Chapter 1-7 Review: Chemistry, Bonds, Reactions, and Body Materials

Atoms and Elements

Analogy: atoms are tiny Legos that build molecules and, in turn, a whole human. The lecture starts with a quick review of basic chemistry.
Periodic table basics:
- Every element has a weight (atomic mass) and a specific number of electrons, protons, and neutrons.
- Protons and neutrons reside in the nucleus; electrons form a cloud around the nucleus.
- Protons are positively charged; electrons are negatively charged.
- In clinical contexts, the term
  proton is often used to refer to a hydrogen ion.
Clinical tie-in: a patient with chronic indigestion might be treated with proton pump inhibitors (PPIs). The class discussed the terminology:
- PPIs are the medication class; the speaker notes the term proton pump inhibitor is commonly used even though the hydrogen ion pump is the actual target.
The four most abundant elements in the human body (accounting for about 95% of body mass):
- Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N).
Atomic models:
- 2D “top-wave” ring models are common for teaching, but in reality atoms are three-dimensional with electron clouds.
Key atomic numbers and masses (as mentioned):
- Atomic number Z is the number of protons. Example: Hydrogen is Z = 1, Helium Z = 2. Carbon’s number is not always remembered by everyone; the speaker humorously notes uncertainty about carbon’s exact number.
- Atomic mass is the sum of protons and neutrons. Example: Carbon-12 is the most common isotope; Carbon-14 is radioactive and used in carbon dating.
Isotopes and medical relevance:
- Isotopes differ in neutron number; stability changes with neutron count.
- Unstable isotopes are radioactive and can be useful in medicine (e.g., imaging) but can be dangerous if misused.
Hydrogen isotopes (examples in the talk):
- Protium (^{1}H), Deuterium (^{2}H), Tritium (^{3}H).
Imaging techniques discussed:
- PET scan uses a radioactive form of glucose; as the tracer is metabolized, radioactive particles are released and detected.
- Barium swallow imaging shows the GI tract using a radiopaque substance (barium) to visualize passage.
- CT/MRI contrast agents:
- Gadolinium-based contrast enhances visibility in imaging; beware of patient allergy history.
Ions and electrolytes:
- An ion is an atom with a different number of p+ and e− (i.e., a net charge).
- Electrolytes are ions that carry charge and are essential for bodily functions (e.g., Na⁺, Ca^{2+}, Cl⁻).
- A cation is positively charged; an anion is negatively charged.
- Mnemonic for calcium as a positive ion: Ca^{2+} -> “calcium two plus” and “cation” help recall.
Electron shells overview:
- Atoms have electron shells; the distribution of electrons affects stability and bonding.
Ionic bonding basics:
- In ionic bonds, electrons are donated from one atom to another rather than shared.
- Example: Sodium chloride, NaCl, where Na donates an electron and becomes Na⁺ while Cl gains an electron and becomes Cl⁻.
- Oxidation and reduction:
- Oxidation: loss of electrons.
- Reduction: gain of electrons.
- Mnemonics: LEO the lion says GER; Oil rig (Oxidation is Loss, Reduction is Gain).
Covalent and hydrogen bonds:
- Covalent bonds: electrons are shared between atoms, forming stable molecules (e.g., O₂, CO₂).
- Double covalent bonds: example O=O in O₂; carbon dioxide (CO₂) as an example of covalent bonding.
- Hydrogen bonds: rely on molecular polarity (one end positive, one end negative) and are common in water molecules; they are relatively weak but highly influential in biology.
- Water and cohesion: hydrogen bonds cause water molecules to stick together (cohesion). A common analogy is the “Mickey Mouse” arrangement of water molecules or the dome of water when a full cup is held without spilling.
- Biological relevance of hydrogen bonding includes properties of water in lungs and the production of surfactant to prevent lung collapse.
Three main types of chemical bonds (recap):
- Ionic bonds (electrons donated)
- Covalent bonds (electrons shared)
- Hydrogen bonds (intermolecular forces due to polarity)

Chemical Reactions and Energy

Factors that affect chemical reactions in the body:
- Acidity (pH), temperature, availability/concentration of reactants, products, size of molecules, and catalysts.
Catalyst:
- A substance that speeds up a chemical reaction by lowering the activation energy, without being consumed.
Activation energy:
- The energy required to start a reaction. Enzymes dramatically lower this barrier.
Synthesis vs decomposition:
- Synthesis (anabolism): building larger molecules from smaller ones (e.g., notebook = note + book).
- Decomposition (catabolism): breaking larger molecules into smaller components (e.g., notebook = book + worm).
Exchange reactions:
- Part of reactants are swapped to form new products.
Energy release/types:
- Exergonic reactions release energy.
- Endergonic reactions require energy input (noted in the body, often related to digestion or metabolism).
Reactants and products terminology:
- Reactants: starting materials in a reaction.
- Products: substances produced by the reaction.
Enzymes:
- Enzymes are molecules that lower activation energy and speed up reactions; they enable biological processes to occur in physiologically relevant timescales.

Organic vs Inorganic Molecules

Organic molecules:
- Carbon-based, often with hydrogen; traditionally defined as carbon-containing compounds.
- The lecture notes that carbon plus hydrogen are typically present in organic molecules; CO₂ is sometimes considered inorganic in strict chemistry contexts.
Inorganic molecules:
- Not carbon-hydrogen based in the same sense; examples include water (H₂O), salts, minerals.
Practical slide examples (bone samples):
- Ground bone: shows inorganic salts only; organic elements (cells, collagen, vessels) are removed during preparation.
- Decalcified bone: pink collagen matrix visible; osteocytes; vessels and nerves present; reveals bone tissue structure.

