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Vermont Tech
CHE1031
Thermochemistry: The VERY Short Story
Covered in Chapter 5 of Brown et al.
Kinetic energy: The energy of an object in motion; proportional to mass & velocity
Ek = (1/2)(m)(v2) = kg-m2/s2 = J
Potential energy: energy stored by physical position, state or chemical bonds
 Results from the attractive & repulsive forces an object experiences relative to other
objects. In chemistry, electrostatic force is the driver of potential energy.
Thermochemistry: study of energy, its transfers and transformations during chemical
reactions.
 Energy: capacity to do work or transfer heat
 Work: moving an object against a force (like gravity)
 Heat: energy causing a change in temperature
 Thermal energy: energy an object or substance possesses as temperature because of
the kinetic energy of its molecules
 Heat flow: heat always moves from warmer objects to cooler objects
Energy is can be transferred or transformed into a variety of forms. For example, kinetic
energy can become potential energy, heat and work.
System vs. surroundings - In a chemistry lab, these terms are defined as follows:
 System: the molecules participating in a chemical reaction (reactants & products)
 Surroundings: everything else, from the water of the aqueous solutions to universe
 Closed system: exchange only energy with the surroundings; artificial but easy to
study
 Open systems: exchange energy & matter with their surroundings
First Law of Thermodynamics: Energy is conserved, neither created nor destroyed but only
transferred or transformed.
 From system to surroundings, or vv
 From kinetic to potential or to work or to heat, etc.
 ∆E = q + w
where q is heat, w is work
 ∆E = Ef – Ei
where Ef is final energy, Ei is initial energy
 Most chemical reactions absorb or release heat, but don’t do work.
Enthalpy (∆H): heat absorbed or released under constant pressure, units of J
 A state function: pathway independent
 ∆H = Hf – Hi = H of products – H of reactants
 ∆H = ∆E + P∆V … for reactions that involve gases
 +∆H = endothermic: system absorbs energy
 - ∆H = exothermic: system releases energy
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Vermont Tech
CHE1031
Thermochemical equation: a balanced chemical equation with an associated enthalpy
change
 Conversion factors are made by combining stoichiometry and enthalpy
Hess’s Law: If a reaction can be broken down into a series of smaller steps, the enthalpy of
the overall reaction is equal to the sum of the enthalpies of the smaller steps.
 If a reaction is reversed, change the sign of its enthalpy but not the magnitude.
 If reactions are multiplied or divided by a fudge factor, multiply or divide the
associated enthalpy.
 How? For each step and the final equation, find a unique chemical and determine
whether the smaller reaction is going in the appropriate direction and whether it
needs to be multiplied or divided.
Calorimetry: experimental measurement of heat flow using an insulated reaction vessel
 Specific heat: energy to increase 1 g of pure substance by 1°C
(J/g-°C)
o water = 4.184 J/g-°C
 Heat capacity: energy to increase an object by 1°C
(J/°C)
 Molar heat capacity: energy to increase 1 mole of pure substance by 1°C (J/mol-°C)
 Units of ∆T can be either °C or °K; each degrees are equivalent
 Assume the density of aqueous solutions is 1 g/mL
∆H = - (mass)(specific heat)(∆T) = (g)(J/g-°C)(°C) = J
To account for heat absorbed by the calorimeter, add a second term: (Ccal)(∆T)
∆H sol’n = - ∆Hrxn
Because energy released by the reaction is absorbed by the
solution; so enthalpy of rxn and sol’n are equal and opposite.
Enthalpy of formation (∆Hf): enthalpy change accompanying formation of 1 mole of
compound from elements in their standard states
 Standard states: physical state at 1 atm, 25°C, 1 M
 ∆Hf in standard state always = zero
(H2, N2, O2, F2, Cl2, Br2, I2)
 Enthalpy of formation equation
 Enthalpy of reaction can be calculated by summing ∆Hf when stoichiometry is
taken into account:
∆Hrxn = (∑n∆Hf products) – (∑n∆Hf reactants)
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