Chapter 10 Energy

10.1 The Nature of Energy

Energy- the ability to do work or produce heat

Potential energy- energy due to position or composition

Kinetic energy- energy due to motion of the object; determined by mass & velocity

Law of Conservation of Energy- energy can be converted from one form to another but can neither be created nor destroyed

Work- force acting over a distance (W = F x d)

State Function- property of a system that changes independently of its pathway

10.2 Temperature & Heat

Temperature- measure of the random motions of the components of a substance

Heat- flow of energy due to temperature difference

Temperature is measured in degrees Celsius in lab; the metric system unit for temperature is Kelvin.

Heat is measured in calories; 1 calorie is the amount of heat necessary to heat 1 gram of water by 1 degree Celsius. The metric system unit for heat is Joules.

1 calorie = 4.184 J

Figure 10.2: Equal masses of hot and cold water.

Figure 10.3: H2O molecules in hot and cold water.

Figure 10.4: H2O molecules in same temperature water.

10.3 Exothermic & Endothermic Processes

System- part of the universe on which we wish to focus attention

Surroundings- everything else in the universe (see next slide)

*Exothermic- heat is evolved; heat exits

*Endothermic- process that absorb energy from the surroundings

The energy gained by the surroundings must be equal to the energy lost by the system.

Figure 10.5: The energy changes accompanying the burning of a

match.

10.4 Thermodynamics

Thermodynamics- study of energy

1st Law of Thermodynamics- the energy of the universe is constant.

E- internal energy; sum of kinetic and potential energies of all particles in the system.

/\E = q + w

q is heat

w is work

Sample problem:

10.5 Measuring Energy Changes

Units of Heat

Calorie- amount of energy needed to raise the temperature of 1 gram of water by one degree Celsius (or 1 Kelvin)

Joule- metric system unit of heat; 1 calorie = 4.184 joule

It takes 4.184 joules to raise the temperature of 1 gram of water by 1 degree Celsius

Example 10.1, p. 295: Express 60.1 cal in joules?

Problem 10.1, p. 295: How many calories of energy correspond to 28.4 J?

Specific heat capacity- amount of energy required to change the temperature of 1 gram of a substance by 1 degree Celsius

Water: specific heat = 1 cal/g oC or 4.184 J/g oC

q = mc/\T (/\T = Tfinal – Tinitial)

Figure 10.6: A coffee-cup calorimeter.

Problem 10.2, p. 296: Calculate joules required to heat 454 grams of water from 5.4oC to 98.6oC.

Note Table 10.1, p. 297

Problem 10.3, p. 299: A 5.63 gram sample of solid gold is heated from 21oC to 32oC. How much energy (in joules and in calories) is required?

Problem 10.4, p. 300: A 2.8 gram sample of pure metal requires 10.1 J of heat to change its temperature from 21oC to 36oC. What is this metal? (Use Table 10.1)

10.6 Thermochemistry (Enthalpy)

Enthalpy- heat that is produced or absorbed in a reaction.

For most reactions, /\Hp = heat

Problem 10.5, p. 302: The reaction that occurs in heat packs used for sports injuries is

4Fe(s) + 3O2(g) 2Fe2O3(s) /\H = -1652 kJ

How much heat is released when 1.00 gram of iron is reacted with excess oxygen gas?

10.7 Hess’s Law

The change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.

1. If a reaction is reversed, the sign of /\H is changed.

2. The magnitude of /\H is directly proportional to the quantities of reactants and products

10.8 Quality v. Quantity of Energy

10.9 Energy and Our World

(read carefully; questions on the test)

Figure 10.7: Energy sources used in the United States.

Figure 10.8: The earth’s atmosphere.

Figure 10.9: The atmospheric CO2 concentration over the past 1000 years.

10.10 Energy as a Driving Force

Energy spread- concentrated energy is dispersed widely

Matter spread- molecules are spread out and occupy a larger volume

Entropy (S) = chaos

The entropy of the universe is increasing. (2nd Law of Thermodynamics)

Example: (NH4)2CO3(s) 2NH3(g) + H2O(g) + CO2(g)

H2O(s) H2O(l) H2O(g)

Figure 10.10: Comparing the entropies of ice and steam.