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

• 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”

• 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• 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?

• 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

E = 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

Oxidation of palmitic acid

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, E = q

• In this case, E = -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

• 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

• When volume is changed against a constant pressure, w = PV

• Assumptions: constant T, gases are ideal, which allows us to use PV = nRT

• w = nRT = -17.3 kJ/mol• SO, under constant pressure q = E + w = E + nRT =

-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

• Because q does not equal E, we need to account for PV work done

• We define a new quantity, enthalpy (H) – H = E + PV

H = E + PV– When the heat of a reaction is measured at

constant pressure, H is determined

E and H measurements are useful for biochemists

• Although oxidation of palmitic acid occurs very differently in the human body than in a calorimeter, the values of E and H 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

E, H, is there a big distinction?

• For most chemical reactions the difference between these two quantities is negligible

• Typically, PV is a tiny quantity

• For instance, it’s about 0.2% difference for palmitic acid oxidation

H 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

The minimal value

of entropy is a

perfect crystal at

absolute zero

Diffusion is an entropy driven process

Increase in entropy can lead to -G

Thermodynamic quantities

H = enthalpy, the heat content of the systemexothermic = negative, endothermic =

positive; Units: Joules/moleS = entropy, randomization of energy and matter;

positive sign indicates increased entropy; Units: joules/mole(K)

G = 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

G = H – TS

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

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

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

G =0 indicates the system is at equilibrium

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 H for this reaction and this term is positive

Entropy favors water over ice

But recall G is also temperature dependent

Entropy and Enthalpy contributions to melting ice

Biochemical reactions can have different contributions

Why is G called “free” energy?

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

TS is amount of energy that is unavailable to do work

G = H – TS

A G Warning!

• You will see many different G’sG – Gibbs Free Energy

G’o or Go – Standard State Free Energy

energy per mole in standard state (1M)

Go – Standard state Free Energy of Activation

enzyme catalysis