Concept 8.2: The free-energy change of a reaction tells us whether or not the reaction occurs spontaneously

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In 1878, J. Willard Gibbs, a professor at Yale, defined a very useful function called

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the Gibbs free energy of a system (without considering its surroundings), symbolized by the letter G.

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is the portion of a system’s energy that can perform work when temperature and pressure are uniform throughout the system, as in a living cell.

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Free energy

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The change in free energy, ▲G , can be calculated for a chemical reaction by applying the following equation:

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▲G=▲H-T▲S

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▲H symbolizes the change in the system’s

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enthalpy (in biological systems, equivalent to total energy)

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▲S is the change in the system’s entropy; and T is the absolute temperature in

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Kelvin (K) units

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Using chemical methods ▲G, we can measure for any

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reaction.

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More than a century of experiments has shown that only processes with a negative ▲G are

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spontaneous

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For ▲G to be negative,

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▲H must be negative (the system gives up enthalpy and H decreases) or T▲S must be positive (the system gives up order and S increases), or both

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When ▲H and T▲S are tallied, ▲G has a negative value for all

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spontaneous processes

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In other words, every spontaneous process decreases the system’s free energy, and processes that have a positive or zero ▲G are

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never spontaneous.

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This information is immensely interesting to biologists, for it allows us to predict which kinds of change can happen without an

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input of energy

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This principle is very important in the study of metabolism, where a major goal is to determine which reactions can supply energy for

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cellular work.

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Another way to think of ▲G is to realize that it represents the difference between the free energy of the final state and the free energy of

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the initial state: ▲G=G final state -G initial state

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▲G can be negative only when the process involves a loss of free energy during the change from initial state to

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final state

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Because it has less free energy, the system in its final state is less likely to change and is therefore

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more stable than it was previously.

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Another term that describes a state of maximum stability is

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equilibrium

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Recall that most chemical reactions are reversible and proceed to a point at which the forward and backward reactions occur at the

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same rate.

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The reaction is then said to be at chemical equilibrium, and there is no further net change in the .

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relative concentration of products and reactants

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As a reaction proceeds toward equilibrium, the free energy of the mixture of reactants and products

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decreases

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For a system at equilibrium, G is at its lowest

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possible value in that system.

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Any change from the equilibrium position will have a positive ▲G and will not be .

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spontaneous

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systems never spontaneously move away from

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equilibrium.

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A process is spontaneous and can perform work only when it is moving

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toward equilibrium.

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Based on their free-energy changes, chemical reactions can be classified as either

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exergonic (“energy outward”) or endergonic (“energy inward”).

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Proceeds with a net release of free energy. Because the chemical mixture loses free energy (G decreases), ▲G is negative for an exergonic reaction.

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exergonic reaction

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▲G as a standard for spontaneity, exergonic reactions are those that occur

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spontaneously

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The magnitude of ▲G for an exergonic reaction represents the maximum amount of work the

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reaction can perform.

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The greater the decrease in free energy, the greater the amount of work that can be

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done.

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cells contain many different molecules that can engage in a variety of

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chemical reactions

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when molecules react, for example when they collide and exchange parts their atoms and bonds are

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rearranged

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reactants are rearranged to form

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products

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reactant and product molecules store

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potential energy in the arrangement of their atoms and bonds

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chemical reactions involve changes in bonding and changes in

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energy

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we can plot potential energy on a graph and compare the potential energy of

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reactants and products

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using molecules as an example, we can image two chemical reactions. molecules AB and CD react to form molecules

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AC and BD or vice versa

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study graph video in animation Exergonic and Endergonic reactions

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chemical reactions that release energy are called

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exergonic reactions

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you may call exergonic reactions

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downhill reactions

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other reactions, they absorb energy from their surroundings

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endergonic reactions are uphill changes

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exergonic reactions occur

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spontaneously or, that is without a net addition of energy

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the potential energy of the molecule

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decrease

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it is easier for a cell to carry out a reaction that does not need

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additional energy input

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a downhill change is easier than an

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uphill change

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endergonic reactions do not occur

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spontaneously

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it is harder for a cell to carry out a reaction that needs additional

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energy input

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learn C6 H12 O6 + 6 O2 → 6 CO2 + 6 H2 O ▲G= -682kcal/mol (-2,870kj/mol)

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For each mole (180 g) of glucose broken down by respiration under what are called

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“standard conditions”

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learn

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1 M of each reactant and product, 25°C, pH 7), 686 kcal (2,870 kJ) of energy is made available for work

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Because energy must be conserved, the chemical products of respiration store 686 kcal less free energy per mole than the

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reactants

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The phrase “energy stored in bonds” is shorthand for the

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potential energy that can be released when new bonds are formed after the original bonds break, as long as the products are of lower free energy than the reactants.

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is one that absorbs free energy from its surroundings

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endergonic reaction

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Because this kind of reaction essentially stores free energy in molecules (G increases), ▲G is

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positive.

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If a chemical process is exergonic (downhill), releasing energy in one direction, then the reverse process must be

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endergonic (uphill), using energy

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If for respiration, which converts glucose and oxygen to carbon dioxide and water, then the reverse process—the conversion of carbon dioxide and water to glucose and oxygen—must be strongly

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endergonic, with ▲G=+680kcal/mol .

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Plants get the required energy—686 kcal to make a mole of glucose—from the environment by capturing

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light and converting its energy to chemical energy

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Because systems at equilibrium are at a minimum of G and can do no work, a cell that has reached metabolic equilibrium is

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dead!

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Like most systems, a living cell is not in

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equilibrium.

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The constant flow of materials in and out of the cell keeps the metabolic pathways from ever reaching

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equilibrium, and the cell continues to do work throughout its life.

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The key to maintaining this lack of equilibrium is that the product of a reaction does not accumulate but instead becomes a reactant in the next step; finally,

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waste products are expelled from the cell.