Energy

Metabolism of Carbohydrates

Organisms break down carbohydrates to produce energy for cellular processes, and photosynthetic plants produce carbohydrates.

Learning Objectives

Analyze the importance of carbohydrate metabolism to energy production

Key Takeaways

Key Points

  • The breakdown of glucose living organisms utilize to produce energy is described by the equation: [latex]{\text{C} }_{ 6 }{\text{H} }_{ 12 }{\text{O} }_{ 6 }+6{\text{O} }_{ 2 }\rightarrow 6{\text{CO} }_{ 2 }+6{\text{H} }_{ 2 }\text{O}+\text{energy}[/latex].
  • The photosynthetic process plants utilize to synthesize glucose is described by the equation: [latex]6\text{CO}_{ 2 }+6{\text{H} }_{ 2 }\text{O}+\text{energy}\rightarrow {\text{C} }_{ 6 }{\text{H} }_{ 12 }{\text{O} }_{ 6 }+6\text{O}_{ 2 }[/latex].
  • Glucose that is consumed is used to make energy in the form of ATP, which is used to perform work and power chemical reactions in the cell.
  • During photosynthesis, plants convert light energy into chemical energy that is used to build molecules of glucose.

Key Terms

  • adenosine triphosphate: a multifunctional nucleoside triphosphate used in cells as a coenzyme, often called the “molecular unit of energy currency” in intracellular energy transfer
  • glucose: a simple monosaccharide (sugar) with a molecular formula of [latex]{\text{C} }_{ 6 }{\text{H} }_{ 12 }{\text{O} }_{ 6 }[/latex]C6H12O6; it is a principal source of energy for cellular metabolism

Metabolism of Carbohydrates

Carbohydrates are one of the major forms of energy for animals and plants. Plants build carbohydrates using light energy from the sun (during the process of photosynthesis), while animals eat plants or other animals to obtain carbohydrates. Plants store carbohydrates in long polysaccharides chains called starch, while animals store carbohydrates as the molecule glycogen. These large polysaccharides contain many chemical bonds and therefore store a lot of chemical energy. When these molecules are broken down during metabolism, the energy in the chemical bonds is released and can be harnessed for cellular processes.

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All living things use carbohydrates as a form of energy.: Plants, like this oak tree and acorn, use energy from sunlight to make sugar and other organic molecules. Both plants and animals (like this squirrel) use cellular respiration to derive energy from the organic molecules originally produced by plants

Energy Production from Carbohydrates (Cellular Respiration )

The metabolism of any monosaccharide (simple sugar) can produce energy for the cell to use. Excess carbohydrates are stored as starch in plants and as glycogen in animals, ready for metabolism if the energy demands of the organism suddenly increase. When those energy demands increase, carbohydrates are broken down into constituent monosaccharides, which are then distributed to all the living cells of an organism. Glucose (C6H12O6) is a common example of the monosaccharides used for energy production.

Inside the cell, each sugar molecule is broken down through a complex series of chemical reactions. As chemical energy is released from the bonds in the monosaccharide, it is harnessed to synthesize high-energy adenosine triphosphate (ATP) molecules. ATP is the primary energy currency of all cells. Just as the dollar is used as currency to buy goods, cells use molecules of ATP to perform immediate work and power chemical reactions.

The breakdown of glucose during metabolism is call cellular respiration can be described by the equation:

[latex]{\text{C} }_{ 6 }{\text{H} }_{ 12 }{\text{O} }_{ 6 }+6{\text{O} }_{ 2 }\rightarrow 6{\text{CO} }_{ 2 }+6{\text{H} }_{ 2 }\text{O}+\text{energy}[/latex]

Producing Carbohydrates (Photosynthesis)

Plants and some other types of organisms produce carbohydrates through the process called photosynthesis. During photosynthesis, plants convert light energy into chemical energy by building carbon dioxide gas molecules (CO2) into sugar molecules like glucose. Because this process involves building bonds to synthesize a large molecule, it requires an input of energy (light) to proceed. The synthesis of glucose by photosynthesis is described by this equation (notice that it is the reverse of the previous equation):

[latex]6\text{CO}_{ 2 }+6{\text{H} }_{ 2 }\text{O}+\text{energy}\rightarrow {\text{C} }_{ 6 }{\text{H} }_{ 12 }{\text{O} }_{ 6 }+6\text{O}_{ 2 }[/latex]

As part of plants’ chemical processes, glucose molecules can be combined with and converted into other types of sugars. In plants, glucose is stored in the form of starch, which can be broken down back into glucose via cellular respiration in order to supply ATP.

