General Biochemistry

• Bioquimica General (Quim 4065) is taught by Professor Vibha Bansal, Ph.D.
• Class hours are Monday and Wednesday from 9:00 - 10:20 AM.
• The course is worth 3 credits.

Biochemistry

• Biochemistry: the branch of organic chemistry focused on matter within living cells known as protoplasm.
• Protoplasm: a complex organic compound mixture with high chemical activity levels.
• Biochemistry seeks to understand how lifeless molecules combine to exhibit life's attributes.

History of Biochemistry

• A history of biochemistry can be found at the link provided and is 6 minutes and 31 seconds long.

Future of Biochemistry

• Biochemistry majors can pursue careers such as college professor, biochemist, pharmaceutical chemist.

Course Objectives

• Course will cover major biological molecules' structures.
• Course will cover biochemical processes in cells like respiration.
• Biochemistry is divided into conformational and informational studies.
• Conformational studies: the study of structure and three-dimensional arrangements of biomolecules.
• Informational studies: the study of language for communication inside and between cells.

Biological Molecules

• Carbohydrates provide energy.
• Sugars and starches are examples of carbohydrates.
• Lipids store energy.
• Fats and oils are examples of Lipids.
• Proteins build cells.
• Amino acids are examples of Proteins.
• Nucleic acids build genetic codes.
• Nucleotides are examples of Nucleic acids.
• DNA comprises bases and pentose sugars.
• RNA comprises bases and pentose sugars.
• Cellulose comes from plants.
• Saturated fats come from animals.
• Unsaturated fats are used to form membranes.
• The peptide bonds link proteins.

Elements in Living Systems

• Elements are divided into macronutrients, trace essential and trace possibly essential.

Common Functional Groups in Biochemistry

• Acyl: C=OR
• Amido: CONH2
• Amino: NH2
• Carbonyl: C=O
• Carboxyl: COOH
• Diphosphoryl: P2O7H4
• Ester: COOR
• Ether: ROR
• Hydroxyl: OH
• Imino: C=NH
• Phosphoryl: PO4H3
• Sulfhydryl: SH

Bioquimica General (Q4065)

• Lectures 1 and 2 cover the topic of Water as the most vital nutrient.
• Readings for lectures can be found in the Textbook Chapter 2, pages 47-74

Vital Role of Water

• The body can lose almost all of its fat.
• The body can lose over half of its protein and survive.
• A 10-15% water loss leads to death.

Importance of Water

• Water molecules display attractive forces through hydrogen bonds.
• Water's solvent properties influence the strength and specificity of bimolecular interactions.
• Water has a slight tendency to ionize.
• Aqueous solutions of weak acids or bases and their salts act as buffers.

Properties of Water via Hydrogen Bonding

• Hydrogen bonding gives water a high melting point, boiling point, and heat of vaporization.

Water as a Solvent

• Water, a polar compound, easily dissolves charged biomolecules like glucose and polar amino acids.
• Nonpolar molecules in water cause water molecules to form a highly ordered cage-like shell.
• Amphipathic compounds form stable structures called 'micelles' in water, including proteins, phospholipids, pigments, sterol, and certain vitamins.

Ionization of Water

• Pure water slightly ionizes; only 2 out of every 10⁹ molecules at 25°C.
• Hydrogen bonding allows for virtually instantaneous hydration of dissociating proteins.

Equilibrium Constant for Ionization of Water

• Ionization can be expressed as an equilibrium constant, measurable by electrical conductivity.

pH Scale

• The pH scale designates H⁺ and OH⁻ concentrations
• Examples of pH values: battery acid (0), gastric fluid (1), lemon juice (2), vinegar (3), beer (5), milk (6), blood (7), and sea water (8)

H+ and OH- concentrations

• In an aqueous solution with H+ concentration of 1 x 10⁻⁶ M, the OH- concentration is 1 x 10⁻⁸ M.

Biological Systems and pH

• pH affects the solubility of many substances.
• pH affects the structure and function of most proteins, including enzymes.
• Many cells and organisms can only survive in a specific pH environment.
• pH is temperature dependent.

Terms

• Terms to define: strong acid, strong base, weak acid, and weak base.
• 1 M HCl has a pH = 0, while 1 M acetic acid only has a pH = 2.4.

