Transcript Enthalpies of Reaction

Thermochemistry
Energy
1st Law of Thermodynamics
Enthalpy / Calorimetry
Hess' Law
Enthalpy of Formation
The Nature of Energy
Kinetic Energy and Potential Energy
• Kinetic energy is the energy of motion:
1
Ek   m  v 2
2
• Potential energy is the energy an object possesses by
virtue of its position.
• Potential energy can be converted into kinetic energy.
Example: a bicyclist at the top of a hill.
The Nature of Energy
Units of Energy
• SI Unit for energy is the joule, J:
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1
2
2
Ek   m  v   2 kg   1 m/s 
2
2
2 -2
 1 kg m s  1 J
sometimes the calorie is used instead of the joule:
1 cal = 4.184 J (exactly)
A nutritional Calorie:
1 Cal = 1000 cal = 1 kcal
Thermochemistry Terminology
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System: part of the universe we are interested in.
Surrounding: the rest of the universe.
Boundary: between system & surrounding.
Exothermic: energy released by system to surrounding.
Endothermic: energy absorbed by system from surr.
• Work ( w ): product of force applied to an
object over a distance.
•Heat ( q ): transfer of energy between two objects
The First Law of Thermodynamics
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Internal Energy
Internal Energy: total energy of a system.
Involves translational, rotational, vibrational motions.
Cannot measure absolute internal energy.
Change in internal energy,
E  Efinal  Einitial
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The First Law of Thermodynamics
Relating E to Heat(q) and Work(w)
• Energy cannot be created or destroyed.
• Energy of (system + surroundings) is constant.
• Any energy transferred from a system must be transferred
to the surroundings (and vice versa).
• From the first law of thermodynamics:
E  q  w
The First Law of
Thermodynamics
First Law of Thermodynamics
• Calculate the energy change for a system
undergoing an exothermic process in which
15.4 kJ of heat flows and where 6.3 kJ of
work is done on the system.
E = q + w
The First Law of Thermodynamics
Exothermic and Endothermic Processes
• Endothermic: absorbs heat from the surroundings.
• An endothermic reaction feels cold.
• Exothermic: transfers heat to the surroundings.
• An exothermic reaction feels hot.
Endothermic Reaction
Ba(OH)2•8H2O(s) + 2 NH4SCN(s) 
Ba(SCN)2(s) + 2 NH3(g) + 10 H2O(l)
The First Law of Thermodynamics
State Functions
• State function: depends only on the initial and final states
of system, not on how the internal energy is used.
Enthalpy
• Chemical reactions can absorb or release heat.
• However, they also have the ability to do work.
• For example, when a gas is produced, then the gas
produced can be used to push a piston, thus doing work.
Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)
• The work performed by the above reaction is called
pressure-volume work.
• When the pressure is constant,
w   PV
Enthalpy
Enthalpy
• Enthalpy, H: Heat transferred between the system and
surroundings carried out under constant pressure.
H  E  PV
• Enthalpy is a state function.
• If the process occurs at constant pressure,
H   E  PV 
 E  PV
Enthalpy
• Since we know that
w   PV
• We can write
H  E  PV
 q P  w  PV
 q P  (  PV )  PV
 qP
• When H is positive, the system gains heat from the
surroundings.
• When H is negative, the surroundings gain heat from
the system.
Enthalpy => Heat of Reaction
Enthalpies of Reaction
• For a reaction:
H  H final  H initial
 H products  H reactants
• Enthalpy is an extensive property (magnitude H is
directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
H = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) H = 1604 kJ
Enthalpies of Reaction
• When we reverse a reaction, we change the sign of H:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g)
H = +802 kJ
• Change in enthalpy depends on state:
H2O(g)  H2O(l) H = -44 kJ
Calorimetry
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Heat Capacity and Specific Heat
Calorimetry = measurement of heat flow.
Calorimeter = apparatus that measures heat flow.
Heat capacity = the amount of energy required to raise
the temperature of an object (by one degree).
Molar heat capacity = heat capacity of 1 mol of a
substance.
Specific heat = specific heat capacity = heat capacity of 1
g of a substance.
q  specific heat  grams of substance  T
Table 5.2: Specific Heats (S) of Some Substances at 298 K
Substance
N2(g)
Al(s)
Fe(s)
Hg(l)
H2O(l)
H2O(s)
CH4(g)
CO2(g)
Wood , Glass
S ( J g-1 K-1 )
1.04
0.902
0.45
0.14
4.184
2.06
2.20
0.84
1.76 , 0.84
If 24.2 kJ is used to warm a piece of aluminum with a mass of 250. g, what is
the final temperature of the aluminum if its initial temperature is 5.0oC?
q  S m T
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Calorimetry
Constant Pressure Calorimetry
• Atmospheric pressure is constant!
