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
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.
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)
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)
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)