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Now what if that same motionless wrecking ball is lifted two stories above a car with a crane? Potential energy is not only associated with the location of matter (such as a child sitting on a tree branch), but also with the structure of matter.
Visit this site and select “A simple pendulum” on the menu (under “Harmonic Motion”) to see the shifting kinetic (K) and potential energy (U) of a pendulum in motion. After learning that chemical reactions release energy when energy-storing bonds are broken, an important next question is how is the energy associated with chemical reactions quantified and expressed?
If energy is released during a chemical reaction, then the resulting value from the above equation will be a negative number. If a chemical reaction requires an input of energy rather than releasing energy, then the ?G for that reaction will be a positive value. An important concept in the study of metabolism and energy is that of chemical equilibrium.
There is another important concept that must be considered regarding endergonic and exergonic reactions.
Why would an energy-releasing, negative ?G reaction actually require some energy to proceed?
The activation energy of a particular reaction determines the rate at which it will proceed. If no activation energy were required to break down sucrose (table sugar), would you be able to store it in a sugar bowl? The ?G of a reaction can be negative or positive, meaning that the reaction releases energy or consumes energy, respectively. Which of the following comparisons or contrasts between endergonic and exergonic reactions is false? Which of the following is the best way to judge the relative activation energies between two given chemical reactions? Explain in your own words the difference between a spontaneous reaction and one that occurs instantaneously, and what causes this difference. Describe the position of the transition state on a vertical energy scale, from low to high, relative to the position of the reactants and products, for both endergonic and exergonic reactions. The transition state is always higher in energy than the reactants and the products of a reaction (therefore, above), regardless of whether the reaction is endergonic or exergonic.
The purpose of this is to give quick reference to information or to use in an emergency (like if your text has accidentally been left under your desk at school).
This is NOT intended to replace reading the text with its excellent photographs, diagrams, charts, and tables. In studying the properties of matter, chemists wish to find out both the physical and the chemical characteristics of a substance. We have found that the change in free energy for a reaction can be used to answer that question. Josiah Willard Gibbs discovered the relationship between entropy and enthalpy called the free energy. As we have seen, a negative free energy change, ΔG is negative, indicates a reaction will proceed spontaneously. However, at room temperature, the reaction proceeds so slowly that sugar is said to be kinetically stable. The relative amounts of solid and liquid can be changed adding or removing a small amount of heat without changing temperature. A chemist studies the structure of matter and the properties which result from this structure.
The arrow means the reaction is read from left to right, but the equation is only partially correct.
Suppose the product of a reversible reaction decomposes faster than the reactants form the product. Reaction rate is usually defined in terms of the rate of disappearance of one of the reactants.
The nature of the reactants involved in a reaction will determine the kind of reaction that occurs.
Ionic reactions (such as double displacement and neutralization reactions) occur almost instantaneously. We can now see why some substances are kinetically stable even though many of their reactions may have negative free energy changes. For example, hydrogen and oxygen can be kept in the same container at room temperature for ages without reacting to form water. If the number of particles per unit volume is increased, the chance of their colliding is also increased. If we keep the concentration of the hydrogen molecules the same, we would expect doubling the concentration of iodine to double the number of collisions between iodine and hydrogen molecules.
If we double the concentration of hydrogen and also doubles the concentration of iodine, there will be four times as many molecules. In this case, the constant k depends upon the size, speed, and kind of molecule involved in the reaction rate constant of the reaction. It should be pointed out here that the actual mechanism of this reaction involves the breaking of the I-I bond before collision. An increase in the pressure on a gas results in a decrease in the volume occupied by the molecules. Chemical reactions which take place at the interface between two phases are called heterogeneous reactions. The reaction takes place on the surface of the zinc which is the interface between the two phases. Reaction rate is determined by the frequency of collision between molecules and increases as the frequency of collision increases. Increased kinetic energy means that more collisions will occur and the reaction rate will increase. However, the increase in reaction rate depends less on the increase in the number of collisions than it does on another factor.
Note from the graph in Figure 23-5 that molecules must collide with a kinetic energy sufficient to react.
Figure 23-6 is a graph which shows the energy changes involved in the reaction of hydrogen with oxygen.
If the molecules have enough kinetic energy to approach close enough to react chemically, they are said to be activated. A substance which increases a reaction rate without being permanently changed is called a catalyst.