H  qP
qrxn  qsoln  specific heat of solution 
 grams of solution   T
Calorimetry
Constant Pressure Calorimetry
qrxn  S  ( total mass of solution)  T
H rxn
qrxn

mol *
* moles of species of interest
Calorimetry
Calorimetry Examples
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1. In an experiment similar to the procedure set out for Part
(A) of the Calorimetry experiment, 1.500 g of Mg(s) was
combined with 125.0 mL of 1.0 M HCl. The initial
temperature was 25.0oC and the final temperature was
72.3oC. Calculate: (a) the heat involved in the reaction
and (b) the enthalpy of reaction in terms of the number of
moles of Mg(s) used. Ans: (a) –25.0 kJ (b) –406 kJ/mol
2. 50.0 mL of 1.0 M HCl at 25.0oC were mixed with 50.0
mL of 1.0 M NaOH also at 25.0oC in a styrofoam cup
calorimeter. After the mixing process, the thermometer
reading was at 31.9oC. Calculate the energy involved in
the reaction and the enthalpy per moles of hydrogen ions
used. Ans: -2.9 kJ , -58 kJ/mol [heat of neutralization
for strong acid/base reactions]
Hess’s Law
• Hess’s law: if a reaction is carried out in a number of
steps, H for the overall reaction is the sum of H for
each individual step.
• For example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
H = -802 kJ
2H2O(g)  2H2O(l)
H= - 88 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
H = -890 kJ
Another Example of Hess’s Law
Given:
C(s) + ½ O2(g)  CO(g)
H = -110.5 kJ
CO2(g)  CO(g) + ½ O2(g)
H = 283.0 kJ
Calculate H for:
C(s) + O2(g)  CO2(g)
Enthalpies of Formation
• If 1 mol of compound is formed from its constituent
elements, then the enthalpy change for the reaction is
called the enthalpy of formation, Hof .
• Standard conditions (standard state): Most stable form of
the substance at 1 atm and 25 oC (298 K).
• Standard enthalpy, Ho, is the enthalpy measured when
everything is in its standard state.
• Standard enthalpy of formation: 1 mol of compound is
formed from substances in their standard states.
Enthalpies of Formation
• If there is more than one state for a substance under
standard conditions, the more stable one is used.
• Standard enthalpy of formation of the most stable form of
an element is zero.
Enthalpies of Formation
Substance
Hof (kJ/mol)
C(s, graphite)
0
O(g)
247.5
O2(g)
0
N2(g)
0
Enthalpies of Formation
Using Enthalpies of Formation to
Calculate Enthalpies of Reaction
• For a reaction
H rxn   n  H f products    m  H f reactants 
• Note: n & m are stoichiometric coefficients.
• Calculate heat of reaction for the combustion of propane
gas giving carbon dioxide and water.
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O()
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O()
∆Hof (kJ/mol):
Foods and Fuels
Foods
• 1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal.
• Energy in our bodies comes from carbohydrates and fats
(mostly).
• Intestines: carbohydrates converted into glucose:
C6H12O6 + 6O2  6CO2 + 6H2O, H = -2816 kJ
• Fats break down as follows:
2C57H110O6 + 163O2  114CO2 + 110H2O, H = -75,520 kJ
Fats contain more energy; are not water soluble, so are good for
energy storage.
Foods and Fuels
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Fuels
Fuel value = energy released when 1 g of substance is
burned.
Most from petroleum and natural gas.
Remainder from coal, nuclear, and hydroelectric.
Fossil fuels are not renewable.
In 2000 the United States consumed 1.03  1017 kJ of
fuel.
• In 2005 the United States consumed 1.05  1017 kJ of energy.
• Hydrogen has great potential as a fuel with a fuel value of
142 kJ/g. [ gasoline ≈ 35 kJ/g ]
Foods and
Fuels
(% for 2000)
Thermochemistry
Ek 
1
 m  v2
2
E  q  w
(kg m 2 s  2  joule)
Energy
1st Law of Thermodynamics
H  qP
qrxn  qsoln  specific heat of solution  grams of solution  T
Enthalpy / Calorimetry
Hess' Law
Enthalpy of Formation
H rxn   n  H f products    m  H f reactants 