• The study of energy and its transformations is known as thermodynamics (Greek: thérme-, “heat”; dy’namis, “power”). This area of study began during the Industrial Revolution in order to develop the relationships among heat, work, and fuels in steam engines.

The Nature of Chemical Energy

• Chemical reactions and energy. Energy changes in chemical reactions can be used to transfer heat or do work.

• Electrostatic potential energy. At finite separation distances the electrostatic potential energy, Eel, is positive for objects with like charges, and negative for objects that are oppositely charged. As the particles move farther apart, their electrostatic potential energy approaches zero.

• Electrostatic potential energy and ionic bonding. As the separation between ions increases, the electrostatic potential energy increases (becomes less negative). As the distance separating the ions goes toward infinity, the electrostatic potential energy goes to zero. In real compounds, repulsions between core electrons place a lower limit on how closely the ions can approach each other.

Energy is released when chemical bonds are formed;

energy is consumed when chemical bonds are broken.

The First Law of Thermodynamics

System and surroundings

• The portion we single out for study is called the system; everything else is called the surroundings.

• Systems may be open, closed, or isolated. An open system is one in which matter and energy can be exchanged with the surroundings. The systems we can most readily study in thermochemistry are called closed systems—systems that can exchange energy but not matter with their surroundings.An isolated system is one in which neither energy nor matter can be exchanged with the surroundings.

Internal Energy

• The internal energy, E, of a system is the sum of all the kinetic and potential energies of the components of the system.

Thermodynamic quantities such as ∆E have three parts:

1. a number

2. a unit

3. a sign

• The symbol ∆ is commonly used to denote change.

• We need to remember, however, that any increase in the energy of the system is accompanied by a decrease in the energy of the surroundings, and vice versa.

• In a chemical reaction, the initial state of the system refers to the reactants and the final state refers to the products.

Relating ∆E to Heat and Work

• The internal energy of a system changes in magnitude as heat is added to or removed from the system or as work is done on or by the system.

• When heat is added to a system or work is done on a system, its internal energy increases. Therefore, when heat is transferred to the system from the surroundings, q has a positive value.

• Sign conventions for heat and work. Heat, q, transferred into a system and work, w, done on a system are both positive quantities, corresponding to “deposits” of internal energy into the system. Conversely, heat transferred out of a system to the surroundings and work done by the system on the surroundings are both “withdrawals” of internal energy from the system.

Endothermic and Exothermic Processes

• When a process occurs in which the system absorbs heat, the process is called endothermic (endo- means “into”).During an endothermic process, such as the melting of ice, heat flows intothe system from its surroundings.
  

• A process in which the system loses heat is called exothermic (exo- means “out of”). During an exothermic process, such as the combustion of gasoline, heat exits or flows out of the system into the surroundings

State Functions

• Internal energy is an example of a state function, a property of a system that is determined by specifying the system’s condition, or state (in terms of temperature, pressure, and so forth). The value of a state function depends only on the present state of the system, not on the path the system took to reach that state.

• Altitude is analogous to a state function because the change in altitude is independent of the path taken. Distance traveled is not a state function.Some thermodynamic quantities, such as E, are state functions. Other quantities, such as q and w, are not.

Enthalpy

• Under conditions of constant pressure, a thermodynamic quantity called enthalpy (from the Greek enthalpein, “to warm”) provides such a function.

• *Enthalpy,*which we denote by the symbol H, is defined as the internal energy plus the product of the pressure, P, and volume, V, of the system: H = E + PV
  

Pressure–Volume Work

The work involved in the expansion or compression of gases is called pressure– volume work (P–V work).

w = -P ∆V

• where P is pressure and ∆V = Vfinal - Vinitial is the change in volume of the system. The pressure P is always either a positive number or zero. If the volume of the system expands, then ∆V is positive as well.

