CSB421_lecture+week+3_AcidsBases_TC_+updated

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• Title: Acids, Bases, Ionisation and pKa

• Course: CSB421: Week 3 | Dr. Yaśmin Antwertinger

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• Part 1: Acids & Bases

• Course: CSB421: Week 3 | Dr. Yaśmin Antwertinger

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• Key Questions:

  • How do you define a base?

  • What is a conjugate acid and conjugate base?

  • How would you define pH?

  • What is a hydronium ion?

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• Definitions:

  • Acid: Donates H+, forming H3O+ (hydronium ion).

  • Base: Accepts H+, reducing concentration of H3O+, or increasing OH- concentration.

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• Proton Jumping in Water:

  • Hydronium ion (H3O+) migrates rapidly, switching partners at a rate of 10^12 per second.

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• Acid-Base Reaction Equation:

  • HA + H2O ↔ H3O+ + A-

  • HA = Acid (proton donor)

  • A- = Base (proton acceptor)

  • pH: A logarithmic measure of [H+] in a solution.

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• Identifying Acids and Bases:

  • Compounds dissociating in water can be classified using their chemical formulas.

    1. Acid: Proton donor.

    2. Base: Proton acceptor.

  • Organic acids and bases have specific functional groups.

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• Identifying Organic Acids:

  • Two functional groups capable of donating protons in organic molecules include phenols.

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• Identifying Organic Acids Continued:

  • Additional specifics on functional groups.

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• Identifying Organic Bases:

  • Discussion of a functional group involved in organic basicity.

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• pH and Characteristics:

  • pH Formula: pH = -log[H+]

  • Pure water has [H+] = 10^-7, thus pH = 7.

  • Acids have high [H+], pH < 7

  • Bases have low [H+], pH > 7

  • Acids produce H+ ions, increasing concentration in solutions.

  • Bases bind H+ ions, reducing solution concentration.

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• Importance of pH:

  • Affects solubility of substances.

  • Influences structure and function of proteins, including enzymes.

  • Many organisms survive only in specific pH environments.

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• Terminology Explanations:

  • Ionised: Has a charge due to H+ loss or gain.

  • Protonated: Contains a proton; can be charged or uncharged.

  • Dissociated: Not attached to proton; can be charged or uncharged.

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• Acid Dissociation Examples:

  • Acids can dissociate to yield H+ and A- (anion); A- is the conjugate base.

  • Discussions on ionisation and dissociation.

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• Conjugate Acid-Base Pairs:

  • An acid becomes a base upon donating a proton, and vice versa for bases.

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• Ionisation and Bases:

  • Bases can ionise by accepting a proton, forming a cation (BH+), the conjugate acid of the base.

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• Acids at Different pH Levels:

  • Acids dissociate more readily in alkaline pHs due to H+ shortage, driving donation of protons.

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• Bases at Different pH Levels:

  • Bases find it easier to accept protons in acidic pHs, where H+ availability is high.

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• Part 2: Ionisation, pH & pKa

• Course: CSB421: Week 3 | Dr. Yaśmin Antwertinger

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• pH vs pKa:

  • pH: General measure of free H+ in solution; depends solely on H+.

  • pKa: Specific number indicating the acidity of a particular molecule; lower pKa indicates stronger acids and their proton donation ability.

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• pKa Significance:

  • Most acids, especially organic ones, are weak and do not fully ionise.

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• pH vs pKa Summary:

  • Lower pH = stronger acid

  • Lower pKa = stronger acid (can have negative or large values).

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• Importance of pKa:

  • pKa indicates necessary pH for a drug to transfer protons, valuable in pharmacology.

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• Environment pH Levels:

  • Plasma/Skin: 7.35 – 7.45

  • Buccal Cavity: 6.2 – 7.2

  • Stomach: 1.0 – 3.0

  • Duodenum: 4.8 – 8.2

  • Jejunum & Ileum: 7.5 – 8.0

  • Colon: 7.0 – 7.5

  • Inflamed tissues: ~5.5

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• pH vs pKa Relationship:

  • Molecule's pKa determines ionisation at specific pH; acetic acid at pKa 4.76 is 50% ionised/unionised.

