Transcript Chapter 6 – Thermochemistry

Thermochemistry
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First Law of Thermodynamics
System, surrounding, and thermodynamic universe
Heat (q), work (w) and internal energy (E)
Calculation of heat gained or lost by system
State Functions
Enthalpy of reactions
Calorimetry; determination of enthalpy of reactions
Hess’s law of enthalpy summation
Energy, Work, Heat & Enthalpy
• Energy – the capacity to do work or to produce heat
• Potential Energy:
– energy associated with the relative position of a substance
in a force field, such as gravitational attraction, chemical
bonds, electrostatic, nuclear force, etc., or on the chemical
composition
• Kinetic Energy:
– energy associated with the translational motion of an
object. Kinetic energy depends on the mass and speed of
the object. EK = ½mv2
Energy, Work, Heat & Enthalpy
• Heat:
– energy associated with temperature change – heat flows
from a hot object to a cold one.
• Heat gained or lost:
q = m.s.Dt
– where m = mass; s = specific heat capacity of the
substance, and Dt = change in temperature
• Enthalpy change (DH):
– Heat gained or lost during a chemical reaction at constant
pressure.
First Law of Thermodynamics
• Energy is not created nor destroyed during chemical or
physical processes
• The change in the internal energy of a system (DE) depends
only on the amount of heat (q) gained or lost by the system and
the work (w) done on or by the system.
DE = q + w
For processes that involve gas expansion or
compression, w = -pDV
DE = q – pDV;
q = DE + pDV
Coffee Cup Calorimeter
Heat Capacity of Calorimeter
• To determine the heat capacity of coffee cup
calorimeter:
– A 25.0-g sample of warm water at 40.0oC was added to a
25.0-g sample of water in a Styrofoam coffee cup
calorimeter initially at 20.0oC. The final temperature of the
mixed water and calorimeter was 29.5oC and the specific
heat capacity of water is 4.184 J/g.oC. Calculate the heat
capacity, Ch, of the calorimeter.
Specific Heat Capacity
• To determine the specific heat capacity of a metal
using coffee cup calorimeter:
– A 55.0-g sample of hot metal initially at 99.5oC was added
to 40.0 g of water in a Styrofoam coffee cup calorimeter.
The water and calorimeter were initially at 21.0oC. If the
final temperature of mixture was 30.5oC, calculate the total
heat lost by metal and the specific heat capacity of the
metal. The specific heat of water is 4.184 J/(g.oC) and heat
capacity of calorimeter is 10.0 J/oC.
Heat of Neutralization
• To determine the molar enthalpy of acid-base reaction using
coffee cup calorimeter.
– 50.0 mL of 2.0 M HCl was reacted with 50.0 mL of 2.0 M
NaOH in a coffee cup calorimeter. The reaction was exothermic,
which caused the temperature of the solution to increase from
22.0oC to 35.6oC. Assume the density of solution as 1.0 g/mL, its
specific heat capacity as 4.18 J/g.oC, and the heat capacity of
calorimeter as 10.J/oC. Calculate the total amount of heat
produced by the reaction. Calculate the enthalpy change (DH, in
kJ/mol) for the following reaction:
• HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Calculation the enthalpy of reaction using
Styrofoam cup calorimeter
• To determine the enthalpy of reaction using coffee cup
calorimeter.
– Suppose 100. mL of 1.0 M HCl solution is placed in a Styrofoam coffee
cup calorimeter. The initial temperature of HCl solution is 22.5oC. A
0.255-g sample of magnesium ribbon is cut to short pieces and added to
the acid solution in which the following exothermic reaction occurred.
• Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
– The heat produced by the above reaction is completely absorbed by the
solution and calorimeter, which attained the highest temperature of
34.2oC. Assume the acid solution has a density of 1.0 g/mL and its
specific heat capacity as 4.0 J/g.oC, and the calorimeter has a heat
capacity of 10. J/oC. Calculate the molar enthalpy change (DH, in
kJ/mol) for the above reaction.
