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Thermochemistry

Thermochemistry


Thermochemistry is the study of the heat released or absorbed as a result of chemical reactions. It is a branch of thermodynamics and is utilized by a wide range of scientists and engineers. For example, biochemists use thermochemistry to understand bioenergetics, whereas chemical engineers apply thermochemistry to design manufacturing plants. Chemical reactions involve the conversion of a set of substances collectively referred to as "reactants" to a set of substances collectively referred to as "products." In the following balanced chemical reaction the reactants are gaseous methane, CH4(g), and gaseous molecular oxygen, O2(g), and the products are gaseous carbon dioxide, CO2(g), and liquid water H2O(l):

CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l)          (1)

Reactions in which a fuel combines with oxygen to produce water and carbon dioxide are called combustion reactions. Because natural gas consists primarily of methane, it is expected that reaction (1) will liberate heat. Reactions that liberate heat are termed exothermic reactions, and reactions that absorb heat are termed endothermic reactions.

The heat associated with a chemical reaction depends on the pressure and temperature at which the reaction is carried out. All thermochemical data presented here are for reactions carried out under standard conditions, which are a temperature of 298 K (24.85°C) and an applied pressure of one bar . The quantity of heat released in a reaction depends on the amount of material undergoing reaction. The chemical formulas that appear in a reaction each represent 1 mole (see article on "Mole Concept") of material; for example, the symbol CH4 stands for 1 mole of methane having a mass of 16 grams (0.56 ounces), and the 2 O2(g) tells us that 2 moles of oxygen are required. Thermochemistry also depends on the physical state of the reactants and products. For example, the heat liberated in equation (1) is 890 kilojoules (kJ); if, however, water in the gas phase is formed, H2O(g), the heat released is only 802 kJ. Reversing a reaction like (l), which liberates heat, yields a reaction wherein heat must be supplied for the reaction to occur. The following reaction absorbs 890 kJ.

CO2(g) + 2 H2O(l) CH4(g) + 2 O2(g)          (2)

Energy and Enthalpy

Thermochemical changes are often discussed in terms of the "system" and the "surroundings." The system is regarded as the reaction products and reactants, whereas the surroundings consist of everything else in the universe. A boundary separates the system from the surroundings. The first law of thermodynamics relates the energy change belonging to a system to the amount of work and heat crossing the boundary. A statement of the first law applied to chemical reactions in which only heat and work cross the boundary is given by the expression:

U products U reactants = ΔU = q + w          (3)

Here U products represents the energy of the products and U reactants represents the energy of the reactants. The heat associated with the reaction is given as q, and w represents work done during the transformation of reactants to products. If the volume of the system changes during the reaction and the applied pressure remains constant, the work carried out is termed pressure-volume work. For example, reaction (2) converts one mole of gas and two moles of liquid to a total of three moles of gas. The volume of the system increases during the reaction because, under standard conditions, a mole of gas occupies more volume than a mole of liquid. The work of expanding a system against atmospheric pressure is experienced when one inflates a balloon, and this work can be shown to be equal to P ΔV. Here P represents the atmospheric pressure and ΔV represents the change in volume of the system.

The first law of thermodynamics also states that U is a state function. State functions are very important in thermodynamics; they depend only on the present state of a system and not on its past history. Neither q nor w are state functions. An understanding of the concept of state function is furthered by considering the example of one's taking a trip from San Diego, California, to Denver, Colorado. The change in altitude that one experiences during this trip does not depend on the route taken and, thus, is similar to a state function. In comparison, the distance traveled between the two cities does depend on the route one follows; similarly, q and w are path-dependent quantities.

If a process such as a chemical reaction is carried out at a constant pressure in a way that involves only pressure-volume work, then PΔV can be substituted for the work term in equation (3). Thus, we have:

ΔU + P ΔV = qp          (4)

The symbol q p represents the heat accompanying a chemical change carried out at constant pressure; in our previous example this would be equivalent to our specifying the exact route of travel between the two cities. The enthalpy of a system H is related to the energy of a system by the expression:

H = U + PV          (5)

For a process or reaction carried out at constant pressure:

ΔH = ΔU + P ΔV = qp          (6)

Enthalpy, like energy, is a state function. Thus, equation (6) shows that, for a reaction carried out at constant pressure, qp depends only on the reactants consumed and the products formed. The enthalpy change associated with a reaction carried out under standard conditions is termed the heat of reaction and is given the symbol ΔH 0, with the superscript denoting standard conditions. Endothermic reactions have a positive ΔH 0 whereas exothermic reactions have a negative ΔH 0. The change in enthalpy accompanying the conversion of reactants to products in a chemical reaction determines the amount of heat liberated or absorbed by the reaction. For a reaction carried out at constant pressure the enthalpy change depends only on the reactants and products.