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• Weakly Acidic Drugs:

  • Completely unionised if pH is more than 2 units below pKa; completely ionised at pH greater than 2 units above.

  • High pH leads to maximum solubility (ionised), low pH is less ionised (better absorption).

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• Weakly Basic Drugs:

  • Opposite rules compared to weakly acidic drugs: unionisation and ionisation levels change with pH.

  • 50% ionisation occurs when pH = pKa.

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• pH vs pKa Dynamics:

  • Effects of pH on dissociation:

    • pH > pKa: Basic conditions, fewer H+, more A-

    • pH = pKa: Neutral conditions

    • pH < pKa: Acidic conditions, lots of H+, less A-

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• Applying pKa:

  • Determines when a drug will give/take a proton based on pH and demonstrates equilibrium.

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• Unionised Form Stability:

  • Above pKa, stable negative ion exists (X-) indicating strong acid characteristics.

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• pKa Equations:

  • pKa can be determined using Henderson-Hasselbalch equation considering concentration of ions.

  • pH = pKa + log[A-]/[HA] for acidic drugs.

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• Henderson-Hasselbalch Equation Explained:

  • pH influences drug ionisation and solubility, aiding in predicting drug absorption across biological membranes.

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• Example Application:

  • pH calculated from a 50:50 solution of acetic acid and acetate gives a pH of 4.76 (equilibrium).

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• Example Outcomes:

  • Variations in acid and conjugate base quantities affect the pH (e.g., increased acidity or basicity).

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• pH Effects on Acid/Base Equilibrium:

  • At different pH values, equilibrium shifts indicating more of the acid or base form.

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• Zwitterions in Biochemistry:

  • Amino acids can exist as zwitterions at neutral pH, containing both positive and negative charges.

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• Importance of Ionisation and pKa:

  • Unionised drugs cross membranes more effectively; ionised forms are more water-soluble, essential for drug administration.

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• pH and Solubility:

  • Weakly acidic drugs see unionised formation with pH decrease; weakly basic drugs with pH increase.

  • Precipitation indicates chemical instability in solutions.

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• pH Influence on Local Anaesthesia:

  • Tissue pH influences anesthetic effects based on local environmental pH conditions.

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• Part 3: Strength of Acids and Bases

• Course: CSB421: Week 3 | Dr. Yaśmin Antwertinger

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• Strength of Acids and Bases:

  • Strong acids like HClO4, 100% ionised in aqueous solution.

  • Weak acids exist as mixtures of HA, A-, and H3O+.

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• Aromatic Ring Influence on Acid Stability:

  • Stability through resonance making it a weaker acid.

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• Factors Influencing Acid Strength:

  • Stability of conjugate base affects acidity; resonance enhances stability thus increasing acid strength.

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• Basicity Explained:

  • Availability of lone pairs impacts base strength; aromatic bases are weaker than aliphatic ones due to electron availability.

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• Comparative pKa Values:

  • Diphenhydramine (pKa = 10.6) as an aliphatic amine vs. an aromatic amine (pKa = 4.6).

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• Part 4: Buffers

• Course: CSB421: Week 3 | Dr. Yaśmin Antwertinger

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• Definition of Buffers:

  • A solution that resists changes in pH, typically a mix of acid and base forms; adjusts to particular pH values.

  • Example: Blood pH ranges from 7.35-7.45.

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• Buffer Action in Acidic Conditions:

  • Ability of bone carbonate to buffer excessive H+ in an acidic environment; affects blood pH.

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• Buffer Significance:

  • Enzyme reactions and cellular functions depend on optimum pH, critical for biological materials.

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• How Buffers Work:

  • Establish equilibrium.

  • Example: Acetate buffer equilibrium shifts to maintain constant [H+].

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• Buffer Mechanisms:

  • Same as above, emphasizing adjustments to H+ addition/removal through equilibrium shifts.

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• Image Credits:

  • Acknowledgment of sources for images and relevant chemistry resources.