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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)