Thermodynamics I
Dr Una Fairbrother
General Biochemistry 2006-07
Module Convenor: Dr Chris Bax,
c.bax@londonmet.ac.uk
Lecture 1: Water (Dr Una Fairbrother).
Lecture 2: Thermodynamics I (Dr Una
Fairbrother).
THERMODYNAMICS
Relationship between heat and movement
BUT better in biology to consider the
relationship between energy and work
For example - can thermodynamics be used
to predict whether any useful work can be
achieved in a chemical reaction?
Does a reaction occur spontaneously or does
energy need to be provided?
DEFINTIONS
SYSTEM
matter within a defined region of the universe
SURROUNDINGS
matter outside that defined region
BOUNDARY
A separator, real or imaginary between system
and surrounding
First Law of Thermodynamics
The total energy of a system and its
surroundings is a constant@ ie energy is
conserved.
DE = EB - EA = Q-W
DE - energy change
EB - energy of system at end of process
EA - energy of system at beginning
Q - heat absorbed by system
W - work done
NOTE about enthalpy
"enthalpy is the amount of energy in a system
capable of doing mechanical work"
Heat content, total heat, enthalpy, H
a thermodynamic quantity equal to the
internal energy of a system plus the product
of its volume and pressure
DH = DE + PDV
H is ENTHALPY, P is pressure, V is volume
change
THERMODYNAMICS
Unfortunately measuring DE (or DH)
does not predict spontaneity of a
reaction.
The next step is to consider another
function of thermodynamics called
ENTROPY (denoted by S) - this is the
measure of randomness or disorder of
system
Second Law of Thermodynamics
A process can occur spontaneously only
if the sum of the entropies of a system
and its surroundings increases@
DS system + DS surroundings > 0
spontaneous reaction)
(for a
Entropy
“the entropy of the universe increases during any spontaneous process”
“universe” just means “the system you’re looking at PLUS its surroundings,
i.e., everything that’s close around it”. System plus surroundings.
Energy spontaneously disperses from being localized to becoming spread
out if it is not hindered from doing so.
A rock will fall if you lift it up and then let go.
Hot frying pans cool down when taken off the stove.
Iron rusts (oxidszes) in the air.
Air in a high-pressure tire shoots out from even a small hole in its side
to the lower pressure atmosphere.
Ice cubes melt in a warm room.
What’s happening in each of those processes?
Energy of some kind is changing from being localized ("concentrated" in
the rock or the pan, etc.) to becoming more spread out
Iron doesn’t have to be hot to
have ‘localized energy’
Iron atoms plus oxygen molecules have more energy
localized within their BONDS than iron rust (iron oxide).
Iron reacts with oxygen releasing energy from higher
energy bonds and form the lower energy bonds in iron
oxide
The difference in energy is dispersed to the
surroundings as ‘heat’ i.e., the reaction is exothermic
and makes molecules in the surroundings move faster
.
Iron spontaneously, but slowly, reacts with oxygen and
each spreads out some of its bond energy to the
surroundings when the iron and oxygen form iron oxide.
System: iron and oxygen, iron oxide.
Surroundings: the nearby air and any moisture or salt in
Motional energy and bond energy
In chemistry the energy that entropy measures as
dispersing is “motional energy”, the translational and
vibrational and rotational energy of molecules
and the DH of phase change energy — both motional
or phase change energy being designated either as
"q" or DH in many equations.
“Bond energy”, the potential energy associated with
chemical bonds that we talked about in the iron
oxidation example,
measured by the potential energy of bond formation
and this is measured by entropy change in
connection with a chemical reaction
FREE ENERGY
Because entropy changes are difficult to
measure a third thermodynamic parameter is
introduced. This is called FREE ENERGY
DG = DH - TDS
DG
DH
DS
T
- free energy change
-enthalpy change
-entropy change
-temperature in Kelvin
Gibbs free energy
The Gibbs free energy was developed in the 1870s by
the American mathematical physicist Willard Gibbs.
A thermodynamic potential which measures the
"useful" work obtainable from an isothermal, isobaric
thermodynamic system.
