Comprehensive study notes: basic chemistry and biochemistry concepts (from transcript)

Catabolic vs. Anabolic Metabolism

• Catabolic reactions: large molecules are broken down; energy is released.
• Anabolic reactions: build up macromolecules; require energy; also referred to as biosynthesis.
• Relationship: anabolism uses energy produced by catabolism to assemble complex molecules.

Basic Chemistry and Biochemistry Foundations

• Chemistry naming note: oxidation and reduction are related concepts but distinct; terminology can be similar but meanings differ.
• Substances can be organized along an umbrella of chemical pathways; keep straight that pathways and names may sound similar but represent different processes.

Periodic Table: Key Elements in Biology

• Major elements commonly discussed:
  • Hydrogen (H), Carbon (C), Nitrogen (N), Oxygen (O) – used most of the time.
  • Others mentioned for examples: Sodium (Na), Potassium (K), Chloride (Cl), Calcium (Ca).
• Elements form atoms, which are the smallest units that retain properties of the element.

Atomic Structure and Subatomic Particles

• Subatomic particles: protons, neutrons, electrons.
• Protons: positive charge; located in the nucleus; each proton ~ 1 atomic mass unit (AMU).
• Neutrons: neutral charge; located in the nucleus; each neutron ~ 1 AMU.
• Electrons: negative charge; orbit the nucleus; mass is so small they are treated as ~0 AMU.
• Nucleus contains protons and neutrons (collectively called nucleons).
• Atomic mass unit (AMU): unit used to express atomic and subatomic particle masses; protons and neutrons each weigh ~1 AMU; electrons negligible by comparison.
• Electron shells: electrons are arranged in shells around the nucleus (simplified model; real behavior is more complex).

Atomic Number, Mass Number, and Isotopes

• Atomic number (Z): number of protons; shown to the left of the element symbol in most periodic tables. It defines the identity of the element.
• Mass number (A): sum of protons and neutrons; A = Z + number of neutrons.
• Isotopes: same number of protons (same Z) but different numbers of neutrons, leading to different mass numbers.
  • Example: Carbon-12 (C-12): 6 protons, 6 neutrons; ~98.9% naturally.
  • Carbon-13 (C-13): 6 protons, 7 neutrons; ~1.1%.
  • Carbon-14 (C-14): 6 protons, 8 neutrons; <0.1%; radioisotope (unstable, emits radiation).
• The average atomic mass reflects the weighted average of isotopes; e.g., carbon’s average mass ~$12.01$, due to isotopic abundances.
• Relationship: mass number A minus atomic number Z gives the number of neutrons: $N = A - Z.$
• In all discussion, protons and neutrons contribute to atomic mass; electrons contribute negligible AMU.

From Atoms to Molecules: Bonds and Bonding

• Molecule: chemical unit formed when two or more atoms are bonded together.
• Compound: substance composed of two or more different elements bonded together.
• Ionic bonding: attraction between oppositely charged ions (cations and anions) due to transfer of electrons.
  • Example: Sodium (Na) donates an electron to Chlorine (Cl) → Na⁺ and Cl⁻; forms NaCl.
  • Ionic bonds are moderately strong and can be broken in water (electrolytes).
  • Ions: positively charged (cation) or negatively charged (anion).
• Covalent bonding: atoms share electrons to achieve stable valence shells; typically stronger than ionic bonds.
  • Single bond: one pair of electrons shared.
  • Double bond: two pairs of electrons shared.
  • Triple bond: three pairs of electrons shared (e.g., nitrogen triple bond in N₂).
  • Polar covalent bonds: unequal sharing of electrons; results in partial charges (one atom slightly negative, the other slightly positive).
  • Example lines: H–O in H₂O is a polar covalent bond; carbon can form multiple covalent bonds to form diverse organic structures.
• Molecular representations: bonds are often shown as lines; single bonds as one line, double bonds as two lines, etc.

Water, Solutions, and the Role of Electrolytes

• Water is a universal solvent; many biological solutes dissolve in water.
• Solute vs solvent: solvent is the dissolving medium (water here); solute is the dissolved substance (e.g., NaCl).
• Hydrophilic substances attract water; hydrophobic substances resist water (e.g., many lipids).
• Electrolytes: substances that dissociate into ions in solution (e.g., salts like NaCl) and are essential for nervous system function and fluid balance.
• Rehydration and electrolyte balance:
  • Sweating depletes electrolytes; fluids with electrolytes help replenish them.
  • Excess electrolytes without adequate water can cause problems; balance with water is important.

