Experiment 01Acid-Base Biophysics • B.Sc. / M.Sc.
Preparation of Buffers & Determination of pKa of Glycine by pH Titration
🎯 Aim: To titrate the simplest amino acid glycine against standardized $0.1,text{N} text{HCl}$ and $0.1,text{N} text{NaOH}$, construct the pH titration curve, identify buffering plateaus, calculate dissociation constants ($text{p}K_1$ and $text{p}K_2$), and determine the isoelectric point ($text{pI}$).
🔬 Theoretical Principle & Dissociation Equilibria
Amino acids are dipolar amphoteric molecules (zwitterions) bearing both acidic carboxyl ($-text{COOH}$) and basic amino ($-text{NH}_3^+$) groups. In strongly acidic solution ($text{pH} < 2$), glycine exists as a fully protonated monocation ($^+text{H}_3text{N}–text{CH}_2–text{COOH}$, net charge $+1$).
As strong base ($text{NaOH}$) is incrementally titrated, the $alpha$-carboxyl group dissociates first with an apparent ionization constant $K_1$ ($text{p}K_1 = 2.34$). When exactly $0.5$ equivalent of base is added, $[^+text{H}_3text{N}–text{CH}_2–text{COOH}] = [^+text{H}_3text{N}–text{CH}_2–text{COO}^-]$, and according to the Henderson-Hasselbalch equation, $text{pH} = text{p}K_1$:
text{pH} = text{p}K_a + log frac{[text{Conjugate Base}]}{[text{Conjugate Acid}]}
At $1.0$ equivalent of $text{NaOH}$, the carboxyl deprotonation is complete, yielding the dipolar zwitterion ($^+text{H}_3text{N}–text{CH}_2–text{COO}^-$, net charge $0$). The pH at which the net electrical charge is zero is the isoelectric point ($text{pI}$), where the molecule displays minimal electrophoretic mobility and lowest solubility:
text{pI} = frac{text{p}K_1 + text{p}K_2}{2} = frac{2.34 + 9.60}{2} = 5.97
Further addition of base titrates the $alpha$-amino group with $text{p}K_2 = 9.60$, yielding the fully deprotonated anion ($text{H}_2text{N}–text{CH}_2–text{COO}^-$, net charge $-1$).
Figure 1.1: Glycine pH Titration Curve & Zwitterion Dissociation States
🧪 Reagents, Materials & Solution Recipes
Reagents & Standard Solutions
- 0.1 M Glycine Solution: Dissolve 0.750 g analytical grade glycine in $100,text{mL} text{dH}_2text{O}$.
- Standardized 0.1 N NaOH: Dissolve 4.0 g NaOH in $1000,text{mL} text{dH}_2text{O}$, standardize against potassium hydrogen phthalate (KHP).
- Standardized 0.1 N HCl: Dilute 8.3 mL concentrated HCl ($37%$) in $1000,text{mL} text{dH}_2text{O}$, standardize against anhydrous $text{Na}_2text{CO}_3$.
- Standard pH calibration buffer tablets (pH 4.00, pH 7.00, pH 10.00).
Apparatus & Glassware
- Digital pH meter with combination glass-reference electrode
- Magnetic stirrer with Teflon stirring bars
- Two 50 mL precision burettes ($0.05,text{mL}$ graduation) & stand
- 150 mL glass beakers, wash bottle with boiled deionized water
📋 Step-by-Step Experimental Protocol
- Electrode Calibration: Calibrate the digital pH meter using freshly prepared standard buffers of pH 4.00, 7.00, and 10.00 at $25^circtext{C}$. Rinse the electrode thoroughly with deionized water and blot dry with lint-free optical wipe.
- Titration Setup: Pipette exactly $25.0,text{mL}$ of 0.1 M glycine solution into a 150 mL beaker. Place a magnetic stir bar in the beaker and mount over a magnetic stirrer set to gentle agitation (avoid frothing or splashing). Immerse the pH electrode securely so the bulb is submerged without touching the stir bar. Record initial pH.
