Chapter 8 Chapter 8

Thermochemistry: Chemical Energy

Energy and Its ConservationEnergy and Its Conservation

Units:

1 Cal = 1000 cal = 1 kcal

1 cal = 4.184 J (exactly)

Energy: The capacity to supply heat or do work.

Kinetic Energy (EK): The energy of motion.

Potential Energy (EP): Stored energy.

Thermal Energy: The kinetic energy of molecular motion and is measured by finding the temperature of an object

Heat: The amount of thermal energy transferred from one object to another as the result of a temperature difference between the two

8.2 Internal Energy and State 8.2 Internal Energy and State FunctionsFunctions

First Law of Thermodynamics: The total internal energy E of an isolated system is constant - The sum of all the kinetic and potential energies of particles making up a substance

E = Efinal - Einitial

Systems and SurroundingsSystems and Surroundings

System = the object in question Surrounding(s) = everything outside the

system When both system and surrounding at same temperature thermal equilibrium When not Heat transfer to surrounding = exothermic (you feel the heat) hot metal!! Heat transfer to system = endothermic (you feel cold) cold metal!!

Internal Energy and State Internal Energy and State FunctionsFunctions

State Function: A function or property whose value depends only on the present state, or condition, of the system, not on the path used to arrive at that state

8.3 Expansion Work8.3 Expansion Work

3CO2(g) + 4H2O(g)C3H8(g) + 5O2(g)

7 mol of gas6 mol of gas

w = F x d = - P∆V

Expansion Work: Work done as the result of a volume change in the system

8.4 Energy and Enthalpy8.4 Energy and Enthalpy

Constant Pressure:

E = q + w

w = work = -PVq = heat transferred

q = E + PV

qP = E + PV

Constant Volume (V = 0): qV = E

Energy and EnthalpyEnergy and Enthalpy

= Hproducts - Hreactants

Enthalpy changeor

Heat of reaction (at constant pressure)

qP = E + PV = H

H = Hfinal - HinitialEnthalpy is a state function whose value depends only on the

current state of the system, not on the path taken to arrive at

that state.

8.6The Thermodynamic 8.6The Thermodynamic Standard StateStandard State

3CO2(g) + 4H2O(g)C3H8(g) + 5O2(g)

3CO2(g) + 4H2O(l)C3H8(g) + 5O2(g) H = -2219 kJ

Thermodynamic Standard State: Most stable form of a substance at 1 atm pressure and at a specified temperature, usually 25 °C; 1 M concentration for all substances in solution.

Standard enthalpy of reaction is indicated by the symbol ΔHo

3CO2(g) + 4H2O(g)C3H8(g) + 5O2(g) H= -2043 kJ

H° = -2043 kJ

Enthalpies of Physical and Enthalpies of Physical and Chemical ChangeChemical Change

Enthalpies of Physical and Enthalpies of Physical and Chemical ChangeChemical Change

Enthalpy of Fusion (Hfusion): The amount of heat necessary to melt a substance without changing its temperature

Enthalpy of Vaporization (Hvap): The amount of heat required to vaporize a substance without changing its temperature

Enthalpy of Sublimation (Hsubl): The amount of heat required to convert a substance from a solid to a gas without going through a liquid phase

Enthalpies of Physical and Enthalpies of Physical and Chemical ChangeChemical Change

2Al(s) + Fe2O3(s)2Fe(s) + Al2O3(s) Ho = +852 kJ

2Fe(s) + Al2O3(s)2Al(s) + Fe2O3(s)Ho = -852 kJ

endothermic

exothermicThe reverse reaction

Applying Stoichiometry to Applying Stoichiometry to Heats of ReactionHeats of Reaction

A propellant for rockets is obtained by mixing the liquids hydrazine, N2H4, and dinitrogen tetroxide, N2O4. These compounds react to give gaseous nitrogen, N2 and water vapor, evolving 1049 kJ of heat at constant pressure when 1 mol N2O4 reacts.

a. Write the thermochemical equation for this reaction

b. Write the thermochemical equation for the reverse of the reaction

c. How much heat evolves when 10.0 g of hydrazine reacts according to the reaction described in (a) ?