Hydrogen Bonding and Water Properties

• Hydrogen bonds: weak interactions between a hydrogen atom bonded to a highly electronegative atom (O, N, or F) and another electronegative atom.
• Water’s hydrogen bonding leads to important properties:
  • Cohesion: water molecules cling to one another, enabling surface tension and droplet formation.
  • Adhesion: water molecules cling to surfaces.
  • Water is polar, enabling it to dissolve many substances (solvent variability).
  • Ice is less dense than liquid water due to open hydrogen-bonded lattice; ice floats on water.
• Water’s heat properties:
  • High heat capacity: resists temperature change; large amounts of energy required to raise its temperature.
  • Heat of vaporization: about $540 ext{ cal/g}$ to vaporize a gram of water; this is why sweating can cool the body as water evaporates.
  • Specific calories in foods: calories as a unit of energy; one gram of water requires $1 ext{ cal}$ to raise by $1^ ext{o}C$; other liquids require less energy.
• Ocean and climate relevance: the ocean stores and slowly releases heat, moderating coastal temperatures.

Dehydration Synthesis vs. Hydrolysis; Enzymes

• Dehydration synthesis (condensation): join monomers by removing water (H₂O) to form a covalent bond; anabolic process.
  • Example: two monomers combining to form a di-mer with removal of water; creates larger macromolecules.
• Hydrolysis: add water to break bonds; catabolic process.
  • One H from water attaches to one monomer, the remaining OH attaches to the other.
• Enzymatic hydrolysis: hydrolysis aided by enzymes; speeds up breakdown of macromolecules.

Inorganic vs. Organic Compounds

• Inorganic compounds: do not contain a carbon-hydrogen (C–H) framework in many molecules; tend to be small and often form ionic bonds.
  • Examples: water (H₂O), molecular oxygen (O₂), carbon dioxide (CO₂), salts, acids, bases.
• Organic compounds: contain carbon and hydrogen; often covalent bonds; include carbohydrates, lipids, proteins, nucleic acids.
• Hydrocarbons: skeletons composed of carbon and hydrogen; form the backbone of many organic molecules and their functional groups.

Organic Macromolecules: Monomers and Polymers

• Four major biological macromolecules: carbohydrates, lipids, proteins, nucleic acids.
• Monomers vs polymers:
  • Carbohydrates: monomers are monosaccharides (e.g., glucose, fructose, galactose); polymers include disaccharides and polysaccharides.
  • Lipids: not polymers in the same way as the above, but can form triglycerides and other lipid structures; energy-dense storage and membrane components.
  • Proteins: monomers are amino acids; polymers are proteins; peptide bonds connect amino acids.
  • Nucleic acids: monomers are nucleotides; polymers are DNA or RNA.
• Note: nucleic acids are polymers; nucleotides are their monomers.

Carbohydrates: Structures, Names, and Roles

• General formula: $CH<em>2O$; when repeated six times: $C</em>6H<em>{12}O</em>6$ (glucose-type formula).
• Monosaccharides: glucose, fructose, galactose – energy sources; body’s preferred energy is glucose.
• Fructose: found in fruits; chemically transformed in metabolism; common form in high-fructose corn syrup (HFCS).
• Disaccharides:
  • Maltose: glucose + glucose.
  • Sucrose: glucose + fructose.
  • Lactose: glucose + galactose.
  • Some sugar configurations (e.g., “superlose”) may not be readily absorbed by the body.
• Polysaccharides: long chains of glucose units; examples include starch (plants), glycogen (animal storage), cellulose (fiber).
• Digestive/biological roles:
  • Energy source and structural components.
  • Starch and glycogen store energy; cellulose is indigestible to humans but serves as dietary fiber.
• Fiber and gut microbiota:
  • Humans cannot digest cellulose; gut bacteria ferment some cellulose, contributing to nutrient availability and gut health.
  • Gut bacteria synthesize vitamin K; newborns receive a vitamin K shot because babies have limited gut bacteria at birth.
• Practical notes:
  • Vitamin K can be synthesized by gut flora; infants often receive a vitamin K injection at birth to prevent deficiency.
• Metabolic notes:
  • Excess glucose is stored as glycogen (in liver and muscles) or converted to fat when storage is full.

Lipids: Structure, Function, and Energy

• Lipids are hydrophobic (mostly nonpolar C–H bonds); dissolve poorly in water.
• Functions and forms:
  • Membrane structure: phospholipid bilayers form cellular membranes; lipids also in organelle membranes (mitochondria, nucleus, etc.).
  • Energy storage: fats (triglycerides) store a large amount of energy; about $9 ext{ kcal/g}$, roughly 3–4 times energy density of carbohydrates per gram.
  • Insulation and padding: fats cushion and insulate organs.
  • Signaling molecules and membranes regulation.
• Triglycerides: fats (solid at room temperature) versus oils (liquid at room temperature).
  • Examples: butter is solid; corn oil, soybean oil, canola oil are liquid.
• Emulsification and micelles:
  • Emulsification increases fat digestion; bile salts from the liver/gallbladder emulsify fats in the small intestine.
  • Micelles are spherical lipid assemblies formed in aqueous environments to solubilize lipids for absorption.
• Bile and emulsification: liver produces bile stored in the gallbladder; bile is released into the duodenum to aid fat digestion.
• Other lipid-related topics mentioned:
  • Waxes act as protective barriers (e.g., earwax) to reduce moisture loss and protect against infection.
  • Emulsification involves polar molecules organizing at oil-water interfaces to stabilize droplets.