Weak Acids, Bases, and Dissociation

• Weak acids and bases have characteristic dissociation constants
• pKa indicates the relative strength of a weak acid or base.

Conjugate Acid-Base Pairs

• Monoprotic acids: Acetic acid (pKa = 4.76), Ammonium (pKa = 9.25)
• Diprotic acids: Carbonic acid (pKₐ = 3.77). Glycine, carboxyl (pKₐ = 2.34)
• Triprotic acids: Phosphoric acid (pKa = 2.14)

Titration Curves

• Titration curves reveal the pKa of weak acids.
• Titration curves can be used for multiple weak acids.

Henderson-Hasselbalch Equation

• Lawrence Joseph Henderson wrote an equation in 1908 to describe the use of Carbonic Acid as a Buffer Solution.
• Karl Albert Hasselbalch re-expressed that formula in logarithmic terms.
• The Henderson-Hasselbalch is used for calculation of pH or relative concentrations of conjugate acid base pairs in Buffered solutions.
• pH = pKa + log([A-]/[HA])

Types of Buffers

• Acidic Buffers: Weak acid and its conjugate base (CH3COOH and CH3COONa)
• Basic Buffers: Weak base and its conjugate acid (NH3 and NH4+)

Acid Buffer

• Acid Buffer (Acetate Buffer): CH3COOH ⇌ H+ + CH3COO¯
• If more H+ is added to this solution, it shifts the equilibrium to the left, absorbing H+, so the [H+] remains unchanged.
• If H+ is removed (e.g. by adding OH-) then the equilibrium shifts to the right, releasing H+ to keep the pH constant

Weak Acid and Conjugate Base

• The weak acid and its conjugate base is needed to be present in significant quantities in order to act like a buffer.

Basic Buffer

• Basic Buffer (Ammonia Buffer): NH4OH ⇌ OH- + NH4+
• If more OH- is added to this solution, it shifts the equilibrium to the left, absorbing OH-, so the [OH-] remains unchanged.
• If OH- is removed (e.g. by adding H+) then the equilibrium shifts to the right, releasing OH-to keep the pH constant

Common Buffering Agents

• Citric Acid: pKa 3.13, 4.76, 6.40, Useful pH range 2.1-7.4
• Acetic Acid: pKa 4.76, Useful pH range 3.8-5.8
• KH₂PO₄: pKa 6.86, Useful pH range 5.9-7.9
• Tris: pKa 8.06, Useful pH range 7.0-9.0
• Ammonium Ion: pKa 9.25, Useful pH range 8.3-10.3

Change in pH Exercises

1. The pKa of acetate buffer is 4.8. Calculate the volume of glacial Acetic Acid (17.6 M) and the mass of Sodium Acetate (MM = 82.0 g/mol) needed to prepare 50 mL of a 0.05 M buffer at pH 4.4.
2. What is the pH of a buffer containing 0.25 M NH3 and 0.45 M NH4Cl?
3. By how much will the pH change if 0.050 mol HCl is added to 1.00 L of the buffer in the previous problem?

Buffering Capacity (β)

• Buffering Capacity (β) refers to the ability of buffer to resist change in pH.
• It is also known as buffer efficiency, buffer index, and buffer value.
• Defined as the number of moles of strong monoprotic acid or base required to produce a one pH unit change in 1.0 L of solution.
• The maximum amount of strong acid or strong base that can be added before the pH changes significantly determines it.

Buffer Capacity

• Buffer Capacity = (number of moles of OH⁻ or H₃O⁺ added) / (ΔpH) (volume of buffer in L)
• It depends on the ionic concentration of buffer.
• It depends on the proximity of pKa to the pH of the solution.
• Buffer concentrations between 0.05 and 0.5 M, and buffer capacities between 0.01 to 0.1 are usually sufficient for pharmaceutical solutions.

Biological Buffers

• Cytoplasm of most cells contains high concentrations of nucleotides and proteins.
• The buffer system (H₂PO₄⁻/HPO₄²⁻) acts in cytoplasm of all cells, pKa = 6.86.
• Weak acids/bases such as Histidine are functional groups of nucleotides and proteins.
• Blood plasma is buffered by the bicarbonate system (H₂CO₃/HCO₃⁻) due to the CO₂ reservoir in the lungs.