6.1 Energy and Energy Transformations
KEY CONCEPTS
By the end of this section, you will be able to do the following:
- Give cellular examples of kinetic and potential energy, and explain why those examples fall into the category of kinetic or potential energy.
- Describe examples of cellular processes that involve transformations of energy (first law of thermodynamics) and increases in entropy (second law of thermodynamics).
- Apply the concept of free energy to determine whether cellular processes (reactions) are endergonic or exergonic, and which of these processes can be used to do cellular work.
- Explain how activation energy limits chemical reaction rates
The challenge for all living organisms is to obtain energy from their surroundings in forms that they can transfer or transform into usable energy to do work. Examples of the types of work that cells need to do include building complex molecules, transporting materials, powering the beating motion of cilia or flagella, contracting muscle fibers to create movement, and reproducing. This chapter will explore the different types of energies, the laws of thermodynamics, and the fundamentals of energy during chemical reactions.
Thermodynamics
Thermodynamics refers to the study of energy and energy transfer involving physical matter. The matter and its environment relevant to a particular case of energy transfer are classified as a system, and everything outside that system is the surroundings. For instance, when heating a pot of water on the stove, the system includes the stove, the pot, and the water. Energy is transferred within the system (between the stove, pot, and water). There are two types of systems: open and closed. An open system is one in which energy and matter can transfer between the system and its surroundings. The stovetop system is open because it can lose heat (a form of energy) and water vapour (a form of matter) into the air. A closed system is one that can transfer energy but not matter to its surroundings. A boiling pot with a lid on would be an example of a closed system as heat can be lost but not matter (water vapour).
Biological organisms are open systems. Energy is exchanged between them and their surroundings, as they consume energy-storing molecules and release energy to the environment by doing work. Like all things in the physical world, energy is subject to the laws of physics. The laws of thermodynamics govern the transfer of energy in and among all systems in the universe. To begin our understanding of thermodynamics, first we need to establish what energy is, and how we can categorize that energy.
Energy Types
We define energy as the ability to do work. As you have likely learned, energy exists in different forms. For example, electrical, light and heat energy are all different energy types. In order to appreciate the way energy flows into and out of biological systems, it is important to understand more about the different energy types that exist in the physical world, including kinetic and potential energy.
Kinetic Energy
The energy that an object has when in motion is referred to as kinetic energy. For example, an airplane in flight has considerable kinetic energy. Moving objects produce kinetic energy as their motion involves the work done by some external force, causing movement. Objects in motion also have the capability of doing work. Think of a wrecking ball; even a slow-moving wrecking ball can do considerable damage to other objects. Rapid molecule movement in the air (which produces heat), cilia and flagella on cells, and the movement of enzymes within a cell all have kinetic energy (Figure 6.2A). Heat energy is a type of kinetic energy due to the random movement of particles within a substance.
Potential Energy
The energy an object has due to its position or structure is referred to as potential energy. Potential energy can be converted into kinetic energy by work done through external forces. Other examples of potential energy include water’s energy held behind a dam, or the energy stored in chemical bonds of molecules that release energy when broken, such as glucose molecules (Figure 6.2B). We associate potential energy not only with the object’s location (such as a child sitting on a tree branch), but also with the object’s structure. A spring on the ground has potential energy if it is compressed; so does a tautly pulled rubber band.
The very existence of living cells relies heavily on structural potential energy. In fact, there is chemical potential energy stored within the bonds of all the food molecules we eat. This is because these bonds can release energy when broken. Another type of potential energy we have already seen is that of concentration or electrochemical gradients across cellular membranes. If there is a higher concentration of particles on one side of a membrane than another, that high concentration of particles contains potential energy (just like the water behind a dam wall) that can be used to do work like secondary active transport.
Link to Learning
Laws of Thermodynamics
Energy Cannot Be Created or Destroyed
To understand how biological organisms transform energy to do work, we must first ensure a good understanding of two laws of thermodynamics. The first law of thermodynamics states that the total amount of energy in the universe is constant. According to the first law of thermodynamics, energy may transfer from place to place or transform into different forms, but it cannot be created or destroyed. The transfers and transformations of energy take place around us all the time. Light bulbs transform electrical energy into light energy. Gas stoves transform chemical energy from natural gas into heat energy. Plants perform one of the most biologically useful energy transformations on earth: that of converting sunlight energy into the chemical energy stored within organic molecules (Figure 6.3). Chemical energy stored within organic molecules such, as sugars and fats, transforms through a series of cellular chemical reactions into energy within ATP molecules. Energy in ATP molecules is easily accessible to do cellular work. For example, the chemical energy in ice cream can ultimately be transformed (via many cellular processes) into the kinetic energy of riding a bicycle (Figure 6.3).