We will discuss two kinds of catalysts: the heterogeneous (or contact) catalyst, and the homogeneous catalyst.
A reaction in which the reactants and catalysts are not in the same state is a heterogeneous reaction.
This attraction weakens the 0-0 bond to the point where the other reactant can break the 0-0 bond. At a given temperature, the rate of a reaction varies directly as the product of the concentrations of the reactants in the slowest step.
The reverse reaction, the collision of two HI molecules to form H2 and 12, does not occur initially because the concentration of HI is zero.
However, as the concentration of HI increases, the reverse reaction, the decomposition of hydrogen iodide, steadily increases. The conditions affecting equilibrium are temperature, pressure, and concentration of either reactants or products. If a system is in equilibrium and a condition is changed, then the equilibrium will shift toward restoring the original conditions.
If stress is put on a reversible reaction at equilibrium, the equilibrium will shift in such a way to relieve the stress.
If the concentration of either of the reactants is increased, the number of collisions between reactant particles will increase. The equilibrium expression for this reaction shows the concentration of ammonia is squared. Pressure would not shift the equilibrium, as the rate in each direction would be affected the same way. Both the forward and reverse reactions at equilibrium are speeded by an increase in temperature. In the example of an airplane in flight, there is a great deal of energy associated with the motion of the airplane. A spring on the ground has potential energy if it is compressed; so does a rubber band that is pulled taut. The change in free energy can be calculated for any system that undergoes such a change, such as a chemical reaction. Standard pH, temperature, and pressure conditions are generally calculated at pH 7.0 in biological systems, 25 degrees Celsius, and 100 kilopascals (1 atm pressure), respectively. Remember that the building of complex molecules, such as sugars, from simpler ones is an anabolic process and requires energy.
Even exergonic reactions require a small amount of energy input to get going before they can proceed with their energy-releasing steps. The source of the activation energy needed to push reactions forward is typically heat energy from the surroundings.
A baby developing from a fertilized egg is an endergonic process; enthalpy decreases (energy is absorbed) and entropy decreases. However, a spontaneous reaction need not occur quickly or suddenly like an instantaneous reaction. Now we can go on to relate this behaviour to the underlying microscopic behaviour of the particles of which the gas is made.Remind your students of the description of pressure arising from bombardment by particles of the walls of the container.
They want to be able to describe the behavior of a given substance with many other substances. Of great importance, especially to the industrial chemist, is the answer to the question: How fast does it react? However, some spontaneous reactions take place so slowly that it takes hundreds of years for any observable change to occur. Therefore, to predict whether a given spontaneous reaction will be useful, we must know the rate at which the reaction occurs as well as at what point equilibrium is established.

If we know the two reaction rates, we can predict whether product or reactant will be in the higher concentration at equilibrium. Reactions with bond rearrangement or electron transfer take longer than reactions without these changes.
They are rapid because ions of one charge are attracted by those of opposite charge and ions collide frequently. Reactions between neutral molecules are slower than ionic reactions because electron transfer and bond rearrangement must occur. Since there are more molecules per unit volume, there has been an increase in concentration. An example of a heterogeneous reaction is zinc (a solid) dissolving in sulfuric acid (a liquid). According to the kinetic theory, the speed (kinetic energy) of molecules increases as the temperature increases. This other factor is the increase in the number of molecules which have reached activation energy. However, the fraction of molecules that have attained the activation energy is greater at the higher temperature T2. This graph can be thought of as a map of the potential energy possessed by the atoms and molecules taking part in a reaction. As two molecules approach, the kinetic energy of motion is transformed into the potential energy of repulsion of electron clouds.
The activated complex may fall back on the left side and break into H2 and 02 molecules, or it may react (fall down on the right side). If the activated complex does react, heat will be released and a molecule of water will be formed. More activated complexes are formed because the number of collisions having the required activation energy increases. If these two gases are brought into contact in the presence of solid vanadium(V) oxide, V205, the reaction is rapid. The V205 lowers the required activation energy by providing a surface on which the activated complex can form. This kind of catalyst does enter into the reaction, but is returned unchanged in a final step of the reaction. If, however, the solution is made acidic, the presence of the acid causes the reaction to proceed readily. Le Chatelier first described the effect of stress (change of conditions) upon systems at equilibrium. If heat is considered to be a product, the addition of heat (a product) will increase the concentration of the product. Reaction rate is the rate of disappearance of one of the reactants of a reaction or the rate of appearance of one of the products.
