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This article incorporates, in modified form, material from Illustrated Guide to Home Chemistry Experiments: All Lab, No Lecture.

 

A buffer solution is a solution of, usually, a weak acid and its conjugate base, or, less commonly, a weak base and its conjugate acid. A buffer solution resists changes in the concentrations of the hydronium ion and hydroxide ion (and therefore pH) when the solution is diluted or when small amounts of an acid or base are added to it. The resistance of a buffer solution to pH change is based upon Le Chatelier’s Principle and the common ion effect.

One common example of a buffer solution is a solution of acetic acid (the weak acid) and sodium acetate (its conjugate base). In solution, acetic acid reaches an equilibrium illustrated by the following equation.

CH3COOH(aq) + H2O(I) ⇌ CH3COO- + H3O+

As we learned earlier in this chapter, acetic acid does not dissociate completely in solution. For example, in a 1 M solution of acetic acid, only about 0.4% of the acetic acid molecules dissociate into hydronium and acetate ions, leaving most of the acetic acid in molecular form. Dissolving sodium acetate in the acetic acid solution forces the equilibrium to the left, reducing the hydronium ion concentration and therefore increasing the pH of the solution.

Consider what happens if you add a small amount of a strong acid or strong base to this buffer solution. Ordinarily, you would expect adding a small amount of a strong acid or base to cause a large change in the pH of a solution. But if you add hydrochloric acid (a strong acid) to the acetate/acetic acid buffer solution, the hydronium ions produced by the nearly complete dissociation of the hydrochloric acid react with the acetate ions to form molecular (non-dissociated) acetic acid. According to Le Chatelier’s Principle, the equilibrium is forced to the left, reducing the concentration of hydronium and acetate ions, and increasing the concentration of the molecular acetic acid in the solution.

The acid dissociation constant for this buffer is:

Ka = [H3O+]·[CH3COO-]/[CH3COOH]

If the buffer solution contains equal amounts of acetic acid and sodium acetate, we can assume that the sodium acetate is fully dissociated and that dissociation of the acetic acid is negligible (because the high concentration of acetate ions from the sodium acetate drives the dissociation equilibrium for acetic acid far to the left). We can therefore assume that the concentrations of CH3COO- and CH3COOH are essentially identical, and simplify the equilibrium equation to:

Ka = [H3O+]

which means that the pH of this buffer solution is equal to the pKa.

If we add hydrochloric acid to the buffer solution, the HCl ionizes completely in solution, yielding hydronium ions and chloride ions. According to Le Chatelier’s Principle, the increase in hydronium ions forces the acetic acid equilibrium to the left, decreasing the concentration of hydronium ions and increasing the concentration of molecular acetic acid. This equilibrium shift changes the effective number of moles of acetic acid and acetate ions, which can be calculated as follows:

final CH3COO- moles = initial CH3COOH moles – initial HCI moles

final CH3COOH moles = initial CH3COO- moles + initial HCI moles

Conversely, if we add sodium hydroxide to the buffer solution, the NaOH ionizes completely in solution, yielding hydroxide ions and sodium ions. According to Le Chatelier’s Principle, the increase in hydroxide ions forces the acetic acid equilibrium to the right, decreasing the concentration of hydroxide ions and increasing the concentration of acetate ions. This equilibrium shift changes the effective number of moles of acetic acid and acetate ions, which can be calculated as follows:

final CH3COO- moles = initial CH3COOH moles + initial NaOH moles

final CH3COOH moles = initial CH3COO- moles – initial NaOH moles

In either case, after you calculate the number of moles of acetic acid and acetate ions, you can use the final volume of the solution to determine the concentration of the acetic acid and acetate ions and plug those values into the Henderson-Hasselbalch equation to determine the new pH.

pH = pK + log10([CH3COO-]/[CH3COOH])

In this lab, we’ll make up a buffer solution of acetic acid and sodium acetate and examine the effects of adding hydrochloric acid and sodium hydroxide to this buffer solution.

Required Equipment and Supplies

  • goggles, gloves, and protective clothing
  • beaker, 150 mL(2)
  • beaker, 250 mL
  • graduated cylinder, 100 mL
  • graduated cylinder, 10 mL
  • pipette, 1.00 mL
  • stirring rod
  • pH meter
  • acetic acid, 1.0 M (100 mL)
  • sodium acetate, 1.0 M (100 mL)
  • hydrochloric acid, 6.0 M (20 mL)
  • sodium hydroxide, 6.0 M (20 mL)
  • distilled or deionized water (boil and cool before use)

All of the specialty lab equipment and chemicals needed for this and other
lab sessions are available individually from Maker Shed or other laboratory
supplies vendors. Maker Shed also offers customized laboratory kits at special
prices, including the Basic Laboratory Equipment Kit, the Laboratory Hardware Kit, the Volumetric Glassware Kit, the Core Chemicals Kit, and the
Supplemental Chemicals Kit.

sciRoomCAUTION2.gif CAUTION

Hydrochloric acid and sodium hydroxide are corrosive. Wear splash goggles, gloves, and protective clothing at all times.