Free Energy Changes in Chemical Reactions

ΔG determines the direction and extent of chemical change.

Learning Objectives

Recall the possible free energy changes for chemical reactions.

Key Takeaways

Key Points

  • If the free energy of the reactants is greater than that of the products, the entropy of the world will increase when the reaction takes place as written, and so the reaction will tend to take place spontaneously.
  • If the free energy of the products exceeds that of the reactants, then the reaction will not take place.
  • An important consequence of the one-way downward path of the free energy is that once it reaches its minimum possible value, net change comes to a halt.
  • In a spontaneous change, Gibbs energy always decreases and never increases.

Key Terms

  • spontaneous change: A spontaneous process is the time-evolution of a system in which it releases free energy (usually as heat) and moves to a lower, more thermodynamically stable energy state.

The Direction and Extent of Chemical Change

ΔG determines the direction and extent of chemical change. Remember that ΔG is meaningful only for changes in which the temperature and pressure remain constant. These are the conditions under which most reactions are carried out in the laboratory. The system is usually open to the atmosphere (constant pressure) and the process is started and ended at room temperature (after any heat that has been added or which was liberated by the reaction has dissipated.)

The importance of the Gibbs function can hardly be over-stated: it determines whether a given chemical change is thermodynamically possible. Thus, if the free energy of the reactants is greater than that of the products, the entropy of the world will increase and the reaction takes place spontaneously. Conversely, if the free energy of the products exceeds that of the reactants, the reaction will not take place.

In a spontaneous change, Gibbs energy always decreases and never increases. This of course reflects the fact that the entropy of the world behaves in the exact opposite way (owing to the negative sign in the TΔS term). Here is an example:

[latex]{\text{H}}_{2}\text{O}(\text {liquid}) \rightarrow {\text{H}}_{2}\text{O} (\text{ice})[/latex]

Water below zero degrees Celsius undergoes a decrease in its entropy, but the heat released into the surroundings more than compensates for this so the entropy of the world increases, the free energy of the H2O diminishes, and the process proceeds spontaneously.

An important consequence of the one-way downward path of the free energy is that once it reaches its minimum possible value, net change comes to a halt. This, of course, represents the state of chemical equilibrium. These relations are summarized as follows:

  • [latex]\Delta G < 0[/latex]: The reaction will occur spontaneously to the right.
  • [latex]\Delta G > 0[/latex]: The reaction will occur spontaneously to the left.
  • [latex]\Delta G = 0[/latex]: The reaction is at equilibrium and will not proceed in either direction.

Conditions for Spontaneous Change

Recall the condition for spontaneous change:

ΔG = ΔH – TΔS < 0

where ΔG = change in Gibbs free energy, ΔH = change in enthalpy, T = absolute temperature, and ΔS = change in entropy

It is apparent that the temperature dependence of ΔG depends almost entirely on the entropy change associated with the process. (it is appropriate to say “almost” because the values of ΔH and ΔS are themselves slightly temperature dependent; both gradually increase with temperature). In particular, notice that in the above equation the sign of the entropy change determines whether the reaction becomes more or less spontaneous as the temperature is raised.

For any given reaction, the sign of ΔH can also be positive or negative. This means that there are four possibilities for the influence that temperature can have on the spontaneity of a process:

Case 1: ΔH < 0 and ΔS > 0

Under these conditions, both the ΔH and TΔS terms will be negative, so ΔG will be negative regardless of the temperature. An exothermic reaction whose entropy increases will be spontaneous at all temperatures.

Case 2: ΔH < 0 and ΔS < 0

If the reaction is sufficiently exothermic it can force ΔG to be negative only at temperatures below which |TΔS| < |ΔH|. This means that there is a temperature defined by [latex]T = \frac{\Delta H}{\Delta S}[/latex] at which the reaction is at equilibrium; the reaction will only proceed spontaneously below this temperature. The freezing of a liquid or the condensation of a gas are the most common examples of this condition.