Calculating the enthalpy of reaction using
Styrofoam cup calorimeter
Reaction: Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g);
Calculations:
Bomb Calorimeter
Heat Capacity of Calorimeter
• To determine the heat capacity of bomb
calorimeter:
– When a 1.200-g sample of glucose, C6H12O6, was
completely combusted in a bomb calorimeter, the
temperature of the calorimeter assembly increased
by 4.48oC. If the combustion of glucose produces
14.0 kJ/g of energy, how much heat energy is
absorbed by the calorimeter. Calculate the heat
capacity, Ch, of the calorimeter. (Assume that all of
the heat produced by the combustion of glucose is
absorbed by the calorimeter.)
Enthalpy changes of exothermic and
endothermic reactions
Heat of Combustion
• To calculate the enthalpy of combustion using bomb
calorimeter
– When a 1.010-g sample of sucrose (cane sugar) is
completely combusted in a bomb calorimeter, the
temperature of the calorimeter was increased by 4.50oC. If
the heat capacity of calorimeter is 3.75 kJ/oC, how much
heat was absorbed by the calorimeter? Calculate the molar
enthalpy of combustion of sucrose according to the
following equation:
• C12H22O11(s) + 12O2(g)  12CO2(g) + 11H2O(l)
Hess’s Law of Enthalpy of Reactions
• According to Hess’s law:
– The net enthalpy change of a given process is
independent of the number of steps taken to
complete the process.
– If a reaction can be broken down into several
steps, then the overall enthalpy for the reaction is
equal to the sum of enthalpies of individual steps.
Hess’s Law of Enthalpy of Summation
• According to Hess’s law:
– For a given reaction, the overall enthalpy change is equal
to the difference between the algebraic sum of enthalpies of
formation of products and the algebraic sum of enthalpies
of formation of reactants.
DHrxn = S(npDHf[products]) – S(nrDHf[reactants])
Enthalpy of Reactions
• Applying Hess’s law to calculate enthalpy of reactions.
• Given:
– C(s) + O2(g)  CO2(g);
DHo = -394 kJ (1)
– CO(g) + ½O2(g)  CO2(g); DHo = -283 kJ (2)
• Calculate DH for the following reaction:
– C(s) + ½O2(g)  CO(g)
Enthalpy Changes in Reactions #1
• Examples:
– C(s) + O2(g)  CO2(g); DH = -394 kJ
– C(s) + ½O2(g)  CO(g); DH = -283 kJ
– C(s) + ½O2(g) ____________ DHf = 0.0 kJ
– CO(g) + ½O2(g) ____________ DHf = -283 kJ
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CO2(g) ____________ DHf = -394 kJ
Enthalpy of Formation
Enthalpy of Combustion
• Calculating enthalpy of combustion of diborane, B2H6, using
the molar enthalpy of formation data:
• Given:
– (1) 4B(s) + 3O2(g)  2B2O3(s);
– (2) 2B(s) + 3H2(g)  B2H6(g);
– (3) 2H2(g) + O2(g)  2H2O(g);
DHf(1) = -2546 kJ
DHf(2) = 36.4 kJ
DHf(3) = -483.7 kJ
• Calculate DH for the following combustion reaction:
– B2H6(g) + 3O2(g)  B2O3(s) + 3H2O(g);
Reaction Enthalpy Diagram
• Reaction: B2H6(g) + 3O2(g)  B2O3(s) + 3H2O(g)
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B2H6(g) ___________ DHf = 36 kJ
• 2B(s) + 3H2(g) + 3O2(g) __________ DHf = 0.0 kJ
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3H2O(g) ____________ DHf = -726 kJ
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B2O2(s) ____________ DHf = -1273 kJ
Heat of Combustion Reaction
• To calculate heat of reaction using molar enthalpy of
formation:
• Given: DHf[C3H8(g)] = -104 kJ/mol;
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DHf[H2O(g)] = -242 kJ/mol, and
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DHf[CO2(g)] = -394 kJ/mol;
• Calculate DH for the following combustion of C3H8(g):
– C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
Enthalpy Diagram
• Reaction: C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
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• 3C(s) + 4H2(g) + 5O2(g) ____________ DHf = 0.0 kJ
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C3H8(g) ____________ DHf = -104 kJ
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4H2O(g) ____________ DHf = -968 kJ
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3CO2(g) ____________ DHf = -1182 kJ
Enthalpy Diagram for the Combustion of
Methane
Enthalpy Diagram for Combustion Reaction
Reactions in Aqueous Solution
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Enthalpy change for reactions in aqueous solution:
