Transcript Chapter 6: Thermochemistry Chemistry 1061: Principles of Chemistry I Andy Aspaas, Instructor

Chapter 6: Thermochemistry
Chemistry 1061: Principles of Chemistry I
Andy Aspaas, Instructor
Thermochemistry
• Thermodynamics: relationships between heat and other
forms of energy
– Thermochemistry: an area of thermodynamics that
involves heat transferred due to a chemical reaction
• Energy: potential or capacity to move matter
– Kinetic energy, Ek: macroscopic energy associated with
an object’s movement
– Potential energy, Ep: macroscopic energy assiciated with
an object’s position in a field of force (only relative values)
– Internal energy, U: microscopic sum of energy contained
in a substance’s particles
– Etot = Ek + Ep + U
Heat of reaction
• System: collection of substances in which the
thermodynamic change is happening
• Surroundings: everything outside the system,
includes the flask, the room, and the universe
• Heat, q: energy that flows into or out of a system
because of a difference in temperature
– Thermal equilibrium: heat flows from areas of
high temperature to areas of low temperature
until the temperatures are equal
Energy added or subtracted from a system
• The sign of q is viewed from the perspective of the
system, not the surroundings
– Exothermic process, the reaction vessel warms,
so energy must have left the system: q is –
– Endothermic process, the reaction vessel cools,
so energy must have been added to the system:
q is +
Enthalpy
• Enthalpy, H: an extensive property (depends on quantity)
which is related to the amount of heat that can be absorbed
or evolved in a chemical reaction
– H is a state function (only depends on present state,
independent of any history)
• ∆H = change in enthalpy = H(products) - H(reactants)
• ∆H = qp (enthalpy change = reaction heat at constant
pressure)
– An exothermic reaction might have qp = -400 kJ
– So, ∆H = -400 kJ
Thermochemical equations
• Thermochemical equation: balanced molar chemical
equation, with enthalpy of reaction given after the equation
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g); ∆H = -368.6 kJ
• 368.6 kJ of heat is evolved when 2 mol Na react with 2 mol
H2O to form 2 mol NaOH and 1 mol H2
• Phase labels are important, ∆H may be different depending
on phase of product
• If a thermochemical equation is multiplied by anything, ∆H is
also multiplied
• If a thermochemical equation is reversed, the sign of ∆H is
also reversed
Stoichiometry and heats of reaction
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l); ∆H = -890.3 kJ
• In excess oxygen, how many kJ of heat can be
obtained from burning 36.0 g CH4?
• Use ∆H as a conversion factor from mol to kJ
Heat capacity
• Heat capacity, C: quantity of heat required to raise
temperature of a substance one degree Celsius
q = C∆t where ∆t = tf - ti
• Specific heat, s: quantity of heat required to raise
temperature of one gram of a substance by one
degree Celsius
q = sm∆t
• Water has a very high specific heat: 4.18 J / (g · °C)
Measuring heat of reaction
• Calorimeter: device used to measure heat of reaction
– Insulated reaction vessel, with thermometer to record
temperature change
– Two nested styrofoam cups works well
• First find amount of heat absorbed by calorimeter
q = sm∆t and q = C∆t
– Heat absorbed by surroundings is opposite of heat given
off by system
qcalorimeter = -qreaction
– Divide qreaction by correct number of moles of limiting
reactant to set up a thermodynamic equation
Hess’s law
• Adding reactions together will also add ∆H for each
reaction
– So, ∆H can be found for reactions where it would
be difficult to measure experimentally
• Remember, multiplying a reaction by a constant also
multiplies the ∆H by that constant, and reversing a
reaction reverses the sign of ∆H
Standard enthalpies of formation
• Standard enthalpy of reaction, ∆H °: ∆H at 25° C
and 1 atm
• Standard enthalpy of formation, ∆Hf°: enthalpy
change for formation of one mole of the substance
in its standard state from its elements
∆H ° =  n ∆Hf°(products) -  n ∆Hf°(reactants)