Transcript Chapter 6 PPT

Thermochemistry
AP Chemistry
Seneca Valley SHS
Chapter 6
6.1 The Nature of Energy
Kinetic and Potential Energy
• Energy: the capacity to do work or
produce heat.
• Kinetic energy: energy an object has due
to its motion
• Potential energy: energy due to position
or composition
2
1
E k  mv
2
6.1 The Nature of Energy
Kinetic and Potential Energy
Potential energy can be converted
into kinetic energy.
Example: a ball of clay dropping
off a building.
6.1 The Nature of Energy
Temperature: the average of energy of
particles in motion.
temp.  heat
Heat: the transfer of energy from a hotter
object to a colder one.
State Function: property of a system that
depends only on its present condition
6.1 The Nature of Energy
Systems and Surroundings
System: part of the universe we
are interested in.
Surroundings: the rest of the
universe.
Exothermic: reaction that
involves the release of heat
to the surroundings
Endothermic: reaction that
absorbs energy from the
surroundings
6.1 The Nature of Energy
Thermodynamics
• Def: the study of energy and its interconversions
First Law of Thermodynamics (Law of Conservation of Energy)
Def: Energy cannot be created nor destroyed in ordinary
chemical reactions
6.1 First Law of Thermodynamics
Internal Energy
• Internal Energy:
total energy of a
system.
• Cannot measure
absolute internal
energy.
• Change in internal
energy, DE = Efinal Einitial
6.1 First Law of Thermodynamics
Relating DE to Heat and Work
•Energy cannot be created or destroyed.
•Energy of (system + surroundings) is constant.
•when a system undergoes a physical or chemical
change, the change in internal energy is given by:
the heat added to or absorbed by the system plus the
work done on or by the system:
DE = q + w
6.1 First Law of Thermodynamics
Relating DE to Heat and Work
Thermodynamic
quantities are measured
from the perspective of
the system
Example: in an
endothermic process
where heat is flowing
into the system, q would
be positive (+)
6.2 Enthalpy
Enthalpy: Heat transferred between the system and
surroundings carried out under constant pressure.
One can only measure the change(D in enthalpy (H):
DH = Hfinal - Hinitial = qP
6.2 Enthalpy
For any reaction: DHrxn = H(products) - H (reactants)
Enthalpy is an extensive property (magnitude DH is
directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
DH = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH = -1604 kJ
When we reverse a reaction, we change the sign of DH:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g)
DH = +802 kJ
Change in enthalpy depends on state:
H2O(g)  H2O(l) DH = -88 kJ
6.2 Calorimetry
Heat Capacity and Specific Heat
Calorimetry = science of measuring of heat flow.
Calorimeter = device that measures heat flow.
Heat capacity = the amount of energy required to raise
the temperature of an object.
Molar heat capacity = heat capacity of 1 mol of a
substance.
Specific heat capacity = heat capacity of 1 g of a subst.
q = (specific heat)  (grams of substance)  T.
6.2 Calorimetry
Constant-Pressure
Calorimetry
Atmospheric pressure is constant!
DH = qP
qrxn = -qsoln =
-(specific heat of solution)  (grams)
of solution)  DT.
6.2 Calorimetry
Constant-Pressure Calorimetry
Problem 1: When a student mixes 50.0 mL of 12.0 M HCl
and 50.0 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 per mole
of water formed, assuming that the calorimeter loses only
a negligible quantity of heat, that the total volume of the
solution is 100.0 mL, the density of the solution is 1.00
g/mL, and its specific heat is 4.18 J/gC.
6.2 Calorimetry
Bomb Calorimetry (Constant-Volume Calorimetry)
The reaction is carried
out under constant vol.
Usually used to study
combustion.
qrxn = -CcalorimeterT.
6.3 Hess’s Law
• Hess’s law: if a reaction is carried out in a number of
steps, DH for the overall reaction is equal to the sum of
DH’s from each individual step.
• For example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
DH = -802 kJ
2H2O(g)  2H2O(l)
DH = -88 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
DH = -890 kJ
6.3 Hess’s Law
Notes on Hess’s Law
1. If a reaction is reversed in order to solve a Hess’s Law problem,
then the sign of the DH of the reaction is also reversed.
2. The magnitude of DH is directly proportional to the quantities
of reactants and products. If the coefficients in an equation are
multiplied by an integer, the DH is also multiplied by that
integer
3. Hess’s Law can be used to calculate the DH in a reaction;
however, the component equations do not necessarily tell us how
the overall reaction occurred.
6.3 Hess’s Law
In the above enthalpy diagram note that:
DH1 = DH2 + DH3
6.3 Hess’s Law
Example (ex. 6.8 pg. 258):
Diborane (B2H6) is a highly reactive boron hydride.
Calculate DH for the synthesis of diborane from its
elements, according to the following equation:
2 B(s) + 3 H2 (g) --> B2H6 (g)
Use the following data to help you:
a)
b)
c)
d)
2 B(s) + 3/2 O2(g) --> B2O3 (s)
B2H6(g) + 3 O2(g) --> B2O3(s) + 3 H2O(g)
H2 (g) + 1/2 O2(g) --> H2O (l)
H2O(l) --> H2O(g)
DH = -1273 kJ
DH = -2035 kJ
DH = -286 kJ
DH =
44 kJ
6.4 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, DHof .
• Degree symbol in DHof indicates standard conditions…
• Standard conditions 1 atm and 25 oC (298 K).
• If there is more than one state for a substance under
standard conditions, the more stable one is used.
• Learn the distinction between standard state and
standard conditions---they are not the same!
6.4 Enthalpies of Formation
• Standard enthalpy of formation of the most stable
form of an element is zero.
6.4 Enthalpies of Formation
Using Enthalpies of Formation to Calculate
Enthalpies of Reaction
For a reaction:
DH f   nDH f (products)   mDHfo (reactants)
6.4 Enthalpies of Formation
Using Enthalpies of Formation to Calculate
Enthalpies of Reaction
Example 1: (Ex. 6.9)Use the standard enthalpies of
formation listed in Table 6.2 to calculate the standard
enthalpy change for the overall reaction that occurs
when ammonia is burned in air to form nitrogen dioxide
and water. This is the first step in the manufacture of
nitric acid.
6.4 Enthalpies of Formation
Using Enthalpies of Formation to Calculate
Enthalpies of Reaction
Example 2: (Ex. 6.10) Using enthalpies of formation,
calculate the standard change in enthalpy for the
thermite reaction, which occurs when a mixture of
powdered aluminum and iron (III) oxide is ignited with
a magnesium fuse:
2 Al(s) + Fe2O3(s) --> Al2O3 (s) + 2 Fe(s)
To save time we will not formally cover 6.5-6.6 in
class (thus you won’t be tested on it!)
Thermochemistry
End of Chapter 6