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Class notes 4BBY1013 (Biochemistry)
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full lecture notes for biochemistry topic energy considerations taken in 2022/2023.
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Lecture 03 – Energy ConsiderationsLearning Outcomes
Demonstrate and understanding of delta H as change in bond energy or enthalpy in a chemical ratio
Explain the concepts of entropy and Hibbs free energy, ‘standard’ free energy changes, and the equation: delta G = delta H
- T(delta S)
Define the terms endothermic, exothermic, endergonic, and exergonic
Explain the relationship between the free energy change delta G, of a reaction and the direction of that reaction
Demonstrate an understanding of the concept of equilibrium in a chemical reaction
Calculate the equilibrium constant of a reaction from the standard free energy change and vice versa
Explain the concept of ‘coupling’ of biochemical reactions and outline the role of ATP as the ‘energy carrier’ of the cell
What determines whether a chemical reaction can take place?
- Energy considerations (thermodynamics)
- Living things must obey thermodynamics laws
First law: Energy cannot be created or destroyed but it can change forms
The study of thermodynamics is the study of energy transformations
What forms of energy are involved in chemical reaction?
- Chemical bond energy (enthalpy of heat content of a molecule) symbol H
- When bonds are made energy is released- the ‘stronger’ the bond, the more energy is released
- Energy is required to break bonds. The stronger the bond the more energy is requires
- We call the energy that is released when a bond is formed, or the energy required to break it, the bond energy.
Enthalpy
The enthalpy change (ΔH) of a chemical reaction is the sum of energy used when bonds are broken during the
reaction and the energy released when new bonds are formed.
ΔH –ve, heat is lost from the molecules and released to the surroundings: The reaction is exothermic.
ΔH +ve, heat is taken up by the molecules, surroundings cool down: The reaction is endothermic.
Enthalpy changes of some representative chemical reactions
Reaction type Reaction ΔH, kJ mol-1
Oxidation glucose + 6O2 → 6CO2 + 6H2O - 2813
Hydrolysis sucrose + H2O → glucose + fructose - 20.1
Hydrolysis glucose-6-phosphate + H2O → glucose + Pi - 12.5
Neutralization NaOH + HCl → NaCl + H2O - 57.7
Reactions with -ve ΔH are more LIKELY to occur.
But SOME reactions with +ve ΔH DO occur.
Enthalpy and Entropy
Reactions with -ve ΔH are more LIKELY to occur.
, But SOME reactions with +ve ΔH DO occur
We also have to consider entropy, S.
Level of disorder, or number of ways something can be arranged
Second law: all processes must increase the entropy of the universe.
Entropy
Living things must struggle against the tendency to disorder (increased entropy)
Reactions where entropy increases (ΔS+ve) are favorable.
Cells don’t defy the 2nd law of thermodynamics
- Cells use energy released by oxidation of foods to maintain organization and do work, but the breakdown of
food also contributes to entropy of universe:
- By release of small molecules such as CO2 into the environment.
- By releasing heat into the environment.
Gibbs free energy change, ΔG
To determine whether a reaction can take place we need to consider both enthalpy and entropy changes.
Not very convenient.
Both are put together in an equation that defines Gibbs 'free' energy - the useful energy 'available' from a
reaction, ΔG.
ΔG = ΔH - T ΔS
T is the temperature in Kelvin.
ΔG has units, kJmol-1
What makes a reaction spontaneous?
The useful energy ‘available’ from a reaction, ΔG
An exergonic reaction: A reaction can only occur spontaneously if ΔG is -ve
An endergonic reaction: A rection cannot occur spontaneously if ΔG is +ve
Note that ‘can’ does not always mean ‘will’ as ΔG does not tell us about the RATE of a reaction.
The value of ΔG changes as a reaction proceeds towards equilibrium
At equilibrium the G = 0
For the energetically favorable reaction Y X
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