Energy

The study of energy and its transformations is known as thermodynamics (thérme- ‘’heat’’, dy’namis ‘’power’’). In this chapter we will examine an aspect of thermodynamics that involves the relationships between chemical reactions and energy changes involving heat. This portion of thermodynamics is called thermochemistry.
The Nature of Energy
Energy is defined as the capacity to do work or totransfer heat. Work is the energy used to cause an object with mass to move, and heat is the energy used to cause the temperature of an object to increase.
* Kinetic energy: the energy of motion. The magnitude of the kinetic energy (Ek) of an object depends on its mass (m) and speed (v). Ek= 12 mv2
* Potential energy: the energy that an object possesses by virtue of its positionrelative to other objects. The potential energy, Ep, is given by the equation Ep=mgh where m is the mass of the object in question, h is the height of the object relative to some reference height, and g is the gravitational constants, 9.8 m/s2
The SI unit for energy is the joule (J).
Ek= 12mv2= 12 (2kg)(1ms)2=1 kg m2/s2 =1 J
A calorie (cal) was originally defined as the amount of energy required toraise the temperature of 1 g of water from 14.5°C to 15.5°C. 1 cal = 4.184 J
The internal energy of a system is the sum of all the kinetic and potential energies of all its components. The first law of thermodynamics states that the change in the internal energy of a system, ∆E, is the sum of the heat, q, transferred into or out of the system and the work, w, done on or by the system:∆E=q+w.
When a process occurs in which the system absorbs heat, the process is called endothermic. During an endothermic process, such as the melting of ice, heat flows into the system from its surroundings. A process in which the system loses heat is called exothermic. During exothermic process, such as the combustion of gasoline, heat exists or flows out of the system and into the surroundings.Enthalpy is a state function because internal energy, pressure, and volume are all state functions. When a change occurs at constant pressure, the change in enthalpy, ∆ H , is given by: ∆H= ∆E+P∆V. When ∆ H is positive, the system has gained heat from the surroundings, which is an endothermic process. When ∆ H is negative, the system has released heat to the surroundings, which is an exothermicprocess.
In a chemical process, the enthalpy of reaction is the enthalpy of the products minus the enthalpy of the reactants: ∆Hrxn=H products-H (reactants). Enthalpies of reaction follow some simple rules:
* The enthalpy of reaction is proportional to the amount of reactant that reacts.
* Reversing a reaction changes the sign of ∆H.
* The enthalpy of reaction depends on the physicalstates of the reactants and products.
The measurement of heat flow is calorimetry; a device used to measure heat flow is a calorimeter. The temperature change of a calorimeter depends on its heat capacity, the amount of heat required to raise its temperature by 1K. The heat capacity for one mole of a pure substance is called its molar heat capacity; for one gram of the substance, we use the termspecific heat. Water has a very high specific heat, 4.18J/gK.
Calorimetry can be used to study the chemical potential energy stored in substances. One of the most important types of reactions studied using calorimetry is combustion, in which a compound (usually an organic compound) reacts completely with excess oxygen.
Many enthalpies of reaction have been measured and tabulated. Hess’ law statesthat if a reaction is carried out in a series of steps, ∆H for the overall reaction will equal the sum of the enthalpy changes for the individual steps.
The energy released when one gram of a material is combusted is often called its fuel value. Because fuel values represent the heat released in combustion, fuel values are reported as positive numbers. The fuel value of any food or fuel can...