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# Introduction to Metabolism Lecture Index Course Resources page
Last revised: Tuesday, February 15, 2000
Ch. 8 in Prescott et al, Microbiology, 4th Ed.
Note: These notes are provided as a guide to topics the instructor hopes to cover during lecture. Actual coverage will always differ somewhat from what is printed here. These notes are not a substitute for the actual lecture!

## Energy and Metabolism

### Energy = capacity to do work

• Measured in calories (heat unit) or joules (work unit). 1 calorie = 4.1840 joules.
• Physicists & chemists use joules; biologists typically use calories. Some biochemistry and microbiology texts use kilocalories, others have converted to kilojoules. I will use kcal. To compare lecture values (kcal) with text values (kJ), multiply by 4.184.
• Example: 40 kjoules = 40/4.184 = 9.56 kcal
• Cells require energy, either as light (phototrophs), inorganic chemicals (chemolithotrophs), or organic chemicals (chemoorganotrophs).

### Metabolism

• = total activity of cell; has two components
1. Anabolism -- used to make new molecules
• assimilative
• biosynthetic
• endergonic
• G > 0; energy consuming
2. Catabolism -- used to obtain energy
• dissimulative
• exergonic
• G < 0; energy producing

### Overview of Catabolism

• electrons from energy source -------> redox carriers --------> terminal electron acceptors:
• some free energy released is trapped in chemically usable forms
• terminal acceptor can be external molecule (oxygen, nitrate, etc.), use membrane-localized electron transfer system (ETS) in respiration (energy yields high)
• or terminal acceptor can be organic molecule derived from energy source during metabolism (pyruvate, etc.) -------> fermentation (energy yields low)

### Thermodynamics = laws governing energy transfer

• Originally from study of heat; later shown to apply universally to all forms of energy transfer
1. Law 1: conservation of energy.
2. Law 2: total amount of entropy (S) increases during energy transfers
• Implications of 2nd Law
• disorder is increasing
• processes go to equilibrium
• heat flows from hot to cold
• diffusion leads to substances becoming uniformly dispersed
• systems far from equilibrium can do useful work; not possible after equilibrium is attained
• Entropy governs availability of energy for useful work.
• Note: biological systems are never at equilibrium (unless dead).

## Free energy: G, Go', relationship to equilibrium

• Free Energy change is a direct measure of the energy that is available to do useful work

### Definition of G:

• G = H - T S; measures free energy change for any reaction, where G = change in free energy, a measure of useful work H = change in heat content S = change in entropy
T = absolute temperature in degrees Kelvin
• Conventions: exergonic reaction have G <0; system can do useful work, will occur spontaneously

### Definition of Go'

• Instead of calculting actual G values, biochemists often use the standard free energy change, Go', to estimate bioenergetic calculations:
• Go' = standard free energy change, 1 M concentrations (the "o"), pH 7 (the ')
• Go' is related to equilibrium
• consider reaction: A + B <=====> C + D
• Any reaction has some equilibrium state; given enough time to reach equilibrium, concentration of products and reactants will remain constant, and can be defined by equilibrium constant Keq:
Keq = [products]/[reactants] = [C][D]/[A][B] Go' = - RT ln Keq

```Sample values of Go':
Keq Go'	     comments
C, D = A, B	              1	          0 Kcal	     no useful work
C, D = 10 x A, B	     10	     - 1.36 Kcal	     useful work
C, D = 100 x A, B	    100	     - 2.73 Kcal	     more useful work
C, D = 1000 x A, B	   1000	     - 4.10 Kcal	     even more useful work
```
• Some examples:
1. fermentation: glucose ---> pyruvate, Go' = - 57 kcal/mole
2. respiration: pyruvate ---> CO2 + H2O; Go' = - 633 kcal/mole
• G can predict reactions
1. if G is negative, reaction is exergonic (spontaneous), will proceed rapidly with release of energy. Can be used for growth.
2. if G is positive, reaction is endergonic, will not occur unless energy is added, it cannot be used for growth.
• relation of G to Go'
• G = Go' + RT lnK, where
R = gas constant (8.29 J/mol/K)
T = absolute temperature
K = equilibrium constant. For a general reaction A + B <==> C + D, then K = [C] [D] / [A] [B]

### Q&A

Q. Do reactions with positive G ever occur?
A. No. But many reactions with positive values of Go' occur in cells.

Q. How is that possible?
A. Remember that G is the sum of two terms (see above): the Go', which can be positive, and the RT lnK term. If the second term is sufficiently negative, it can overcome a strongly postive Go' term.

Q. Wait a minute. Isn't the Go' itself related to the equilibrium constant?
A. Yes, it's related to the equilibrium constant under standard conditions, which include everything being at 1 molar concentrations and pH 7. But no cell ever comes close to 1 molar concentrations. In fact, cells often keep the concentrations of many reactants at very small levels; as soon as a compound is produced in one reaction, it is immediately used up in another reaction. Most cell chemicals occur in mMolar amounts or less.

A consequence of this is that the RT lnK term of the G reaction (above) can be modified quickly in the cell. If the cell needs to make this a very large negative number, the concentration of product needs to be kept very small relative to the concentration of reactant. Since the log of 1 is zero, the log of numbers smaller than 1 is a negative number, which makes the RT lnK term negative. The smaller the ratio of [C] [D] / [A] [B], the more negative that number becomes.

