Lecture 28: Energy calculations and Metabolism

Energy and Metabolism

Instructor and Contact Information

  • Instructor: Dr. Katrine Wallis

  • Email: Katrine.wallis@warwick.ac.uk

  • Location: LF 130 Cellular and Molecular Biology

  • Office: D134

Recap from Last Lecture

System and Surroundings

  • System: The specific part of the universe being studied or observed.

  • Surroundings: Everything outside the system, including the environment and other factors that may influence the system.

Gibbs Free Energy (ΔG)

  • Role: A thermodynamic quantity that helps predict the spontaneity of a reaction, determining whether a reaction can occur naturally without external input.

  • Equation: ΔG = ΔH – TΔS

  • Variables:

    • ΔG: Change in Gibbs Free Energy

    • ΔH: Change in enthalpy, reflecting the total heat content

    • ΔS: Change in entropy, associated with disorder or randomness in the system

    • T: Absolute temperature measured in Kelvin

Reactions Close to Equilibrium

  • Forward and Backward Reactions: The same enzyme catalyzes both the forward and reverse reactions, indicating that they are reversible under certain conditions.

  • Example Reaction:

    • Glucose-6-phosphate ↔ Fructose-6-phosphate

  • Equilibrium Condition:

    • If ΔG = 0, this indicates that there is no overall change in the concentration of reactants and products, affirming a state of dynamic equilibrium.

  • Concentration Effects:

    • Increased levels of glucose-6-phosphate ([G-6-P]) lead to a negative change in ΔG (indicating the spontaneity of glycolysis) and favor the formation of fructose-6-phosphate. Conversely, increased concentrations of fructose-6-phosphate ([F-6-P]) result in a positive change in ΔG, indicating the reverse reaction (gluconeogenesis) becomes more favorable.

Reactions with Large ΔG

  • Examples of Reactions Far From Equilibrium:

    • Reaction: Glucose + ATP → Glucose-6-phosphate + ADP

    • Implication: A negative ΔG (<0) signifies that these reactions are typically considered irreversible under physiological conditions, meaning they proceed in a single direction under most cellular environments.

Mass Action Ratio

  • Definition: The mass action ratio (q) describes the ratio of products to reactants in a chemical reaction, giving insight into how far the system is from equilibrium.

    • Equilibrium Condition: At equilibrium, the mass action ratio q equals the equilibrium constant K_D.

  • Example Calculation: For the reaction from F-1,6-bP to G3P + DHAP, the mass action ratio can be expressed as:

    • q = [C][D]/[A][B]

Standard Free Energy

  • Equation:

    • ΔG = ΔG° + RT ln q

  • Units: Measured in J/mol or kcal/mol

  • Variables:

    • R: The gas constant

    • T: Absolute temperature in Kelvin

    • ΔG°: Standard Free Energy change at 1M concentration at 25°C and 1 atm pressure

    • ΔG°’: Standard Free Energy change at pH 7, 25°C, and 1 atm.

  • Dependencies: The value of ΔG varies based on the nature of the reaction and the concentrations of reactants and products involved.

Equilibrium Conditions

  • Condition: When ΔG = 0, this denotes that there is no net change in free energy, indicating that the rates of formation and degradation are balanced.

  • Example: The rate of formation equals the rate of degradation of F-1,6-bP.

Equilibrium Constant and Free Energy

  • Relationship Established:

    • ΔG = ΔG°’ + RT ln q

    • When ΔG = 0, this indicates a direct relationship with the equilibrium constant (K_eq).

  • Equation for K_eq:

    • ln K_eq = -ΔG°’/RT

  • Example Calculation of K_eq:

    • To derive K_eq, use K_eq = e^(-ΔG°’/2.5) where R = 8.31 x 10^-3 kJ mol-1 K-1 and T = 298 K.

Changes in Free Energy and K_eq

  • Significance of ΔG°’: A small change in the standard free energy change (ΔG°’) can lead to significant variations in the equilibrium constant (K_eq).

  • Example Values of ΔG°’:

    • ΔG°’ = -5.7 kJ/mol leads to K_eq = 10

    • ΔG°’ = -11.5 kJ/mol leads to K_eq = 100

    • ΔG°’ = -17.3 kJ/mol leads to K_eq = 1000

Example Reaction

  • Specific Conversion: Dihydroxyacetone-phosphate is converted to glyceraldehyde-3-phosphate, showcasing the type of metabolic conversions that are crucial in cellular respiration.

Calculation of ΔG

  • At Equilibrium:

    • G3P/DHAP Ratio = 0.0475

    • Equations for ΔG calculation include:

    • ΔG = ΔH – TΔS

    • q = [C][D]/[A][B]

    • R = 8.31 x 10^-3 kJ mol-1 K-1 and T = 298 K

Calculate ΔG°’

  • Similar Parameters: The calculation of ΔG°' follows similar conditions and parameters as that used on the previous page.

