12-08: Energy & Metabolism

Kinetic & Potential Energy

• Metabolism: the sum of all chemical reactions needed to sustain life
• Energy is the ability to do work (movement)
• Types of energy:
  • Kinetic: motion
  • Potential (chemical potential energy): energy contained/stored in chemical bonds within a molecule
• First law of thermodynamics: energy cannot be created or destroyed - it can only be converted from one form to another
  • Some reactions produce or take up more energy than they require

Metabolism

• Metabolism = anabolism + catabolism

  • Anabolism: builds up
  • Catabolism: breaks down
  • The sum of all chemical reactions

• In a reaction, bonds between reactants are broken down and bonds between products are formed

  • Energy is absorbed when reactant bonds break and energy is needed
  • Energy is released when product bonds form

  

  

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Anabolic Reactions

• Build complex substances from smaller subunits

• Overall require the input of energy to occur

• Endergonic reactions: a need for energy

• E.g. photosynthesis

  

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Catabolic Reactions

• Breakdown of complex substances

• Overall release energy

• Exergonic reactions: release of energy

• Energy can be used for other jobs

• E.g. cellular respiration

  

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Exothermic vs. Endothermic

• Enthalpy (∆H): a measure of the energy in a system, related to the amount of heat released or absorbed by a reaction

• Exothermic reactions: ∆H < 0 (negative) – releases more thermal energy than they absorb

  • Products have less potential energy than reactants

• Endothermic reactions: ∆H > 0 (positive) – absorbs more thermal energy than it releases

• Chemical reactions need activation energy for a reaction to happen

  

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Energy Loss

• Second law of thermodynamics: in every energy conversion, some energy becomes unusable, thus increasing the entropy of the universe

• Heat is the typical form of loss of usual energy – it contributes to the increasing disorder in the surroundings

• Entropy (∆S): measure of randomness/disorder

  • +∆S: more disorder
  • -∆S: less disorder

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Gibbs Free Energy

• The most useful kind of energy

• In a chemical change, since some energy is lost to entropy, the usable remainder is called Gibbs Free Energy

• Exergonic reactions release free energy (-∆G) and are spontaneous – i.e. don’t need certain conditions to occur, will continue as long as there is sufficient reactant w/o a continuous input of free energy

• Endergonic reactions absorb free energy (+∆G) and are not spontaneous – i.e. they absorb more free energy than it produces and it won’t happen with a continual source because of this

  • Spontaneous: don’t need a continual source of energy, once they hit their activation energies they will occur (e.g. once the match is lit, the fire continues)

  

Spontaneous Reactions

• Reactions that release free energy (-∆G) are spontaneous

• Spontaneous reactions are those that will continue on their own once started (after activation energy)

• Total energy change/enthalpy(∆H) and total disorder change/entropy (∆S) both play a key role in determining whether or not a reaction will be spontaneous (along with temperature)

  ∆G = ∆H --- T∆S

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Metabolic Pathways

• Catabolic pathways

  • Break down complex molecules into simple ones
  • Release free energy – spontaneous (-∆G)

• Anabolic pathways

  • Build more complex molecules from simple ones
  • Absorb free energy – non spontaneous (+∆G)

• In order for non spontaneous reactions to occur, they must be coupled with spontaneous ones

  

  ATP Hydrolysis

  • Hydrolysis: breaking something down (catabolic)

• Energy for endergonic reactions in cells is mostly provided by coupling the reaction with the hydrolysis of ATP

  • When water reacts with it, it breaks it into ADP and the phosphate bonds to a reactant thus making it more reactive, provides energy

  

  

Coupled Reactions

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