Q. What's the bottom line here?
A. If reactions have negative G values, they're very useful to a cell -- they can release energy, they can happen spontaneously. All the cell needs to do is make an enzyme to speed up such reactions. On the other hand, if needed reactions have positive G values, they're not going to happen. In order to make such reactions occur, the cell has to change the G to a negative value, either by coupling such reactions with exergonic reactions (we'll see examples of that later), or manipulating concentrations of reactants and products to force the reaction to have a negative value. That works surprisingly well for intermediary metabolites (chemicals used only to get from one molecule to another), but not for desired end products, which by definition must accumulate in large quantity.

## Biological Oxidations

• Redution-Oxidation (Redox) reactions are often associated with Energy transfer in cells
• oxygen rarely involved
• oxidation reaction always coupled with a reduction reaction.
• Definition: loss of electron = Oxidation.
• gain of electrons = Reduction
• Mnemonic: LEO the lion says GER
• Examples of oxidation reactions:
1. Fe++ ---> Fe+++ + e- -- occurs in cytochromes
2. succinic acid ----> fumaric acid + [2H]
• Note: this is oxidation, also dehydrogenation reaction, since 2H = 2 H+ + 2 e-.
• Oxidation reactions are always accompanied by reduction reactions:
• Example of a reduction reaction:

## Redox Carriers

• NAD+ is one of a small number of biomolecules that function as redox carriers; alternately get reduced, then oxidized. See text Fig. 4.8
• View NAD in oxidized and reduced forms
• Note: very small concentrations of NAD+ in cell; so must continually be recycled from (red) to (ox) state and back. Like an electron shuttle.
• Other redox carriers
• Ubiquinone (Coenzyme Q) -- carries 2H
• Heme groups (in cytochromes) -- carries single electron
• View FAD and heme structures

## Use of Redox Tower to measure Go', predict reactions

• Redox tower (see handout) provides a measure of electron donating capacity. Each substance has its own potential to give up electrons, some more (high on tower), some less (low on tower).
• View Redox tower diagram
• At top, good energy sources, very reduced, easily give up electrons.
• At bottom, good electron acceptors, very oxidized, readily accept electrons.
• Most compounds can either accept or donate electrons, depending on what chemicals they are paired with.
• For any pair, the member higher on the tower will tend to give up electrons (and be oxidized) to the lower member (which will be reduced).
• But note: actual G for a reaction depends not only on the Go', but also on actual reactant concentrations; if concentrations are favorable, it is possible to transfer electrons in opposite direction, as long as overall G is negative. (See above)
• Use tower to determine amount of Energy available from any pair or redox reactions.
• Go' = (- Eo') n F, where Eo' = (Eo' acceptor - Eo' donor), n = # of electrons transferred, and F = Faraday constant, 23 Kcal/mole
• Example: for H2 + O2 ---> H2O
• Eo' = + 0.82 - (-0.43) v. = 1.25 v.; n = 2
• Go' = (- 1.25 v. )(2)(23) Kcal = - 57.5 Kcal/mole
• Therefore this is a spontaneous reaction, liberates sizable free energy

## Use of ATP to store Energy

• Certain molecules store Energy for cell use. Exs: ATP, GTP
• Interact with structure of ATP (Note: you need the Chime plug-in to be able to view this structure)
• note: highly ionizable Đ charges in phosphate groups repel each other -- therefore easy to push one or two Pi groups away from rest of molecule
• ATP + H2O ---> ADP + Pi Go' = -7300 Kcal/mole -- very exergonic
• ADP + H2O ---> AMP + Pi Go' = -6800 Kcal/mole -- very exergonic
• Note: these reactions don't actually occur in cell -- no enzymes to carry them out!
• ATP Synthesis
• To make ATP, must supply more than 7300 Kcal
• Major problem for every cell is to find ways to make ATP
• ATP used as energy donor for many synthetic (anabolic) reactions, to link with other reactions so that overall G is negative, and reactions proceed spontaneously.
• Note: also GTP, CTP, UTP present. Pools of nucleotide phosphate groups can be interchanged by appropriate enzymes.

## Activation Energy and the role of enzymes

• Activation Energy = amount of energy put into a reaction in order to break bonds holding reactants together, allow new partnerships to form
• Example: Nitrogen fixation
N2 + H2 ----> NH3 Go' = - 19 kcal/mole (spontaneous)
• note: dinitrogen is held together by triple bond, one of strongest bonds known. To break it requires 225 kcal/mole.
• In Industry: uses 400 atmosphere pressure, 450 deg C -----> 3% of natural gas consumption in U.S. each year
• In bacteria: cells use enzyme (nitrogenase) ambient temperatures, 1 atm.

## Enzymes:

• protein catalysts -- lower activation energy of reaction
• speed up reactions by up to 1010 times; if enzyme is not present, reaction will not occur at rates useful to cell
• contain highly 3-dimensionally stereospecific active sites, only bind one or a very few similar substrates
• View pdb file of lysozyme bound to peptidoglycan substrate
• direct reaction to desired end products
• usually many enzymes organized in sequence
• each enzyme different. Mol. Wts. from 10,000 to 1,000,000, unique for each enzyme
• most enzymes contain prosthetic groups covalently bound; Exs: folic acid, riboflavin, FAD, heme groups, Zinc, Molybdenum, etc.
• can react with coenzymes; ex. NAD, Coenzyme A (small molecules, not covalently bound)

## Mechanisms of energy release: overview

1. Fermentation -- oxidation of an organic compound in the absence of external electron acceptor (no oxygen required). Uses SLP (substrate-level phosphorylation)
2. Respiration -- oxidation of an organic compound where oxygen is the final electron acceptor. Uses ETS (electron transport system) as well as SLP
3. Anaerobic respiration (unique to bacteria) -- oxidation of organic compounds where an external substrate other than oxygen serves as final electron acceptor. Exs: nitrate, sulfate, carbon dioxide

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