Hess's Law

Because enthalpy is a state function, the heat associated with a reaction does not depend on whether the reaction proceeds from reactants to products in a series of steps or in a single step. This is the basis for Hess's law, which states that if two reactions are combined to yield a third reaction, the sum of the ΔH 0s for the first two reactions is equal to the ΔH 0 for the third. For example, consider the conversion of gaseous methane to liquid methanol:

CH4(g) + 1/2 O2(g) CH3OH (l)          (7)

and the subsequent combustion reaction:

CH3OH(l) + 3/2 O2(g) 2 H2O(l) + CO2(g)          (8)

Combining reactions (7) and (8) by adding them together gives reaction (1). Thus, the ΔH 0 for combined reactions (7) and (8) must equal 890kJ. If the ΔH 0 for reaction (8) is known to be 681 kJ, then the ΔH 0 for reaction(7) can be calculated by Hess's law to equal 209 kJ. Born-Haber cycles represent an application of Hess's law to reactions associated with the formation of salts, such as potassium chloride. Born-Haber cycles can be used to determine the enthalpy change accompanying the breakup of the potassium chloride lattice into isolated potassium and chlorine ions.

see also Energy; Heat; Physical Chemistry; Temperature; Thermodynamics.

Michael Eastman

Bibliography

Atkins, Peter, and de Paula, Julio (2002). Physical Chemistry, 7th edition. New York: W. H. Freeman.

Chang, Raymond (2002). Chemistry, 7th edition. Boston: McGraw-Hill.

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thermochemistry

ther·mo·chem·is·try / ˌ[unvoicedth]ərmōˈkemistrē/ • n. the branch of chemistry concerned with the quantities of heat evolved or absorbed during chemical reactions. DERIVATIVES: ther·mo·chem·i·cal / -ˈkemikəl/ adj.

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Thermochemistry

Thermochemistry

History

Thermodynamics and thermochemistry

Change

Energy

Measurement of thermal energy

Enthalpy

Entropy

Gibbs free energy

Resources

Thermochemistry is the part of thermodynamics that studies the relationship between heat and chemical reactions. The word thermodynamics is derived from the Greek words that mean heat and power. Thermodynamics is studied and has applications in all the sciences.

History

french chemist Antoine Laurent Lavoisier (1743 1794) and French mathematician Pierre Simon de Laplace (17491827) are considered to have established the field of thermochemistry around 1780, when the two scientists showed that the heat produced in a particular reaction equals the heat absorbed in the opposite reaction. Sixty years later, Swiss-Russian chemist Henri Hess (18021850) showed (in what is now called Hess Law) that the amount of heat produced in a reaction is the same whether it is produced in a series of steps or as the result of one step.

Thermodynamics and thermochemistry

Thermochemistry is a very important field of study, because it helps to determine if a particular reaction will occur and if it will release or absorb energy as it occurs. It is also possible to calculate how much energy a reaction will release or absorb; this information can be used to determine if it is economically viable to use a particular chemical process. Thermochemistry, however, does not predict how fast a reaction will occur.

In order to understand the terminology of thermochemistry it is first necessary to define the world as viewed by thermodynamics. The chemical reaction being studied is considered the system. For instance, if an acid is being mixed with a base, the acid, the base, any water used to dissolve them, and the beaker in which they are all held are considered the system. Everything else that is not part of the system is considered the surroundings. This definition includes everything from the countertop on which the beaker is held to the planets in outer space. The system and surroundings together form the universe. From this wide set of definitions, it is easy to understand why the system is the only part of any interest. The surroundings are too complex to be considered.

Change

any process that involves a chemical reaction involves change. Sometimes the change occurs on its own. Such a process is called spontaneous. If a change does not occur on its own, it is called non-spontaneous. A spontaneous change may not occur immediately. For example, if a barrel of fuel is left alone, it will remain as fuel indefinitely. However, if a match is used to ignite the fuel, it will burn spontaneously until all the reactants (air, fuel) are completely consumed. In this instance, the spontaneous process required a small amount of energy to be added to the system before a much larger amount of energy could be released. However, once started, it proceeded without assistance. An electrolysis reaction, in which electricity is passed through water to dissociate it into hydrogen and oxygen, is not considered spontaneous because the reaction stops if the electricity is removed. An electrolysis reaction is a non-spontaneous process. How is it possible to determine if a process is spontaneous or non-spontaneous without actually mixing the chemicals together? Two factors in combination are used to determine whether a process occurs spontaneously or not. These factors are energy and disorder.

