Biology and Thermodynamics Lecture Vocabulary

1. Fundamentals of Chemistry and Atomic Structure

• Elements

  • An element is a pure substance that cannot be broken down into any simpler substance by ordinary chemical means.

  • Protons determine the atomic number and chemical identity of an element.

  • Neutrons act as nuclear glue to hold protons together against electrostatic repulsion.

  • Isotopes: Atoms of the same element containing differing numbers of neutrons.

  • Chemically identical to normal atoms, but some are radioactive.

  • Example: Radium ($Ra$) sits in the same periodic table column as Calcium ($Ca$) and Barium ($Ba$). Because elements in the same column exhibit similar chemical behaviors, radium gets incorporated into bone tissue when ingested.

• Molecules and Chemical Bonds

  • A molecule consists of multiple atoms linked together via chemical bonds.

  • Bonds Types:

  • Covalent Bonds: Strong links formed by shared electrons; essential structural component of molecules.

  • Ionic Bonds: Strong electrostatic interactions, easily dissolved in polar solvents like water.

  • Hydrogen Bonds: Intermediate-strength interactions key to polar liquid properties.

  • Van der Waals Forces: Weak intermolecular attractions.

2. Properties of Water

• High Heat Capacity: Stores heat effectively, requiring $5$ times more heat energy to raise its temperature compared to metals.

• Low Density of Ice: Solid water (ice) is less dense than liquid water and forms on the surface of water bodies.
  
  

  • Serves as an insulating layer that prevents lakes and ponds from freezing solid from the bottom up, preserving aquatic life.

• High Heat of Vaporization: Requires $540JS_NOP_QUOTEcal\,g^{-1}$ during phase transition from liquid to gas without raising temperature.
  
  

  • Provides evaporative cooling mechanisms (sweating in humans, panting in dogs). Cooling efficiency decreases at high humidity levels.

• Cohesion and Adhesion:
  
  

  • Cohesion: Attraction between water molecules themselves.

  • Adhesion: Attraction between water molecules and other polar surfaces.

  • Together with capillary action and surface evaporation, adhesion allows tall trees (up to $380\,ft$) to transport water upwards without mechanical pumps.

• High Polarity: Excellent solvent for polar substances (e.g., sugars, salts); insoluble with nonpolar compounds (e.g., oils, lipids).

3. $pH$ and Acid-Base Buffers

• $pH$ Definition:

  • Mathematical definition:
        $pH = -\log[H^+]$

  • Concentration $[H^+]$ is measured in moles per liter ($mol\,dm^{-3}$).

  • A Mole: Avogadro's number ($1\,mol = 6.022 \times 10^{23}$ particles).

• $pH$ Scale Characteristics:

  • Neutral pure water has $[H^+] = 10^{-7}\,mol\,dm^{-3}$, giving a $pH = 7.0$.

  • Scale ranges from $0$ to $14$. Because it is logarithmic, a change of $1.0\,pH$ unit represents a $10\text{-fold}$ change in hydrogen ion concentration.

  • Acids: Dissociate in water to increase $[H^+]$, lowering $pH$ below $7.0$.

  • Bases: Decrease $[H^+]$ or increase $[OH^-]$, raising $pH$ above $7.0$.

• Biological Buffer Systems:

  • Buffers consist of weak acids and their conjugate bases in roughly equal concentrations to resist changes in $pH$.

  • Bicarbonate Buffer System: Primary short-term buffer in blood maintain blood $pH$ around $7.4$:
        $CO_2 + H_2O \rightleftharpoons H_2CO_3 \rightleftharpoons H^+ + HCO_3^-$

  • Increased respiratory $CO_2$ shifts equilibrium right, increasing $[H^+]$ and lowering blood $pH$. Sensory receptors detect low $pH$, triggering the urge to breathe.

  • Hyperventilation expels $CO_2$, shifting equilibrium left, decreasing $[H^+]$, and raising blood $pH$.

4. Radiometric Dating

• Carbon-14 Dating:
  
  

  • Measures the ratio of radioactive Carbon-14 ($^{14}C$) to stable Carbon-12 ($^{12}C$) in biological remains.

  • Half-life of $^{14}C$ is $5730$ years, decaying exponentially over time.

  • Effective for dating organic materials up to approximately $50,000$ years old (e.g., Ötzi the Iceman dated to roughly $5800$ years ago).

5. Laws of Thermodynamics and Energy

• Energy: The capacity to do work.

  • Kinetic Energy: Energy of motion.

  • Potential Energy: Stored energy based on position or chemical structure.

• First Law of Thermodynamics:

  • Energy cannot be created or destroyed; the total amount of energy in a closed system remains constant.

  • Energy can change forms (e.g., potential to kinetic), but energy conversions always release some energy as non-useful heat due to friction.

• Second Law of Thermodynamics:

  • The total entropy ($S$) or disorder of a closed system increases over time.

  • Living organisms are highly organized (low entropy) systems. To maintain order against thermodynamic decay, organisms disorganize complex food molecules (catabolism) to harness energy for structural upkeep (anabolism).

6. Chemical Reactions and Activation Energy

• Reaction Types:

  • Endergonic Reactions: Require net energy input; products hold more energy than reactants.

  • Exergonic Reactions: Net release of energy; products hold less energy than reactants.

• Activation Energy:

  • The initial energy input required to destabilize chemical bonds and initiate a reaction.

  • Explains why reduced organic matter (wood, living tissue) does not spontaneously combust despite Earth's atmosphere containing $21\%$ oxygen ($O_2$). Biological enzymes lower activation energy barriers to allow controlled metabolic reactions.