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Module 1 and 2 of 'Biochemistry' (AB_1137)
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A detailed elaboration of module 1 and 2 (and a bit of 3) of the course 'Biochemistry' (AB_1137). It helped me pass the course with a 7.
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BIOCHEMISTRYBiomedical Sciences
VU Amsterdam
Module I Basic Thermodynamics and Interactions
Chapter 4 Water, Acids, Bases and Buffers
Chapter 5 Structures of the Major Compounds of the Body
THERMODYNAMICS
Thermodynamics is the reasoning about the forces that allow flow of material and, in general,
the forces that allow life to organize itself in ever more structures. It is a theory that links
various forms of energy and how their conversion can be used to do work.
System = the object of investigation (engine, organism, cell, molecule, etc.).
Environment = everything that is outside the system.
→ Thermodynamics describes the exchange of energy between the system and its
environment.
Variables = parameters that influence the behaviour of the system.
State variables = parameters that describe the state in which the system can be/is found,
irrespective of how it got there. Examples: temperature, position, volume, pressure, etc.
Intensive Variables Extensive Variables
(average out) (add up)
Temperature Energy
Pressure Volume
Concentration Particles/molecules
Often, one can recognize pairs of state variables. A difference in an intensive variable (of the
pair) lead to the exchange of an extensive variable (to reach equilibrium).
FIRST LAW OF THERMODYNAMICS
Energy is conserved; energy can only be transformed
𝛥𝑈 = 𝑞 – 𝑤
U internal energy of a system
q heat
w work
* heat flows from the system to the environment. This flux of heat away from the system is
indicated as a negative heat flow. A positive heat flow means that heat is put into the system.
* a decrease of internal energy leads to positive work.
,Types of internal energy
- Translation energy – rate of movement in space;
- Rotation energy – molecules can rotate;
- Vibration energy – movement of atoms within a molecule;
- Binding energy – energy in chemical bonds between atoms (electrons);
- Potential energies caused by intermolecular interactions (H-binds, Vanderwaals forces);
- Electron energies – energies of electrons within an atom.
Biology generally works with a constant temperature, constant pressure and a (slowly)
varying volume (growth).
𝛥𝑈 = 𝛥ₑ𝑈 = 𝑞 – 𝑝𝛥𝑉
* subscript ‘e’ is used to emphasize that this is an exchange process; heat is exchanged from
one object to another.
Since systems have constant temperatures, energy is being exchanged between the system and
the environment in the form of heat exchange. This heat is:
𝑞ₚ = 𝛥𝑈 + 𝑝𝛥𝑉 ≡ 𝛥𝐻
* subscript ‘p’ is used to emphasize that we work under constant pressure.
ΔH is the change in enthalpy H. It is defined as 𝐻 = 𝑈 + 𝑝𝑉. Enthalpy is the heat that is
released under constant pressure. In many biological processes, volume work is negligible and
𝛥𝑈 ≈ 𝛥𝐻. However, in some processes volume work cannot be ignored.
* volume work at constant volume and fluctuating pressure is
𝑤 = 𝑉 ∙ 𝛥𝑝 and 𝛥𝑈 = 𝑞 − 𝑉𝛥𝑝.
* volume work at constant temperature, constant pressure and slowly fluctuating volume is
𝑤 = 𝑝∆𝑉 and ∆𝑈 = 𝑞ₚ − 𝑝𝛥𝑉.
Enthalpy (H) = the heat that is added to or produced by the system at constant pressure
(enthalpy change). It is a very useful parameter in biology and chemistry, where volumes
change, but temperatures and pressures do not change.
SECOND LAW OF THERMODYNAMICS
The total entropy of an isolated system can never decrease/must increase. The universal
driving force behind reaching equilibriums is probability: a state in which molecules/energy
quanta are spread around many different places is much more probable than a state in which
they are confined in a specific spot.
Multiplicity (W/Ω) = the number of microscopic arrangements that have the same
macroscopic appearance. The higher the number (W), the more probable the state.
, Equilibrium = the most probable (macroscopic) state. Individual particles will still move, but
there are no net. displacements.
* each microscopic state is equally probable.
Entropy (S) = probability in thermodynamics
𝑆 = 𝑘𝐵 ∙ ln(𝑊)
𝛥𝑆 = 𝑆𝑒𝑛𝑑 – 𝑆𝑏𝑒𝑔𝑖𝑛
S entropy
kB Boltzmann constant (1,380649 × 10−23 J K-1)
W multiplicity of a system
ln(𝑥) natural logarithm/log 𝑒 (𝑥)
* concentration differences and pressure differences (𝑊 = 𝑛 + 1) can be considered as
driving forces for macroscopic average movement of molecules/probability drive.
* equilibrium is achieved when there is no thermodynamic driving force anymore.
* highly structured systems (such as all molecules localized in one place) have low
multiplicity and hence low entropy.
At the molecular level, energy is quantitated (defined packages of fixed energy). Energy can
jump to a higher energy quantum; quantitation increases multiplicity and thus entropy.
Multiplicity of a total system
𝑊ₜₒₜ = 𝑊𝐴 ∙ 𝑊𝐵
A temperature difference is a probability drive: it indicated how much entropy can be gained
if energy (in the form of heat) is exchanged.
𝑞
∆ₑ𝑆 =
𝑇
* subscript ‘e’ is used to emphasize that this is an exchange process: heat is exchanged from
one object to another.
* maximal entropy is achieved in the most probable state where the probability drive is zero
(equilibrium).
Upon exchange of heat, the sum of the entropies of the biological system and the environment
must increase. If there is exchange of heat – because we are in thermal contact of heat to keep
the system at constant T – there is exchange of entropy and we cannot be certain anymore that
the system will evolve to the most probable state; only 𝑠𝑦𝑠𝑡𝑒𝑚 + 𝑒𝑛𝑣𝑖𝑟𝑜𝑛𝑚𝑒𝑛𝑡 will.
→ Isolated system: equilibrium when 𝑆𝑠𝑦𝑠 is maximal.
→ Open system (system at constant temperature): equilibrium when 𝑆𝑠𝑦𝑠 + 𝑆𝑒𝑛𝑣 is maximal.
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