Weak Acid – Strong Base Titration
A titration is the combination of two chemicals that will react. Often, we will titrate to the endpoint with a standardized solution (called the titrant) in order to find the concentration of an unknown solution. For example, we might titrate an unknown acid with a titrant of known concentration, like titrating HCl of unknown concentration with NaOH of known concentration. We know a neutralization occurs in an acid-base titration and an equivalence point is reached when the reactants are stoichiometrically related.
HCl(aq) + NaOH(aq) H2O + NaCl(aq) (1)
Since this reaction is a 1:1 mole ratio, at the endpoint (equivalence point) we know the moles of the HCl and NaOH that reacted will be equal. There may be more than one equivalence point for a diprotic or triprotic acid or a base that produces more than one hydroxide ion. Ideally, we do this with an indicator that changes color at the equivalence point so that the endpoint will coincide with the equivalence point.
Titration is a volumetric technique based on a precise measurement of volume of the solution of known molarity (standard solution or titrant) which enables one to calculate number of moles of titrant reacted and to proceed with further solution stoichiometry calculation to obtain quantitative data needed about an unknown solution (analyte). We will perform a titration slightly differently. We will begin with a sample of a weak acid and add strong base (see equation 2), measuring the pH as we go. From these data we will create a titration curve.
HC2H3O2 (aq) + NaOH(aq) H2O(l) + NaC2H3O2 (aq) (2)
A titration curve (Figure 1) monitors a titration as it proceeds with added titrant (base in this case) by graphing pH vs volume of titrant added. The curve shows several very important features that give us information. We can find the equivalence point, when the added titrant neutralizes the analyte. As you can see, this will sometimes be a neutral pH (Figure 1, i and ii) and sometimes not (Figure 1, iii). The equivalence point is at the inflection point of the curve, where we can calculate the moles of titrant added to find the moles of analyte.
Figure 1: Titration curves of (i) strong acid with a strong base, (ii) strong base with a strong acid, and (iii) a weak base with a strong acid. From Chemistry: A Molecular Approach, 5th Edition. Nivaldo Tro, Pearson Education, Inc.
Let’s say, in the titration of a weak base with a strong acid, we were titrating 100. mL of ammonia (weak base) with 0.100M HCl (strong acid). We see from the graph (Figure 1, iii) that the equivalence point occurs after 25 mL of acid has been added. In order to find the original concentration of the base we would do the following calculations.
Moles of acid added at equivalence point = (volume)acid x (M)acid (3)
Moles acid = (0.025L)(0.100M)
Moles acid = 0.0025 moles
At the equivalence point we know
Initial moles unknown base = moles acid added
moles base = 0.0025 moles
And the original base concentration would be
Mbase = = 0.025 M NH3
Because this is a weak base, the pH at the equivalence point is slightly acidic. At the equivalence point all of the weak base will be reacted and converted to the conjugate acid. Consequently the presence of the weak acid will make the solution acidic. The same reasoning applies to the titration of a weak acid with a strong base and the conjugate weak base that is present at the equivalence point.
We also see the half equivalence point (labeled in Figure 1, iii) which is in the buffer region and is where the solution contains an equimolar amount of a weak base analyte and its conjugate weak acid. At this point
pH = pKa (of the conjugate acid) (4)
pOH = pKb (of the weak base) (5)
There are several different types of calculations (Figure 2) we must do at different times of the titration (different regions of the graph). Here we will see the titration of a weak acid with a strong base.
Figure 2: Calculation regions of a weak acid/strong base titration.
Initial pH, weak acid only region (no titrant added)
In this region there is only the initial substance (weak acid in Figure 2) before any base is added. The pH will be the pH of the weak acid solution.
Buffer region (before equivalence point)
Here we have a mix of weak acid and conjugate base. We would determine this with an SCF table (stoichiometry) and an ICE table or the Henderson Hasselbalch equation.
We would do an SCF table (stoichiometry) and then an ICE table, since we are left with a weak base.
All the acid is neutralized to produce the conjugate acid at the equivalence point, but this weak base is insignificant compared to the excess strong base (titrant). We calculate the amount of excess base with an SCF table (stoichiometry) and then find pH or pOH from that concentration.
In addition to using the pH meter to measure the pH, we will also use indicator paper. Indicator paper is a piece of paper that is impregnated with an indicator or a mixture of indicators that change color depending on the degree of protonation. At a low pH, the indicator (a weak acid) contains a proton. As the pH increases, the indicator loses the proton. The key to these molecules is the weak acid form is a different color than the conjugate base. Litmus paper is an example of indicator paper, but we will use universal indicator paper.
Prelab: Show all you’re CALCULATIONS in each problem
Figure 3: A Titration curve for the titration of 20. mL of acetic acid with NaOH.
1) Predict (ICE table) the pH of 50. mL of 0.10 M HC2H3O2(aq), abbreviated as HA(aq) in further text, (Ka=1.8 x 10-5). This is the initial pH on a titration curve, if this solution is titrated with NaOH(aq). Use an arrow to point to and label this point on the curve above (Figure 3).
2) Predict the pH of a solution if 10. mL of 0.10M NaOH is added to the solution in #1. In order to solve this problem, you must calculate how many moles of NaOH are being added, then react it at 100% (SCF table) to find out how much HA(aq) and A-1(aq) is present after the reaction, and finally calculate the pH using the Henderson-Hasselbalch equation or by using the Ka (ICE table and be aware that there is a new volume). Use an arrow to point to and label this point on the curve above (Figure 3).
3) Predict the pH of a solution if 15.0 mL of 0.10M NaOH is added to the solution in #1. In order to solve this problem, you must calculate how many moles of NaOH are being added, then react it at 100% (SCF table) to find out how much HA and A-1 is present, and finally calculate the pH (Henderson Hasselbalch or ICE table). Use an arrow to point to and label this point on the curve above (Figure 3).