Note that it is never appropriate to say that cells “create” or “generate” energy; for example it is incorrect to say that mitochondria generate energy for cells. This is because energy cannot be created, only transformed, as per the first law of thermodynamics. It is appropriate to say that mitochondria generate (or synthesize) ATP, but not that they generate energy.
None of the energy transfers that we have discussed, along with all energy transfers and transformations in the universe, is completely efficient. In every energy transfer, some amount of energy is lost to the environment in a form that is unusable. In most cases, this form is heat energy, which cannot be used to do work. For example, when an airplane flies through the air, it loses some of its energy as heat energy due to friction with the surrounding air. This friction actually heats the air by temporarily increasing air molecule speed. Likewise, some energy is lost as heat energy during cellular metabolic reactions. This is good for warm-blooded creatures like us, because heat energy helps to maintain our body temperature. Note that none of this “lost” energy is destroyed (as per the first law of thermodynamics) nor does it leave the universe. That “lost” energy is just no longer available to the open system (biological organism) of interest.
Entropy: The Drive Towards Disorder
A living cell’s primary tasks of obtaining, transforming, and using energy to do work may seem simple. However, the second law of thermodynamics explains why these tasks are harder than they appear. An important concept in physical systems is that of order and disorder (or randomness). The more energy that a system loses to its surroundings, the less ordered and more random the system. Scientists refer to the measure of randomness or disorder within a system as entropy. High entropy means high disorder and low energy. For example, a gas has higher entropy than a liquid or solid because the gas particles can move around more freely (with less order) than those in a liquid or solid (Figure 6.4). To better understand entropy, think of a student’s bedroom. If no energy or work were put into it, the room would quickly become messy. It would exist in a very disordered state, one of high entropy. Energy must be put into the system, in the form of the student doing work and putting everything away, in order to bring the room back to a state of cleanliness and order. For a biological example, the broken-down sugar molecules during digestion are in a disordered state in the body. Cells must use energy to form more ordered structures (like organelles) from these molecules to use in the body. Cells constantly need to gain energy from their environment to keep up the “fight” against entropy.
Free Energy
Cells need to perform an enormous amount of work in order to maintain their essential functions and avoid falling into a disordered state. One concept that is crucial for understanding this notion is that of free energy. But what exactly is free energy? Free energy is defined as the capacity of a system’s energy that is available to do work. We usually call this free energy Gibbs free energy (abbreviated with the letter G) after Josiah Willard Gibbs, the scientist who developed the measurement. According to the second law of thermodynamics, the amount of energy available to do work decreases as energy is transformed one form to another, resulting in entropy of a closed system always increasing over time. Whenever energy is converted within a cell, such as during a chemical reaction, it is possible to determine whether the energy transformation leads to an increase or decrease in the amount of available free energy by calculating delta G (∆G).
Every chemical reaction involves a change in free energy, ∆G. We can calculate the change in free energy for any system that undergoes such a change, such as a chemical reaction. To calculate ∆G, we use this formula: ΔG = ΔH − TΔS, where ∆H is the change in enthalpy, which is the total energy in the system, T refers to the absolute temperature in Kelvin (°C + 273) and ∆S is the change in entropy. It is important to note that from the formula we can decipher that increases in entropy often result in negative ΔG values, whereas decreases in entropy often result in positive ΔG values. This will become relevant in the section below on exergonic and endergonic processes.
Some technical notes about ∆G values: We express a chemical reaction’s standard free energy change as an amount of energy per mole of the reaction product (either in kilojoules or kilocalories, kJ/mol or kcal/mol; 1 kJ = 0.239 kcal) under standard pH, temperature, and pressure conditions. We generally calculate standard pH, temperature, and pressure conditions at pH 7.0 in biological systems, 25 °C, and 100 kPa (1 atm pressure). Cellular conditions vary considerably from these standard conditions, and so standard calculated ∆G values for biological reactions will be different inside the cell.
Exergonic and Endergonic Processes
We call reactions that have a negative ∆G and release free energy exergonic, and this released free energy can be used to perform work in the cell. Think: exergonic means energy is exiting the process. If energy is released during a chemical reaction, then the resulting value from the above equation will be a negative number, ∆G < 0. A negative ∆G also means that the reaction’s products have less free energy than the reactants, as free energy was released during the reaction (Figure 6.5). We also refer to these reactions as spontaneous reactions because they can occur without adding energy into the system. Contrary to the everyday use of the term, a spontaneous reaction is not one that suddenly or quickly occurs. Rusting iron is an example of a spontaneous reaction that occurs slowly, little by little, over time. Understanding which chemical reactions are spontaneous and release free energy is extremely useful for biologists, as these reactions can be harnessed to perform work inside the cell.
