Bioenergetics and Metabolism PDF

# Bioenergetics and Metabolism Almost all cellular activities, including movement, membrane transport, and the synthesis of cell constituents, require energy. Consequently, the generation and utilization of metabolic energy is fundamental to all of cell biology. All cells use ATP as their source of...

# Bioenergetics and Metabolism Almost all cellular activities, including movement, membrane transport, and the synthesis of cell constituents, require energy. Consequently, the generation and utilization of metabolic energy is fundamental to all of cell biology. All cells use ATP as their source of metabolic energy, and the mechanisms that cells use for the generation of ATP, either from the breakdown of organic molecules or from photosynthesis, are discussed in this chapter. In addition, this chapter presents an overview of the network of chemical reactions that constitute the metabolism of the cell and are responsible for the synthesis of major cell constituents, including carbohydrates, lipids, proteins, and nucleic acids. ## 3.1 Metabolic Energy and ATP ### Learning Objectives You should be able to: - Interpret the first and second laws of thermodynamics. - Explain how changes in Gibbs free energy determine the direction of chemical reactions. - Summarize the role of ATP in cell physiology. Energy is a basic requirement for many tasks that a cell must perform, so a large portion of the cell's activities is devoted to obtaining energy from the environment and using that energy to drive energy-requiring reactions. As discussed in the previous chapter, enzymes control the rates of virtually all chemical reactions within cells and therefore determine what chemical reactions can take place under the mild conditions of living cells. However, enzymes only control reaction rates; the equilibrium position of chemical reactions is not affected by enzymatic catalysis. Rather, the laws of thermodynamics govern chemical equilibria and determine the energetically favorable direction of all chemical reactions. Many of the reactions that must take place within cells are energetically unfavorable, and are therefore able to proceed only at the cost of additional energy input. Consequently, cells must constantly generate and expend energy derived from the environment. ### The laws of thermodynamics The first and second laws of thermodynamics are particularly applicable to the behavior of cells. The first law is conservation of energy. It states that the total energy of a system and its surroundings remains constant. In other words, energy can neither be created nor destroyed. However, energy can be converted from one form to another. For example, chemical energy can be converted to mechanical energy, such as the contraction of muscle. The second law of thermodynamics states that the degree of disorder in a system (entropy) increases over time. Thus, a system will change spontaneously to a state of greater entropy. Cells intuitively appear to violate the second law, because the molecules of living cells are more ordered (have lower entropy) than their precursors. For example, proteins are ordered polymers of amino acids, so protein synthesis involves a reduction rather than an increase in entropy. However, the cell is not an isolated system, so changes in entropy within a cell are balanced by changes in its environment. For example, formation of a peptide bond releases energy in the form of heat, which increases the thermal motion (and degree of disorder) of surrounding molecules. In order for a biochemical reaction that lowers entropy to occur, the heat released by the reaction must be sufficient to generate a greater increase in entropy in the surroundings. The energetics of biochemical reactions are best described in terms of the thermodynamic function called Gibbs free energy (G), named for Josiah Willard Gibbs. The change in free energy (ΔG) of a reaction combines the effects of changes in enthalpy (the heat released or absorbed during a chemical reaction, ΔH) and entropy (the degree of disorder resulting from a reaction, ΔS) according to the equation (T is the absolute temperature): $AG=AH-TAS$ All chemical reactions spontaneously proceed in the energetically favorable direction, accompanied by a decrease in free energy (ΔG < 0). For example, consider a hypothetical reaction in which A is converted to B: $A \rightarrow B$ If ΔG < 0, this reaction will proceed in the forward direction, as written. If ΔG > 0, however, the reaction will proceed in the reverse direction, and B will be converted to A. The ΔG of a reaction is determined not only by the intrinsic properties of reactants and products, but also by their concentrations and other reaction conditions (e.g., temperature). It is thus useful to define the free-energy change of a reaction under standard conditions. (Standard conditions are considered to be a 1-M concentration of all reactants and products, and 1 atm of pressure). The standard free-energy change (ΔG°) of a reaction is directly related to its equilibrium position because the actual ΔG is a function of both ΔG° and the concentrations of reactants and products. For example, consider the reaction $A \rightarrow B$ The free-energy change can be written as follows: $AG=AG°+ RT In [B]/[A]$ where R is the gas constant. At equilibrium, ΔG = 0 and the reaction does not proceed in either direction. The equilibrium constant for the reaction (K = [B]/[A] at equilibrium) is thus directly related to ΔG° by the above equation, which can be expressed as follows: $0=AG°+RT In K$ or $AG°=-RT In K$ If the actual ratio [B]/[A] is greater