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 calorimeterq = 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)