Thermochemistry

1

• Energy is the capacity to do work
• Radiant energy comes from the sun and is earths
  primary energy source
• Thermal energy is the energy associated with the
  random motion of atoms and molecules
• Chemical energy is the energy stored within the
  bonds of chemical substances
• Nuclear energy is the energy stored within the
  collection of neutrons and protons in the atom
• Potential energy is the energy available by
  virtue of an objects position

6.1
2
Energy Changes in Chemical Reactions
Heat is the transfer of thermal energy between
two bodies that are at different temperatures.
Temperature is a measure of the thermal energy.
greater thermal energy
6.2
3
Thermochemistry is the study of heat change in
chemical reactions.
The system is the specific part of the universe
that is of interest in the study.
open
closed
isolated
energy
nothing
mass energy
Exchange
6.2
4
Exothermic process is any process that gives off
heat transfers thermal energy from the system
to the surroundings.
Endothermic process is any process in which heat
has to be supplied to the system from the
surroundings.
6.2
5
Exothermic
Endothermic
6.2
6
Thermodynamics is the study of the
interconversion of heat and other kinds of energy.
State functions are properties that are
determined by the state of the system, regardless
of how that condition was achieved.
energy
, pressure, volume, temperature
DE Efinal - Einitial
Potential energy of hiker 1 and hiker 2 is the
same even though they took different paths.
6.3
7
First law of thermodynamics energy can be
converted from one form to another, but cannot be
created or destroyed.
DEsystem DEsurroundings 0
or
DEsystem -DEsurroundings
Exothermic chemical reaction!
6.3
8
Another form of the first law for DEsystem
DE q w
DE is the change in internal energy of a system
q is the heat exchange between the system and the
surroundings
w is the work done on (or by) the system
w -PDV when a gas expands against a constant
external pressure
6.3
9
A sample of nitrogen gas expands in volume from
1.6 L to 5.4 L at constant temperature. What is
the work done in joules if the gas expands (a)
against a vacuum and (b) against a constant
pressure of 3.7 atm?
w -P DV
P 0 atm
W -0 atm x 3.8 L 0 Latm 0 joules
P 3.7 atm
w -3.7 atm x 3.8 L -14.1 Latm
6.3
10
Enthalpy and the First Law of Thermodynamics
DE q w
At constant pressure
q DH and w -PDV
DE DH - PDV
DH DE PDV
6.4
11
Enthalpy (H) is used to quantify the heat flow
into or out of a system in a process that occurs
at constant pressure.
DH H (products) H (reactants)
DH heat given off or absorbed during a reaction
at constant pressure
Hproducts lt Hreactants
Hproducts gt Hreactants
DH lt 0
DH gt 0
6.4
12
Thermochemical Equations
Is DH negative or positive?
System absorbs heat
Endothermic
DH gt 0
6.01 kJ are absorbed for every 1 mole of ice that
melts at 00C and 1 atm.
6.4
13
Thermochemical Equations
Is DH negative or positive?
System gives off heat
Exothermic
DH lt 0
890.4 kJ are released for every 1 mole of methane
that is combusted at 250C and 1 atm.
6.4
14
Thermochemical Equations

• The stoichiometric coefficients always refer to
  the number of moles of a substance

• If you reverse a reaction, the sign of DH changes

• If you multiply both sides of the equation by a
  factor n, then DH must change by the same factor
  n.

6.4
15
Thermochemical Equations

• The physical states of all reactants and products
  must be specified in thermochemical equations.

