# Chapter 6 Outline

## 6.2 The Nature of Energy: Key Definitions

There are two basic types of energy. Energy is often described as the potential to do work. There are many different forms of energy. These are classified as either :
1. Kinetic energy. The energy of motion. A car moving at 55 mph has more kinetic energy than a sponge at 5 mph. That is why the car hurts more when it hits you. Kinetic energy depends on the mass and the velocity (speed) of an object. KE = mass * (velocity)2
2. Potential Energy. This is the energy that something has stored. The classic example is a rock on the top of a hill. Think about the rock rolling down the hill and releasing energy by smashing Willey E. Coyote. The bigger the rock, the more potential energy it has. The higher the hill the more potential energy the rock has. After the rock smashes Willey, it has no potential energy left. If Willey pushes the rock back up to the top of the hill the rock has more potential energy (and Willey has less because he is tired)

## 6.3 The First Law of Thermodynamics: There is No Free Lunch

1. The First law of thermodynamics "Energy can be neither created nor destroyed"
2. To work thermodynamics problems scientists typically define a "system". This is just the part of the universe that we are paying attention to. In the Willey E. Coyote example the system is the rock and I can describe the energy of the rock.
3. Heat and work. The energy of a system is divided into heat and work. The sign conventions are kind of tricky here (for positive and negative energy, heat, or work). Remember we are talking about the "system". As the rock falls, the system loses energy so Delta Energy is negative. This section introduces some new terms. They include:
4. /\ (an e-mail delta, different from the ^ symbol used for a superscript). This symbol is used to mean "change". It is important to keep in mind that we can only measure change. In the above example I talked about the potential energy of a rock as it falls from a hill. All I can really measure is the change in energy as it falls from the top of the hill to the bottom. The total, absolute energy would also have to include things like how fast the earth is moving through space, and other things that can not be determined.
5. State functions. (This is not the same as an elegant dinner at the governor's mansion). The idea here is that you can describe the "state" a system is in. Typically in chemistry we describe the "initial state" (the reactants), and the "final state" (the products). Thermochemistry is about measuring the energy difference between the initial and final states. In the above example, the initial state is "Rock on top of hill" and the final state is "Rock on Willey E. Coyote". From an energetics standpoint it does not matter how he gets down.

## 6.4 Quantifying Heat and Work

One of the underlying concepts of thermodynamics is that energy is required to change temperature. This is frequently used because we can measure temperature change. If we know the energy required to change the temperature of an object, we can use temperature change as an indirect measure of energy. This is important because we can not measure energy directly. Use the units of specific heat capacity (J g-1 K-1) and heat capacity (J K-1) to keep track of the problems. Units are key to double checking thermodynamics problems.

E = mass * /\T * C (where C is the specific heat capacity)

Key points to remember when working these problems are that 1) this only works for temperature CHANGE, 2) you have one equation with four variables, rearrange to solve what you need. This is the relationship between energy and temperature change. USE IT.

## 6.5 Measuring E for Chemical Reactions: Constant-Volume Calorimetry

• Energy = Heat + Work
• One type work is as pressure * /\ Volume.
• If only P /\V work is done, then the change in heat (absorbed or released by the system) is called ENTHALPY (H). Remember that: /\H = /\E + P*/\V
• In this section we will work with reactions that do not produce any change in volume. This means /\V = 0 so that /\E = q.

• ## 6.6 Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure

1. Energy = Heat + Work
2. One type work is as pressure * /\ Volume.
3. If only P /\V work is done, then the change in heat (absorbed or released by the system) is called ENTHALPY (H). Remember that: /\H = /\E + P*/\V
4. In this section we will work with reactions that do not produce much change in volume. This means /\V = 0 so that /\H =/\E. (We will come back to this assumption in more detail next semester). This semester we will talk about /\H, the enthalpy of reaction. Remember that /\H is the change in energy for a "system"

## 6.7 Constant-Pressure Calorimetry: Measuring Hrxn

This section describes how /\H is determined experimentally. The basic idea is to measure the heat released or absorbed by a reaction. The amount of heat is also /\H (remember this assumes no PV work). to measure heat you need two things, the heat capacity and the temperature change. Heat capacity is kind of like density. the heat capacity is the amount of energy required to change the temperature by 1 degree.

## 6.9 Enthalpies of Reaction from Standard Heats of Formation

"This is the energy required to make one mole of a compound from the elements in their standard states" This is just an easy way to use Hess's law. You can use tables of "heat of formation" to make a series of reactions to use for Hess's law. Think of these as building blocks. In these problems you "un-form" the reactants to give the energy to change all the elements to their standard states (use the enthalpy of formation to determine the amount of energy required). Then you "form" the products from the elements in their standard states (use /\H formation to determine the amount of energy required). Add up all these reaction steps to get the total reaction and add up all the energy steps to get the total energy