Chemical Composition of the Body - Chapter 2 (Flashcards)

Atomic and Molecular Foundations of the Body

• Core idea: The body is built from atoms; structure dictates function across all levels of organization.
• 2-3 statement: A 150 lb male is composed of roughly 10^28 atoms; all biological structures are atom-based. This emphasizes that life emerges from arrangement, not from a different kind of matter.
• Einstein-level perspective (Schrödinger): The events in space and time inside a living organism can be accounted for by physics and chemistry, a question that fascinated Schrödinger (1944).

Levels of Biological Organization

• Atomic level: Atoms form molecules and macromolecules.
• Subatomic level: Protons, neutrons, and electrons; fundamental particles underpin chemistry.
• Molecular level: Molecules build cellular and tissue structures.
• Cellular level: Cells are the basic units of life.
• Tissular level: Tissues are groups of cells with common function.
• Organ level: Organs consist of multiple tissues.
• Organ system level: Organ systems coordinate to sustain life.
• System level and organism level: Integrated, functioning whole.
• Key idea: From one fertilized egg, all ten organ systems emerge; all life starts from a single cell.

Common Elements and Body Composition

• Major elements in the body (percent of total body mass):
  • Oxygen (O): $65.0\%$
  • Carbon (C): $18.5\%$
  • Hydrogen (H): $9.5\%$
  • Nitrogen (N): $3.2\%$
• Significance of O, C, H, N:
  • Oxygen: Part of water; major component of organic molecules; used to generate ATP (energy storage in cells).
• Lesser elements: Ca, P, K, S, Na, Cl, Mg, Fe, etc., in smaller but essential amounts (percentages vary; see table in transcript).
• Trace elements: Aluminum (Al), Boron (B), Chromium (Cr), Cobalt (Co), Copper (Cu), Fluorine (F), Iodine (I), Manganese (Mn), Molybdenum (Mo), Selenium (Se), Silicon (Si), Tin (Sn), Vanadium (V), Zinc (Zn); required in small amounts for various biological roles.
• Essential elements note: Only about 25 elements are known to be essential for human structure and function; Hydrogen is the most abundant atom, while calcium is the most abundant mineral.
• Boron is specifically highlighted as noteworthy in discussions of essential trace elements.

Atomic Structure and Isotopes

• Atomic constituents:
  • Protons (p+), in the nucleus
  • Neutrons (n°), in the nucleus
  • Electrons (e¯), in electron shells/clouds
• Two common representations of atomic structure:
  • Electron cloud model
  • Electron shell model
• Mass number (A) and atomic number (Z):
  • Atomic number Z = number of protons
  • Mass number A = Z + N (where N = number of neutrons)
  • Atomic mass is the weighted average of stable isotopes in daltons
• Example shells (first 4 shells shown in transcript):
  • 1st shell: can hold up to 2 electrons
  • 2nd shell: can hold up to 8 electrons
  • 3rd shell and beyond: fill according to energy levels and orbital type
• Isotopes: Atoms of the same element can have different mass numbers due to different neutron numbers (e.g., ^12C, ^13C, ^14C).

Periodic Table and Electron Orbitals

• The periodic table organizes elements by recurring chemical properties, driven by electron arrangements in outer orbitals.
• Periodic trends arise from outer-shell electron configurations (block structure: s, p, d, f blocks; G-block discussed conceptually in transcript).
• Janet left-step periodic table: a variant arrangement of blocks; predictions about future elements (e.g., 119–120 with s-type outer orbitals; 121 with a new g-type block).
• Electron configurations (illustrative examples from the first 18 elements):
  • Hydrogen (Z=1): 1s^1
  • Helium (Z=2): 1s^2
  • Lithium (Z=3): [He] 2s^1
  • Carbon (Z=6): [He] 2s^2 2p^2
  • Nitrogen (Z=7): [He] 2s^2 2p^3
  • Oxygen (Z=8): [He] 2s^2 2p^4
• Orbital theory for nitrogen: each orbital holds a maximum of 2 electrons; 1s^2, 2s^2, 2p^3 (five orbitals in 2p have up to 6 electrons; nitrogen fills partially).
• First 18 elements table (electron arrangements) summarized: how electrons populate 1s, 2s, 2p, 3s, and 3p orbitals across the period 1–18; note periodicity and blocks.

Atomic Numbers, Mass Numbers, and Isotopes

• Mass number vs atomic number recap:
  • Atomic number (Z): number of protons
  • Mass number (A): number of protons + neutrons
  • Atomic mass: average mass of all stable atoms of an element in daltons
• The concept of isotopes: varieties of an element with differing neutron counts, altering mass but not chemical identity.