than the equilibrium ratio (K), ΔG > 0 and the reaction proceeds in the reverse direction (conversion of B to A). On the other hand, if the ratio [B]/[A] is less than the equilibrium ratio, ΔG < 0 and A is converted to B. The standard free-energy change (ΔG°) of a reaction therefore determines its chemical equilibrium and predicts in which direction the reaction will proceed under any given set of conditions. For biochemical reactions, the standard free-energy change is usually expressed as ΔG', which is the standard free-energy change of a reaction in aqueous solution at pH = 7 - approximately the conditions within a cell. ### The role of ATP Many biological reactions (such as the synthesis of macromolecules) are thermodynamically unfavorable (ΔG > 0) under cellular conditions. In order for such reactions to proceed, an additional source of energy is required. For example, consider the reaction $A \rightarrow B \Delta G= +10 kcal/mol$ The conversion of A to B is energetically unfavorable, so the reaction proceeds in the reverse rather than the forward direction. However, the reaction can be driven in the forward direction by coupling the conversion of A to B with an energetically favorable reaction, such as: $C \rightarrow D \Delta G = -20 kcal/mol$ If these two reactions are combined, the coupled reaction can be written as follows: $A + C \rightarrow B + D \Delta G = -10 kcal/mol$ The ΔG of the combined reaction is the sum of the free-energy changes of its individual components, so the coupled reaction is energetically favorable and will proceed as written. Thus, the energetically unfavorable conversion of A to B is driven by coupling it to a second reaction associated with a large decrease in free energy. Enzymes are responsible for carrying out such coupled reactions in a coordinated manner. The cell uses this basic mechanism to drive the many energetically unfavorable reactions that must take place in biological systems. Adenosine 5'-triphosphate (ATP) plays a central role in this process by acting as a store of free energy within the cell (Figure 3.1). The bonds between the phosphates in ATP are known as high-energy bonds because their hydrolysis is accompanied by a relatively large decrease in free energy. There is nothing special about the chemical bonds themselves; they are called high-energy bonds only because a large amount of free energy is released when they are hydrolyzed within the cell. In the hydrolysis of ATP to ADP plus phosphate (P), ΔG°' = -7.3 kcal/mol. Recall, however, that ΔG°' refers to "standard conditions" in which the concentrations of all products and reactants are 1 M. Actual intracellular concentrations of P; are approximately 10-2 M, and intracellular concentrations of ATP are higher than those of ADP. These differences between intracellular concentrations and those of the standard state favor ATP hydrolysis, so for ATP hydrolysis within a cell, ΔG is approximately -12 kcal/mol. Alternatively, ATP can be hydrolyzed to AMP plus pyrophosphate (PP). This reaction yields about the same amount of free energy as the hydrolysis of ATP to ADP. However, the pyrophosphate produced by this reaction is then itself rapidly hydrolyzed, with a ΔG similar to that of ATP hydrolysis. Thus, the total free-energy change resulting from the hydrolysis of ATP to AMP is approximately twice that obtained by the hydrolysis of ATP to ADP. For comparison, the bond between the sugar and phosphate group of AMP, rather than having high energy, is typical of covalent bonds; for the hydrolysis of AMP, ΔG° = -3.3 kcal/mol. Because of the accompanying decrease in free energy, the hydrolysis of ATP can be used to drive other energy-requiring reactions within the cell. For example, the first reaction in glycolysis (discussed in the next section) is the conversion of glucose to glucose-6-phosphate. The reaction can be written as follows: Glucose + Phosphate -> Glucose-6-phosphate + H₂O Because this reaction is energetically unfavorable as written (ΔG°' = +3.3 kcal/mol), it must be driven in the forward direction by being coupled to ATP hydrolysis (ΔG° = -7.3 kcal/mol): ATP + H₂O -> ADP + HPO₄^2- The combined reaction can be written as follows: Glucose + ATP -> Glucose-6-phosphate + ADP The free-energy change for this reaction is the sum of the free-energy changes for the individual reactions, so for the coupled reaction, ΔG = -4.0 kcal/mol, favoring glucose-6-phosphate formation. Other molecules, including other nucleoside triphosphates (e.g., guanosine 5'-triphosphate, GTP), also have high-energy bonds and can be used as ATP is used to drive energy-requiring reactions. For most reactions, however, ATP provides the free energy. The energy-yielding reactions within the cell are therefore coupled to ATP synthesis, while the energy-requiring reactions are coupled to ATP hydrolysis. The high-energy bonds of ATP thus play a central role in cell metabolism by serving as a usable storage form of free energy to drive energy-requiring reactions. Hydrolysis of ATP can also drive the energy-requiring transport of molecules against a concentration gradient (active transport, see Figure 2.40) and cell movements, such as muscle contraction and the movements of chromosomes during mitosis (discussed in Chapter 14). ## 3.1 Review The behavior of cells is governed by the first and second laws of thermodynamics. Gibbs free energy combines the effects of entropy and enthalpy to predict the direction of biochemical reactions, which proceed in the energetically favorable direction. ATP serves as a store of free energy, which can be used to drive energy-requiring reactions within cells.