Four factors influence reaction rate: nature of reactants, concentration, temperature, and catalysis.
Increasing the concentration of a reactant increases the rate of reaction by increasing the number of collisions.
A catalyst is a substance which causes an increase in reaction rate without being permanently changed. A reaction in which the reactants and catalyst are not in the same phase is a heterogeneous reaction.
At a given temperature, the rate of a single-step reaction varies directly as the product of the concentrations of the reactants.
Le Chatelier's principle states that if an equilibrium system is subjected to stress, then the equilibrium will shift to relieve the stress.
Equilibrium constants for gas phase reactions can be expressed in terms of the partial pressures of the reactants and products. Energy conversion – any form of energy can be converted into any other form of energy (ex.
When potential energy decreases, kinetic energy increases and when kinetic energy increases, potential energy decreases (ex. For example, electrical energy, light energy, and heat energy are all different types of energy.
To calculate ?G, subtract the amount of energy lost to entropy (denoted as ?S) from the total energy change of the system. It is important to note that cellular conditions vary considerably from these standard conditions, and so standard calculated ?G values for biological reactions will be different inside the cell. A negative ?G also means that the products of the reaction have less free energy than the reactants, because they gave off some free energy during the reaction.
They can proceed in both directions, releasing energy into their environment in one direction, and absorbing it from the environment in the other direction ([link]). These reactions have a net release of energy, but still require some energy in the beginning. Heat energy (the total bond energy of reactants or products in a chemical reaction) speeds up the motion of molecules, increasing the frequency and force with which they collide; it also moves atoms and bonds within the molecule slightly, helping them reach their transition state. Objects that are not in motion may have the potential to do work, and thus, have potential energy.
Sand art being destroyed is an exergonic process; there is no change in enthalpy, but entropy increases.
It may occur over long periods due to a large energy of activation, which prevents the reaction from occurring quickly. If we now add a piece of metal which has chilled to -20oC, we may be surprised to find that the ice grows larger Evidently the process can go either way. Actually, the study of chemical reactions is the study of the breaking and forming of chemical bonds. A reaction goes to completion when all of one the reactants is used completely and the reaction stops. The hydrogen iodide is decomposing rapidly, more rapidly than H2 and I2 can combine to produce it.
However, some collisions do have enough energy to cause changes in the electron clouds of the colliding molecules. If the activation energy is high, few collisions have enough energy to form the activated complex. Consequently, unless very strong reaction conditions are used, the reactions do not take place. If, however, the mixture is heated to 800oC, or a flame or spark is introduced, a violent reaction occurs.
Actual experiment confirms that the rate of reaction varies directly as the concentration of iodine.
There will be four times as many molecules per unit volume and four times as many collisions. Because of the unusual properties of surface molecules, their bonds are more easily broken.
As more and more kinetic energy is transformed into potential energy, the line of the graph, which represents potential energy, rises. The effect of raising the temperature is to produce more activated complexes through collision. It changes the reaction mechanism in such a way that the activation energy required is less than in the uncatalyzed reaction. In the process of adsorbing a molecule, such as 02, the catalytic surface attracts the 02 molecule. It forms an intermediate compound or compounds which react more readily than the uncatalyzed reactants. As a result, the observation of reaction rates has given scientists an insight into the mechanisms of reactions. Increasing the pressure of a gas or the surface area of a heterogeneous reactant has the same effect on the reaction rate.
There are more frequent collisions and more of the collisions involve sufficient energy to form the activated complex.
While these are all familiar types of energy that one can see or feel, there is another type of energy that is much less tangible.
The suspended wrecking ball has energy associated with it that is fundamentally different from the kinetic energy of objects in motion.
On a chemical level, the bonds that hold the atoms of molecules together have potential energy.
Free energy is called Gibbs free energy (abbreviated with the letter G) after Josiah Willard Gibbs, the scientist who developed the measurement.
