Download unit 6 power point File - District 196 e

Chapter 6
Thermochemistry: pp. 235-274
6.1
The Nature of Energy
• Energy
– Capacity to do work or produce heat.
– 1st Law of Thermodynamics: Energy can not be created
or destroyed, but converted to different forms.
– The energy of the universe is constant.
– Potential Energy: Energy due to position or composition
– Kinetic Energy: Energy due to motion; depends on mass
and velocity
6.1
The Nature of Energy
Lower Potential
Energy
Add Energy
Higher Potential
Energy
6.1
The Nature of Energy
• Energy
• Heat involves the transfer of energy between two objects due
to a temperature difference.
• Work: Force acting over a distance
• Energy is a state function: Work and Heat are not
– State Function: Property that does not depend in any
way on the system’s past or future (only depends on
present state).
6.1
The Nature of Energy
• Chemical Energy
• System – part of the universe where a reaction is happening.
• Surroundings – everything else in the universe
– Endothermic Reaction
• Energy flows into the system, absorbs energy from
surroundings.
– Exothermic Reaction
• Energy flows out of the system, surroundings gain energy.
– Energy gained by surroundings must be equal to the
energy lost by the system (opposite is true also).
6.1
The Nature of Energy
• Reality Check
– Is the freezing of water an exothermic or endothermic
process?
• Why?
– Define the system, surroundings, and direction of energy
transfer:
• Methane is burning in a Bunsen burner in a laboratory.
6.1
The Nature of Energy
• Internal Energy
– Internal energy, E, of a system is the sum of the kinetic
and potential energies of all the “particles” in the
system.
– A change to the internal energy of a system is described
as:
E = q + w
• Where q is heat and w is work.
6.1
The Nature of Energy
• Internal Energy
– Sign reflects the system’s point of view.
q
+
+
-
w
+
+
-
What we observe
Endothermic, surroundings do work
Endothermic, system does work
Exothermic, surroundings do work
Exothermic, system does work
6.1
The Nature of Energy
• Work
– Work = P × A × Δh = PΔV




P is pressure.
A is area.
Δh is the piston moving a distance.
ΔV is the change in volume.
6.1
The Nature of Energy
• Work
– For an expanding gas, ΔV is a positive quantity because
the volume is increasing. Thus ΔV and w must have
opposite signs:
w = - PV
6.1
The Nature of Energy
• Reality Check
– Determine the sign of E for each of the following with
listed conditions
• An endothermic process that performs work.
– l work l > heat
– l work l < heat
Δ E = negative
Δ E = positive
• Works is done on a gas and the process is exothermic.
– l work l > heat
– l work l < heat
Δ E = positive
Δ E = negative
6.2
Enthalpy and Calorimetry
• Change in Enthalpy
– State Function
– ΔH = q at constant pressure
– ΔH = Hproducts - Hreactants
6.2
Enthalpy and Calorimetry
• Reality Check
Consider the combustion of propane:
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l)
ΔH = –2221 kJ
Assume that all of the heat comes from the
combustion of propane. Calculate ΔH in which
5.00 g of propane is burned in excess oxygen at
constant pressure.
-252 kJ
6.2
Enthalpy and Calorimetry
• Calorimetry
– Science of measuring heat
– Specific heat capacity:
• The energy required to raise the temperature of one gram of a
substance by one degree Celsius.
– Molar heat capacity:
• The energy required to raise the temperature of one mole of
substance by one degree Celsius.
6.2
Enthalpy and Calorimetry
• Simple Calorimeter
• If two reactants are
mixed and the solution
gets warmer an
exothermic reaction is
happening.
• An endothermic
reaction cools the
solution.

6.2
Enthalpy and Calorimetry
• Calorimetry
Heat = s  m  T
– Where,
• s = specific heat capacity (J/oCg)
• m = mass (g)
• ΔT = change in temperature (oC)
6.2
Enthalpy and Calorimetry
• Reality Check
A 100.0 g sample of water at 90°C is added to
a 100.0 g sample of water at 10°C.
The final temperature of the water is:
a) Between 50°C and 90°C
b) 50°C
c) Between 10°C and 50°C
6.2
Enthalpy and Calorimetry
• Reality Check
You have a Styrofoam cup with 50.0 g of water at 10.C. You
add a 50.0 g iron ball at 90.C to the water. (sH2O = 4.18
J/°C·g and sFe = 0.45 J/°C·g)
The final temperature of the water is:
a) Between 50°C and 90°C
b) 50°C
c) Between 10°C and 50°C
Calculate the final temperature of the water.
18oC
6.3
Hess’ Law
• Hess’ Law
– In going from a particular set of reactants to a
particular set of products, the change in enthalpy
is the same whether the reaction takes place in
one step or in a series of steps.
6.3
Hess’ Law
N2(g) + 2O2(g) → 2NO2(g)
•
ΔH1 = 68 kJ
This reaction also can be carried out in two distinct steps,
with enthalpy changes designated by ΔH2 and ΔH3.
N2(g) + O2(g) → 2NO(g)
2NO(g) + O2(g) → 2NO2(g)
N2(g) + 2O2(g) → 2NO2(g)
ΔH1 = ΔH2 + ΔH3 = 68 kJ
ΔH2 = 180 kJ
ΔH3 = – 112 kJ
ΔH2 + ΔH3 = 68 kJ
6.3
Hess’ Law
• Principle of Hess’ Law
6.3
Hess’ Law
Characteristics of Hess’ Law
•
•
If a reaction is reversed, the sign of ΔH is also
reversed.
The magnitude of ΔH is directly proportional to the
quantities of reactants and products in a reaction. If
the coefficients in a balanced reaction are
multiplied by an integer, the value of ΔH is
multiplied by the same integer.
6.3
Hess’ Law
• Reality Check
– Calculate the ΔH for the reaction
N2H4 (l) + O2 (g)  N2 (g) + 2H2O (l)
2NH3 (g) + 3N2O(g)  4N2 (g) + 3H2O(l)
ΔH = -1010. kJ
N2O(g) + 3H2 (g)  N2H4 (l) + H2O(l)
ΔH = -317 kJ
2NH3 (g) + ½ O2 (g)  N2H4 (l) + H2O(l)
ΔH = -143 kJ
H2 (g) + ½ O2 (g)  H2O (l)
ΔH = -286 kJ
6.4
Standard Enthalpy of Formation
Standard Enthalpy of Formation (ΔHof)
• Change in enthalpy that accompanies the formation
of one mole of a compound from its elements with
all substances in their standard states.
6.4
Standard Enthalpy of Formation
Conventional Definitions of Standard States
• For a Compound
 For a gas, pressure is exactly 1 atm.
 For a solution, concentration is exactly 1 M.
 Pure substance (liquid or solid)
• For an Element
 The form [N2(g), K(s)] in which it exists at 1 atm
and 25°C.
 Heat of formation is zero.
6.4
Standard Enthalpy of Formation
• Calculating Enthalpy of formation
H°rxn = npHf(products) - nrHf(reactants)
• Remember,
– Elements enthalpy of formation is zero.
– Don’t forget to multiply by coefficients from the
balanced equation.
6.4
Standard Enthalpy of Formation
Reality Check
Calculate H° for the following reaction:
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Given the following information:
Hf° (kJ/mol)
Na(s)
0
H2O(l)
–286
NaOH(aq)
–470
H2(g)
0
H° = –368 kJ