ΔG Calculation In Vivo

  • Concentration Measurements: Concentrations measured as follows:

    • [DHAP] = 2 x 10^-4 M; [G3P] = 3 x 10^-6 M.

  • Equations for ΔG calculation are identical to those used in prior calculations.

Energy Storage and Transfer

  • Importance of Coenzyme A: Recognized as a critical cofactor in metabolic pathways, particularly in the transfer of acetyl groups.

  • ATP: Identified as the primary energy carrier in cells, often referred to as 'energy-rich phosphate bonds' as stated by Friz Lippman, highlighting its essential role in energy transfer.

Energy of ATP Hydrolysis

  • Hydrolysis Reactions:

    • ATP + H2O → ADP + Pi: ΔG°’ = -30.5 kJ/mol

    • ADP + Pi → ATP + H2O: ΔG°’ = 30.5 kJ/mol

    • ATP Conversion to AMP + PPi: ΔG°’ = -45.6 kJ/mol

  • Interconversion Process: Adenylate Kinase facilitates the reaction:

    • ATP + AMP ↔ 2 ADP

  • Other Important Phosphate Sources: Creatine phosphate, PEP, and 1,3-bisphosphoglycerate are also vital energy-rich compounds used in energy transfer.

Energy Released During Reactions

  • General Observation: Most physiological reactions do not reflect equilibrium; instead, they operate under non-equilibrium conditions, allowing for the release of energy.

  • Example Calculation:

    • Reaction: ATP + H2O → ADP + Pi results in ΔG estimated to be higher than standard conditions due to the influence of non-equilibrium states.

ATP’s Energy Bonds

  • Common Misconception: ATP is often considered to possess inherently 'high energy' bonds; however, the actual energy is derived from its movement away from equilibrium rather than the structural properties of the bond themselves.

ATP Functions

  • Diverse Roles of ATP:

    • ATP plays a fundamental role in various biological processes, including:

      • Energy production through light (photosynthesis) and biochemical reactions (respiration).

      • Synthesis of vital macromolecules such as DNA, RNA, and proteins.

      • Cellular movement, activation of transport mechanisms for molecules, and neurotransmission.

ATP Usage

  • Physiological Example: During a cheetah's sprinting at 110 km/h, it utilizes 55 grams of ATP per second.

  • Short-Timed Sprint Impact: In just 10 seconds, a cheetah can deplete 0.55 kg of ATP, corresponding to roughly 3% of its total body weight.

  • Human ATP Needs: An average human requires approximately 70 kg of ATP per day to sustain basic physiological functions and activities.

Effect of Coupling Reactions on K_eq

  • Illustration of Coupled Reactions:

    • Reaction 1: Glucose + Pi → Glucose-6-phosphate + H2O: ΔG°’ = +14.0 kJ/mol

    • Reaction 2 (ATP Hydrolysis): ATP + H2O → ADP + Pi: ΔG°’ = -30.5 kJ/mol

    • Net Reaction Analysis: Combining these reactions yields a net change in free energy: ΔG°’ = -16.5 kJ/mol, demonstrating how coupling can favorably shift the equilibrium.

Redox Reactions in Energy Storage

  • Central Reactions:

    • The oxidation of glucose during glycolysis results in ATP generation, a crucial energy production line in metabolism.

    • The oxidation of Acetyl-CoA in the citric acid cycle leads to the production of reduced coenzymes NADH and FADH2.

    • Finally, oxidative phosphorylation uses these reduced compounds to synthesize ATP efficiently.

Overview of Redox Reactions

  • Details of Redox Reactions:

    • Key aspects of these reactions include:

      • Oxidation: The process of donating electrons.

      • Reduction: The process of accepting electrons, establishing a fundamental basis for metabolic energy transformations.

Redox Potential

  • Definition: Describes the inherent tendency of a chemical compound to accept electrons, important in determining the direction of electron flow in biochemical reactions.

  • Example Half Reactions:

    • Comparisons between compounds like ferredoxin and NAD+ demonstrate differing potentials that guide their respective reduction and oxidation behaviors.

Example of Redox Reaction Calculation

  • Overall Reaction:

    • NADH + ½ O2 → H2O + NAD+

    • Calculation of free energy released during this electron transfer process is critical for understanding its efficiency.

Free Energy and Redox Potentials

  • Equation:

    • ΔG°’ = -nFΔE°’

  • Example: For the reaction NADH + ½ O2 → H2O + NAD+, with ΔE°’ calculated to be 1.14 V, resulting in ΔG°’ = -220 kJ/mol, indicating a highly exergonic reaction.

Summary

  • Key Points:

    • Gibbs free energy is fundamentally related to the spontaneity and specific conditions of chemical reactions.

    • ATP serves as the main energy carrier, facilitating the execution of endergonic reactions essential for cellular function and metabolism.

    • The concept of redox potential is crucial for understanding electron transfer processes and is vital for various metabolic pathways.