Energy

energy is a state function. There are a number of different forms of energy, which is the ability to do work. Work is done anytime a force is applied to make an object move. There is energy of motion, called kinetic energy, and energy of position or stored energy, called potential energy. Potential and kinetic energy are interconvertible; that is, one form can change to the other. Different types of energy include thermal energy, electrical energy, radiant energy, chemical energy, mechanical energy, and nuclear energy. One type of energy can be converted to another. However, energy can neither be created nor destroyed. It is always conserved. For example, passing electrical energy through a tungsten filament converts it to light energy. All the electrical energy is not converted to light however. Some of it is converted to thermal energy, which is why a light bulb becomes hot after some time.

In most chemical reactions, chemical energy is converted to some other, more useful form of energy. For example, in a flashlight, chemical energy from the batteries is converted to electrical energy. In a car, chemical energy from the combustion of the fuel is converted into mechanical energy. Thermochemistry concerns itself with the relation between chemical reactions and thermal energy. Thermal energy is the energy of motion of particles such as atoms, molecules, or ions. Thermal energy depends on the quantity of a substance present and is thus known as an extensive property. The thermal energy provided by a drop of water is much less than that provided by a pot full of water. Temperature, however, is a property that is not dependent on the quantity of substance. The temperature of a drop of boiling water is the same as that of a pot of boiling water. Heat is the transfer of thermal energy that occurs between two objects when they are at different temperatures. If the two objects are at the same temperature, no thermal energy is transferred and no heat is felt. That is how people can tell if an object is hot by touching it. When heat is released from the system in a chemical reaction, the reaction is said to be exothermic. When heat is absorbed by the system, the reaction is said to be endothermic. In an endothermic reaction, the surroundings provide the heat for the reaction; in an exothermic reaction, the surroundings are heated by the reaction. For this reason, it is accepted that exothermic quantities are negative quantities, since the system is losing energy, and endothermic quantities are positive quantities, since the system is gaining energy.

Measurement of thermal energy

How can thermal energy be measured? One way is to measure a quantity called the specific heat. The specific heat of a substance is the amount of thermal energy required to heat one gram of that substance by one degree. Once again, the specific heat of a substance is an intensive property, meaning that it does not depend on the amount of substance present. The specific heat of a drop of water and a pan of water are the same. When one multiplies the mass of an object by its specific heat, it is possible to calculate the heat capacity of that object. Heat capacity is an extensive property, meaning that it is dependent on the quantity of substance present. The heat capacity of a drop of water is much, much less than that of a lake. The specific heat of water is unusually high compared to many other substances. This fact has an important impact on humans. Cities located near huge bodies of water tend to have more moderate climates. Such cities are cooler in the summer, as large water bodies take a long time to absorb the heat of the summer sun. These cities are warmer in the winter, as the water slowly releases the heat it had absorbed during the summer. Since human bodies are composed largely of water, people are able to maintain a fairly constant body temperature, in spite of outside temperature fluctuations. Even so, one important fact arises from all the information so far. Scientists cannot measure an absolute value for energy. They can, however, measure energy differences.

Enthalpy

the mathematical representation for thermal energy contains many terms. Scientists can, however, simplify it based on how experiments are performed. Most chemical reactions take place under atmospheric pressure, which is (for the most part) constant. When a thermal energy change is measured under constant pressure conditions, it is called a change in enthalpy. The symbol for enthalpy is H. Since only a difference in enthalpy can be measured, the difference is called delta H. When describing different kinds of changes, scientists can indicate the difference as part of the name by using subscripts after the H. However, reactions can be done under different pressure and temperature conditions. For the sake of uniformity, a standard state is defined as the state of a substance at one atmosphere pressure. For solids and liquids, the standard state is the pure solid or liquid at one atmosphere pressure. For gases, the standard state is the ideal gas at a partial pressure of one atmosphere. Once the standard state is defined, some simplifications can be made. The enthalpy of formation of an element in its standard state is zero. If there is more than one form of the element under the defined conditions, the most stable form is given an enthalpy of formation of zero. Carbon, for example, has two forms at one atmosphere, graphite and diamond. Graphite is more stable and is assigned an enthalpy of formation of zero. Diamond does not have an enthalpy of formation of zero.