Reactions that have a negative ?G and consequently release free energy are called exergonic reactions. On the other hand, the catabolic process of breaking sugar down into simpler molecules releases energy in a series of exergonic reactions. The same is true for the chemical reactions involved in cell metabolism, such as the breaking down and building up of proteins into and from individual amino acids, respectively. This small amount of energy input necessary for all chemical reactions to occur is called the activation energy (or free energy of activation) and is abbreviated EA ([link]). For example, when a glucose molecule is broken down, bonds between the carbon atoms of the molecule are broken.
For this reason, heating up a system will cause chemical reactants within that system to react more frequently. Molecules also have potential energy because the breaking of molecular bonds has the potential to release energy. Exergonic reactions are said to be spontaneous, because their products have less energy than their reactants. A ball rolling downhill is an exergonic process; enthalpy decreases (energy is released), but there is no change in entropy.
Note that this use of the word reversible is different from the use in Chapter 18 concerning work. There is little chance of three or more particles colliding with the proper position and energy to cause a reaction.
The series of reaction steps that must occur for a reaction to go to completion is called the reaction mechanism.

Thus, with the increase in number of activated complexes, the number that will react will also increase. Platinum and other finely divided metals and metallic oxides are common examples of this kind of catalyst. However, the net result to the system as a whole is to shift the equilibrium toward the right.
Doubling the pressure then should increase the rate of the forward reaction by 23 x 2, or 16 times!
One easy way to predict the shift in an equilibrium subjected to a temperature change is to consider heat as a reactant or product. In the Haber process, the reverse action (the decomposition of ammonia) is favored by the addition of heat. This form of energy results from the fact that there is the potential for the wrecking ball to do work. Remember that anabolic cellular pathways require energy to synthesize complex molecules from simpler ones, and catabolic pathways release energy when complex molecules are broken down. Recall that according to the second law of thermodynamics, all energy transfers involve the loss of some amount of energy in an unusable form such as heat, resulting in entropy. Like the example of rust above, the breakdown of sugar involves spontaneous reactions, but these reactions don’t occur instantaneously. Reactants within a closed system will undergo chemical reactions in both directions until a state of equilibrium is reached. Additionally, the burning of many fuels, which is strongly exergonic, will take place at a negligible rate unless their activation energy is overcome by sufficient heat from a spark. Living cells depend on the harvesting of potential energy from molecular bonds to perform work.
The products of endergonic reactions have a higher energy state than the reactants, and so these are nonspontaneous reactions. We conclude, then, that the reaction rate varies directly as the product of the concentrations of hydrogen and iodine. Increasing the surface area increases the number of surface molecules in the same space (increases concentration). The acid lowers the required activation energy by attacking the oxygen atom linking the two sugar molecules. In order to appreciate the way energy flows into and out of biological systems, it is important to understand more about the different types of energy that exist in the physical world. The fact that energy can be released by the breakdown of certain chemical bonds implies that those bonds have potential energy. Gibbs free energy specifically refers to the energy associated with a chemical reaction that is available after entropy is accounted for. These reactions are also referred to as spontaneous reactions, because they can occur without the addition of energy into the system. This state of equilibrium is one of the lowest possible free energy and a state of maximal entropy.
However, to get them into a state that allows the bonds to break, the molecule must be somewhat contorted.
Once reactants have absorbed enough heat energy from their surroundings to reach the transition state, the reaction will proceed.
Once they begin to burn, however, the chemical reactions release enough heat to continue the burning process, supplying the activation energy for surrounding fuel molecules. However, all reactions (including spontaneous -?G reactions) require an initial input of energy in order to reach the transition state, at which they’ll proceed. Because this type of energy refers to the potential to do work, it is called potential energy. In fact, there is potential energy stored within the bonds of all the food molecules we eat, which is eventually harnessed for use.
In other words, Gibbs free energy is usable energy, or energy that is available to do work.
Understanding which chemical reactions are spontaneous and release free energy is extremely useful for biologists, because these reactions can be harnessed to perform work inside the cell. Later sections will provide more information about what else is required to make even spontaneous reactions happen more efficiently. Energy must be put into the system to push the reactants and products away from a state of equilibrium.
Like these reactions outside of cells, the activation energy for most cellular reactions is too high for heat energy to overcome at efficient rates.
The free energy of a system changes during energy transfers such as chemical reactions, and this change is referred to as ?G.