Isotopes in Practice and Isotope Notation

• Examples of isotopic forms used in research and medicine:
  • Deuterium (D, ^2H): two neutrons and one proton in the nucleus; heavier than protium (^1H)
  • Tritium (^3H): two neutrons and one proton plus one more neutron compared to protium
  • ^13C and ^14C used for tracing and radiometric dating; percentage abundances of carbon isotopes in biology and research contexts.
• Applications include detecting metabolic pathways, dating, and studies of biochemical processes via isotope labeling (e.g., [13C]-cyanocobalamin in vitamin B-12 studies).

Atomic Bonding and Ions

• Ionization: loss or gain of electrons to form ions; ions are charged species which may be cations (+) or anions (-).
• Cations vs Anions:
  • Cations are positively charged (loss of electrons)
  • Anions are negatively charged (gain of electrons)
• In physiology, hydrogen ions (H+) are treated as cations with positive charge.
• The concept of solubility, electrolytes, and hydration shells in aqueous biology is tied to ionization and ionic interactions.

Molecules, Compounds, and Bonding

• Molecule: two or more atoms sharing electrons via covalent bonds (e.g., O2, N2, H2O, CO2).
• Compound: substance containing atoms of two or more different elements (e.g., H2O, CO2, NaCl).
• Covalent bonds: sharing of electrons; can be nonpolar or polar depending on electronegativity differences.
• Ionic bonds: electrostatic attraction between oppositely charged ions; typically weaker than covalent bonds.
• Free radicals: atoms/molecules with unpaired electrons in the outer shell; highly reactive and can propagate oxidative damage (examples: hydroxyl radical •OH, nitric oxide •NO).
• Strength ranking of bonds (relative): Covalent > Ionic > Hydrogen > van der Waals; overall, covalent bonds are the strongest, van der Waals are the weakest.

Covalent Bonding and Electronegativity

• Bond formation depends on octet rule (outermost shell tends toward eight electrons).
• Electronegativity: measure of a atom’s ability to attract electrons in a covalent bond; fluorine (F) is the most electronegative element.
• Polar vs Nonpolar Covalent Bonds:
  • Nonpolar covalent bonds: electrons shared equally; little charge separation; often found in molecules with similar electronegativities; example: H2, O2, many hydrocarbon chains
  • Polar covalent bonds: electrons unevenly shared; results in partial charges; contribute to solubility in water and interactions with polar environments.
• Water as a polar molecule: promotes dissolution of polar/ionic substances; “like dissolves like.”
• Ionic interactions form salts like NaCl in crystals; ions can be surrounded by hydration shells in solution (solvation/hydration).
• Hydrogen bonds: weaker than covalent bonds but numerous and critical for biological structure (shape of proteins and DNA).
• Molecular geometry (VSEPR): shapes such as tetrahedral, pyramidal, and bent are dictated by electron pair repulsion around central atoms; shapes influence reactivity and interactions (e.g., protein folding).
• Phospholipids: amphipathic molecules with polar (hydrophilic) heads and nonpolar (hydrophobic) tails; essential for cell membranes; phospholipid bilayer arrangement shows polar heads facing aqueous environments and nonpolar tails forming the interior.

Macromolecules and their Building Blocks

Carbohydrates

• General role: energy storage and structure; contain carbon, hydrogen, and oxygen in specific ratios (CH2O unit common).
• Monosaccharides: glucose (C6H12O6), fructose, galactose; ribose and deoxyribose for nucleic acids (RNA and DNA).
• Monosaccharides are aldehyde or ketone sugars with multiple hydroxyl groups; ribose is the sugar in RNA; deoxyribose lacks one hydroxyl group compared to ribose.
• Disaccharides: formed by dehydration synthesis; example: sucrose (glucose + fructose) with water released; hydrolysis breaks disaccharides back into monosaccharides.
• Polysaccharides: glycogen (animal starch); starch (plant storage); cellulose (plant structural polymer; non-digestible by humans).
• Chemical formula examples:
  • Glucose: $C<em>6H</em>{12}O_6$
  • Sucrose: $C<em>{12}H</em>{22}O<em>{11} + H</em>2O$ (dehydration synthesis forms the glycosidic linkage; hydrolysis adds water)

Lipids

• Classes (as summarized in transcript): five sub-classes (not all named in excerpt, but triglycerides and phospholipids are discussed).
• General features: predominantly hydrogen and carbon; nonpolar covalent bonds; fats provide ~$9\ \text{kcal/g}$ energy.
• Triglycerides: major storage form of lipids in the body.
• Phospholipids: amphipathic; form the bilayer of cell membranes; contain a polar head (phosphate group) and nonpolar tails.
• Lipid saturation: saturated fats have all single bonds (fewer kinks; typically solid at room temperature), unsaturated fats contain one or more cis double bonds (kinks; typically liquid at room temperature).
• Structural sterols: cholesterol is a steroid with a multi-ring structure; steroids include testosterone, estradiol, and cortisol.