Enthalpy has a special property. Its value is determined based on the initial state of the system and the final state of the system. It does not depend on how the system gets from the initial state to the final state. A function that has this property is called a state function. The fact that enthalpy is a state function makes it possible to calculate enthalpies for some compounds without having to measure them experimentally. By combining different reactions with known enthalpies, it is possible to calculate the unknown enthalpy. Hess Law summarizes this observation by stating that the thermal energy absorbed or released in a change is the same whether the change occurs in a single step or in multiple steps.

Those enthalpies that cannot be calculated using Hess law can be measured experimentally. An apparatus called a calorimeter is used to measure the quantity of thermal energy gained or lost in a chemical change. A simple calorimeter can be constructed using two nested styrofoam cups with lids and a thermometer. A more complex type of calorimeter is the bomb calorimeter, which measures thermal energy changes under constant volume conditions.

Entropy

as mentioned much earlier, two quantities determine whether a reaction will be spontaneous or not, the thermal energy and disorder. Disorder is also known as entropy. Entropy is given the symbol S. In general, entropy always has a tendency to increase. In other words, the universe has a tendency to move towards disorder. When disorder increases, scientists say entropy increases. An increase in entropy is assigned a positive sign. When order increases, scientists say entropy decreases. A decrease in entropy is assigned a negative sign. Entropy only has a zero value if one considers a perfect crystal at absolute zero. Since it is not possible to reach absolute zero, no substance has a value of zero entropy.

KEY TERMS

Enthalpy The measurement of thermal energy under constant pressure conditions.

Entropy The measurement of a tendency towards increased randomness and disorder.

Equilibrium The conditions under which a system shows no tendency for a change in its state. At equilibrium the net rate of reaction becomes zero.

Gibbs free energy Mathematically equal to the change in the enthalpy minus the product of the temperature and the change in the entropy. Used to determine if a process is spontaneous or not.

Heat The transfer of thermal energy that occurs between two objects when they are at different temperatures.

Surroundings Everything that is not part of the system.

System The materials pertinent to the reaction being studied.

Thermal energy The total amount of energy contained within any body as a consequence of the motion of its particles.

Gibbs free energy

Certain processes that release a great deal of energy are not spontaneous, even though it would seem that they should be. Similarly, certain processes that greatly increase disorder are not spontaneous, although it would seem that they should be. If scientists mathematically manipulate the expressions for enthalpy and entropy, it is possible to define a new quantity called the Gibbs free energy. The Gibbs free energy, sometimes simply called free energy, equals the change in the enthalpy minus the product of the temperature and the change in the entropy. The term free energy should not be misunderstood. As stated earlier, energy can neither be created nor destroyed. This energy does not come free of cost. The term free in free energy is better interpreted as available. The free energy can be used to predict if a process is spontaneous or not. If the free energy is negative, the process is spontaneous. If the free energy is positive, the process is not spontaneous. A non-spontaneous process can sometimes be made spontaneous by varying the temperature. If the free energy is zero, the process is at equilibrium, meaning that the forward rate of the reaction equals the reverse rate of the reaction.

Resources

BOOKS

Brock, William Hodson. Norton History of Chemistry. New York: Norton, 2000.

Burshtein, A. I. Introduction to Thermodynamics and Kinetic Theory of Matter. Weinheim, Germany: Wiley-VCH, 2005.

Roy, Bimalendu Naravan. Fundamentals of Classical and Statistical Thermodynamics. West Sussex, UK, and New York: John Wiley, 2002.

Tro, Nivaldo J. Introductory Chemistry. Upper Saddle River, NJ: Pearson Education, 2006.

Turns, Stephen R. Thermodynamics: Concepts and Applications. New York: Cambridge University Press, 2006.

Rashmi Venkateswaran

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Thermochemistry

Thermochemistry

Thermodynamics and thermochemistry

The word thermodynamics is derived from the Greek words that mean "heat" and "power." Thermodynamics is studied and has applications in all the sciences. Thermochemistry is the part of thermodynamics that studies the relationship between heat and chemical reactions . Thermochemistry is a very important field of study because it helps to determine if a particular reaction will occur and if it will release or absorb energy as it occurs. It is also possible to calculate how much energy a reaction will release or absorb and this information can be used to determine if it is economically viable to use a particular chemical process. Thermochemistry, however, does not predict how fast a reaction will occur.

