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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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