Buffers CHEM 41 Intarmed 2030 PDF

Buffers CHEM 41 INTARMED 2030 | Prof. Cruz/Genato/Lim | LU2 SEM 1 | SY. 2024-2025 Given that HA is a strong acid, the reaction would be:...

Buffers CHEM 41 INTARMED 2030 | Prof. Cruz/Genato/Lim | LU2 SEM 1 | SY. 2024-2025 Given that HA is a strong acid, the reaction would be: HA(aq)+H2O(l) → H3O+(aq)+A−(aq) TABLE OF CONTENTS I. Acids and Bases In the reaction, H2O acts as a base that reacts with II. pH acid HA. The conjugate base of acid HA is A− while the A. Ka (Acid Dissociation Constant) conjugate acid of H2O is the H3O+. B. pKa The ionization of a strong acid — HCl for this C. Henderson-Hasselbalch Equation particular example — looks like this: III. Buffer Preparation HCl → H++Cl− IV. Titration Curves a. Titration Setup On the other hand, the ionization of a strong base — b. Parts of a Titration Curve KOH for this particular example — looks like this: c. General Trends KOH → K++OH− The “→” signifies complete dissociation I. ACIDS AND BASES Table 2. Examples of strong acids and bases DEFINITION ______________________________________________________ Strong Acids Strong Bases Chemicals that interact with one another — forming salt and H2O Hydrochloric Acid Lithium Hydroxide Can be classified/defined in three ways: (HCl) (LiOH) ○ Arrhenius definition Focus: increase in the concentration of a Hydrobromic Acid Sodium Hydroxide specific ion in an aqueous solution (HBr) (NaOH) ○ Bronsted-Lowry definition Hydroiodic Acid Potassium Hydroxide Focus: capability of a species to (HI) (KOH) accept/donate a proton (H+) Applicable to solvents other than H2O ○ Lewis definition WEAK ACIDS AND BASES ______________________________________________________ Focus: capability of a species to accept/donate an electron pair Only partial ionization/dissociation into its electrolyte The most general definition as it is components in aqueous solution since some initial applicable to species that do not have H+ acid/base does not dissociate Conjugate acid-base pair is observed Table 1. Summary for the definitions of acids and bases Definition Acids Bases The reaction below shows the dissociation of acetic acid (CH3COOH): Arrhenius Increased Increased CH3COOH + H2O ⇌ H3O+ + CH3COO− [H3O+] in [OH-] in ○ Weak acid: CH3COOH aqueous aqueous ○ Conjugate base: CH3COO− (one H removed) solution solution On the other hand, the reaction below shows the Bronsted-Lowry* H+ donor H+ acceptor dissociation of ammonia (NH3): NH3 + H2O ⇌ OH− + NH4+ Lewis Electron pair Electron pair ○ Weak base: NH3 acceptor donor ○ Conjugate acid: NH4+ (one H added) *For Chem 41, focus on the Bronsted-Lowry definition. The “⇌” signifies reaction reversibility and equilibrium Organic acids and bases are weak acids and bases STRONG ACIDS AND BASES ______________________________________________________ 100% ionization/dissociation into its electrolyte components in aqueous solution CHEM 41 LU 2 SEM 1 | IMED 2030 Page 1 of n CUADRO, JGT; NUCUP, ZJP; PAZ, ERB; ANG, EMS; BALASANOS, KMB; MERCADO, AIS; REYES, JCDC Buffers CHEM 41 INTARMED 2030 | Prof. Cruz/Genato/Lim | LU2 SEM 1 | SY. 2024-2025 Table 3. Examples of weak acids and bases + − [𝐻3𝑂 ][𝐴 ] 𝐾𝑎 = [𝐻𝐴] Weak Acids Weak Bases Higher Ka Formic Acid Ammonia ○ Higher tendency of weak acid to dissociate (HCOOH) (NH3) ○ Relatively more acidic Lower Ka Acetic Acid Trimethyl Ammonia ○ Lower tendency of weak acid to dissociate (CH3COOH) N(CH3)3 ○ Relatively less acidic Hydrocyanic Acid Pyridine (HCN) (C5H5N) pK ______________________________________________________ a The negative logarithm of Ka II. pH More convenient way to express acidity 𝑝𝐾𝑎 = − 𝑙𝑜𝑔(𝐾𝑎) Quantitative