Case 3: ΔH > 0 and ΔS > 0

This is the reverse of the previous case; the entropy increase must overcome the handicap of an endothermic process so that TΔS > ΔH. Since the effect of the temperature is to “magnify” the influence of a positive ΔS, the process will be spontaneous at temperatures above [latex]T = \frac{\Delta H}{\Delta S}[/latex]. (Think of melting and boiling. )

Case 4: ΔH > 0 and ΔS < 0

With both ΔH and ΔS working against it, this kind of process will not proceed spontaneously at any temperature. Substance A always has a greater number of accessible energy states, and is therefore always the preferred form.

Internal Energy and Enthalpy

The enthalpy of reaction measures the heat released/absorbed by a reaction that occurs at constant pressure.

Learning Objectives

Review enthalpy of reaction

Key Takeaways

Key Points

  • At constant volume, the heat of reaction is equal to the change in the internal energy of the system.
  • At constant pressure, the heat of reaction is equal to the enthalpy change of the system.
  • Most chemical reactions occur at constant pressure, so enthalpy is more often used to measure heats of reaction than internal energy.

Key Terms

  • enthalpy: In thermodynamics, a measure of the heat content of a chemical or physical system.
  • internal energy: A property characteristic of the state of a thermodynamic system, the change in which is equal to the heat absorbed minus the work done by the system.
  • first law of thermodynamics: Heat and work are forms of energy transfer; the internal energy of a closed system changes as heat and work are transferred into or out of it.

In thermodynamics, work (W) is defined as the process of an energy transfer from one system to another. The first law of thermodynamics states that the energy of a closed system is equal to the amount of heat supplied to the system minus the amount of work done by the system on its surroundings. The amount of energy for a closed system is written as follows:

[latex]\Delta U = Q - W[/latex]

In this equation, U is the total energy of the system, Q is heat, and W is work. In chemical systems, the most common type of work is pressure-volume (PV) work, in which the volume of a gas changes. Substituting this in for work in the above equation, we can define the change in internal energy for a chemical system:

[latex]\Delta U=Q-P\Delta V[/latex]

Internal Energy Change at Constant Volume

Let’s examine the internal energy change, [latex]\Delta U[/latex], at constant volume. At constant volume, [latex]\Delta V=0[/latex], the equation for the change in internal energy reduces to the following:

[latex]\Delta U = Q_V[/latex]

The subscript V is added to Q to indicate that this is the heat transfer associated with a chemical process at constant volume. This internal energy is often very difficult to calculate in real life settings, though, because chemists tend to run their reactions in open flasks and beakers that allow gases to escape to the atmosphere. Therefore, volume is not held constant, and calculating [latex]\Delta U[/latex] becomes problematic. To correct for this, we introduce the concept of enthalpy, which is much more commonly used by chemists.

Standard Enthalpy of Reaction

The enthalpy of reaction is defined as the internal energy of the reaction system, plus the product of pressure and volume. It is given by:

[latex]H=U+PV[/latex]

By adding the PV term, it becomes possible to measure a change in energy within a chemical system, even when that system does work on its surroundings. Most often, we are interested in the change in enthalpy of a given reaction, which can be expressed as follows:

[latex]\Delta H = \Delta U +P\Delta V[/latex]

When you run a chemical reaction in a laboratory, the reaction occurs at constant pressure, because the atmospheric pressure around us is relatively constant. We will examine the change in enthalpy for a reaction at constant pressure, in order to see why enthalpy is such a useful concept for chemists.

Enthalpy of Reaction at Constant Pressure

Let’s look once again at the change in enthalpy for a given chemical process. It is given as follows:

[latex]\Delta H=\Delta U + P\Delta V[/latex]

However, we also know that:

[latex]\Delta U=Q-W=Q-P\Delta V[/latex]

Substituting to combine these two equations, we have:

[latex]\Delta H=Q-P\Delta V+P \Delta V=Q_P[/latex]

Thus, at constant pressure, the change in enthalpy is simply equal to the heat released/absorbed by the reaction. Due to this relation, the change in enthalpy is often referred to simply as the “heat of reaction.”

Enthalpy: An explanation of why enthalpy can be viewed as “heat content” in a constant pressure system.