Transcript Chapter5_20130809103108

BIOENERGY. The sugars in sugarcane, produced from
CO2, H2O, and sunshine via photosynthesis, can be
converted into ethanol, which is used as an alternative to
gasoline. In certain climate, such as that in Brazil, the
sugarcane crop replenishes itself rapidly, making canebased ethanol a sustainable fuel source.
INTRODUCTION
 ENERGY AND CHEMISTRY
 Energy is necessary for all life.
 Plants, sunlight, and photosynthesis
 Human beings derive energy from plants and animals
 Energy changes accompany chemical reactions
 All of the energy we use are derived from chemical
reactions
 Thermodynamics is the study of energy and its
transformations
5.1 THE NATURE OF ENERGY
 DEFINITION OF ENERGY
 Energy is the capacity to do work or to transfer heat
 Work is the energy used to cause an object with mass to
move against a force
 Heat is the energy used to cause the temperature of an
object to increase
5.1 THE NATURE OF ENERGY
 KINETIC ENERGY AND POTENTIAL ENERGY
 Objects that have mass and motion possess kinetic energy
 Potential energy is the stored energy that arise from the
attractions and repulsions an object experiences in relation
to other objects
5.1 THE NATURE OF ENERGY
 UNITS OF ENERGY
 The SI unit of energy is the joule (J)
 An older, non-SI unit is still in widespread use:
the calorie (cal), the amount of energy required
to raise the temperature of 1 g of water by 1 °C
1 cal = 4.184 J
5.1 THE NATURE OF ENERGY
 SYSTEM AND SURROUNDINGS
• The system is the portion we single
out for study and typically includes
the molecules we want to study
(here, the hydrogen and oxygen
molecules).
• The surroundings are everything
else (here, the cylinder and piston).
• Open systems: matter & heat
exchangeable
• Closed systems: only heat
exchangeable
• Isolated systems: no exchange
5.1 THE NATURE OF ENERGY
 TRANSFERRING ENERGY: WORK AND HEAT
 Energy used to move
an object over some
distance is work.
 w=Fd
where w is work, F is
the force, and d is the
distance over which the
force is exerted.
5.1 THE NATURE OF ENERGY
 TRANSFERRING ENERGY: WORK AND HEAT
 Energy can also be
transferred as heat.
 Heat flows from
warmer objects to
cooler objects.
5.1 THE NATURE OF ENERGY
 TRANSFERRING ENERGY: WORK AND HEAT
5.2 THE FIRST LAW OF THERMODYNAMIC
 ENERGY IS CONSERVED
 Energy is neither created nor destroyed.
 In other words, the total energy of the universe is
a constant; if the system loses energy, it must be
gained by the surroundings, and vice versa.
5.2 THE FIRST LAW OF THERMODYNAMIC
 INTERNAL ENERGY
 The internal energy (E) of a
system is the sum of all kinetic
and potential energies of all
components of the system
E = Ek + Ep
5.2 THE FIRST LAW OF THERMODYNAMIC
 INTERNAL ENERGY
 By definition, the change in internal energy, E, is the
final energy of the system minus the initial energy of the
system:
E = Efinal − Einitial
5.2 THE FIRST LAW OF THERMODYNAMIC
 INTERNAL ENERGY
• If E > 0, Efinal > Einitial
– Therefore, the system
absorbed energy from
the surroundings.
5.2 THE FIRST LAW OF THERMODYNAMIC
 INTERNAL ENERGY
• If E < 0, Efinal < Einitial
– Therefore, the system
released energy to the
surroundings.
5.2 THE FIRST LAW OF THERMODYNAMIC
 RELATING E TO WORK AND HEAT
• When energy is
exchanged between
the system and the
surroundings, it is
exchanged as either
heat (q) or work (w).
• That is, E = q + w.
5.2 THE FIRST LAW OF THERMODYNAMIC
 RELATING E TO WORK AND HEAT
5.2 THE FIRST LAW OF THERMODYNAMIC
 ENDOTHERMIC AND EXOTHERMIC PROCESSES
5.2 THE FIRST LAW OF THERMODYNAMIC
 STATE FUNCTION
 Consider chemical changes of H2O:
 The value of a state function (internal energy, E, in
this case) depends only on the present state of the
system, not on the path the system took to reach
that state
5.2 THE FIRST LAW OF THERMODYNAMIC
 STATE FUNCTION
• However, q and w are not
state functions.