In order to understand the terminology of thermochemistry it is first necessary to define the world as viewed by thermodynamics. The chemical reaction being studied is considered to be the "system." For instance, if an acid is being mixed with a base, the acid, the base, any water used to dissolve them and the beaker in which they are all held are considered the system. Everything else that is not part of the system is considered to be the "surroundings." This includes everything from the countertop on which the beaker is held to the planets in outer space . The system and surroundings together form the "universe." From this wide set of definitions, it is easy to understand why the system is the only part of any interest to us. The surroundings are too complex to be considered.


Change

Any process that involves a chemical reaction involves change. Sometimes the change occurs on its own. Such a process is called spontaneous. If a change does not occur on its own, it is called non-spontaneous. A spontaneous change may not occur immediately. For example, if a barrel of fuel is left alone, it will remain as fuel indefinitely. However, if a match is used to ignite the fuel, it will burn spontaneously until all the reactants (air, fuel) are completely consumed. In this instance, the spontaneous process required a small amount of energy to be added to the system before a much larger amount of energy could be released. However, once started, it proceeded without assistance. An electrolysis reaction, in which electricity is passed through water to dissociate it into hydrogen and oxygen , is not considered spontaneous because the reaction stops if the electricity is removed. An electrolysis reaction is a non-spontaneous process. How is it possible to determine if a process is spontaneous or non-spontaneous without actually mixing the chemicals together? There are two factors whose combination determines whether a process occurs spontaneously or not. These factors are energy and disorder.


Energy

Energy is a state function. There are a number of different forms of energy, which is the ability to do work . Work is done anytime a force is applied to make an object move. There is energy of motion , called kinetic energy and energy of position or stored energy, called potential energy. Potential and kinetic energy are interconvertible; that is, one form can change to the other. Different types of energy include thermal energy, electrical energy, radiant energy, chemical energy, mechanical energy and nuclear energy. One type of energy can be converted to another. However, energy can neither be created nor destroyed. It is always conserved. For example, passing electrical energy through a tungsten filament converts it to light energy. All the electrical energy is not converted to light however. Some of it is converted to thermal energy, which is why a light bulb becomes hot after some time.

In most chemical reactions, chemical energy is converted to some other, more useful form of energy. For example, in a flashlight, chemical energy from the batteries is converted to electrical energy. In a car, chemical energy from the combustion of the fuel is converted into mechanical energy. Thermochemistry concerns itself with the relation between chemical reactions and thermal energy. Thermal energy is the energy of motion of particles such as atoms , molecules or ions. Thermal energy depends on the quantity of a substance present and is thus known as an extensive property. The thermal energy provided by a drop of water is much less than that provided by a pot full of water. Temperature , however, is a property that is not dependent on the quantity of substance. The temperature of a drop of boiling water is the same as that of a pot of boiling water. Heat is the transfer of thermal energy that occurs between two objects when they are at different temperatures. If the two objects are at the same temperature, no thermal energy is transferred and no heat is felt. That is how we can tell if an object is hot by touching it. When heat is released from the system in a chemical reaction, the reaction is said to be exothermic. When heat is absorbed by the system, the reaction is said to be endothermic . In an endothermic reaction, the surroundings provide the heat for the reaction while in an exothermic reaction, the surroundings are heated by the reaction. For this reason it is accepted that exothermic quantities are negative quantities, since the system is losing energy and endothermic quantities are positive quantities since the system is gaining energy.


Measurement of thermal energy

How can thermal energy be measured? One way is to measure a quantity called the specific heat. The specific heat of a substance is the amount of thermal energy required to heat one gram of that substance by one degree. Once again, the specific heat of a substance is an intensive property, meaning that it does not depend on the amount of substance present. The specific heat of a drop of water and a pan of water are the same. When we multiply the mass of an object by its specific heat, it is possible to calculate the heat capacity of that object. Heat capacity is an extensive property, meaning that it is dependent on the quantity of substance present. The heat capacity of a drop of water is much, much less than that of a lake . The specific heat of water is unusually high compared to many other substances. This fact has an important impact on us. Cities located near huge bodies of water tend to have more moderate climates. Such cities are cooler in the summer as large water bodies take a long time to absorb the heat of the summer sun , and these cities are warmer in the winter, as the water slowly releases the heat it had absorbed during the summer. Since our bodies are composed largely of water, we are able to maintain a fairly constant body temperature, in spite of outside temperature fluctuations. Even so, one important fact arises from all the information so far. We cannot measure an absolute value for energy. We can, however, measure energy differences.