We call reactions that have a positive ∆G and consume free energy endergonic or non-spontaneous. If a chemical reaction requires an energy input rather than releasing energy, then the ∆G for that reaction will be a positive value, ∆G > 0. In this case, the products have more free energy than the reactants (Figure 6.5). Thus, we can think of the reactions’ products as energy-storing molecules. An endergonic reaction will not take place on its own without adding free energy.
For an example of endergonic and exergonic reactions, let’s consider the synthesis and breakdown of the food molecule, glucose. Remember that building complex molecules, such as sugars, from simpler ones requires energy. Therefore, the chemical reactions used to make glucose (e.g., via photosynthesis) are endergonic reactions and have a positive ΔG. Alternatively, the process of breaking glucose down into simpler molecules releases energy in a series of exergonic reactions, with an overall negative ΔG that is equal to but the opposite sign from the ΔG of glucose synthesis. Sugar breakdown involves spontaneous reactions, but these reactions do not occur instantaneously. Later sections will provide more information about what else is needed to make even spontaneous reactions happen more efficiently.
Cells Do Work to Avoid Chemical Equilibrium
An important concept in studying metabolism and energy is that of chemical equilibrium. Most chemical reactions are reversible. They can proceed in both directions, releasing energy into their environment in one direction, and absorbing it from the environment in the other direction. The same is true for the chemical reactions involved in cellular metabolism, such as the breaking down and building up of proteins into and from individual amino acids, respectively. Reactants within a closed system will undergo chemical reactions in both directions until they reach a state of equilibrium, which is one of the lowest possible free energy states and a state of maximal entropy. To push the reactants and products away from a state of equilibrium requires energy. Either reactants or products must be added, removed, or changed. If a cell were a closed system, its chemical reactions would reach equilibrium, and it would die because there would be insufficient free energy left to perform the necessary work to maintain life. In a living cell, chemical reactions are constantly moving towards equilibrium, but never reach it. This is because a living cell is an open system. Materials pass in and out, the cell recycles the products of certain chemical reactions into other reactions, resulting in never reaching chemical equilibrium. In this way, living organisms are in a constant energy-requiring, uphill battle against equilibrium and entropy. This constant energy supply ultimately comes from sunlight, which produces nutrients in the photosynthesis process.
Activation Energy
There is another important concept that we must consider regarding chemical reactions. Even though exergonic reactions have a net release of energy, they still require a small amount of initial energy input before they can proceed with their energy-releasing steps. We call this small amount of energy input necessary for all chemical reactions to occur the activation energy (or free energy of activation) abbreviated as EA (Figure 6.6). The activation energy of a particular reaction determines the rate at which it will proceed. The higher the activation energy, the slower the chemical reaction. Iron rusting illustrates an inherently slow reaction. This reaction occurs slowly over time because of its high EA. Additionally, burning many fuels, which is strongly exergonic, will take place at a negligible rate unless sufficient heat from a spark overcomes their activation energy. However, once they begin to burn, the chemical reactions release enough heat, and this heat energy can supply the activation energy for burning of the surrounding fuel molecules. Like these reactions outside of cells, the activation energy for most cellular reactions is too high for heat energy to overcome at efficient rates. In other words, for important cellular reactions to occur at appreciable rates (number of reactions per unit time), their activation energies must be lowered.
Why would an energy-releasing, negative ∆G reaction require some energy to proceed? The reason lies in the steps that take place during a chemical reaction. During chemical reactions, certain chemical bonds break and new ones form. For example, when a glucose molecule breaks down, bonds between the molecule’s carbon atoms break. Since these are energy-storing bonds, energy is released when the bond is broken. However, to get them into a state that allows the bonds to break, the molecule must be somewhat contorted. A small energy input is required to achieve this contorted state. This contorted state is the transition state (Figure 6.6), and it is a high-energy, unstable state. For this reason, reactant molecules do not last long in their transition state, but very quickly proceed to the chemical reaction’s next steps.
Links to Learning
Watch an animation of the move from free energy to transition state from Shomu’s Biology.
A couple useful video from Crash Course Biology to clarify the concepts (don’t get overwhelmed by all the formulas in the videos):