A speeding bullet, a walking person, the rapid movement of molecules in the air (which produces heat), and electromagnetic radiation like light all have kinetic energy. Objects transfer their energy between kinetic and potential in the following way: As the wrecking ball hangs motionless, it has 0 kinetic and 100 percent potential energy. An important distinction must be drawn between the term spontaneous and the idea of a chemical reaction that occurs immediately. This contorted state is called the transition state, and it is a high-energy, unstable state. In other words, in order for important cellular reactions to occur at appreciable rates (number of reactions per unit time), their activation energies must be lowered ([link]); this is referred to as catalysis. Even if the derivation is not required, you will have to explain the terms used, and show that the equation is plausible. Once it is released, its kinetic energy begins to increase because it builds speed due to gravity.
The type of potential energy that exists within chemical bonds, and is released when those bonds are broken, is called chemical energy ([link]). Contrary to the everyday use of the term, a spontaneous reaction is not one that suddenly or quickly occurs. 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 work needed to maintain life. For this reason, reactant molecules don’t last long in their transition state, but very quickly proceed to the next steps of the chemical reaction. Point out that the quantities on the left are macroscopic, while those on the right are microscopic.Stress the underlying assumptions of the model.
The rusting of iron is an example of a spontaneous reaction that occurs slowly, little by little, over time.
In a living cell, chemical reactions are constantly moving towards equilibrium, but never reach it.
Important macromolecules, such as proteins, DNA, and RNA, store considerable energy, and their breakdown is exergonic. The gas:has zero volume at zero temperature so the volume of the actual molecules is negligiblehas zero pressure at zero temperature so thermal energy is its source of kinetic energyhave atoms or molecules which behave as elastic spheres with no long-range intermolecular forcesThere are a number of points to look out for in the derivation. The release of energy is brought about by breaking the molecular bonds within fuel molecules. Whether the reaction is exergonic or endergonic determines whether the products in the diagram will exist at a lower or higher energy state than both the reactants and the products.
If cellular temperatures alone provided enough heat energy for these exergonic reactions to overcome their activation barriers, the essential components of a cell would disintegrate. Students will need convincing about the idea of averaging the square of a velocity; the average velocity is zero, because particles with opposite velocities cancel out.
Just before it hits the ground, the ball has nearly lost its potential energy and has near-maximal kinetic energy. Materials pass in and out, the cell recycles the products of certain chemical reactions into other reactions, and chemical equilibrium is never reached. However, regardless of this measure, the transition state of the reaction exists at a higher energy state than the reactants, and thus, EA is always positive. Other examples of potential energy include the energy of water held behind a dam ([link]), or a person about to skydive out of an airplane.
In this way, living organisms are in a constant energy-requiring, uphill battle against equilibrium and entropy.
The square of the velocity arises because v increases both the momentum change and the frequency of impact. This constant supply of energy ultimately comes from sunlight, which is used to produce nutrients in the process of photosynthesis.
All of this will rely on having covered momentum and impulse thoroughly.If necessary try a demonstration to help convince the students.
This idea follows directly from a statistical analysis of elastic collisions but the maths involved is too advanced for the post-16 level. However the idea may be demonstrated by computer models that have programmed into them only the laws of dynamics for collisions, and nothing about thermodynamics. Calculate the rms speed for oxygen molecules at room temperature; ask the class to repeat the calculation for nitrogen, or for other molecules in air. All stages increase internal energyJoulea€™s paddle wheel experiment where instead of heating increasing the internal energy, mechanical working is converted directly to internal energy.
The temperature is raised and the equivalence of heating and doing mechanical work as two means of transferring energy is demonstrated. This is leading towards the first law of thermodynamicsCooling something down by evaporation. Compare when it is on a surface that then supplies further heat energy to keep the temperature constant (people sweat to lose heat energy) and evaporation when there is no such reheating, so that the liquid cools (puddles evaporate away totally).
A drop of isopropyl alcohol on the back of the hand demonstrates the cooling effect clearly a€“ it feels cold as it evaporatesYou are now in a position to talk about internal energy as the kinetic and potential energy of the molecules. Potential has to be in there because energy must be supplied to melt a solid or boil a liquid; this does not become the kinetic energy of the particles, as the temperature stays constant.
This leads to explanations of everything from evaporation to chemical reaction rates.All of the above are quite subtle, sophisticated ideas that require the students to absorb definitions of closely related concepts.

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