Introduction to Biochemistry: Course Overview and Foundations

Course Introduction and Instructor Information

This course provides an introduction to the fundamental principles of biochemistry, focusing on the chemistry of life and the molecular processes that govern living organisms. The instructor is Dr. Ryan Myers ([e-mail]), with office hours available MWF 3:30-4:30 and WF 1-2pm by appointment.

Instructor Background

Dr. Myers has a research background in:

• Ion Pairing: Study of interactions between divalent metal cations and phosphate groups, and the influence of cations on water structure around functional groups.

• Polymer Hydrophobicity: Investigation of hydrophobic and hydrophilic interactions in polymers and their impact on structure and function.

• Coil Structure: Research on the behavior of ammonium ions and their effect on polymer structure.

Course Structure and Grading

• Exams: 30% (two exams, 15% each)

• Cumulative Final Exam: 20%

• Project: 20%

• Homework/Quizzes: 10%

• Notebook: 10%

• Attendance and Participation: 10%

Exams will cover material from lectures and assigned readings. Homework and quizzes are designed to reinforce understanding and track progress. A handwritten notebook defining all bolded terms in each chapter is required. Participation includes attendance, engagement in class discussions, and presentations.

Project Requirements

• Choose an interesting chemistry question (e.g., "Does aspartame impact the body the same as natural sugar?")

• Submit a proposal by September 8th

• Write a 4-5 page paper (double spaced, 12 pt Times New Roman)

• Discuss relevant scientific literature

• Present findings in a 15-minute presentation (+5 minutes for questions)

Textbook

• Biochemistry, 9th edition by Campbell and Farrell

Learning Objectives (Chapters 1 & 2)

• Define biochemistry and its scope

• Identify and name biological functional groups

• Discuss the thermodynamic possibility of life

• Explain the essential role of water in life on Earth

Foundations of Biochemistry

What is Biochemistry?

Biochemistry is the study of the chemical processes and substances that occur within living organisms. It bridges biology and chemistry, focusing on the molecular mechanisms that underlie life.

• Key Biomolecules: Proteins, nucleic acids, carbohydrates, and lipids

• Life Processes: All life is regulated by similar biomolecules, following the rules of chemistry and physics

• Molecular Pathways: Life follows specific, studyable pathways at the molecular level

Chemical Foundations

• Organic Chemistry: The study of carbon-containing compounds and their derivatives

• Functional Groups: Specific groups of atoms within molecules that determine their chemical properties and reactions (e.g., hydroxyl, carboxyl, amino, phosphate)

Origins of Life

• After the Big Bang, the universe contained mainly hydrogen, helium, and lithium

• Heavier elements formed through stellar processes

• Early Earth's atmosphere: NH3, H2S, CO, CO2, H2, N2, H2O

• Miller-Urey Experiment: Demonstrated that organic molecules (e.g., amino acids) could form under prebiotic conditions

Thermodynamics in Biochemistry

Basic Principles

• First Law: Energy cannot be created or destroyed, only transformed

• Second Law: The entropy (disorder) of the universe always increases

• Life is thermodynamically possible because Earth is not a closed system; local order can increase as long as overall disorder increases

Gibbs Free Energy

• Predicts spontaneity of reactions based on enthalpy and entropy changes

$\Delta G = \Delta H - T\Delta S$

• Spontaneous (exergonic): $\Delta G < 0$ (energy released)

• Non-spontaneous (endergonic): $\Delta G > 0$ (energy required)

• Equilibrium: $\Delta G = 0$

Water: The Solvent of Life

Properties of Water

• Principal component of most cells

• Polar molecule: uneven distribution of charge, leading to hydrogen bonding

• Excellent solvent for polar and ionic substances

Hydrogen Bonding

• Occurs when a hydrogen atom covalently bonded to an electronegative atom (e.g., O, N, F) interacts with another electronegative atom

• Responsible for many of water's unique properties

Acid-Base Chemistry in Biology

Definitions

• Brønsted Acid: Proton donor

• Brønsted Base: Proton acceptor

• Strong Acids: Completely dissociate in water

• Weak Acids: Partially dissociate

Acid Ionization Constant

• Measures acid strength

$K_a = \frac{[H^+][A^-]}{[HA]}$

• Larger $K_a$ indicates a stronger acid

• $pK_a = -\log_{10} K_a$

pH and Autoionization of Water

• pH measures the concentration of H+ ions in solution

$pH = -\log_{10} [H^+]$

• At 25°C, pH < 7 is acidic, pH = 7 is neutral, pH > 7 is basic

• Autoionization: $2H_2O \rightleftharpoons H_3O^+ + OH^-$

Henderson-Hasselbalch Equation

• Relates pH, pKa, and the ratio of conjugate base to acid

$pH = pK_a + \log \left( \frac{[A^-]}{[HA]} \right)$

Titration Curves and Buffers

• Buffers resist changes in pH by using a weak acid and its conjugate base

• Buffering range is typically ±1 pH unit from the pKa

• Biological buffers include phosphate (intracellular) and carbonic acid (extracellular)

Buffer Systems in the Human Body

Carbonic Acid Buffer System

• Maintains blood pH through the equilibrium:

$CO_2 (g) \rightleftharpoons CO_2 (aq)$ $CO_2 (aq) + H_2O (l) \rightleftharpoons H_2CO_3 (aq)$ $H_2CO_3 (aq) \rightleftharpoons H^+ (aq) + HCO_3^- (aq)$ Net: $CO_2 (g) + H_2O (l) \rightleftharpoons H^+ (aq) + HCO_3^- (aq)$

• pH of blood is calculated as:

$pH = 6.1 + \log \left( \frac{[HCO_3^-]}{0.03 \times pCO_2} \right)$

Acid-Base Disorders

• Respiratory Acidosis: Caused by hypoventilation or decreased lung function

• Respiratory Alkalosis: Caused by hyperventilation

• Metabolic Acidosis: Increased acid production or decreased acid excretion

• Metabolic Alkalosis: Elevated plasma bicarbonate (e.g., excessive antacid use or vomiting)

Buffer Preparation Considerations

• Required pH

• Compatibility with experiment

• Potential for precipitation

• Availability and ease of preparation

Summary Table: Common Buffer Systems

Buffer System	Location	pKa Range	Example
Phosphate	Intracellular	6.8 - 7.4	H2PO4- / HPO42-
Carbonic Acid	Extracellular (blood)	6.1	H2CO3 / HCO3-
Protein	Intracellular	Varies	Histidine residues

Key Terms

• Biochemistry: The study of chemical processes in living organisms

• Functional Group: Specific group of atoms responsible for characteristic reactions

• Buffer: Solution that resists changes in pH

• pH: Measure of hydrogen ion concentration

• pKa: Negative logarithm of the acid dissociation constant

• Gibbs Free Energy: Thermodynamic quantity predicting reaction spontaneity