6.2: Lab-Acid Equilibria
- Page ID
- 342657
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Learning Objectives
Goals:
- Determine an equilibrium constant from experimental data
- Understand the effect of dilution on percent ionization for a weak acid
- understand the common ion effect and LeChatlier's principle
By the end of this lab, students should be able to:
- Apply the concept of equilibria and equilibrium constants to weak acids (and bases)
- Design and run an experiment using a pH meter to determine the equilibrium constant for a weak acid
- Use an equilibrium constant and pH meter to determine the concentration of an unknown weak acid or base
- Predict the effect on the pH of and acid (or base) by adding the salt of that acid (or base).
Prior knowledge:
Concurrent Reading:
- Section 15.3: Determining an Equilibrium Constant
- Section 15.6: Disturbing Equilibrium (LeChatlier's Principle)
Safety
- Emergency Preparedness
- Eye protection is mandatory in this lab, and you should not wear shorts or open toed shoes.
- Any spills should be cleaned up immediately
- Minimize Risk
- Label all containers
- Keep work station clear and neat
- Recognize Hazards
- Acetic Acid PubChem LCSS
- All solutions should be considered harmful and care should be taken to avoid contact with your skin or other body tissues.
- In event of contact with reagents you should flush contacted area with water and notify instructor immediately.
Figure \(\PageIndex{1}\):OSHA's Chemical Safety Board release on July 27, 2021 fatal accident in Texas involving acetic acid. Now, this incident involved 100,000 pounds of the acid being released. Glacial acetic acid (anhydrous pure acetic acid) is very dangerous and we are using dilute aqueous solutions. Acetic acid is the common ingredient in vinegar, which you probably have in your kitchen, but you must alert your instructor if you have a spill (Copyright; screen capture by Bob Belford, CC0)
Equipment and materials needed
Burets | 4-250 mL Erlenmeyer flask | 50 mL vol flask |
1-25 mL pipet | 2-10 mL pipets | 50 mL beaker |
3-100 mL vol flasks | 2.00 M Acetic Acid | 2.00 M Sodium Acetate |
pH meter | 3-buffers at pH=4,7,10 | Acetic acid unkown concentration |
Background
Before proceeding read sections 15.2: Equilibrium Constants and 15.3: Determining an Equilibrium Constant. The generic expression for an equilibrium constant for the reaction \[aA+bB \rightleftharpoons cC+dD \] is: \[ K =\dfrac{[C]^c[D]^d}{[A]^a[B]^b}\]
Rice Diagram
We typically use a RICE diagram to describe the chance in concentration as a reaction proceeds from initial to final (equilibrium) concentrations. For the above generic reaction, the RICE diagram is:
Reactants | aA + | bB | ⇌ | cC + | dD |
---|---|---|---|---|---|
Initial | [A]Initial | [B]Initial | [C]Initial | [D]Initial | |
Change | -ax | -bx | +cx | +dx | |
Equilibrium | [A]Eq | [B]Eq | [C]Eq | [D]Eq |
noting that:
\[ \left [ A\right ]_{Eq}=\left [ A\right ]_{Initial}-ax \\ \left [B\right ]_{Eq}=\left [ B\right ]_{Initial}-bx \\ \left [ C\right ]_{Eq}=\left [C\right ]_{Initial}+cx \\ \left [D\right ]_{Eq}=\left [D\right ]_{Initial}+dx \label{15.3.3} \]
where,
x=extent of reaction
and
\[K=\frac{\left [C\right ]_{eq}^{c}\left [D\right ]_{eq}^{d}}{\left [A\right ]_{eq}^{a}\left [B\right ]_{eq}^{b}}\]
We will use the equilibria of a weak acid as our example because we can measure the hydronium ion concentration with a pH meter.
pH and Concentration
In this lab we will use the pH meter to measure the hydronium ion concentration. From section 4.5 (gen chem 1) we know
pH = -log[H3O+]
[H3O+] = 10-pH = 1/10pH
Weak Acid Equilibria
From section 3.5.1.2: Acid-Base Reactions we know that the Bronsted-Lowry definition of an acid is a proton donor, which can generically be represented as HA, and that in aqueoous solutions the proton is donated to water
\[HA(aq)+H_2O(l)⇌H_3O^+(aq)+A^-(aq)\]
At first glance this gives an equilibrium constant of
\[K=\frac{[H_{3}O^{+}]_{eq}[A^{-}]_{eq}}{[HA]_{eq}[H_{2}O]_{eq}}\]
but water is in excess (section 15.2.5) and so this reduces to
\[K_a=\frac{[H_{3}O^{+}]_{eq}[A^{-}]_{eq}}{[HA]_{eq}}\]
where we ignore the water and write the equation as:
\[HA \rightleftharpoons H^+ + A^-\]
Which results in the following RICE diagram.