Enthalpy Change

• ∆H = ∆(E + PV) = ∆E + P∆V (constant pressure)

• The change in internal energy is equal to the heat gained or lost at constant volume, and the change in enthalpy is equal to the heat gained or lost at constant pressure.

Enthalpies of Reaction

• The enthalpy change that accompanies a reaction is called either the enthalpy of reaction or the heat of reaction and is sometimes written ∆Hrxn, where “rxn” is a commonly used abbreviation for “reaction.”

• The negative sign for ∆H tells us that this reaction is exothermic. Balanced chemical equations that show the associated enthalpy change in this way are called thermochemical equations.

GUIDELINES IN USING THERMOCHEMICAL EQUATIONS AND ENTHALPY DIAGRAMS:

1. Enthalpy is an extensive property.

2. The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to ∆H for the reverse reaction.

3. The enthalpy change for a reaction depends on the states of the reactants and products.

Calorimetry

• The measurement of heat flow is calorimetry; a device used to measure heat flow is a calorimeter.

Heat Capacity and specific Heat

• The temperature change experienced by an object when it absorbs a certain amount of heat is determined by its heat capacity, denoted C.

• The heat capacity of one mole of a substance is called its molar heat capacity, Cm. The heat capacity of one gram of a substance is called its specific heat capacity, or merely its specific heat, Cs.

• Specific heat = (quantity of heat transferred)/(grams of substance) * (temperature change)
  

Constant-Pressure Calorimetry

• Calorimetry methods and equipment vary by procedure. Pressure can be controlled to directly measure ∆H in various processes, including those in solution. In ordinary chemistry labs, a "coffee-cup" calorimeter is used to demonstrate calorimetry principles. The reaction takes place under atmospheric pressure because the calorimeter is open.

• Coffee-cup calorimeter. This simple apparatus is used to measure temperature changes of reactions at constant pressure.

Bomb Calorimetry (Constant-Volume Calorimetry)

• An important type of reaction studied using calorimetry is combustion, in which a compound reacts completely with excess oxygen. Combustion reactions are most accurately studied using a bomb calorimeter. The substance to be studied is placed in a small cup within an insulated sealed vessel called a bomb. The bomb, which is designed to withstand high pressures, has an inlet valve for adding oxygen and electrical leads for initiating the reaction.

• The heat released when combustion occurs is absorbed by the water and the various components of the calorimeter (which all together make up the surroundings), causing the water temperature to rise. The change in water temperature caused by the reaction is measured very precisely.

• To calculate the heat of combustion from the measured temperature increase, we must know the total heat capacity of the calorimeter. Because reactions in a bomb calorimeter are carried out at constant volume, the heat transferred corresponds to the change in internal energy, ∆E, rather than the change in enthalpy, ∆H.

Hess’s Law

• Hess’s law states that if a reaction is carried out in a series of steps, ∆H for the overall reaction equals the sum of the enthalpy changes for the individual steps.

• This law is a consequence of the fact that enthalpy is a state function. Hess’s law provides a useful means of calculating energy changes that are difficult to measure directly.

• Because H is a state function, for a particular set of reactants and products, ∆H is the same whether the reaction takes place in one step or in a series of steps.

Enthalpies of Formation

• The enthalpy change associated with this process is called the enthalpy of formation (or heat of formation), ∆Hf, where the subscript f indicates that the substance has been formed from its constituent elements.

• The standard enthalpy change of a reaction is defined as the enthalpy change when all reactants and products are in their standard states.

• The standard enthalpy of formation of a compound, ∆Hf°, is the change in enthalpy for the reaction that forms one mole of the compound from its elements with all substances in their standard states.

• The standard enthalpy of formation of the most stable form of any element is zero because there is no formation reaction needed when the element is already in its standard state.

• The symbol Σ (sigma) means “the sum of,” and n and m are the stoichiometric coefficients of the relevant chemical equation.