Proteins

• Proteins are the most versatile macronutrients; ~100,000 different proteins are produced from the same genome through transcription/translation and folding.
• Amino acids: 20 standard amino acids are the building blocks of proteins; some are hydrophobic, some hydrophilic, with charged side chains contributing to protein folding and function.
• Polypeptides: sequences of amino acids linked by peptide bonds; hydrolysis can break them into shorter peptides or individual amino acids.
• Protein conformation: final 3-D shape (conformation) arises from primary structure (amino acid sequence) and higher-order structures.
  • Primary structure: sequence of amino acids from amino (N) to carboxyl (C) terminus.
  • Secondary structure: alpha helices and beta sheets stabilized by hydrogen bonds between backbone atoms.
  • Tertiary structure: overall 3-D folding driven by interactions among side chains (hydrogen bonds, ionic bonds, hydrophobic interactions, van der Waals).
  • Quaternary structure (when present): assembly of multiple polypeptide subunits (e.g., hemoglobin is a tetramer).
• Disulfide bonds: covalent linkages between cysteine residues can strongly influence folding (strongest among the discussed forces for maintaining structure).
• Protein examples and domains: SH2 domains, globular enzymes, structural proteins like collagen, transport proteins like hemoglobin, enzymes like catalase and lysozyme, membrane proteins like porins, etc.

Nucleic Acids

• Nucleic acids are polymers made of nucleotides; nucleotides account for about 2% of body weight and store, express, and transmit genetic information.
• Backbone: alternating sugar (pentose) and phosphate groups with attached nitrogenous bases.
• Bases (purines vs pyrimidines):
  • Purines: Adenine (A) and Guanine (G)
  • Pyrimidines: Cytosine (C), Thymine (T, in DNA), Uracil (U, in RNA)
• DNA specifics:
  • Double helix with two strands held together by hydrogen bonds between complementary bases.
  • Genome contains about 20,000 genes; laid out across ~3×10^9 base pairs per haploid genome; two copies per nucleus give ~6×10^9 base pairs per diploid cell.
• RNA specifics:
  • Generally single-stranded; sugars differ (ribose vs deoxyribose); Uracil replaces Thymine.
  • Three major forms in the excerpt: messenger RNA (mRNA), transfer RNA (tRNA), and ribosomal RNA (rRNA).
• DNA-RNA comparison (highlights):
  • DNA uses deoxyribose; RNA uses ribose.
  • DNA bases: A, G, C, T; RNA bases: A, G, C, U.
  • DNA is double-stranded; RNA is typically single-stranded.

Acids, Bases, and pH

• Acids dissociate to release hydrogen ions (H⁺); bases accept hydrogen ions.
• Acidity is defined by the free hydrogen ion concentration; higher [H⁺] means greater acidity.
• pH scale:
  • pH = $-\log_{10}[H^+]$
  • Physiological pH of the body: approx. $7.2 - 7.4$; values outside this range can be life-threatening (e.g., pH = 6.8 is incompatible with life).
• pH visualization:
  • The neutral point is pH 7.0; acidic conditions have pH < 7.0; basic/alkaline conditions have pH > 7.0.
  • In blood, venous vs arterial pH can differ; extreme deviations can cause acidosis or alkalosis and can be fatal.

Solutions and Concentrations

• Concentration concept:
  • One mole of a substance in one liter yields 1 M (1 mol/L) solution.
  • Avogadro’s number: $N_A = 6.022\times 10^{23}$ entities per mole.
• In biology, concentrations of solutes in body fluids are typically in the millimolar (mM) range; many solute concentrations fall well below 1 M (common clinical ranges are mM).
• Hydration shells: ions in solution are typically surrounded by solvent molecules forming solvation shells.