4) Predict the pH of a solution if a total of 50. mL of 0.10M NaOH is added to the solution in #1. In order to solve this problem, you must calculate how moles of NaOH are being added, then react it at 100% (SCF table) to find out how much HA and A-1 is present, if any. Once you know what species are present in the solution at this point, you will have to write the corresponding base dissociation of that given species with water, set appropriate ICE table, and calculate pH in the solution using Kb of the base conjugate of the weak acid initially present in the solution (described in question 1). Use an arrow to point to and label this point on the curve above (Figure 3).
5) Predict the pH of a solution if a total of 90. mL of 0.10M NaOH is added to the solution in #1. In order to solve this problem, you must calculate how moles of NaOH are being added, then react it at 100% (SCF table) to find out how much excess NaOH is present, and finally calculate the pH. Use an arrow to point to and label this point on the curve above (Figure 3).
Safety: Wear safety goggles.
Chemical Disposal: All chemicals in this experiment can be disposed of in the sink. pH paper can be disposed of in the trash.
0.1 M NaOH
250 mL beaker
Glass stirring rod
Calibrating the volume of a drop of solution:
- Prefill a 10 mL graduated cylinder with approximately 2 mL of water. Carefully record the volume to the appropriate number of significant figures. Use a plastic disposable pipet from the kit to add 25 drops of water. Be careful to make each drop consistent. Measure and record the final volume. Repeat adding an additional 25 drops, measure, and record the final volume. Repeat a third time, adding an additional 25 drops, measure, and record the final volume.
0 (original volume)
- Calculate the volume of 1 drop of solution () for each trial and report the average. Show your work and report your answer here.
- Measure approximately 5 mL of vinegar in the graduated cylinder. It does not have to be exactly 5 mL, but you must record the exact volume and add this the 250 mL beaker. Add approximately 50 mL of water to the beaker and mix with the stirring rod.
- Measure and record the pH of the solution with the pH meter. This is your initial pH. Measure and record the pH of the solution with pH paper. Dip your glass rod in the solution, remove it and touch it to a piece of pH paper. Compare the color of the paper with the chart on the container. Dispose of the paper in the garbage.
- Measure 5.0 mL of 0.1M NaOH in the 10 mL graduated cylinder. Record the volume in your data table. If it is not precisely 5.0 mL record the precise amount. Add it to the solution of vinegar (described in step 3 above) and stir the mixture thoroughly.
- Measure and record the pH with both the pH meter and the pH paper.
- Continue adding NaOH and measuring the pH by adding 2 more 5.0 mL aliquots (samples) of NaOH. Stir, measure the pH with the pH meter and the pH paper, and record your results after each addition.
- Measure 1.0 mL of NaOH with the pipet and add it to the solution. Then stir, measure the pH with the pH meter and the pH paper, and record your results.
The graduated pipet is marked to measure 0.25, 0.50, 0.75, and 1.0 mL. When you use the pipet, be sure the tip of it is below the level of the liquid when you are sucking it up and be sure there are no air bubbles in the pipet. Suck up the solution so it is above the 1.0 mL line, remove it from the solution, and squirt out the solution so it is at the line. Make this amount as precise as possible.
- Continue adding 1.0 mL aliquots of NaOH until the pH is close to 6.0. Stir, measure the pH with the pH meter and the pH paper, and record your results after each addition. Go to the next step when the pH is at or above 6.0.
- Begin adding the NaOH dropwise, 10 drops at a time. Record the number of drops you add in the first column of the data table. Using your average drop volume, calculate the volume you added and record that to the volume column in the data table. Stir, measure the pH, and record your results.
- Continue this addition of 10 drops at a time, stirring, measuring and recording, until the pH is greater than 10.0.
- Add 5 more 1.0 mL 0.1 M NaOH aliquots (measured with the pipet). Stir, measure the pH with the meter and the paper, and record your results after each addition.
- Add 1 more 5.0 mL 0.1M NaOH aliquot. Stir, measure the pH, and record your results.
Drops added, when appropriate
Volume Base added this step
Total Volume Base Added
Drops added, when appropriate
Volume Base added this step
Total Volume Base Added
- Draw a titration curve for the titration on the attached graph paper or with the scatter plot with smooth lines and markers using Excel (see Figure 1 and Figure 2). Make the graph as large as possible.
- Label the following regions on the titration curve: Weak Acid, Buffer, Equivalence point, Strong Base.
- Label the titration curve with the species (HC2H3O2, C2H3O21-, NaOH) present in each of the following regions you labeled in question 2.
- Why are the calculations done differently in the different regions you labeled in Question 2? Be specific by referring to the species present in each region.
- What is the volume of base added to reach the equivalence point (the inflection point of the curve)? ___________
- What is the initial concentration of the Acetic Acid solution (the vinegar you began with)? Keep in mind that for a monoprotic acid the moles of titrant at the equivalence point equals the moles of analyte. Using this information and the initial vinegar volume, you can calculate the concentration.
- What is the volume of base added and pH of the half titration point? _______________
- How does this pH compare to the pKa of acetic acid? Is this what you would expect?
- Since you know the pH at the equivalence point, you know the H+ concentration and can use an ICE table to calculate the initial concentration of acetic acid a different way. What was the initial concentration of acetic acid in the vinegar solution using this calculation?
- Do your calculated initial concentrations (questions 6 and 9) agree? Why or why not? If they are different, which do you think is more accurate and why?
- Compare the results of the pH readings from the pH strips with those from the pH meter. Why might they be similar or different. Which tool is better for the creation of a titration curve and why?