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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.

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slide1
Ch. 16

Reaction

Energy

thermochemistry
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
thermochemistry1
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
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
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
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

enthalpy1
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
Exothermic Rxns

C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O + (-2043KJ)

16-8

exothermic rxns1
Exothermic Rxns

C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O + (-2043KJ)

In this reaction, energy is _______. ΔH is ___________becauseproducts have a ________ value for H than the reactants. ΔH is always negative for exothermic reactions

16-9

endothermic rxns
Endothermic Rxns

C(s) + H2O(g) + 113KJ → CO(g) + H2(g)

16-10

endothermic rxns1
Endothermic Rxns

C(s) + H2O(g) + 113KJ → CO(g) + H2(g)

  • In this reaction, energy is __________.ΔHis _________ because the products have a __________ value for H than the reactants. ΔH is always positive for endothermic reaction.

16-11

enthalpy2
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

enthalpy3
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
Hess’s Law
  • Rules for applying Hess’s Law:

1) If you ________________________, you must multiply the ΔHby 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 ΔHis also ___________.

16-14

hess s law1
Hess’s Law

Ex. CalculateΔHofor 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 law2
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

practice1
Practice

2) Calculate the enthalpy of rxn for the combustion of methane, CH4, to form CO2(g) and H2O(l).

16-17

practice3
Practice

Example: 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

practice5
Practice

3) Calculate the ΔHofof butane, C4H10

16-21

spontaneous reactions
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 = - Spontaneous

ΔG = + Not spontaneous

ΔG = 0 Equilibrium

16-23

example1
Example

Example: 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

practice7
Practice

4) 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