fill it with deionized water and return it to the buret rack. Use of the pH Meter Your textbook discusses the theory of electrical pH measurements in section 23G. According to the equation on page 617 the EMF of a glass electrode assembly is related to the pH of the solution in which it is immersed by the equation

    \[pH = \dfrac{\Delta \epsilon - constant}{ 0.0592} \label{1}\]

    The constant 0.0592 is a lumping of natural constants:

    \[0.0592 = 2.30259 \left(\dfrac{RT}{F}\right) \label{2}\]

    In Equation \(\ref{2}\), R = 8.314 Joules•mol-1•Kelvin-1; T = 298.16 K ; F = 96,493 coulombs•equivalent-1 and 2.30259 is the conversion factor from natural logarithms to base 10 logs. Note that the readout of a pH meter depends on temperature. Since there is no guarantee that an individual glass electrode follows Equation \(\ref{1}\) exactly, even at 298.16 K, we will calibrate each meter-electrode combination so that it reads the correct pH of two buffer solutions having accurately known values near 4 and 10. Microprocessors in the pH meter accurately calibrate the response of thermometer for the temperature dependence.

    Operation of the pH Meter - Standardization

    1. Plug the lead from the combination electrode into the input marked "pH/mV" on the back of the meter.
    2. Turn on the meter by pressing the button with a vertical line in a circle, then press the button marked "C" to clear the meter. The display should show "Clr, auto".
    3. The electrode should be gently washed with a stream of water from your wash bottle and blotted dry with a tissue. Do not wipe the electrode or you may leave a static charge on it that can seriously distort the meter reading.
    4. Immerse the electrodes into the first buffer solution. Very gently stir the electrode to ensure that all bubbles have been removed from the electrode surface. During the measurement at least the bottom 3/4" of the electrode must be immersed in the solution. Some of the electrodes have a vent hole near the top that must be uncovered.
    5. Press the "pH" button followed by the "STD" button. After the "eye" stops flashing, the display should read the pH of the buffer solution you are using along with the "1 STD" sign.
    6. Rinse and blot the electrode. Place the electrode in buffer 2 and press the "STD" button. The display should read the pH of the buffer along with the "1 STD" and "2 STD" signs. The meter is now standardized and you should never hit the "STD" button again. If you need to re-standardize, then you will need to press the "C" button, so that the display shows "Clr, auto", and continue from step 3 of this procedure.
    7. The electrode should be washed with a stream of water from your wash bottle and blotted dry with a tissue. Do not wipe the electrode or you may leave a static charge on it that can seriously distort the meter reading. 8. Gently place the electrode into the solution to be measured. Measure the pH by touching the "pH" button.

    Operation of the pH Meter: Experiment

    1. Standardize the pH meter first. Place the electrode in the solution to be titrated. Press the “auto” button to allow continuous pH monitoring.
    2. Monitor the pH as you add HCl to the solution. There is no pH lock in this mode so read the value at the same time to add the solution of HCl.
    3. Record the pH as a function of the volume of HCl added to the solution.

    Labs 2/3: Standardization of HCl & Potentiometry: Determination of an Unknown Soda Ash is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.

    • Was this article helpful?