• Whether the battery is
shorted out or is
discharged by running the
fan, its E is the same.
– But q and w are different
in the two cases.
5.3 ENTHALPY
 ENTHALPY
• Consider a chemical reaction:
w = -PV
5.3 ENTHALPY
 ENTHALPY
 Enthalpy is a measure of the total energy of a system
 Enthalpy accounts for heat flow in processes occurring at
constant pressure when no forms of work are performed
other than P-V work
 Enthalpy (H) is expressed by:
H = E + PV
(state function)
 When the system changes at constant pressure, the
change in enthalpy, H, is
H = (E + PV) = E + PV
= (qp + w) − w = qp
(state function?)
 H equals the heat gained or lost at constant P
5.3 ENTHALPY
 ENTHALPY
 When H is positive:
endothermic process
 When H is negative:
exothermic process
 At constant V, the change in E is equal to the heat
gained or lost
 At constant P, the change in H is equal to the heat
gained or lost
 The difference between E and H is the amount of P-V
work done by the system when the process occurs at
constant P, −PV
 In many reactions, V is close to zero, which makes
PV (the difference between E and H) small
 Generally H can be used to measure energy changes
during most chemical processes
5.3 ENTHALPY
(a) H > 0, endothermic
(b) H < 0, exothermic
5.4 ENTHALPIES OF REACTION
 ENTHALPY OF REACTION
thermochemical equation
enthalpy diagram
5.4 ENTHALPIES OF REACTION
This tragedy, in Lakehurst, New Jersey, on May 6,
1937, led to the discontinuation of hydrogen as a
buoyant gas in such craft
5.4 ENTHALPIES OF REACTION
 ENTHALPY OF REACTION
 Enthalpy is an extensive property.
for 1 mol of CH4
 H for a reaction in the forward direction is equal
in size, but opposite in sign, to H for the reverse
reaction.
 H for a reaction depends on the state of the
products and the state of the reactants.
5.4 ENTHALPIES OF REACTION
5.4 ENTHALPIES OF REACTION
5.5 CALORIMETRY
 HEAT CAPACITY AND SPECIFIC HEAT
 Heat capacity, C
• The amount of heat required to raise temperature
by 1 K (˚C)
 Molar heat capacity, Cm
• The heat capacity of one mole of a substance
 Specific heat, Cs
• The heat capacity of one gram of a substance
5.5 CALORIMETRY
 HEAT CAPACITY AND SPECIFIC HEAT
5.5 CALORIMETRY
Sample Exercise 5.5
Relating Heat, Temperature Change, and Heat Capacity
(a) How much heat is needed to warm 250 g of water (about 1 cup)
from 22 ˚C (about room temperature) to near its boiling point, 98 ˚C?
The specific heat of water is 4.18 J/g-K.
(b) What is the molar heat capacity of water?
5.5 CALORIMETRY
 “Are you running a fever?”
 Difficult to maintain T
 Maintaining T in our body
• High heat capacity of water
• Optimal for muscle function and
biochemical rxns in our body at 35.8-37.2 ˚C
• Hypothalamus controls body T
• Radiation, convection, and evaporation
(humidity and water replenishment)
• Shivering and reddish skin
5.5 CALORIMETRY
 CONSTANT-PRESSURE CALORIMETRY
 Calorimetry is the
measurement of heat flow
 Consider a system:
• A reaction in a coffee-cup
• System: reactants and
products
• Surrounding: water and
calorimeter
5.5 CALORIMETRY
Sample Exercise 5.6
Measuring ΔH Using a Coffee-Cup Calorimeter
When a student mixes 50 mL of 1.0 M HCl and 50 mL of 1.0 M NaOH in a
coffee-cup calorimeter, the temperature of the resultant solution increases
from 21.0 ˚C to 27.5 ˚C. Calculate the enthalpy change for the reaction in
kJ/mol HCl, assuming that the calorimeter loses only a negligible quantity
of heat, that the total volume of the solution is 100 mL, that its density is
1.0 g/mL, and that its specific heat is 4.18 J/g-K.
