Chapter 6

Thermochemistry

6.1: I. Nature of Energy• A. Energy (E): capacity for work or producing heat

• B. Law of Conservation of Energy: can’t be created or destroyed only changed to different forms

• C. Potential Energy (PE): due to position or composition, Ex. Ball held in the air

• D. Kinetic Energy (KE): due to motion

• E. Activation Energy (Ea): The energy that must be overcome for a chemical rxn to occur

II. Energy Transfer• A. Frictional heating: loss of energy as heat

by friction against surface• B. Temperature: property of random motion of

particles in a substance• C. Heat (q): transfer of energy due to temp.

difference between objects• D. Work (w): force acting over a distance

Ex. Lifting a dumbbell*** Energy can be transferred as heat or

work***

III. Chemical Energy• A. For energy transfer we need to identify parts

• B. “System”: part we focus on

• C. “Surroundings”: everything around system

• D. This reaction between Na and H2O would

be your system

• E. The rest of the universe is the surroundings

IV. Endothermic

• C. The reaction has absorbed energy as reactants become products

• A. Endothermic: the system absorbs energy• B. (“Endo” = in, “thermic” = heat)

V. Exothermic

• A. Exothermic: when system loses energy • (“Exo”= Exit, “Thermic” = heat)• B. The reaction has released energy as

reactants become products

VI. 1st Law of Thermodynamics• A. Energy of the “universe” is constant

• B. Internal energy of a system (E) is the sum of the kinetic and potential energy

• C. ***∆ Before anything means how much that value is changing***

• D. ∆E = q + w (q= heat, w= work)

• E. ∆E = - for an exothermic process, energy lost

• F. ∆E = + for endothermic, energy gained

VII. Gas Energy• A. Related to compression/expansion of gas

• B. Gas Work (W) = Force (F) x distance (∆h)

• C. F = Pressure (P) x Area (A)

• D. W = P x A x ∆h

• E. A x ∆h = ∆V (∆V + expanding, - compressed)

• F. W is released (-) for expanding gas, work added (+) to compress a gas

• G. Since W and ∆V have opposite signs, must be a “–” sign in front of pressure

• H. W = - P ∆V

6.2: I. Enthalpy and Calorimetry• A. Enthalpy (H): E + PV

• B. Enthalpy change is equal to heat (q) at constant pressure

• C. ∆H = - for exothermic reaction

• D. ∆H = + for endothermic reaction

II. Calorimetry• A. Measure of heat change in a system

• B. Heat capacity (C): heat absorbed/increase in temperature

• C. Specific heat capacity (C): heat capacity per gram of substance (J/gºC) or (J/g•K)

• D. Molar heat capacity (C): (J/mole ºC) or (J/mole K)

III. Constant Pressure Calorimetry• A. Allows us to measure heat change or

enthalpy change

• B. ∆Q or ∆H = C x m x ∆T

Constant Volume Calorimetry• C. Bomb calorimeter: container where substance is ignited to measure energy of combustion

• D. ∆E = q + w, w = P ∆V, ∆V = 0

• E. ∆E = q

6.3: I. Hess’s Law• A. States that the heat (or enthalpy) for

reactions yielding a net reaction can be added to yield the heat or enthalpy of the net reaction

• B. Ex. H for 1 mole H2O(s) to H2O(g)

H2O(s) H2O(l) H = 6.01KJ/mole

H2O(l) H2O(g) H = 40.7 KJ/mole

II. Characteristics of Enthalpy Change• A. If a rxn reversed,

sign of q or ∆H switches (because direction of heat flow reverses)

• B. Magnitude of q or ∆H is directly proportional to the quantities of reactants and products in a reaction (if double rxn, then double heat)

6.4: I. Standard Enthalpy of Formation• A. ΔH for formation of 1 mole of a compound

from its elements in their standard states

• B. Std. conditions are 1 atm, 1 M, and the states found at 25 ºC

• C. ∆H = Sum ∆Hf (products) - Sum ∆Hf (reactants)