Everyday Biochemical Implications and Real-World Relevance

• Isotopes in research and medicine (e.g., ^13C-labeled compounds) are used to trace metabolic pathways (e.g., measuring vitamin B-12 bioavailability with [^13C]-cyanocobalamin).
• Carbon isotope ratios (e.g., ^12C/^13C) are used in forensic and athletic doping contexts to differentiate synthetic from natural sources (e.g., synthetic testosterone distinguishable by isotope ratios).
• Biomolecule polarity and solubility affect distribution in the body: polar/ionized molecules are hydrophilic and dissolve in blood and intracellular fluids; nonpolar molecules are hydrophobic and tend to associate with lipid environments (e.g., membrane cores).
• The principle of amphipathicity explains soap action: amphipathic molecules interact with both water and oils to emulsify fats.

Practical Formulas and Key Numbers (for quick recall)

• Atomic mass and isotopes: A = Z + N; atomic mass is the weighted average of stable isotopes in daltons.
• Avogadro’s concept: $N_A = 6.022\times 10^{23}$ particles per mole; 1 mol of any substance contains this number of entities.
• Molarity: $1~\mathrm{M} = 1~\mathrm{mol/L}$; 1 mol/L of glucose contains the same number of solute particles as any other substance at 1 M (Avogadro’s number).
• Energy content of lipids: $9\ \mathrm{kcal/g}$ (lipids yield more energy per gram than carbohydrates or proteins).
• Carbohydrate formula examples:
  • Glucose: $\mathrm{C<em>6H</em>{12}O_6}$
  • Sucrose: $\mathrm{C<em>{12}H</em>{22}O<em>{11}} + H</em>2O$ (dehydration synthesis forms glycosidic bonds; hydrolysis adds water)
• Nucleotides and nucleic acids: nucleotides are the building blocks; DNA typically contains about $3\times 10^9$ base pairs per haploid genome; diploid cells have ~6×10^9 base pairs in total.
• ATP: an energy-carrying molecule; hydrolysis releases energy to perform work; ATP is not stored long-term and must be regenerated continuously.

Quick Connections to Foundational Principles

• Chemistry underpins biology: atoms -> molecules -> macromolecules -> cells -> tissues -> organs -> systems.
• Emergent properties: function arises from the arrangement and interactions of atoms and molecules, not from the atoms themselves.
• Structure governs function across scales: molecular shapes (VSEPR, lipid bilayers, protein folding) determine biological activity.
• Hydration and solvation are central to biochemistry: ions and polar molecules interact with water to enable transport, signaling, and metabolism.
• Isotopes and labeling enable tracing of biochemical pathways and dating techniques, linking chemistry to medicine and biology.

Ethical, Philosophical, and Practical Implications

• The Schrödinger question about physics and chemistry explaining life touches on fundamental questions about reductionism in biology and the nature of life.
• Isotope labeling and carbon dating have profound implications for archaeology, forensics, and medical diagnostics.
• Understanding molecular interactions informs drug design, nutrition, and public health; e.g., water solubility, lipid solubility, and pH management are critical in pharmacology and physiology.
• The ethics of manipulating isotopes and genetic material (DNA/RNA) require careful consideration of safety, consent, and societal impact.

Summary Reference Points (recap)

• Body composition is dominated by a few major elements, with numerous trace elements playing specialized roles.
• Atoms organize into molecules, which form macromolecules like carbohydrates, lipids, proteins, and nucleic acids.
• The macromolecules have defined basic building blocks (monosaccharides, fatty acids/glycerol, amino acids, nucleotides) and follow fundamental bonding principles (covalent, ionic, hydrogen bonds).
• Water’s polarity and hydrogen bonding are central to molecular interactions, solubility, and biological structure.
• pH and acidity regulate biochemical processes; physiological pH is tightly regulated for life.
• ATP provides energy for cellular processes and is regenerated continually.
• The genome’s structure (DNA) and transcriptional/translational processes (RNA) underlie inheritance and protein synthesis.

Appendix: Key Notations to Remember

• $N_A = 6.022\times 10^{23}$
• $\mathrm{M} = \mathrm{mol/L}$
• $\mathrm{pH} = -\log_{10}[\mathrm{H^+}]$
• Glucose: $\mathrm{C<em>6H</em>{12}O_6}$
• Sucrose: $\mathrm{C<em>{12}H</em>{22}O_{11}}$ + H2O (dehydration synthesis/hydrolysis context)
• Lipids energy: $9\ \mathrm{kcal/g}$
• DNA base-pairs per diploid nucleus: ~$6\times 10^9$ base pairs
• ATP: energy-carrying molecule; energy released by hydrolysis