Enthalpy

The mathematical representation for thermal energy contains many terms. We can, however, simplify it based on how we perform the experiments. Most chemical reactions take place under atmospheric pressure , which is (for the most part) constant. When a thermal energy change is measured under constant pressure conditions, it is called a change in "enthalpy." The symbol for enthalpy is H. Since only a difference in enthalpy can be measured, the difference is called "delta H." When describing different kinds of changes, we can indicate the difference as part of the name by using subscripts after the H. However, reactions can be done under different pressure and temperature conditions. For the sake of uniformity, a standard state is defined as the state of a substance at one atmosphere pressure. For solids and liquids, the standard state is the pure solid or liquid at one atmosphere pressure. For gases, the standard state is the ideal gas at a partial pressure of one atmosphere. Once the standard state is defined, some simplifications can be made. The enthalpy of formation of an element in its standard state is zero. If there is more than one form of the element under the defined conditions, the most stable form is given an enthalpy of formation of zero. Carbon , for example, has two forms at one atmosphere, graphite and diamond . Graphite is more stable and is assigned an enthalpy of formation of zero. Diamond does not have an enthalpy of formation of zero.

Enthalpy has a special property. Its value is determined based on the initial state of the system and the final state of the system. It does not depend on how the system gets from the initial state to the final state. A function that has this property is called a "state function." The fact that enthalpy is a state function makes it possible to calculate enthalpies for some compounds without having to measure them experimentally. By combining different reactions with known enthalpies, it is possible to calculate the unknown enthalpy. Hess' Law summarizes this observation by stating that the thermal energy absorbed or released in a change is the same whether the change occurs in a single step or in multiple steps.

Those enthalpies that cannot be calculated using Hess' law can be measured experimentally. An apparatus called a calorimeter is used to measure the quantity of thermal energy gained or lost in a chemical change. A simple calorimeter can be constructed using two nested styrofoam cups with lids and a thermometer . A more complex type of calorimeter is the bomb calorimeter, which measures thermal energy changes under constant volume conditions.


Entropy

As mentioned much earlier, two quantities determine whether a reaction will be spontaneous or not, the thermal energy and disorder. Disorder is also known as entropy . Entropy is given the symbol S. In general, entropy always has a tendency to increase. In other words, the universe has a tendency to move towards disorder. When disorder increases, we say entropy increases. An increase in entropy is assigned a positive sign. When order increases, we say entropy decreases. A decrease in entropy is assigned a negative sign. Entropy only has a zero value if we consider a perfect crystal at absolute zero . Since it is not possible to reach absolute zero, no substance has a value of zero entropy.


Gibbs' free energy

Certain processes that release a great deal of energy are not spontaneous, even though it would seem that they should be. Similarly, certain processes that greatly increase disorder are not spontaneous, although it would seem that they should be. If we mathematically manipulate the expressions for enthalpy and entropy, it is possible to define a new quantity called the Gibbs' free energy. The Gibbs' free energy, sometimes simply called free energy, equals the change in the enthalpy minus the product of the temperature and the change in the entropy. The term free energy should not be misunderstood. As stated earlier, energy can neither be created nor destroyed. This energy does not come free of cost. The term 'free' in "free energy" is better interpreted as available. The free energy can be used to predict if a process is spontaneous or not. If the free energy is negative, the process is spontaneous. If the free energy is positive, the process is not spontaneous. A non-spontaneous process can sometimes be made spontaneous by varying the temperature. If the free energy is zero, the process is at equilibrium, meaning that the forward rate of the reaction equals the reverse rate of the reaction.

Resources

books

Oxtoby, Gillis, and Nachtrieb. Principles of Modern Chemistry. 4th ed. Saunders College Publishing, 1999.

Silberberg. Chemistry, the Molecular Nature of Matter andChange. 2nd Ed., McGraw-Hill, 2000.


Rashmi Venkateswaran

KEY TERMS

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Enthalpy

—The measurement of thermal energy under constant pressure conditions.

Entropy

—The measurement of a tendency towards increased randomness and disorder.

Equilibrium

—The conditions under which a system shows no tendency for a change in its state. At equilibrium the net rate of reaction becomes zero.

Gibbs' free energy

—Mathematically equal to the change in the enthalpy minus the product of the temperature and the change in the entropy. Used to determine if a process is spontaneous or not.

Heat

—The transfer of thermal energy that occurs between two objects when they are at different temperatures.

Surroundings

—Everything that is not part of the system.

System

—The materials pertinent to the reaction being studied.

Thermal energy

—The total amount of energy contained within any body as a consequence of the motion of its particles.

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