# Buffer Solution Demonstration: Acetic Acid/Acetate vs. Water Compare pH

This demonstration focuses on imparting the concept of a buffer solution to students.  Solutions having a pH of  4, 5, 6, 7, 8, 9, 10 are placed in separate labeled Erlenmeyer flasks. Universal indicator is added to each solution.  [pH meter is optional]

Distilled water is adjusted to have a pH = 7.0.  Half of this solution is placed in beaker "A" . Beaker "A" and its contents are on a stir plate with a magnetic stir bar in the water.  Universal indicator solution is added. Next, the demonstrator builds the buffer solution in front of the students starting with deionized water in Beaker "B" and then adding acetic acid.   Beaker "B" and its contents are on a stir plate with a magnetic stir bar in the solution.  The instructor asks students to estimate the pH of the acetic acid solution and to write the equilibrium equation representing this weak acid system.

CH3COOH(aq) + H2O(l) <=> H3O+(aq)CH3COO(aq)    Ka = 1.8 x 10-5

Universal indicator is added to the acetic acid solution and to the students observe the indicator in the acetic acid solution has a color corresponding to acidic.

The instructor asks students what will happen to the pH of the acetic acid solution when sodium acetate is added.   Students should invoke LeChatilier's Principle and the common ion effect.  The acetate ion is the common ion.  When added to the acetic acid system at equilibrium, the acetate will react with some of the hydronium ions, causing the equilibrium to a shift to the left.  Since the hydronium ion concentration decreases, the pH should increase (become less acidic).

Next, solid sodium acetate is added to the acetic acid solution until the color of the indicator in the solution is "green" corresponding to pH = 7.   Both the water and the acetic acid/acetate solution have the same color and therefore both solutions have same pH.

Small quantities of 010 M HCl and then 0.10 M NaOH are added to water.  Students observe the color of the indicator in the water changes dramatically. Small quantities of 010 M HCl and then 0.10 M NaOH are added to the acetic acid/acetate solution. The color of the indicator in the buffer solution slightly changes.  Water is not a buffer solution and the acetic acid/acetate solution is a buffer solution.

An acidic buffer is a solution of a weak acid (acetic acid) and its conjugate base pair (sodium acetate) that prevents the pH of a solution from changing drastically through the action of each component with incoming acid or base.

Acetic acid in the buffer solution will react with the addition of sodium hydroxide, NaOH

CH3CO2H(aq) + OH-(aq) –> CH3CO2-(aq) + H2O(l)       K=1.8 x 109

The acetate anion in the buffer solution will react with the addition of hydrochloric acid, HCl

CH3CO2-(aq) + H3O+(aq) –> CH3CO2H(aq) + H2O(l)     K=5.6 x 104

Curriculum Notes

A Class activity accompanies this interactive demonstration.  A computer animation showing a dynamic representation of the interactions of weak and and conjugate base in the the acid-base reactions at the particle level can accompany this activity.  When instructors expect their students to think of buffers in microscopic or symbolic terms and to relate the different representations of buffers with each other, they need to show those representations and their connections with the students explicitly. Students do not generally make these connections on their own or think in terms of "molecular scenes".

Clicker questions asking students to predict what will occur can accompany this activity. Quiz questions assess students understanding of buffer solutions are available.  Sample lecture notes accompany this activity.

Suggested Computer Animations

The following animations present a simplified representation of a molecular view of what occurs when acid or base is added to a buffer system.

http://www.chembio.uoguelph.ca/educmat/chm19104/chemtoons/chemtoons7.htm

https://video.search.yahoo.com/yhs/search;_ylt=AwrUily3FcBa4lIAvYU2nIlQ?...

Student Difficulties

1. Buffer solutions are commonly viewed by students as static systems instead of dynamic equilibria systems.

2. Students have a difficult time interpreting chemical formulas confidently.  Students have difficulty with identifying if chemical formula represents a weak or strong acid, weak or strong base, acidic or basic ionic salt. The chemical formulas for salts of conjugate bases are particularly hard for students to interpret.

3.  Students have a difficult time understanding and predicting if a a soluble ionic salt will generate an acidic, basic, or neutral solution when dissolved in water.

4. Students have a difficult time writing an equilibrium chemical equation representing a buffer system.  Students not understand the relationship between a weak acid and its conjugate base.  Students assume that any two chemicals that are mixed will react together and students will write an equation for the chemical reaction between the weak acid and its conjugate base.

HA  + A- <=>

5. Student difficulties in understanding buffer conceptually are related to their inability to visualize buffers on the microscopic scale. Students have a difficult representing (drawing) a buffer solution using a "picture diagram" or "molecular scene" and have difficulty interpreting a "molecular scene" representing of a buffer . Students have difficulty relating the macroscopic, microscopic and symbolic representations of buffers.  Students who can draw and interpret "molecular scenes" of buffer solutions exhibit and demonstrate a better conceptual understanding of buffers compared to students who cannot do so.

