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KNOWL Flashcards: Energy & Thermodynamics

1. Definition of Energy

Q: What is energy?
A: The capacity to do work (to cause change).


2. Potential vs. Kinetic Energy

Q: Compare potential and kinetic energy. Classify different energy forms.
A:

  • Potential Energy: Stored energy.

    • Examples: Chemical energy (bonds & atoms), Concentration gradient (across membranes).

  • Kinetic Energy: Energy of motion.

    • Examples: Electrical, Radiant, Thermal, Motion energy.


3. Cellular Energy Needs

Q: Why do cells need energy? What are types of cellular work?
A: Cells require energy for:

  • Synthetic Work – Biosynthesis of macromolecules.

  • Mechanical Work – Movement (e.g., muscle contraction, flagella motion).

  • Concentration Work – Active transport across membranes.

  • Electrical Work – Ion gradients (e.g., nerve signaling).

  • Heat Generation – Regulating temperature.

  • Light Generation – Bioluminescence.


4. Energy-Converting Organisms

Q: Define autotrophs, heterotrophs, phototrophs, and chemotrophs.
A:

  • Autotrophs – Make organic molecules from inorganic ones.

  • Heterotrophs – Obtain organic molecules from others.

  • Phototrophs – Use light energy to make chemical energy.

  • Chemotrophs – Obtain energy by oxidizing chemical bonds.


5. Energy & Matter in the Biosphere

Q: Explain "Energy flows, matter cycles." Why must energy be replenished?
A:

  • Energy flows from the sun → captured by photoautotrophs → transferred in food chains → lost as heat.

  • Matter cycles between phototrophs & heterotrophs (e.g., carbon cycle).

  • Energy must be replenished because heat loss increases entropy.


6. Definition of Thermodynamics

Q: What is thermodynamics?
A: The study of energy transformations and how energy is used for work.


7. Laws of Thermodynamics

Q: State and explain the first and second laws of thermodynamics.
A:

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

  • Second Law: Every energy transfer increases entropy (disorder).


8. Free Energy & Spontaneity

Q: Define free energy and spontaneous processes.
A:

  • Free energy (G): Usable energy for work.

  • Spontaneous process: Occurs without added energy (∆G < 0).


9. Delta G & Equilibrium

Q: What does ∆G tell us? What happens when ∆G = 0?
A:

  • ∆G tells us if a reaction is spontaneous.

  • When ∆G = 0, the system is at equilibrium, meaning no net work can be done.


10. Gibbs Free Energy Equation

Q: Interpret the equation ∆G = ∆H - T∆S.
A:

  • ∆G: Free energy change (spontaneity).

  • ∆H: Enthalpy (heat content).

  • T: Temperature (Kelvin).

  • ∆S: Entropy (disorder).

  • If ∆G < 0: Spontaneous (exergonic).

  • If ∆G > 0: Non-spontaneous (endergonic).


11. Equilibrium & Work

Q: What happens when ∆G = 0?
A: The system is at equilibrium, the most stable state, where no work can be done.


12. Free Energy Diagrams & Reaction Spontaneity

Q: How do free energy diagrams determine spontaneity?
A: If final energy is lower than initial energy (∆G < 0), the reaction is spontaneous.


13. Endergonic vs. Exergonic Reactions

Q: Define and compare exergonic & endergonic reactions.
A:

  • Exergonic: Releases energy (∆G < 0).

  • Endergonic: Requires energy (∆G > 0).


14. Spontaneous Changes

Q: Provide examples of spontaneous processes. How does initial state compare to final state?
A:

  • Examples: Jumping off a diving board, diffusion, chemical reactions.

  • Initial state has more energy than the final state.


15. Entropy & Thermodynamics

Q: What is entropy, and how does it relate to thermodynamics?
A:

  • Entropy (S): Measure of disorder.

  • The Second Law of Thermodynamics states that entropy increases with energy transformations.


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