environmental cycles of metabolism n.
Skip this Video
Loading SlideShow in 5 Seconds..
Environmental Cycles of Metabolism PowerPoint Presentation
Download Presentation
Environmental Cycles of Metabolism

Loading in 2 Seconds...

play fullscreen
1 / 23

Environmental Cycles of Metabolism - PowerPoint PPT Presentation

  • Uploaded on

Environmental Cycles of Metabolism. Carbon is fixed (incorporated) by autotrophs (CO 2 ) and heterotrophs (complex such as carbohydrates) Nitrogen (N 2 ) is solely introduced into biological systems through microbes Also phosphate cycle, sulfur cycle, etc. Modes of metabolism.

I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.
Download Presentation

PowerPoint Slideshow about 'Environmental Cycles of Metabolism' - brett-gould

An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.

- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript
environmental cycles of metabolism
Environmental Cycles of Metabolism
  • Carbon is fixed (incorporated) by autotrophs (CO2) and heterotrophs (complex such as carbohydrates)
  • Nitrogen (N2) is solely introduced into biological systems through microbes
  • Also phosphate cycle, sulfur cycle, etc.
modes of metabolism
Modes of metabolism
  • Catabolism – nutrient breakdown
  • Anabolism – macromolecule synthesis
  • Both are linked via carriers of chemical energy NADH, ATP, NADPH, FADH2
  • These sources of chemical energy allow cells to perform “work” (synthesis, etc…)
consider the cell a system
Consider the cell a “system”
  • Isolated system – cannot exchange energy or matter with its surroundings (not a cell)
  • Closed system – can exchange energy, but not matter with its surroundings (still not a cell)
  • Open system – can exchange energy and matter in and out (A Cell!)
internal energy is a state function
Internal energy is a state function
  • The thermodynamic state is defined by prescribing the amounts of all substances present, and two of these variables: temperature (T), Pressure (P), and Volume (V) of the system.
  • The internal energy (E) of the system reflects all of the kinetic energy of motion, vibration, and rotation and all of the energy contained within chemical bonds and non-covalent interactions
how do cells make and use chemical energy
How do cells make and use chemical energy?
  • Bioenergetics must follow the laws of thermodynamics
  • First Law: the total amount of energy in the universe remains constant; energy may change form or location, but cannot be created or destroyed.
  • Second Law: Entropy is always increasing
first law of thermodynamics
First Law of Thermodynamics

DE = q – w

q = heat; positive q indicates heat is absorbed by the system, negative q indicates heat given off by system

w = work; positive w means the system is doing work, negative w means work is being done on the system

a bomb calorimeter allows reactions to be carried out at constant volume
A “bomb” calorimeter allows reactions to be carried out at constant volume
  • Because the reaction in (a) is carried out at constant V, no work is done on the surroundings
  • Therefore, DE = q
  • In this case, DE = -9941.4 kJ/mole
  • The negative sign indicates the reaction releases energy stored in chemical bonds and transfers heat to the surroundings
reactions at constant pressure
Reactions at constant pressure
  • In reaction (b), the reaction proceeds at 1 atm pressure
  • The system is free to expand or contract, the final state has contracted because the amount of gas has changed from 23 moles to 16
  • The decrease in volume means that work has been done on the system by the surroundings
pv work appears as extra heat released
PV work appears as extra heat released
  • When volume is changed against a constant pressure, w = PDV
  • Assumptions: constant T, gases are ideal, which allows us to use PV = nRT
  • w = DnRT = -17.3 kJ/mol
  • SO, under constant pressure q = DE + w = DE + DnRT = -9941.4 kJ/mol – 17.3 kJ/mol

= -9958.7 kJ/mol – In (b) the surroundings can do work on the system, this (PV) work looks like extra heat

most biochemical reactions occur under constant pressure not constant volume
Most biochemical reactions occur under constant pressure, not constant volume
  • Because q does not equal DE, we need to account for PV work done
  • We define a new quantity, enthalpy (H)
    • H = E + PV


    • When the heat of a reaction is measured at constant pressure, DH is determined
d e and d h measurements are useful for biochemists
DE and DH measurements are useful for biochemists
  • Although oxidation of palmitic acid occurs very differently in the human body than in a calorimeter, the values of DE and DH are the same regardless of the pathway
  • Average human expends ~6000 kJ or roughly 1500 kcal for bodily function, with exercise that figure easily doubles
d e d h is there a big distinction
DE, DH, is there a big distinction?
  • For most chemical reactions the difference between these two quantities is negligible
  • Typically, PDV is a tiny quantity
  • For instance, it’s about 0.2% difference for palmitic acid oxidation

DH is generally considered a direct measure of the energy change in a process and is the heat evolved in a reaction at constant P

entropy and the second law
Entropy and the second law

The minimal value

of entropy is a

perfect crystal at

absolute zero

thermodynamic quantities
Thermodynamic quantities

DH = enthalpy, the heat content of the system

exothermic = negative, endothermic = positive; Units: Joules/mole

DS = entropy, randomization of energy and matter; positive sign indicates increased entropy; Units: joules/mole(K)

DG = Gibbs Free energy, amount of energy that is available to do work at constant T and P; Units: Joules/mole

Note 1 calorie = 4.184 Joule

gibbs helmholtz equation
Gibbs-Helmholtz equation


Positive DG is endergonic, requires energy for reaction to occur, this is unfavorable

Negative DG is exergonic, releases energy, this is a favorable process; spontaneous but not necessarily rapid

A decrease in energy (-DH) and/or increase in entropy (+DS) make favorable processes

DG =0 indicates the system is at equilibrium

thermodynamics of melting ice
Thermodynamics of melting ice

Ice is a crystal lattice held together by H-bonds, bonds must be broken to form water

Energy for breakage of H-bonds is almost entirely the DH for this reaction and this term is positive

Entropy favors water over ice

But recall DG is also temperature dependent

why is d g called free energy
Why is DG called “free” energy?

DG represents the portion of an energy change DH that is available or free to do useful work.

TDS is amount of energy that is unavailable to do work


a d g warning
A DG Warning!
  • You will see many different DG’s

DG – Gibbs Free Energy

DG’o or DGo – Standard State Free Energy

energy per mole in standard state (1M)

DGo – Standard state Free Energy of Activation

enzyme catalysis