Substitutions and Modifications

  • You may substitute foam cups or similar containers for the beakers.
  • You may substitute a calibrated beral pipette for the 1.00 mL transfer pipette. (To calibrate the beral pipette, count the number of drops required to reach 10.00 mL in a graduated cylinder, and then calculate the number of drops per mL.)
  • If you do not have a pH meter, you can substitute pH test paper, which is sufficiently accurate for this lab session.
  • You can make up 1.0 M acetic acid by diluting 5.7 mL of concentrated (glacial) acetic acid to 100 mL. Alternatively, you may use distilled white vinegar, which is close to 1 M straight out of the bottle.
  • You can make 100 mL of 1.0 M sodium acetate by dissolving 8.20 g of anhydrous sodium acetate or 13.61 g of sodium acetate trihydrate in some distilled water and then making up the solution to 100 mL.
  • You can make 20 mL of 6.0 M hydrochloric acid by mixing 10.00 mL of concentrated (37%, 12 M) hydrochloric acid with 10.00 mL of water. If you have hardware-store muriatic acid (31.45%, 10.3 M), mix 11.65 mL of the acid with 8.35 mL of water.
  • You can make 20 mL of 6.0 M sodium hydroxide by dissolving 4.80 g of sodium hydroxide, stirring constantly, in 90 mL of water and making up the solution to 100 mL. (Caution: this reaction is extremely exothermic.)

Procedure

  1. If you have not already done so, put on your splash goggles, gloves, and protective clothing.
  2. Make up the buffer solution by mixing 100 mL of 1.0 M acetic acid and 100 mL of 1.0 M sodium acetate in the 250 mL beaker.
  3. Transfer 100 mL of the buffer solution to one of the 150 mL beakers, and 100 mL of distilled water to the second 150 mL beaker.
  4. Read and follow the directions for your pH meter with respect to calibrating it, rinsing the electrode between measurements and so on.
  5. Use the pH meter or pH test paper to measure the pH of the buffer solution, and record the observed value on Line A in the second and fourth columns of Table 11-3.
  6. Use the pH meter to measure the pH of the distilled water, and record the observed value on Line A in the third and fifth columns of Table 11-3.
  7. Use the 1.00 mL pipette to transfer 1.00 mL of 6.0 M hydrochloric acid to the beaker that contains the buffer solution and another 1.00 mL of hydrochloric acid to the beaker that contains the distilled water. Stir or swirl the beakers to mix the solutions thoroughly. (If you use only one stirring rod, rinse it thoroughly before using it in the other beaker.)
  8. Use the pH meter to determine the pH value for the buffer solution, and record that value on Line B in the second column of Table 11-3.
  9. Use the pH meter to determine the pH value for the water solution, and record that value on Line B in the third column of Table 11-3.
  10. Repeat steps 7 through 9, adding 1.00 mL of hydrochloric acid each time until you have added a total of 10.00 mL of hydrochloric acid to the buffer solution and water solution.
  11. Rinse the beakers and pipette thoroughly.
  12. Transfer the remaining 100 mL of buffer solution to one of the 150 mL beakers, and 100 mL of distilled water to the second 150 mL beaker.
  13. Repeat steps 7 through 11 using the 6.0 M sodium hydroxide solution.

Table 11-3. Observe the characteristics of a buffer solution – observed data

Acid/Base Added Buffer + HCI Water + HCI Buffer + NaOH Water + NaOH
A. 0.00 mL
___.___pH
___.___pH
___.___pH
___.___pH
B. 1.00 mL
___.___pH
___.___pH
___.___pH
___.___pH
C. 2.00 mL
___.___pH
___.___pH
___.___pH
___.___pH
D. 3.00 mL
___.___pH
___.___pH
___.___pH
___.___pH
E. 4.00 mL
___.___pH
___.___pH
___.___pH
___.___pH
F. 5.00 mL
___.___pH
___.___pH
___.___pH
___.___pH
G. 6.00 mL
___.___pH
___.___pH
___.___pH
___.___pH
H. 7.00 mL
___.___pH
___.___pH
___.___pH
___.___pH
I. 8.00 mL
___.___pH
___.___pH
___.___pH
___.___pH
J. 9.00 mL
___.___pH
___.___pH
___.___pH
___.___pH
K. 10.00 mL
___.___pH
___.___pH
___.___pH
___.___pH

Disposal

All of the solutions used in this laboratory can be flushed down the drain with plenty of water.

Review Questions

Q1: How did the pH change compare in the buffer solution versus water?

Q2: Calculate the expected pH of the buffer solution. How does the calculated value correspond to the observed value?

Q3: Graph the amounts of added acid and base against the pH of the buffer solution. At what point does the buffer begin to lose effectiveness? Why?

Q4: You have on hand 0.5 M solutions of acetic acid and sodium acetate. You want to make up 100 mL of a pH 5.0 buffer solution to calibrate your pH meter. Calculate the amounts of acetic acid solution and sodium acetate solution required to make up this buffer solution.

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