Standard condition for gas: P = 1 atm at 25oC
Standard condition for solution: 1 M at 25oC
Under standard condition: DHf[H3O+(aq)] = 0.0 kJ
Given:
 DHf[Mg2+(aq)] = -467 kJ/mol;
 DHf[OH-(aq)] = -230 kJ/mol, and DHf[Mg(OH)2(s)] = -925 kJ/mol,
• Calculate DH for the reaction:
– MgCl2(aq) + 2NaOH(aq)  Mg(OH)2(s) + 2NaCl(aq)
Enthalpy of Ionic Reactions
• Given the following enthalpy of formation (in kJ/mol):
• DHf[Ba2+(aq)] = -537.6; DHf[CO32-(aq)] = -677.1;
DHf[BaCO3(s)] = -1219; DHf[BaSO4(s)] = -1465;
DHf[OH-(aq)] = -230;
DHf[H2SO4(aq)] = -909.3 and
DHf[H2O(l)] = - 286 kJ/mol
• Calculate enthalpy changes for the following reactions in
aqueous solution:
• (1) BaCl2(aq) + Na2CO3(aq)  BaCO3(s) + 2NaCl(aq)
• (2) HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
• (3) Ba(OH)2(aq) + H2SO4(aq)  BaSO4(s) + 2H2O(l)
Global Energy Resources
• Biomass (mainly wood) – major sources of energy in many
under-developed countries;
• Coal was once the major source of energy in U.S.A. and
industrialized European countries;
• Petroleum replaces coal in the middle of 20th Century as the
major source of energy for power plants and transportation;
• Hydroelectric power and nuclear energy are used in certain
developed countries. Geothermal energy is used as a secondary
source of energy
• Solar energy is a secondary source of energy, but mainly for
household heating.
Comparison of Enthalpy of Combustion
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Energy (kJ/g fuel)
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Hydrogen gas (H2)
Natural gas (CH4)
Gasoline
Crude petroleum
Animal fat
Coal
Charcoal
Ethanol
Methanol
Paper
Dry biomass (wood)
120
50
48
43
38
29
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27
20
20
16
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Petroleum and Natural Gas
• Origin of petroleum and natural gas:
– most likely from fossilized remains of marine organisms that
lived approximately 500 millions years ago
• Petroleum –
– thick, dark liquid composed of mixture of hydrocarbons
– Composition varies from one location to another, but mostly
hydrocarbon compounds containing C5 to > C25
• Natural Gas
– Consists mostly methane (CH4, >90%) and some ethane (C2H6),
propane (C3H8) and butane (C4H10)
Petroleum Refining
• Petroleum refining
– Fractional distillation of crude petroleum yields the following
fractions:
• Gasoline (C5 – C10);
• Kerosene & jet fuel (C10 – C18);
• Diesel fuel, heating, and lubricating oil (C15 – C25),
• Asphalt (>C25)
– More gasoline is produced by pyrolytic (high temperature)
cracking of larger HC compounds (> C25)
Coal
• Formed from fossilized plant remains that have been subjected
to high temperature and pressure for many millions years
• Coal matures through 4 stages:
– Lignite, subbituminous, bituminous, and anthracite;
• Composition by mass%:
– Lignite: 71% C, 4% H, 23% O, 1% N, and 1% S;
– Subbituminous: 77% C, 5% H, 16% O, 1% N, and 1% S;
– Bituminous: 80% C, 6% H, 8% O, 1% N, and 5% S;
– Anthracite: 92% C, 3% H, 3% O, 1% N, and 1% S.
• The relative carbon content increases and those of hydrogen and
oxygen decrease as coal matures.
Coal as Energy Source
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Coal furnishes about 23% of energy needs in U.S.A.
Underground coal mining is dangerous
Strip mining destroys lands and the environments
Coal contains sulfur and burning coal causes severe air
pollution due to:
– Air particulate matters, CO2, CO, and SO2 (from sulfur in coal)
– In atmosphere, SO2 is oxidized to SO3, which yields acid rain
(SO3 forms H2SO4 when mixed with rain water)
Processing Coal
• Coal gasification - converting coal into gaseous fuel
– Treating coal with air and steam at high temperature produces
mixture of CO, H2, and CH4. Some CO2 and SO2 are also formed
– Mixture of CO and H2 is also called syngas
• Reactions in coal gasification:
– C(s) + O2(g)  CO2(g);
DH = -394 kJ
– C(s) + ½O2(g)  CO(g);
DH = -111 kJ
– C(s) + H2O(g)  CO(g) + H2(g); DH = 131 kJ
– C(s) + 2H2(g)  CH4(g);
DH = -75 kJ
• Note: both exothermic and endothermic reactions occur. An energy balance
can be maintained by controlling the temperature, the rate of coal feed, and
the flow of air and steam.
