Thermochemistry:

• Energy
• Kinetic & Potential
• First Law of Thermo
• internal energy, heat & work
• endothermic & exothermic processes
• state functions
• Enthalpy
• Enthalpies of Reaction

• Calorimetry
• heat capacity and specific heat
• constant-pressure calorimetry
• bomb calorimetry (constant-volume calorimetry)
• Hess's Law
• Enthalpies of Formation
• for calculation of enthalpies of reaction
• Foods and Fuels

• Energy
• work is a form of energy w = F x d
• energy is the capacity to do work or transfer heat
• Kinetic Energy

energy of motion E = ½ mv2

• potential energy

energy of position

applies to electrostatic energy

applies to chemical energy (energy of bonds)


• energy units

one joule = energy of a 2 kg mass moving at 1 m/s

E = ½ mv2 (½)(2 kg) (m/s)2 = kg m2/s2 = 1 J

1 cal = 4.184 J 1 kcal = 1 food calorie (Cal)


• Systems & Surroundings
• system -- chemicals in the reaction

closed system can exchange energy (but not matter) with its surroundings

• surroundings -- container & all outside environment

• First Law of Thermo.

• Energy is always conserved

any energy lost by system, must be gained by surroundings

• Internal Energy -- total energy of system

combination of all potential and kinetic energy of system

incl. motions & interactions of of all components

we measure the changes in energy
DE = Efinal - Einitial


+ D E = Efinal > Einitial system has gained E from
surroundings

• - D E = Efinal < Einitial system has lost E to
  surroundings
• Relating D E to heat and work

D E = q + w q is positive if heat goes from
surroundings to system

w is positive if work is done on
system by surroundings





• Endothermic

system absorbs heat or heat flows into the system

• Exothermic

system gives off heat or heat flows out of the system

• State Function

a property of a system that is determined by specifying its condition or state (T, P, etc.)

internal energy is a state function, \ DE depends only on Efinal & Einitial


• Enthalpy
• for most reactions, most of the energy exchanged is in the form of heat, that heat transfer is called enthalpy, H
• enthalpy is a state function
• like internal energy, we can only measure the change in enthalpy, DH

DH = qp when the process occurs under constant pressure
DH = Hfinal - Hinitial = qp

• - DH Þ exothermic process
• +DH Þ endothermic process

• enthalpy change for forward rxn is equal in magnitude but opposite in sign for the reverse rxn

CH4(g) ® C(s) + 2H2(g)                 D H = +74.8 kJ/mol

C(s) + 2H2(g) ® CH4(g)                 D H = - 74.8 kJ/mol


• enthalpy change for a reaction depends n the state of the reactants and products

C(g) + 2H2(g) ® CH4(g)        D H = -793.2 kJ/mol

2H2(g) + O2(g) ® 2H2O(g)     D H = -486.6 kJ/mol

2H2(g)+ O2(g) ® 2H2O(l)      D H = -571.7 kJ/mol



• Calorimetry

experimental determination of D H using heat flow

• heat capacity

measures the energy absorbed using temperature change

the amount of heat required to raise its temp. by 1 K

molar heat capacity -- heat capacity of 1 mol of substance


• Practice Ex. 5.3:

Calculate the quantity of heat absorbed by 50.0 kg of rocks if their temp. increases by 12.0 °C if the specific heat of the rocks is 0.82 J/gK.

S.H. x g x D T = joules


What unit should be in the solution?                  joules -- quantity of heat


0.82 J x 50.0 x 10³ g x 12.0 K = 4.9 x 10⁵ J
g K


• Constant-Pressure Calorimetry
• D H = qp at constant pressure as in coffee cup calorimeter

heat gained by solution = qsoln

\ qsoln = (S.H.soln)(gsoln)(DT)

heat gained by solution must that which is given off by reaction

\ qrxn = - qsoln = - (S.H.soln)(gsoln)(DT)


• Practice Ex. 5.4:

When 50.0 mL of 0.100 M AgNO₃ and 50.0 mL of 0.100 M HCl are mixed in a c.p. calorimeter, the temp. of the mixture increases from 22.30°C to 23.11°C. Calculate D H for this reaction, assuming that the combined solution has a mass of 100.0 g and a S.H. = 4.18 J/g °C.

AgNO3(aq) + HCl(aq) ® AgCl(s) + HNO3(aq)

qsoln = 4.18 J x 100.0 g soln x 0.81°C = 3.39 x 10² J
          g ^oC

qrxn = - qsoln = - 3.39 x 102 J = - 68,000 J    or    - 68 kJ/mol
                       0.00500 mol             

• Practice Ex. 5.5:

A 0.5865 g sample of lactic acid, HC₃H₅O₃, is burned in a calorimeter with C = 4.812 kJ/°C. Temp. increases from 23.10°C to 24.95°C. Calculate heat of combustion per gram and per mole. D T = +1.85°C

qrxn = - (4.812 kJ/°C) (1.85°C) = - 8.90 kJ per
                                                       0.5865 g lactic acid
-8.90 kJ = - 15.2 kJ/g
0.5865 g

- 15.2 kJ x 90 .1 g = - 1370 kJ/mol
   1 g            1 mol

• Hess's Law
• rxns in one step or multiple steps are additive because they are state functions

eg.


CH4(g) + 2O2(g) ® CO2(g) + 2H2O(g) D H = - 802 kJ


CH4(g) + 2O2(g) ® CO2(g) + 2H2O(l) D H = - 890 kJ


• Practice Ex. 5.6:
• Calculate D H for the conversion of graphite to diamond:

Cgraphite ® Cdiamond


Cgraphite + O2(g) ® CO2(g)            D H = -393.5 kJ

Cdiamond + O2(g) ® CO2(g)           D H = -395.4 kJ


Cgraphite + O2(g) ® CO2(g)            D H = -393.5 kJ

CO2(g) ® Cdiamond + O2(g)           D H = 395.4 kJ



• Enthalpies of Formation
• enthalpies are tabulated for many processes

vaporization, fusion, formation, etc.


• enthalpy of formation describes the change in heat when a compound is formed from its constituent elements, DHf
• standard enthalpy of formation, DHfo, are values for a rxn that forms 1 mol of the compound from its elements under standard conditions, 298 K, 1 atm
• elemental forms

eg. C(s) graphite, Ag(s) , H2(g) , O2(g) , etc.


DHf^o, for any element is = 0


• used for calculation of enthalpies of reaction, DHrxn
  
• DHrxn = S DH prod - S DH react

• Practice Ex. 5.9:

Given this standard enthalpy of reaction, use the standard enthalpies of formation to calculate the standard enthalpy of formation of CuO(s)

CuO(s) + H₂(g) ® Cu(s) + H₂O(l)             DH^o = -130.6 kJ


DHrxn = S DH _f ^o prod - S DH _f ^o react


-130.6 kJ = [(0) + (-285.8)] - [(CuO) + (0)]

DH_f^o CuO = -155.2 kJ/mol