Ch. 16 Reaction Energy
-
Upload
octavius-ferrell -
Category
Documents
-
view
37 -
download
0
description
Transcript of Ch. 16 Reaction Energy
Ch. 16
Reaction
Energy
Thermochemistry
• __________________: the study of the transfers of energy as heat that accompany chemical reactions and physical changes.
• ______________: an instrument to measure the energy absorbed or released as heat in a chemical or physical change.
• ______________: a measure of the average kinetic energy of the particles in a sample of matter. ________________
• ____________: the SI unit of heat as well as other forms of energy. 16-2
Thermochemistry• _______ = N x m = kg x m2
s2
• _______: the energy transferred between samples of matter because of a difference in their temperatures.
• _____________: the amount of energy required to raise the temperature of one gram of a substance by degree (C or K)
_________________________________ heat = specific x mass x change in
heat temp.
16-3
Example
q = cp x m x ΔT Example: A 4.0g sample of glass was heated from 274K to 314 K,
and was found to have absorbed 32 J of energy as heat. What is the specific heat of this glass, and how much energy would be gained with a temp. change of 314k to 344K?
________________________
cp = 0.20 J/gK
________________________)q = 24 J
16-4
Practice
q = cp x m x ΔT
1) Determine the specific heat of a material if a 35 g sample absorbed 96 J as it was heated from 293 K to 313 K.
16-5
Enthalpy• ________________: the amount of energy absorbed
by a system as heat during a process at constant pressure.
_________________________
• ________________________: the quantity of energy transferred as heat during a chemical rxn.
• ________________________: an equation that includes the quantity of energy released or absorbed as heat during the reaction.
16-6
Enthalpy
• ________________________: energy is released during the rxn.
• _____________________: energy is absorbed during the rxn.
• The quantity of energy ______________ is proportional to the amount of reactants.
16-7
Exothermic Rxns
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O + (-2043KJ)
16-8
Exothermic Rxns
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O + (-2043KJ)
In this reaction, energy is _______. ΔH is ___________because products have a
________ value for H than the reactants. ΔH is always negative for exothermic reactions
16-9
Endothermic Rxns
C(s) + H2O(g) + 113KJ → CO(g) + H2(g)
16-10
Endothermic Rxns
C(s) + H2O(g) + 113KJ → CO(g) + H2(g)
• In this reaction, energy is __________. ΔH is _________ because the products have a __________ value for H than the reactants. ΔH is always positive for endothermic reaction.
16-11
Enthalpy
• _____________________: the enthalpy change that occurs when one mole of a compound is formed from its elements in their standard state at STP. (standard temp and pressure, 0oC and 1 atm.)
• ______ = standard enthalpy of a rxn. • ______ = standard enthalpy of formation. (elements in
their standard state have ΔHof = 0, compounds with
positive values are unstable)
16-12
Enthalpy
• ______________________: the energy change that occurs during the complete combustion of one mole of a substance.
• ___________: the overall enthalpy change in rxn is equal to the sum of enthalpy changes for the individual steps in the process.
16-13
Hess’s Law
• Rules for applying Hess’s Law: 1) If you ________________________, you must
multiply the ΔH by the same coefficient:CH4(g) + 2O2(g) CO2(g) + 2H2O(g)
ΔH = -802 kJ
2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(g) ΔH = -1604 kJ
2) If an equation is ___________, the sign of ΔH is also ___________.
16-14
Hess’s Law
Ex. Calculate ΔHo for NO(g) + ½O2(g) → NO2(g) from th enthalpy data found in Appendix A-14. Solve by combining known eq.
Rxn1) ½N2(g) + ½O2(g) → NO(g) ΔHof = 90.29 kJ
Rxn2) ½N2(g) + O2(g) → NO2(g) ΔHof = 33.2 kJ
NO(g) + ½O2(g) → NO2(g)
________________________.
16-15
Hess’s Law
Rxn1) NO(g) → ½N2(g) + ½O2(g) ΔHof = - 90.29 kJ
(reversed)
Rxn2) ½N2(g) + O2(g) → NO2(g) ΔHof = 33.2 kJ
NO(g) + ½O2(g) → NO2(g)
ΔHo = (-90.29 kJ) + (33.2 kJ)
= -57.1 kJ 16-16
Practice2) Calculate the enthalpy of rxn for the combustion of methane,
CH4, to form CO2(g) and H2O(l).
16-17
Practice
16-18
PracticeExample: Calculate the enthalpy of formation of pentane, C5H12
5C(s) + 6H2(g) → C5H12(g) ΔHof = ?
Rnx1) C5H12(g) + 8O2(g) → 5CO2(g) + 6H2O(l) ΔHoc = -3535.6
kJ
Rxn2) C(s) + O2(g) → CO2(g) ΔHof = -393.5 kJ
Rxn3) H2(g) + ½O2(g) → H2O(l) ΔHof = -285.8 kJ
_____________________________________________________________________________________
16-19
Practice
16-20
Practice3) Calculate the ΔHo
f of butane, C4H10
16-21
Practice
16-22
Spontaneous Reactions• _____________ a measure of the degree of randomness of the
particles, such as molecules, in a system.
• ____________________ a combined enthalpy-entropy function.
• _______________________ the difference between the change in enthalpy and the product of the kelvin temp. and the entropy change.
________________________
ΔG = - SpontaneousΔG = + Not spontaneousΔG = 0 Equilibrium
16-23
ExampleExample: For the rxn NH4Cl(s)→NH3(g) + HCl(g), ΔHo = 176 kJ/mol
and ΔSo = 0.285 kJ/molK. Calculate ΔGo, is the rxn spontaneous at 298.15k?
ΔGo = ΔHo – TΔSo
= (176 kJ/mol) - (298.15k)(0.285 kJ/molK)
= 91 kJ/mol
16-24
Practice4) For the rxn Br2(l)→Br2g) ΔHo = 31 kJ/mol and ΔSo = 93 J/molK.
At what temp. will this rxn be spontaneous?
16-25
Ch. 16The End!