ExampleExample An LP gas tank in a home barbeque contains 13.2 kg of

propane, C3H8. Calculate the heat (in kJ) associated with the complete combustion of all of the propane in the tank

C3H8(g) + CO2(g) 3 CO2(g) + 4 H2O(g) ΔHo = -2044kJ

8.6 Calorimetry8.6 Calorimetry The process of measuring

heat transfer in at constant pressure (ΔH).

chemical/physical process qrxn + qsol = 0

qrxn = -qsol

Rrxn = system

Ssoln = surrounding

What you’ll do in lab Heat given off by rxn Measured by

thermometter Figure out qrxn indirectly

Calorimetry and Heat Calorimetry and Heat CapacityCapacityMeasure the heat flow at constant volume (E).

Calorimetry and Heat Calorimetry and Heat CapacityCapacity

Specific Heat: The amount of heat required to raise the temperature of 1 g of a substance by 1 °C.

q = (specific heat) x (mass of substance) x T

Molar Heat Capacity (Cm): The amount of heat required to raise the temperature of 1 mol of a substance by 1 °C.

q = (Cm) x (moles of substance) x T

Heat Capacity (C): The amount of heat required to raise the temperature of an object or substance a given amount.

q = C x T

Calorimetry and Heat Calorimetry and Heat CapacityCapacity

ExampleExample How much heat must be added to a 8.21 g sample of

gold to increase its temperature by 6.2 oC? The specific heat of gold is 0.13 J/goC.

ExampleExample When 1.045 g of CaO is added to 50.0 mL of water at

25.0oC in a calorimeter, the temperature of the water increases to 32.2 oC. Assuming that the specific heat of the solution is 4.18 J/goC and that the calorimeter itself absorbes a negligible amount of heat, calculate ∆H in kilojoules for the reaction

CaO(s) + H2O(l) Ca(OH)2(aq)

8.98.9Hess’s LawHess’s Law

8.9 Hess’s Law8.9 Hess’s Law

Haber Process:

Multiple-Step Process - Given

N2H4(g)2H2(g) + N2(g)

Hess’s Law: The overall enthalpy change for a reaction is equal to the sum of the enthalpy changes for the individual steps in the reaction.

2NH3(g)3H2(g) + N2(g) H°total = ???

H°1 = 95.4 Kj

2NH3(g)N2H4(g) + H2(g)

H°total =2NH3(g)3H2(g) + N2(g)

H°2 = -187.6 kJ

ExampleExample Find ΔHo

rxn for the following reaction

C(s) + H2O(g) CO(g) + H2(g) Horxn = ?

Use the following reactions with known H’s

C(s) + O2(g) CO2(g) ΔHo = -393.5 kJ

2CO(g) + O2(g) 2CO2(g) Δ Ho = -566.0kJ

2H2 (g) + O2(g) 2H2O (g) Δ Ho = -483.6 kJ

Standard Heats of FormationStandard Heats of FormationStandard Heat of Formation (Ho

f ): The enthalpy change for the formation of 1 mol of a substance in its standard state from its constituent elements in their standard states

Hof = -74.8 kJCH4(g)C(s) + 2H2(g)

Standard states

1 mol of 1 substance

Standard Heats of FormationStandard Heats of Formation

cC + dDaA + bB

ReactantsProducts

Ho = Hof (Products) - Ho

f (Reactants)

Ho = [c Hof (C) + d Ho

f (D)] - [a Hof (A) + b Ho

f (B)]

Standard Heats of FormationStandard Heats of FormationUsing standard heats of formation, calculate the standard enthalpy of reaction for the photosynthesis of glucose (C6H12O6) and O2 from CO2 and liquid H2O.

C6H12O6(s) + 6O2(g)6CO2(g) + 6H2O(l)

Ho = ?

ExampleExample Calculate the enthalpy change for the following

reaction

3 NO2(g) + H2O(l) 2 HNO3(aq) + NO(g)

*Use standard enthalpies of formation

An Introduction to EntropyAn Introduction to EntropySpontaneous Process: A process that, once started, proceeds on its own without a continuous external influence

Entropy (S): The amount of molecular randomness in a system

An Introduction to EntropyAn Introduction to Entropy

Spontaneous processes are• favored by a decrease in H (negative H).• favored by an increase in S (positive S).

Nonspontaneous processes are• favored by an increase in H (positive H).• favored by a decrease in S (negative S).

ExampleExample Predict whether ΔSo is likely to be positive or negative

for each of the following reactions

a. H2C=CH2(g) + Br2(g) BrCH2CH2Br(l)

b. Consider the following figures