Chapter 6

Energy is the ability to do work, which is the ability to move matter. It takes many forms, which can be converted into each other, frequently quite easily. Radiant energy – energy from the sun Thermal energy – Energy associated with the movement of atoms and molecules (basically heat energy)

Chemical energy – Energy stored by the bonds of chemical substances Potential energy – Energy available by virtue of an object’s position Electrical energy and light energy are also common forms.

A basic principle in science is that energy cannot be created nor destroyed, only converted from one to the other. This is called the Law of Conservation of Energy. (Actually, as a result of Albert Einstein, we now believe that the Laws of Conservation of Energy and of Mass are not 2 separate laws, but rather parts of one Law, the Law of Conservation of Energy + Mass.)

In order to study energy changes during chemical reactions (called Thermochemistry), we have to understand that heat is always transferred from a hot object to a cooler object. Heat is a quantity; basically, for our purposes right now, thermal energy. The terms hot and cold refer to the temperature which measures intensity of light. Think of a very intense flashlight bulb that provides a lot of light to a small area, compared to a much less intense large light bulb, that provides less light but over a much larger area.

We also have to be careful in defining what we are actually studying, called the system. Everything else in the universe is the surroundings. If the system releases heat to the surroundings, that is an exothermic process. Conversely, if the system picks up (absorbs) heat from the surroundings, that is an endothermic process. Examples: touching a hot stove, putting an ice cube in a warm drink, sweating.

To help determine heat flow into and out of a system, chemists define a term called enthalpy (H). The actual enthalpy of a substance can’t be measured but the change in enthalpy,  H, can be measured. The enthalpy of a reaction is the difference between the enthalpies of products – reactants, in other words,  H. This represents the heat absorbed or given off during a reaction.  H can be positive or negative. Positive is for endothermic and negative is for exothermic.

Thermochemical Equations: Lists the chemical equation followed by the  H value for that exact equation and amounts. CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2(l)  H = kJ  This means that kJ is released to the surroundings when 1 mole of CH 4 (g) and 2 moles of O 2 (g) react completely. If less is reacted, than  H will be proportionally smaller and if more is reacted, then  H will be proportionally larger.

 If the reaction is run in reverse, then the  H value remains the same, only opposite in sign.  It is important to specify the state of all substances when writing thermochemical equations.

We have already studied specific heat. A related property is Heat Capacity, which is the amount of heat required to raise the temperature of a given quantity of the substance by one degree Celcius. Heat capacity depends on the amount of the substance, while specific heat is a constant fixed property for that substance.

Enthalpy of Formation – (  H f ) is the  H when a substance is formed from its elements in the normal states. Standard Enthalpy of Formation (  H 0 f ) is the heat change (  H) when a substance is formed from its elements in their normal, standard states (1 atm and 25  C) We can use these values, which can be obtained from reference sources, to easily calculate the standard enthalpy of reaction (  H 0 rxn ). Note: The  H 0 f of an element in its normal, standard state is = 0.  H 0 rxn =  [  H 0 f (products) -  (  H 0 f (reactants)] Let’s do example 6.10 on page 250.

A second way to determine  H 0 rxn makes use of a discovery called Hess’s Law – When reactants are converted to products,  H is the same whether the reaction takes place in one step or in a series of steps. Let’s do example 6.9 on page 249