ACC BIOL 1406 Lab Manual Hays

ACC BIOL 1406 Lab Manual: Lab 3 - Understanding pH and Buffers

Objectives

• After completing this exercise, students should be able to:

  • Define and correctly use terms:

  • Dissociation

  • Acid

  • Base

  • pH

  • Buffer

  • Buffering range

  • Buffering capacity

  • Explain how solutions with different pH values compare regarding their H+ and OH- concentrations.

  • Calibrate and use a pH meter to measure the pH of a solution.

  • Plot and interpret a pH titration curve to determine the buffering range and buffering capacity of a buffered solution.

Prelab: Understanding pH and Buffers

A. Dissociation of Water

• Water is fundamental to life chemistry; when mixed with ionic or polar solutes, water molecules are attracted to them.

• Dissociation refers to the process where molecules held by ionic bonds are separated into oppositely charged ions.

• Water dissociates into:

  • $H_2O <br>ightleftharpoons H^+ + OH^-$

• The double-headed arrow indicates that the reaction is reversible.

• Pure water has equal concentrations of hydrogen ions and hydroxide ions:

  • $[H^+] = [OH^-] = 1 imes 10^{-7} ext{ M}$

• Hydrogen ions (H+) have several names including hydronium ion and proton.

B. Acids and Bases

• Definitions are often contextual, but for this lab:

  • Bronsted-Lowry Definition:

  • Acid: A proton donor.

  • Base: A proton acceptor.

• Example Reaction:

  • $HCl + NH<em>3 ightarrow Cl^- + NH</em>4^+$

  • Here, HCl (acid) dissociates and donates a proton to NH3 (base), forming Cl- and NH4+.

• **Effects of Acids and Bases on pH: **

  • Acids increase [H+] when dissociated, e.g., $HCl <br>ightarrow H^+ + Cl^-$.

  • Bases either:

  • Release $OH^-$ (e.g., sodium hydroxide $NaOH$ dissociates easily into $Na^+$ and $OH^-$).

  • Combine with H+ ions (e.g., $NH<em>3$ combines with H+ to form $NH</em>4^+$).

• Strength of Acids/Bases:

  • Strong acids and bases (HCl, H2SO4, NaOH, KOH) fully dissociate.

  • Weak acids and bases (acetic acid, citric acid, NH3) dissociate slower, leading to smaller changes in [H+] and [OH-].

C. pH

• Definition: pH is the negative base-10 logarithm of the hydrogen ion concentration:

  • $pH = - ext{log} [H^+]$

• pH scale ranges from 0 (very acidic) to 14 (very basic).

• Neutral solutions have:

  • $pH = 7$

• Inverse Relationship:

  • An acid has [H+] > $1 imes 10^{-7} ext{ M}$ and pH < 7.

  • A base has [H+] < $1 imes 10^{-7} ext{ M}$ and pH > 7.

• Example Comparison:

  • Solution A: pH 3 ($[H^+] = 1 imes 10^{-3} ext{ M}$) vs. Solution B: pH 5 ($[H^+] = 1 imes 10^{-5} ext{ M}$).

  • Solution A is a stronger acid than Solution B by a factor of 100.

• To calculate [H+] from pH, reverse the logarithm:

  • Example: pH 10 corresponds to $[H^+] = 1 imes 10^{-10} ext{ M}$.

D. Buffers

• Definition: A buffer or buffer system is a mixture of molecules that stabilizes pH by resisting changes when acids or bases are added.

• Mechanism:

  • Buffers release or bind H+ to maintain a constant pH.

  • Consist generally of a weak acid (releases H+) and a weak base (binds H+).

• Buffering Actions:

  • Addition of base: The weak acid releases H+, combining with extra OH- to form water.

  • Addition of acid: The weak base combines with H+ from added acid, reducing free H+.

• Buffering Range: The specific pH range in which a buffer is effective (e.g., a buffer may be effective between pH 2-6).

• Buffering Capacity: The amount of acid or base that can be added to a buffer without exceeding its capacity to maintain pH.

E. pH Titration Curves

• Definition: A pH titration curve graphs pH changes as acid or base is added to a solution.

• Axes:

  • X-axis represents the volume of acid/base added.

  • Y-axis shows the resulting pH.

• Use of titration curves can help determine:

  • Buffering Range: The pH range where the buffer resists pH changes, typically found where the curve is horizontal.

  • Buffering Capacity: Indicated by the relative length of the horizontal region on the curve.

F. Making Buffered Solutions

• Buffered Solution Definition: A solution containing both acid and base components necessary for maintaining pH.

• Laboratory Context: Buffered solutions can include numerous nutrients, such as vitamins and amino acids in addition to the buffer constituents.

• Experiment Procedure:

  • Prepare three solutions, all having the same sucrose concentration, with one unbuffered and two buffered (acetate buffer, bicarbonate buffer).

G. Calculations for Preparing Solutions

• Using the parallel dilution equation:

  • $C<em>1 V</em>1 = C<em>2 V</em>2$

• For the unbuffered solution and the two buffered solutions, calculate:

  • Volumes needed for:

  • Unbuffered sucrose solution $(20 ext{ mM})$

  • Acetate buffer approximately pH 4.8 (20 mM sucrose + 40 mM acetic acid + 60 mM sodium acetate)

  • Bicarbonate buffer approximately pH 9.8 (20 mM sucrose + 22 mM sodium carbonate + 28 mM sodium bicarbonate)

• Remember to account for the total desired volume and include units in the answers.