Coal Gasification Process
• Coal + Steam + Air
Heat
• CH4 + CO, CO2, H2, H2O
(+ sulfur containing impurities)
Separate
Treatment
• CO + H2O  CO2 + H2
• CO + 3H2  CH4 + H2O
• After removal of CO2 and H2O, the remaining mixture contains CH4
and syngas (CO + H2)
Coal to Coal Slurry
• Coal is pulverized and mixed with water to form a
thick slurry
• Coal slurry burns like residual oil, with less CO and
SO2 produced
Hydrogen Fuel
• Combustion of H2:
– H (g) + ½O2(g)  H2O(l);
DH = -286 kJ
– H2(g) + ½O2(g)  H2O(g); DH = -242 kJ
• Combustion of H2 produces 2.5 times more
energy per gram than natural gas
• Combustion of H2 only produces water.
• However, production, storage and
transportation of the gas pose major problems
Problems in Hydrogen Production
• H2 does not exist in the free form (like N2 or O2)
• All existing methods to produce H2 gas involve endothermic
reactions or high temperature:
– Steam reformation:
• CH4(g) + H2O(g)  CO(g) + 3H2(g);
DH = 206 kJ
– Electrolysis: H2O(l)  H2(g) + ½O2(g);
DH = 286 kJ
– Thermochemical decomposition of H2O:
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2HI(g)  H2(g) + I2(s);
(425oC)
2H2O(l) + SO2(g) + I2(s)  H2SO4(aq) + 2HI(aq);
(90oC)
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H2SO4(aq)  SO2(g) + H2O(l) + ½O2(g); (825oC)
• Net Reaction: H2O(l)  H2(g) + ½O2(g);
Problems in Storage & Transportation of H2
• H2 decomposes to H-atoms on metal surface
• H-atoms are small enough to be absorbed into the metal
lattices and makes the metal to become brittle
• H2 requires a much larger container for storage and
transportation
• Storage or transportation of H2 under high pressure (say as
liquefied gas) poses high explosion hazard
• Energy produced per unit volume by H2 is only one-third that
produced by natural gas under similar conditions
• An alternative method suggested for H2 storage/transportation
is to convert it into solid metal hydrides, MH2.
Alternative Fuels
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Hydrogen gas is most efficient and clean fuel
It produces the most energy per gram,
Transportation and storage are difficult
It requires a much large fuel tank than gasoline;
May be compressed into liquid form, but poses explosion
hazard
• An alternative way is to convert it into a solid metal hydride
(MH2) – one that is able to release H2 when heated.
• Most likely the hydrogen will be used to power fuel cells that
will be installed in automobiles
Alternative Fuels
• Oil shale
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Shale rocks contain complex HC called kerogen
Huge deposits in western states, especially Colorado;
Kerogen is not fluid like petroleum - cannot be pumped;
Rocks containing fuel must be heated to >250oC to
decompose the kerogen into smaller HC molecules;
– Process produces large quantities of waste rocks – a
negative environmental impact.
Alternative Fuels
• Ethanol:
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Produced by fermentation of sugar (from sugar cane or corn)
Pure ethanol produces about 27 kJ/g of energy
Use as fuel supplement
Added to gasoline as gasohol (which contains ca. 10% ethanol)
Pure ethanol not suitable for motor fuel in USA, especially in
winter, because it does not vaporize easily in cold climate.
– Pure ethanol is widely used as motor fuel in Brazil where the
climate is warmer.
Alternative Fuels
• Methanol:
– Methanol produces about 20 kJ/g of energy
when completely combusted
– It has been used in race cars as a mixture of
85% methanol and 15% gasoline.
– California is evaluating methanol as motor
fuel
– Arizona and Colorado are considering similar
step.
New Energy?