Reaction | \(HA\) | \(H^+\) | \(A^-\) |
---|---|---|---|
Initial | [HA]i | 0 | 0 |
Change | -x | +x | +x |
Equilibrium | [HA]i-x | x | x |
Noting that \(x=10^{-pH}\) (at equilibrium) and substituting, gives\[K_a =\frac{(10^{-pH})^2}{[HA]_i-10^{-pH}}\]
So with a pH meter we can study equilibrium problems involving acids (and bases) and by measuring the pH we can
- Calculate Ka knowing [HA]i (solutions 1-3)
- Calculate Ka knowing [HA]i and [NA]i (solution 4)
- Calculate [HA]i knowing Ka (solution 5 - the unknown)
Common Ion Effect
In solution 4 we are mixing equal volumes of 0.1M acetic acid (weak acid) and 0.1M sodium acetate (strong salt section 3.4.2), which have the common ion acetate. Before mixing we would expect little of the acid to dissociate, and all the salt to be dissociated. So before mixing we expect
\[\begin{align}HC_2H_3O_2 &\rightleftharpoons H^+ + C_2H_3O_2^- \\ NaC_2H_3O_2 &\rightarrow Na^+ + C_2H_3O_2^- \end{align}\]
According to LeChatlier's principle (section 15.6), a system shifts its equilibrium to consume a chemical species involved in a reaction if they are added to a system at equilibrium, and so if we add sodium acetate to the weak acidic acid we would expect the reaction to proceed to the left, and thus the common ion inhibits the ionization.
Reaction | \(HA\) | \(H^+\) | \(A^-\) |
---|---|---|---|
Initial | [HA]i | 0 | [NaA]i |
Change | -x | +x | +x |
Equilibrium | [HA]i-x | x | [NaA]i+x |
\[K_a=\frac{x\left ( [NaA]_i+x] \right )}{[HA]_i-x} \; \\ \; \\ \; =\frac{x\left ( [NaA]_i] \right )}{[HA]_i}\]
This is saying that the common ion (A-) added by the salt pushes the equilibrium to the left and inhibits the ionization of the acid, driving the pH up. We note here a specific case where if the acetate ion concentration equals the acetic acid concentration they cancel and you have a quick way to measure the equilibrium constant.
\[K_a=\frac{x\left (\cancel{[NaA]_i]} \right )}{\cancel{[HA]_i}} = x = 10^{-pH}\]
So we have a very easy way to measure the acid ionization constant of a weak acid.
Experimental Procedures
In this lab we are going to measure the pH of five solutions. Three are solutions of acetic acid (HA), one of an unknown and one is an equimolar solution of acetic acid and its salt.
Preparing Reagents
- Solution 1: Calculate the volume of 2.00 M acetic acid that will make 100 ml of 0.100 M acetic acid, transfer that volume to a 100 mL volumetric flask and dilute to volume with DI water. You will measure the pH of this solution, and use it to make three other solutions.
- Solution 2: Using a volumetric pipette transfer 25 mL of solution 1 to a 50 mL Volumetric flask and dilute to volume.
- Solution 3: Using a volumetric pipette transfer 10 mL of solution 1 to a 100 mL Volumetric flask and dilute to volume
- Solution 4: Using volumetric pipettes transfer 10 mL of solution 1 (0.100 M acetic acid) and 10 mL of 0.100 M sodium acetate to a small beaker
- Solution 5: Unknown (add enough to a small beaker to submerge the pH probe tip).
Measuring pH
- Set up pH meter as instructed in the general information page, section 0.4.4 (pH Meter under Instrumentation)
- Test pH meter with buffers and if required, calibrate.
- Measure pH of each of the 5 solutions and record values in the data sheet.
Data Analysis
Cover Page
As always, fill out the cover page of your Google Workbook
Solutions
The material in light brown represents data or values you used in making the solutions (like the concentration of a reagent you diluted, and the volume you diluted it to). The stuff in blue represents calculations you do with the data. Some of these you will use in other sheets of this workbook.
![clipboard_ee04610c5ea609aa5be8f13381f9fdcba.png](https://chem.libretexts.org/@api/deki/files/392048/clipboard_ee04610c5ea609aa5be8f13381f9fdcba.png?revision=1)
Remember, you can use formulas in worksheets. For example you wish to find the initial volume of concentrated acetic acid from the dilution formula. So you know that
\[M_iV_i=M_fV_f\] and can solve for the volume you need with the spreadsheet, and then use that volume in the lab. So in the cell B4 you can type
=(Z1*Z2)/Z3
where Z1, Z2 and Z3 are the cells of your sheet that have the data you wish to use in your calculation.
Data Analysis
Once again, you can use functions or work the data up by hand.
![clipboard_ea328e16292b7d899455bf544318aa2b4.png](https://chem.libretexts.org/@api/deki/files/392050/clipboard_ea328e16292b7d899455bf544318aa2b4.png?revision=1)
- Row 2 for solutions 1-4 are the molarities of the solutions created in the solutions worksheet (you can cut and paste values).
- Row 3 has the pH values for each of the solutions you measured and recorded in your data sheet.
- Calculate hydronium ion concentration from pH (weak acid equilibria above).
- For solutions 1-3
- Calculate acetate ion concentration from hydronium ion concentration (Table \(\PageIndex{3}\) above)
- Calculate undissociated acid concentration at equilibrium from the initial concentration and the extent of reaction (Table \(\PageIndex{3}\) above)
- Calculate Ka from equilibrium acid, hydronium and acetate concentrations
- Calculate percent ionization from initial acid concentration and extent of reaction (x) in (Table \(\PageIndex{3}\)
- For solution 4 insert values but calculate Ka from just the pH measurement (see eq. 6.2.13)
- The unknown is acetic acid and you are using the pH meter to determine its concentration. Use the average Ka from the four solutions you prepared.