How much heat is evolved when 266 g of white
phosphorus (P4) burn in air?
266 g P4
6470 kJ
6.4
16
A Comparison of DH and DE
DE DH - PDV
At 25 0C, 1 mole H2 24.5 L at 1 atm
PDV 1 atm x 24.5 L 2.5 kJ
DE -367.5 kJ/mol 2.5 kJ/mol -370.0 kJ/mol
6.4
17
The specific heat (s) of a substance is the
amount of heat (q) required to raise the
temperature of one gram of the substance by one
degree Celsius.
The heat capacity (C) of a substance is the
amount of heat (q) required to raise the
temperature of a given quantity (m) of the
substance by one degree Celsius.
C m x s
Heat (q) absorbed or released
q m x s x Dt
q C x Dt
Dt tfinal - tinitial
6.5
18
How much heat is given off when an 869 g iron bar
cools from 940C to 50C?
s of Fe 0.444 J/g 0C
Dt tfinal tinitial 50C 940C -890C
q msDt
869 g x 0.444 J/g 0C x 890C
-34,000 J
6.5
19
Constant-Volume Calorimetry
qsys qwater qbomb qrxn
qsys 0
qrxn - (qwater qbomb)
qwater m x s x Dt
qbomb Cbomb x Dt
Reaction at Constant V
DH qrxn
No heat enters or leaves!
6.5
20
Constant-Pressure Calorimetry
qsys qwater qcal qrxn
qsys 0
qrxn - (qwater qcal)
qwater m x s x Dt
qcal Ccal x Dt
Reaction at Constant P
DH qrxn
No heat enters or leaves!
6.5
21
Chemistry in Action
Fuel Values of Foods and Other Substances
1 cal 4.184 J
1 Cal 1000 cal 4184 J
22
Because there is no way to measure the absolute
value of the enthalpy of a substance, must I
measure the enthalpy change for every reaction of
interest?
The standard enthalpy of formation of any element
in its most stable form is zero.
6.6
23
6.6
24
Hesss Law When reactants are converted to
products, the change in enthalpy is the same
whether the reaction takes place in one step or
in a series of steps.
(Enthalpy is a state function. It doesnt matter
how you get there, only where you start and end.)
6.6
25
6.6
26
Calculate the standard enthalpy of formation of
CS2 (l) given that
1. Write the enthalpy of formation reaction for
CS2
2. Add the given rxns so that the result is the
desired rxn.
6.6
27
Benzene (C6H6) burns in air to produce carbon
dioxide and liquid water. How much heat is
released per mole of benzene combusted? The
standard enthalpy of formation of benzene is
49.04 kJ/mol.
6.6
28
In class problems

• How much heat (in kJ) is evolved with 5.00 g of
  aluminum reacts with
• a stoichiometric amount of Fe2O3?
• 2 Al (s) Fe2O3 (s) -------? 2 Fe (s)
  Al2O3 (s) DH -852 kJ
• 2. The specific heat of copper metal is 0.385
  J/gK. How many joules of
• heat are necessary to raise the temperature of a
  1.42 kg block of copper
• from 25 oC to 88.5 oC? What is the molar heat
  capacity of copper metal?
• 3. A calorimeter contained 75 g of water at 16.95
  oC. A 93.3 g sample of
• iron at 65.58 oC was placed in it, giving a final
  temp.of 19.68 oC for the
• system. Calculate the heat capacity of the
  calorimeter. Specific heats
• are 4.184 J/goC for water and 0.444 J/goC for Fe.

29

• 4. A coffee cup calorimeter having a heat
  capacity of 472 J/oC is used
• to measure the heat evolved when the following aq
  solutions, both
• initially at 22.6 oC, are mixed 100 g of
  solution containing 6.62 g
• of Pb(NO3)2, and 100 g of solution containing
  6.00 g of NaI. The
• final temperature is 24.2 oC. Assume that the
  specific heat of the
• mixture is the same as that for water, 4.184
  J/goC. The reaction is
• Pb(NO3)2 (aq) 2 NaI (aq) -----? PbI2 (s) 2
  NaNO3 (aq)
• Calculate the heat evolved in the reaction.
• Calculate DH for the reaction (per mole of
  Pb(NO3)2 consumed).
• 5. A 0.187 g sample of benzene, C6H6, was burned
  in a bomb calor-
• imeter whose heat capacity had been determined to
  be 1.34 kJ/oC. The
• temperature of 2000 g of water rose from 21.92 oC
  to 23.67 oC.
• Determine DE in J/g benzene then in kJ per mole
  of benzene. The
• specific heat of water is 4.184 J/goC.

30

• Methane, the main constituent of natural gas,
  burns in oxygen to
• yield carbon dioxide and water
• CH4 (g) 2O2 (g) -------? CO2 (g)
  2H2O (l)
• Using the following information, calculate the
  DHo (in kJ) for the
• reaction
• CH4 (g) O2 (g) -------? CH2O (g) H2O (g)
  DHo -284 kJ
• CH2O (g) O2 (g) -------? CO2 (g) H2O (g)
  DHo -518 kJ
• H2O (l) ---------? H2O (g)
  DHo 44 kJ
• 7. Calculate DHo for the fermentation of glucose
  to make ethanol
• according to the following equation
• C6H12O6 (s) -------? 2 C2H5OH (l) 2 CO2 (g)
• useful info
• substance DHfo (kJ/mol)
• glucose -1260
• ethanol -277.7
• CO2 -393.5