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CHE1031
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Thermochemistry Superconcepts
This material is presented in chapter 5 of Brown et al. 12/e
Superconcepts
1. Energy is the ability to move an object against a force, to raise the temperature of an
object, or the potential to do either.
2. Unless they involve the use or production of gases, chemical reactions involve potential
energy and heat, but not kinetic energy.
3. Energy is conserved, but is transferred or transformed by chemical reactions. And as
energy is transferred some is always lost.
Concepts
a. Kinetic energy is the energy of objects in motion while potential energy is energy
stored by position or in chemical bonds.
b. In thermochemistry, the system is the chemical reaction itself (only the reactants &
products). The surroundings are everything else in the universe.
c. First law of thermodynamics = energy is conserved, neither created nor destroyed;
only transferred or transformed.
d. Enthalpy = heat flow at constant pressure (= ΔH or qp)
e. Thermochemical equations are balanced chemical equations with associated
enthalpy changes.
f. Hess’s law says that when a chemical reaction can be seen as the sum of a series of
smaller reactions, the enthalpy of that chemical reaction is equal to the sum of the
enthalpies of the summed reactions.
g. Calorimetry = the experimental determination of heat flow.
h. Enthalpy of formation = amount of energy to form 1 mole of product from
components in their standard states (J/mol or kJ/mol).
Details
i.The unit of energy is a joule (J) = (kg)(m2)/s2
[1 kJ = 1000 J]
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ii.Kinetic energy = (1/2)(mass)(velocity )
iii.Heat always flows uni-directionally: from hot to cold
iv.Can you describe a scenario in which potential energy is transformed into kinetic
energy, heat, and work?
v.Closed systems exchange only energy with their surroundings while open systems
exchange both energy and matter with their surroundings.
vi.Energy lost by the system is gained by the surroundings, & vice versa.
a. ΔE = q + w [ΔE = internal energy; q = heat; w = work]
b. ΔE = Ef – Ei [Ef = final energy; Ei = initial energy]
vii.Exothermic reactions occur when the system loses energy to the surroundings (ΔH =
negative). Endothermic reactions gain energy from their surroundings (ΔH =
positive).
viii.Enthalpy is extensive, reversible (sign convention), and dependent on physical state.
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ix.To solve Hess’s problems: 1) Look for reagents unique to each reaction & the overall
reaction; 2) place them on the same side of the reaction; 3) multiply or divide to
achieve the proper number of molecules.
x.Specific heat = the amount of heat required to raise the temperature of 1 gram of a
pure substance by 1 degree C.
xi.Molar heat capacity = “
“ one mole of a pure substance “
.
xii.Heat capacity = the amount of heat required to raise the temperature of an object by
1 degree C.
xiii.ΔH = (mass)(specific heat)(ΔT) = (g)(J/g/°C)(°C) = J
Or, since calorimeters absorb some of the heat:
ΔH = (mass)(specific heat)(ΔT) + (Ccal)(ΔT)
[Ccal = heat cap of calorimeter]
xiv.Standard state is the physical state of any element at 1 atm pressure, 25°C or 298°K.
xv.The enthalpy of a reaction can be calculated from enthalpies of formation of its
products and reactants:
ΔHrxn = nH°f(products) - nH°f(reactants)
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