When a system evolves from a well-defined initial
state to a well-defined final state, the Gibbs free
energy DG equals the work exchanged by the system
with its surroundings, less the work of the pressure
forces, during a reversible transformation of the
system from the same initial state to the same final
state.
How do you use DG?
A reaction can only occur spontaneously if DG is
negative .
If DG is 0 then the reaction is at equilibrium.
A reaction cannot normally occur if DG is positive
( an input of free energy from another reaction is
required to drive it).
Consider the first step of
glycolysis
Glucose + Pi
Glucose 6-phosphate
DG = +13.8 kJ mol-1
Not spontaneous
If the reaction is COUPLED to
the hydrolysis of ATP
ATP + H2O
+
Pi DG= -30.5 kJ mol-1
and
Glucose 6-phosphate DG=+13.8 kJ
ADP
Glucose + Pi
mol-1
Glucose + Pi + ATP + H2O Glucose 6-phosphate +
ADP + Pi
DG = -16.7 kJ mol-1
The net reaction is spontaneous and therefore
permits the phosphorylation of glucose
Relationship between free energy
(G) and the equilibrium constant
(Keq)
A+B
C+D
DG = DGo + RT ln [C][D]
[A][B]
NOTE: ln = loge
DG0 = standard free energy change [ie change
under standard conditions - A,B,C,D present at
1.0M, 298.15 K, atmospheric pressure of 101,235
Pa]
R = gas constant 8.314 JK-1mol-1
T = temperature in K
DG0 and DG0’
At equilibrium: DG = 0, Keq = [C][D]
[A][B]
So: DG = DGo + RT ln [C][D]
[A][B]
Becomes: 0 = DGo + RT ln Keq
Therefore: DGo = -RT ln Keq
In biochemistry there is a further standard
free energy change which occurs at pH=7 this is given the symbol DGo’
Reducing agents
A reducing agent is defined as a
substance that will donate an electron
and become oxidised
eg
Fe2+ Fe3+ + eAn oxidising agent is able to accept an
electron and becomes reduced
eg
Fe3+ + e- Fe2+
HALF-REACTIONS
Reactions showing the electrons being
accepted (or donated) but where the electron
donor (or acceptor) is not shown are called
HALF-REACTIONS
eg H+ + e- H (sometimes written as 1/2 H2 )
Eo’ of -0.42 V
E0 , which is the reduction potential,
a measure of the tendency of H+ to accept
electrons
Half reaction (written as a reduction) E0’ (at pH 7.0), V
Fe3+ + eFe2+
0.77
Dehydroascorbic acid + 2H+ + 2 e0.06
ascorbic acid
Ethanal + 2H+ + 2 eethanol
-0.16
NAD+ + 2H+ + 2 eNADH + H+
0.32
Two half reactions can be added
together to obtain the full reaction
When any two half reactions are
coupled the half reaction with the more
positive reduction potential will proceed
as written (ie as a reduction) driving the
other half reaction backwards (ie as an
oxidation)
Free energy is related to
reduction potential as follows
DG0’ = - n F DE0’
n = number of electrons transferred
F = Faraday’s constant (96,496 JV-1)
DE0’= [E0’of half reaction containing the oxidising
agent] - [E0’of half reaction containing reducing
agent]
[Ethanal + 2H+ + 2 e- ethanol] -[NAD+ + 2H+ + 2 eNADH + H+]
NOTE :
overall the reaction is
Ethanal + NADH + H+ ethanol + NAD+
DE0’ = -0.16 -(-0.32) V
= +0.16V
Therefore
substituting in DG0 = - n FDE0’
DG0’ = -(2)(96,496)(0.16) joules (per mole)
= -30 kJmol-1
Summary
1st Law of Thermodynamics: enthalpy
2nd Law of Thermodynamics: entropy
FREE ENERGY, DG
Relationship between free energy and the equilibrium
constant (Keq)
DG0 and DG0’
Reducing agents
HALF-REACTIONS:Two half reactions can be
added together to obtain the full reaction
Free energy is related to reduction potential
Reading list
Stryer, 5th edition Principles of Biochemistry
http://www.entropysite.com/students_approach.html