- Acid Titration (pK1 Region): Fill burette 1 with standardized 0.1 N HCl. Add HCl in increments of $0.5,text{mL}$, stirring for 15 seconds after each addition until the pH stabilizes. Record pH after each increment until pH drops to $approx 1.50$.
- Base Titration (pK2 Region): Repeat the setup with a fresh $25.0,text{mL}$ aliquot of 0.1 M glycine. Fill burette 2 with standardized 0.1 N NaOH. Add NaOH in increments of $0.5,text{mL}$, recording pH after each addition until pH reaches $approx 12.0$.
- Plotting & Calculations: Plot pH (Y-axis) versus volume of 0.1 N NaOH / HCl added (X-axis). Determine the midpoints of the two buffering plateaus ($text{p}K_1$ and $text{p}K_2$) and the inflection point ($text{pI}$).
📊 Representative Titration Data Matrix
| Added 0.1 N Reagent (mL) | Equivalent of Acid / Base | Measured pH | Predominant Ionic Species | Net Electrical Charge |
|---|---|---|---|---|
| 12.5 mL HCl | 0.5 Eq HCl Added | 2.34 (pK1) | $[^+H_3N–CH_2–COOH] = [^+H_3N–CH_2–COO^-]$ | +0.5 (Buffering Zone 1) |
| 0.0 mL (Initial) | 0.0 Eq Added | 5.97 (pI) | $^+H_3N–CH_2–COO^-$ (Zwitterion) | 0.0 (Isoelectric Point) |
| 12.5 mL NaOH | 0.5 Eq NaOH Added | 9.60 (pK2) | $[^+H_3N–CH_2–COO^-] = [H_2N–CH_2–COO^-]$ | −0.5 (Buffering Zone 2) |
| 25.0 mL NaOH | 1.0 Eq NaOH Added | 11.30 | $H_2N–CH_2–COO^-$ (Fully Deprotonated) | −1.0 (Anion) |
⚠️ Practical Precautions & Errors
- Boiled, cooled deionized water must be used to eliminate dissolved atmospheric $text{CO}_2$, which forms carbonic acid and distorts high-pH readings.
- Ensure the magnetic stir bar does not collide with the fragile glass bulb of the pH electrode.
- Wait 15–20 seconds after each titrant addition for complete ionic equilibrium before reading the meter.
💬 Viva-Voce Questions & Model Answers
Q1: What is a zwitterion and why does glycine have zero net charge at its isoelectric point?
A zwitterion (dipolar ion) is an electrically neutral molecule carrying equal numbers of positively ($-text{NH}_3^+$) and negatively ($-text{COO}^-$) charged functional groups. At pH 5.97, exactly 100% of glycine molecules carry a $+1$ charge on the amino group and a $-1$ charge on the carboxyl group, giving a net electrical charge of zero.
Q2: Why does glycine have two distinct buffer zones on its titration curve?
Glycine possesses two ionizable groups with distinct $text{p}K_a$ values. Maximum buffering capacity occurs within $text{pH} = text{p}K_a pm 1$. Thus, glycine acts as an effective buffer in the acidic range ($text{pH } 1.34–3.34$) centered at $text{p}K_1 = 2.34$, and in the alkaline range ($text{pH } 8.60–10.60$) centered at $text{p}K_2 = 9.60$. It has virtually zero buffering capacity near physiological pH 7.4.
Q3: How would the titration curve of glutamic acid differ from that of glycine?
Glutamic acid is a dicarboxylic acidic amino acid possessing three ionizable groups: $alphatext{-COOH}$ ($text{p}K_1 = 2.19$), side-chain $gammatext{-COOH}$ ($text{p}K_R = 4.25$), and $alphatext{-NH}_3^+$ ($text{p}K_2 = 9.67$). Its titration curve displays three distinct buffering plateaus and an acidic isoelectric point calculated as $text{pI} = (text{p}K_1 + text{p}K_R)/2 = (2.19 + 4.25)/2 = 3.22$.