measurement of the acidity or basicity of a substance or + − [𝐻3𝑂 ][𝐴 ] Value ranges from 0–14 𝑝𝐾𝑎 = − 𝑙𝑜𝑔( ) [𝐻𝐴] Resulting pH value is dependent on the concentration of H+ in the aqueous solution since: Higher pKa pH = −log[H+] ○ Relatively less acidic Lower pKa Conversely: ○ Relatively more acidic [H+] = 10-pH Henderson-Hasselbalch Equation ______________________________________________________ Higher [H+] = Lower pH = more acidic Figure 1. The pH scale Acidic solutions: pH < 7 Neutral solutions: pH = 7 Basic solutions: pH > 7 pH of a solution can be qualitatively determined by using a litmus paper ○ Acidic substances: blue litmus paper turns red Figure 2. Derivation of ○ Basic substances: red litmus paper turns blue Henderson-Hasselbalch Equation A pH meter can also be used to quantitatively measure the actual pH of the solution Allows the pH to be computed given that the values for pKa, weak acid concentration [HA], and conjugate K a (Acid Dissociation Constant) ______________________________________________________ base concentration [A−] are already determined − [𝐴 ] 𝑝𝐻 = 𝑝𝐾𝑎 + 𝑙𝑜𝑔( [𝐻𝐴] ) Quantitative measurement of how strong an acid tends to dissociate into its components in an aqueous solution III. Buffer Preparation Given the reaction HA(aq)+H2O(l) ⇌ H3O+(aq)+A−(aq) : CHEM 41 LU 2 SEM 1 | IMED 2030 Page 2 of n CUADRO, JGT; NUCUP, ZJP; PAZ, ERB; ANG, EMS; BALASANOS, KMB; MERCADO, AIS; REYES, JCDC Buffers CHEM 41 INTARMED 2030 | Prof. Cruz/Genato/Lim | LU2 SEM 1 | SY. 2024-2025 1 pH unit above and below the midpoint pH of 4.76. DEFINITION ______________________________________________________ In this zone, a given amount of H+ or OH- added to the system has much less effect on pH than the same Buffers are mixtures of equal concentrations of amount added outside the zone. weak acids (proton donor) and its conjugate base Buffering power of the system is maximal at (proton acceptor) midpoint of the buffering region These are aqueous systems that tend to resist ○ pH changes least on addition of H+ or OH changes in pH when small amounts of acid (H+) - or base (OH-) are added. RELEVANCE IN BIOLOGY ______________________________________________________ Almost every biological process is pH dependent; a PROCESS OF BUFFERING small change in pH produces a large change in the rate of the process. Buffering results from two reversible reaction ○ The enzymes that catalyze cellular reactions equilibria occurring in a solution of nearly equal contain ionizable groups with characteristic pKa concentrations of a proton donor and its conjugate values. proton acceptor. Biological buffers regulate the constancy of pH in biological systems ○ Example: Cells maintain a constant cytosolic pH near pH 7 to keep biomolecules in their optimal ionic state IMPORTANT TO NOTE ______________________________________________________ pH of the buffer system does change slightly when a small amount of H+ or OH- is added, but this change is very small compared with the pH change that would result if the same amount of H+ or OH- were added to pure water or to a solution of the salt of a strong acid and strong base without buffering power. Figure X. The acetic acid-acetate pair as a buffer system Whenever H+ or OH- is added to a buffer, the result is Titration Curves and Buffers a small change in the ratio of the relative concentrations of the weak acid and its anion and thus a small change in the pH. ○ Decrease in concentration of a component is balanced by an increase in the other The sum of the buffer components does not change, only their ratio Buffering action is simply the consequence of two