Proteins and Nucleic Acids: Building Blocks and Polymers

• Proteins:
  • Monomer: amino acid; general structure includes an amino group, a carboxyl group, and an R group (side chain) attached to a central carbon.
  • Polymers formed by peptide bonds linking amino acids; sequence determines protein structure and function.
  • The term peptide is often associated with protein-related sequences or fragments.
• Nucleic acids:
  • Monomer: nucleotide; polymer: nucleic acid (DNA or RNA).
  • Nucleotides consist of a sugar, a phosphate group, and a nitrogenous base.
• Carbohydrates, lipids, proteins, nucleic acids are the major macromolecules that constitute living systems and their polymers.

Water-Biochemistry-Physiology Connections and Real-World Context

• Water as a solvent, reactant, and medium for biochemical reactions (e.g., dehydration synthesis and hydrolysis).
• pH and buffers:
  • pH scale runs from 0 to 14; acidic solutions have pH < 7; basic (alkaline) solutions have pH > 7; neutral water is pH ~ 7.
  • Acids release hydrogen ions (H⁺) in water; bases release OH⁻ or accept H⁺.
  • The human body maintains pH homeostasis using buffers (e.g., blood buffers); stomach has its own buffering in digestion.
  • Common example: stomach acid (gastric acid) has pH around $pH \,\approx \,2$.
• Practical implications:
  • Consuming highly basic water can raise stomach pH temporarily; body buffers respond to restore pH, but excessive alkalinity can disrupt digestion.
  • Household bleach and some other oxidizers are caustic and dangerous if ingested; do not rely on them as a pH-balancing strategy.
• pH and physiology continue to be important in physiology courses; blood is slightly basic; buffers prevent large pH shifts.
• Environmental considerations: acid rain from atmospheric pollutants can alter natural buffering systems; buffers help mitigate pH changes in the environment and body fluids.

Quick Recap: Key Concepts and Connections

• Energy metabolism: catabolism releases energy; anabolism builds macromolecules using that energy.
• Atomic structure basics: Z (protons), A (mass number), N (neutrons); isotopes affect atomic mass; carbon isotopes illustrate natural abundance and radioactivity.
• Bonding: ionic vs covalent; polar covalent bonds introduce partial charges; hydrogen bonding explains many properties of water and molecular interactions.
• Water properties drive biology: solvent capability, cohesion/adhesion, high heat capacity, and high heat of vaporization underpin thermoregulation and metabolism.
• Macromolecules: carbohydrates, lipids, proteins, and nucleic acids form the basis of life; monomers and polymers define their assembly.
• Digestion chemistry: dehydration synthesis builds; hydrolysis breaks; enzymes accelerate hydrolysis.
• Inorganic vs. organic chemistry: carbon-containing compounds (organic) dominate biology; water, salts, acids, bases illustrate inorganic chemistry.
• Practical biology connections: glucose as primary energy source; starch/glycogen store energy; cellulose/fiber supports digestion via gut microbiota; gut bacteria synthesize vitamin K.
• Lipids provide energy storage, membranes, and protection; emulsification and micelles are essential for fat digestion; waxes protect moisture.
• pH health implications: balance of acids, bases, and buffers maintains physiological homeostasis; improper pH can influence digestion and overall health.

Key Equations and Values (LaTeX)

• Mass number and neutrons: $N = A - Z$
• General carbohydrate formula: $CH<em>2O$; hexose example: $C</em>6H<em>{12}O</em>6$
• Energy densities (approximate): fats $\approx 9\ \text{kcal/g}$; carbohydrates $\approx 4\ \text{kcal/g}$ (contextual note: typical values discussed in class).
• Water vaporization energy: $540\ \text{cal/g}$
• pH concept: $pH\in[0,14],\ pH<7\rightarrow\text{acidic},\ pH>7\rightarrow\text{basic},\ pH=7\rightarrow\text{neutral}$
• Stomach pH: $pH\approx 2$
• Carbonic-isotope example masses: $^{12}\text{C},\^{13}\text{C},\^{14}\text{C}$; typical natural abundances: $^{12}\text{C}:\approx98.9\%,$ $^{13}\text{C}:\approx1.1\%,$ ^{14}\text{C}:<0.1\%.