Water, Solutions, and Body Composition

Water in the body:
- Approximately 70% of body weight is water; technically inorganic; it serves multiple roles.
- Roles of water:
- Lubricant and cushion for tissues.
- Heat sink: distributes heat throughout the body to maintain relatively uniform temperature.
- Universal solvent: enables many biochemical reactions.
Types of mixtures in biology:
- Solutions, colloids, suspensions.
- Hydrophilic vs hydrophobic substances explain interactions with water.
Dehydration and hydration concepts:
- Dehydration synthesis (condensation): two molecules combine and release a water molecule, forming a larger molecule.
- Byproduct: water, produced during synthesis.
- Example: forming a polymer from monomers.
- Hydrolysis: water is added to break bonds, separating molecules.
Water-derived questions in labs:
- Dehydration synthesis and hydrolysis are common topics for physiology labs and exams.

pH, Acids, Bases, and Buffers

pH scale:
- Measures acidity or basicity (alkalinity).
- Lower pH = more acidic; higher pH = more basic; neutral pH = 7.
Buffers:
- Buffers are weak acids or bases that resist changes in pH by shifting equilibrium.
- In blood, carbonic acid/bicarbonate buffering system is key:
- Carbon dioxide interacts with water to form carbonic acid, which can dissociate to bicarbonate and a hydrogen ion.
- The buffering equilibrium helps maintain blood pH around the narrow range of about 7.35 ext{ to } 7.45.
Carbonic acid/bicarbonate system:
- Chemical equilibrium example:
  \mathrm{CO2 + H2O \rightleftharpoons H2CO3 \rightleftharpoons H^+ + HCO_3^-}
- Carbon dioxide (CO₂) originates from cellular respiration; carbonic acid/bicarbonate buffering relates to respiratory control and pH sensing by receptors in the brain and arteries.
Body-specific pH values:
- Blood: pH \approx 7.35 - 7.45
- Urine: typically pH \approx 4.5 - 8 (varies with elimination of acids/bases and dietary influences)
- Stomach: highly acidic, pH \approx 1.5 - 3.5 (gastric acid) to aid protein digestion
Clinical contexts:
- Acidosis and alkalosis refer to states outside normal pH ranges; more detailed AP2 topics cover these conditions.

Electrolytes, Ions, and Body Fluids

Electrolytes:
- Salts dissolve in water to form ions (cations and anions).
- Common bodily electrolytes include Na⁺, K⁺, Ca^{2+}, Cl⁻, and others.
- Electrolyte imbalances can disrupt many physiological systems; frequent in elderly patients on supplements like potassium.
Dissociation in water:
- Water’s polarity surrounds ions; the negative end of water around positive cations and the positive end around negative anions facilitates dissociation of salts into √
  ext{Salt} \rightarrow \text{Cation}^+ + \text{Anion}^-

Practical Imaging and Contrast Agents

PET scans use radiolabeled compounds (e.g., radioactive glucose) to visualize metabolic activity; energy from decay traces the metabolic pathway.
Contrast agents:
- Gadolinium-based contrast is used to improve visibility in CT/MRI scans.
- Always review patient allergy history before administering contrast.

Practical Topics in Anatomy and Lab Techniques

Bone slide preparation differences:
- Ground bone: shows inorganic salts; no cells, collagen, vessels.
- Decalcified bone: reveals collagen, osteocytes, vessels, nerves, and bone tissue organization.

Key Takeaways and Connections

Three primary bond types shape molecular behavior in the body:
- Ionic bonds (electrons donated); example: NaCl; oxidation vs reduction concepts summarized by mnemonics.
- Covalent bonds (electrons shared); examples: O₂, CO₂.
- Hydrogen bonds (polarity-driven, weaker but crucial for properties of water and biological structures).
Water plays critical roles: solvent, temperature distribution (heat sink), lubrication; also participates in dehydration synthesis and hydrolysis reactions.
The body maintains tight pH control via buffers (notably the carbonic acid/bicarbonate system) to keep blood within a narrow pH window; disruptions can cause acidosis/alkalosis.
Electrolytes (Na⁺, K⁺, Ca²⁺, Cl⁻, etc.) dissociate in water and are essential for muscle, nerve, and fluid balance; imbalances can have systemic effects.
Organic vs inorganic distinction helps frame how the body stores energy and builds tissues; bones illustrate the importance of both inorganic minerals and organic collagen in tissue structure.
Instructors often tailor depth to exam scope: expect definitions, basic concepts, and practical applications rather than exhaustive chapter detail.

Quick Reference Equations and Items (LaTeX)

Atomic number and mass concepts:
- Z = \text{number of protons}
- A = Z + N
Isotopes (examples):
- ^{12}\mathrm{C}, \quad ^{14}\mathrm{C}
- Hydrogen isotopes: ^{1}\mathrm{H}, \; ^{2}\mathrm{H}, \; ^{3}\mathrm{H}
Ions and charges:
- \mathrm{Na^+}, \; \mathrm{Ca^{2+}}, \; \mathrm{Cl^-}
Covalent bonds (example): \mathrm{O2}, \; \mathrm{CO2} (double bonds in O=O and C=O contexts)
Hydrogen bonding and cohesion illustrations are conceptual rather than closed-form equations.
Acid–base buffering equilibrium:
- \mathrm{CO2 + H2O \rightleftharpoons H2CO3 \rightleftharpoons H^+ + HCO_3^-}
Blood pH window: 7.35 \leq pH \leq 7.45
Stomach pH range: 1.5 \leq pH \leq 3.5
Urine pH range: 4.5 \leq pH \leq 8
Activation energy concept (enzymes lower it):
- Without enzyme: high energy barrier; with enzyme: lower barrier and faster rate (qualitative, not a fixed number).