reversible reactions taking place simultaneously and reaching their points of equilibrium as governed by the equilibrium constants, KW and Ka Figure X. The titration curve of acetic acid Buffering region of the acetic acid-acetate buffer pair IV. Titration Curves is defined by the relatively flat zone extending about CHEM 41 LU 2 SEM 1 | IMED 2030 Page 3 of n CUADRO, JGT; NUCUP, ZJP; PAZ, ERB; ANG, EMS; BALASANOS, KMB; MERCADO, AIS; REYES, JCDC Buffers CHEM 41 INTARMED 2030 | Prof. Cruz/Genato/Lim | LU2 SEM 1 | SY. 2024-2025 Graphs that show the change in pH of the solution with respect to the amount of titrant added Titration Setup ○ Burette: Holds the titrant (acid or base) and allows controlled addition to the solution. ○ Erlenmeyer Flask: Contains the analyte, the solution of unknown concentration. ○ Analyte: The substance being titrated, often an acid or base of unknown concentration. ○ Titrant: A solution of known concentration (acid or base) added to the analyte to reach the equivalence point. ○ Indicator Solution (Optional): A chemical added to the analyte that changes color near the equivalence point. ○ pH Meter or pH Probe: Monitors real-time pH changes as the titration proceeds. ○ Magnetic Stirrer: Ensures the analyte and titrant are well mixed during the titration process. ○ Distilled Water (Optional): Added to dilute the analyte or titrant to improve accuracy and visibility of changes. General Trend Parts of a Titration Curve MONOPROTIC TITRATION ______________________________________________________ 1. Buffering Region: ○ The flat part of the curve where the solution resists changes in pH ○ Occurs when the solution contains a significant amount of both weak acid and its conjugate base 2. Equivalence Point/Stoichiometric Point: ○ The point where the number of moles of titrant equals the number of moles of analyte ○ Sharp rise (or drop) in pH, depending on whether titrating an acid or base ○ solution only contains salt and water Strong acid titrated with strong base: 3. Half-Equivalence Point: ○ S-shaped curve ○ Occurs when half of the analyte has been ○ abrupt increase in pH once the equivalent point neutralized is reached ○ pH equals the pKa of the weak acid or pKb of the Strong base titrated with strong acid: weak base ○ S-shaped curve ○ abrupt decrease in pH once the equivalent point is reached EXAMPLE OF TITRATION CURVE ______________________________________________________ CHEM 41 LU 2 SEM 1 | IMED 2030 Page 4 of n CUADRO, JGT; NUCUP, ZJP; PAZ, ERB; ANG, EMS; BALASANOS, KMB; MERCADO, AIS; REYES, JCDC Buffers CHEM 41 INTARMED 2030 | Prof. Cruz/Genato/Lim | LU2 SEM 1 | SY. 2024-2025 Weak acid titrated with strong base: ○ Curve Shape: Gradual increase in pH as conjugate base forms. ○ Buffer Region: A buffer is formed before the equivalence point due to the weak acid and its Example: titration of the triprotic acid H3PO4 with NaOH conjugate base. ○ Equivalence Point: Occurs at pH >7, as the LAB ACTIVITY 1 strong base dominates after neutralizing the weak acid. PROCEDURE ______________________________________________________ Weak base titrated with strong acid: 1. Wash all glassware with soap and tap water. Then ○ Curve Shape: Gradual decrease in pH due to the rinse with distilled water thrice. formation of a conjugate acid. 2. Refer to the pKa table for calculating the quantity ○ Buffer Region: A buffer forms before the of reagents needed for your buffers. equivalence point as the weak base is converted 3. Consider the individual concentrations of salt and to its conjugate acid. acid needed to obtain a suitable buffer capacity. ○ Equivalence Point: Occurs at pH