Bond Enthalpies

• The bond enthalpy is the enthalpy change, ∆H, for the breaking of a particular bond in one mole of a gaseous substance. It is easiest to determine bond enthalpies from simple reactions where only one bond is broken, such as the dissociation of Cl2(g).

• It is straightforward to assign bond enthalpies for a reaction involving the breaking of a bond in a diatomic molecule: the bond enthalpy is simply equal to the enthalpy of the reaction.

• The bond enthalpy is always a positive quantity because energy is required to break chemical bonds. Conversely, energy is always released when a bond forms between two gaseous atoms or molecular fragments. The greater the bond enthalpy, the stronger the bond.

Bond Enthalpies and the Enthalpies of Reactions

• Because enthalpy is a state function, we can use average bond enthalpies to estimate the enthalpies of reactions in which bonds are broken and new bonds are formed.

Our strategy for estimating reaction enthalpies is a straightforward application of Hess’s law.

1. We supply enough energy to break those bonds in the reactants that are not present in the products. The enthalpy of the system is increased by the sum of the bond enthalpies of the bonds that are broken.

2. We form the bonds in the products that were not present in the reactants. This step releases energy and therefore lowers the enthalpy of the system by the sum of the bond enthalpies of the bonds that are formed.

• It is important to remember that bond enthalpies are derived for gaseous molecules. In solids, liquids, and solutions, intermolecular forces between different molecules must also be taken into account.

Foods and Fuels

• The energy released when one gram of any substance is combusted is the fuel value of the substance. The fuel value of any food or fuel can be measured by calorimetry.

Foods

• Most of the energy our bodies need comes from carbohydrates and fats.

• The carbohydrates known as starches are decomposed in the intestines into glucose.

• Glucose is soluble in blood, and in the human body it is known as blood sugar. Because carbohydrates break down rapidly, their energy is quickly supplied to the body.

• However, the body stores only a very small amount of carbohydrates.

• The body uses the chemical energy from foods to maintain body temperature, to contract muscles, and to construct and repair tissues.

• Fats are well suited to serve as the body’s energy reserve for at least two reasons: (1) They are insoluble in water, which facilitates storage in the body, and (2) they produce more energy per gram than either proteins or carbohydrates, which makes them efficient energy sources on a mass basis.

• The combustion of carbohydrates and fats in a bomb calorimeter gives the same products as when they are metabolized in the body.

• The metabolism of proteins produces less energy than combustion in a calorimeter because the products are different.

• Proteins are used by the body mainly as building materials for organ walls, skin, hair, muscle, and so forth.

• The amount of energy our bodies require varies considerably, depending on such factors as weight, age, and muscular activity.

Fuels

• Coal, petroleum, and natural gas, which are the world’s major sources of energy, are known as fossil fuels. All have formed over millions of years from the decomposition of plants and animals and are being depleted far more rapidly than they are being formed.

• Natural gas consists of gaseous hydrocarbons, compounds of hydrogen and carbon.

• Petroleum is a liquid composed of hundreds of compounds, most of which are hydrocarbons, with the remainder being chiefly organic compounds containing sulfur, nitrogen, or oxygen.

• Coal, which is solid, contains hydrocarbons of high molecular weight as well as compounds containing sulfur, oxygen, or nitrogen.

Other Energy sources

• Nuclear energy is the energy released in either the fission (splitting) or the fusion (combining) of atomic nuclei.

• Fossil fuels and nuclear energy are nonrenewable sources of energy—they are limited resources that we are consuming at a much greater rate than they can be regenerated.

• Renewable energy sources include solar energy from the Sun, wind energy harnessed by windmills, geothermal energy from the heat stored inside Earth, hydroelectric energy from flowing rivers, and biomass energy from crops and biological waste matter.

• Electricity is typically generated by using a moving gas or liquid to turn a turbine that is connected to a generator, although there are some exceptions like solar cells.

• Fulfilling our future energy needs will depend on developing technology to harness solar energy with greater efficiency. Solar energy is the world’s largest energy source.