6. Students confuse the concentration of hydronium ion in the buffer solution (used to calculate the pH) with the initial concentration of weak acid (i.e. 1.0 M) - used in the Henderson-Hasselbalch equation. Students are not able  to distinguish between or relate the weak acid component of the buffer and the hydrogen ions that determine the pH of the solution.

7.  Students believe that all buffers have a pH = 7, neutral.

8. Students do not conceptual relate that fact that pH is a logarithmic scale and the log of the concentration of the H3O+ ions in solution (with acknowledgement of the role that an activity coefficient plays) determines the acidity or alkalinity of the solution. Students’ inability to understand logarithmic functions (base 10) has consequences for their understanding of buffers.

9. Students have difficulty differentiating between Ka and pKa or pH and [H+].

10. Because students do not have a good conceptual understanding of buffers and because they try to approach buffer problems from a purely mathematical perspective, some students believe that there is only one way to solve a particular type of buffer problem.

11. Some students believe the strength of the buffer is determined by the strength of its component acid and base: a buffer made from a strong acid and a strong base would be stronger (have a higher buffer capacity) than a buffer made from a weak acid and a weak base.

12.  Some students have the idea that buffers have an unlimited ability to resist pH changes.  In fact, buffer solutions have a finite capacity to resist pH changes.  In an acidic buffer solution it is the number of moles of weak acid and number of moles of its conjugate base that determine the extent to which a buffer can neutralize added acid or base.

13. Students have a difficulty explaining what occurs when strong acid or a strong base is added to a buffer system.  Students need to incorporate in their explanation, chemical equations, particle diagrams, and written explanations.

Learning Objectives

1. Identify two components of an acidic buffer solution and explain the function of each component.

2.  Write the chemical equilibrium equation representing and acetic acid-sodium acetate buffer system.

3.  Write appropriate chemical equations and explain how one component of a buffer system reacts when acid is added, and the other component reacts when base is added.  Show that these reactions only slightly increase or decrease the pH of the solution.

4. Write appropriate chemical equations and explain how water reacts when acid is added, and when base is added.  Show that these reactions  increase or decrease the pH of the solution.

5.  Write appropriate chemical equations and explain why the concentrations of the two buffer components must be high to minimize the change in pH due to the addition of small quantities of acid,H3O+, or base OH-.

6.  Explain why the best pH of an acidic buffer system is ±1 pH of the pKa of the weak acid.

AP Chem Learning Objective

The student can identify a solution as being a buffer solution and explain the buffer mechanism in terms of the reactions that would occur on addition of acid or base.

One day of lead time is required for this project.
Discussion

Ask students to predict the pH of this system.  Nearly all students will say, with confidence, acidic.  Add some universal indicator solution and students will observe the color of the indicator corresponds to an acidic solution.  Tell students the acetic acid is the "Big Dog" in the yard.  The "Big Dog" establishes the equilibrium system.

Next ask students what will happen to the pH of the solution when some  solid sodium acetate is added to the acetic acid solution?  Tell students that the sodium acetate is the "Little Dog".  Remind students they have observed the common ion effect and they know Le Chatelier's Principle.  Students should predict that adding some acetate ions will shift the equilibrium to the left, decreasing the [H+] concentration, thus increasing the value of pH.

Ask students to write the equilibrium equation representing this system.  Some students will attempt to write

HA  + A<=>

Students will need to be convinced that the equation representing the equilibrium system remains

HA  + H2<=> H+ + A

Acetic acid is the "Big Dog" and the "Big Dog" establishes the equilibrium in the yard.

Preparation

3M or 6 M acetic acid, 50 g of solid sodium acetate, DI water, three large beakers, two stir plates, two magnetic stir bars, universal indicator solution, 0.10 M HCl. 0.10 M NaOH   Solutions pH = 4, 5, 6, 7, 8, 9, 10 in labeled Erlenmeyer flasks with universal indicator in each solution.   A pH meter is optional.

Footnotes

References

Orgill, M.K., & Sutherland, A. (2008). Undergraduate chemistry students’ perceptions of and misconceptions about buffers and buffer problems. Chemistry Education Research and Practice, 9, 131–143.

Drechsler, M., & Schmidt, H. J. (2005), Textbooks’ and Teachers’ Understanding of Acid-Base Models Used in Chemistry Teaching, Chemistry Education Research and Practice, 6 (1), 19-35.

Summerlin, L.; Borgford, C.; Ealy, J. Chemical Demonstrations: A Sourcebook for Teachers; Volume 2; 1987; p. 172-173.

Donahue, C.J.; Panel, M.G. (1985). Buffer capacity of various acetic acid-sodium acetate systems: A lecture experiment.  Journal of Chemical Education62(4) p 337. DOI: 10.1021/ed062p337.

Paul G. Hobe Jr. (1979).  Buffer effect demonstration on the overhead projector. J. Chem. Educ.56 (1), p 47.

James C. Chang (1976). A buffer solution and its action.  J. Chem. Educ.197653 (4), p 228.

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