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!