Hmol = -2.7 kJ/0.050 mol = -54 kJ/mol
5.5 CALORIMETRY
 CONSTANT-VOLUME CALORIMETRY
 Bomb calorimeter
• Designed for measuring T
by combustion reactions
• Inlet valve: O2 supply
• Electrical contact: initiation
• Surrounding: water and
calorimeter
• The heat capacity of the
calorimeter is measured by
combusting a standard
sample
Sample Exercise 5.7
Measuring qrxn Using a Bomb Calorimeter
5.5 CALORIMETRY
Methylhydrazine (CH6N2, MW 46.1 g/mol) is used as a liquid rocket fuel. The
combustion of methylhydrazine with oxygen produces N2(g), CO2(g), and H2O(l):
2 CH6N2(l) + 5 O2(g) → 2 N2(g) + 2 CO2(g) + 6 H2O(l)
When 4.00 g of methylhydrazine is combusted in a bomb calorimeter, the
temperature of the calorimeter increases from 25.00 ˚C to 39.50 ˚C. In a separate
experiment the heat capacity of the calorimeter is measured to be 7.794 kJ/ ˚C.
Calculate the heat of reaction for the combustion of a mole of CH6N2.
5.6 HESS’S LAW
 COMBUSTION OF CH4(g)
 Consider the combustion reaction of CH4(g)
 We can think of the reaction as a two-step
process
5.6 HESS’S LAW
 HESS’S LAW
 If a reaction is carried out in a series of steps, H
for the overall reaction will be equal to the sum of
the enthalpy changes for the individual steps
 Based on the fact that
enthalpy is a state
function
 Useful for calculating
H that are difficult to
measure directly
(Ex. Carbon to CO)
5.6 HESS’S LAW
5.6 HESS’S LAW
5.6 HESS’S LAW
5.6 HESS’S LAW
5.7 ENTHALPIES OF FORMATION
 An enthalpy of formation, Hf, is defined as
the enthalpy change for the formation of a
compound from its constituent elements
 The conditions of T, P, and state (aq, s, l,…)
should be defined.
 Standard enthalpy change, H°
• H when all reactants and products are in their
standard state (1 atm, 273 K or 25 ˚C)
5.7 ENTHALPIES OF FORMATION
 STANDARD ENTHALPY OF FORMATION
 The standard enthalpy of formation of a compound, H°f,
is the enthalpy change for the reaction that forms one
mole of the compound from its elements in their standard
states:
 By definition, H°f of the most stable form of any element
is zero
The enthalpy change of the reaction
5.7 ENTHALPIES OF FORMATION
 STANDARD ENTHALPY OF FORMATION
5.7 ENTHALPIES OF FORMATION
 CALCULATION OF REACTION ENTHALPIES
5.7 ENTHALPIES OF FORMATION
 CALCULATION OF REACTION ENTHALPIES
5.7 ENTHALPIES OF FORMATION
1 mol
5.7 ENTHALPIES OF FORMATION
Calculate H°f of CaCO3(s).
Calculate H°f of CuO(s).
5.8 FOODS AND FUELS
 FOODS
 Fuel value: the energy released when 1 g of a material is
combusted
 The fuel value of carbohydrates: 17 kJ/g (4 kcal/g)
 The fuel value of fats: 38 kJ/g (9 kcal/g)
 The fuel value of proteins: 17 kJ/g (4 kcal/g)
5.8 FOODS AND FUELS
 FOODS
5.8 FOODS AND FUELS
 FUELS
EX. 5.106
The hydrocarbons acetylene (C2H2) and benzene (C6H6) have the same
empirical formula. Benzene is an “aromatic” hydrocarbon, one that is
unusually stable because of its structure.
(a) By using the data in Appendix C, determine the standard enthalpy
change for the reaction 3C2H2(g) → C6H6(l).
(b) Which has greater enthalpy, 3 mol of acetylene gas or 1 mol of liquid
benzene?
(c) Determine the fuel value in kJ/g for acetylene and benzene.
(a) H° = -631.3 kJ (b) 3 mol acetylene
(c) 50 kJ/g (acetylene) and 42 kJ/c (benzene)
EX. 5.100
How many grams of methane CH4(g) must be combusted to heat 1.00 kg
of water from 25 ˚C to 90 ˚C, assuming H2O(l) as a product and 100%
efficiency in heat transfer?
4.90 g of CH4(g)