1. Ch 17 Thermochemistry

2. What is the difference between heat and temperature?
HEAT is energy that transfers from one object/substance to another
TEMPERATURE is a measure of the amount of energy an object/substance has

3. A. Energy Transformations
Energy that is stored in the chemical bonds of a substance is called CHEMICAL POTENTIAL ENERGY
Heat ALWAYS flows from hot to cold

4. B. Endothermic and Exothermic Processes
System=the reaction
Surroundings=everything around the reaction
Surroundings
Law of Conservation of Energy- energy can be neither created nor destroyed
System

5. B. Endothermic and Exothermic Processes
Endothermic Process- heat is absorbed from the surroundings
Endo = Into
Endothermic processes are represented by a positive “q”
HEAT

6. B. Endothermic and Exothermic Processes
Exothermic process- heat is released into the surroundings
Exo = Exit
Exothermic processes are represented by a negative “q”
HEAT

7. C. Measuring Heat Flow Two Common Units 1J = 0.2390 cal
Joule
calorie
1J = cal
4.184 J = 1 cal
1Calorie = 1 kilocal = 1000 cal

8. D. Heat Capacity and Specific Heat
Heat Capacity depends on:
The mass of the object
The chemical composition of the object
“the amount of heat needed to increase the temperature of an object by 1 oC
Specific heat capacity- amount of heat needed to raise the temperature of 1g by 1 oC

9. D. Heat Capacity and Specific Heat
C = q / (m X ΔT)
C =Specific Heat
q = heat (joules or calories)
m = mass (grams)
ΔT = change in temperature
The change in temperature can be measure in Kelvin or degrees Celsius

10. C= 849J/ (95.4g)(23.0 oC) = 0.387 J/(g· oC)
The temperature of a 95.4g piece of copper increases from 25.0 oC to 48.0 oC when it absorbs 849 J of heat. What is the specific heat of copper?
Known:
m= 95.4g
q=849 J
ΔT= =23.0 oC
Work:
C= 849J/ (95.4g)(23.0 oC)
= J/(g· oC)

11. E. Calorimetry Measures the heat flow into or out of a system
Heat released by the system is equal to heat absorbed by the surroundings
ENTHALPY: (H) the heat constant of a system at constant pressure

12. E. Calorimetry The terms heat and enthalpy change are interchangeable
q = ΔH
qsys = ΔH = -m x C x ΔT
Negative enthalpy = exothermic
Positive enthalpy = endothermic

13. When 25. 0mL of water containing HCl at 25. 0 oC is added to 25
When 25.0mL of water containing HCl at 25.0 oC is added to 25.0mL of water containing NaOH at 25.0 oC in a calorimeter a rxn occurs. Calculate the enthalpy change (in kJ) during the rxn if the highest temperature observed was 32.0 oC. Assume all densities =1g/mL KNOWN: Cwater=4.18J/g oC V=25.0mL+25.0mL ΔT=7 oC Density= 1g/mL ΔH = ?

14. m= (50mL) x (1. 00g/mL) = 50g ΔT= TF – Ti = 32. 0 – 25. 0 =7
m= (50mL) x (1.00g/mL) = 50g ΔT= TF – Ti = 32.0 – 25.0 =7.0 ΔH= -mCΔT= -(50.0g)(4.18J/goC)(7.0oC) ΔH= J = -1500J = -1.5kJ

15. F. Thermochemical Equations
In a thermochemical equation, the enthalpy of change for the reaction can be written as either a reactant or a product
Endothermic (positive ΔH)
2NaHCO kJ Na2CO3 + H2O + CO2
Exothermic (negative ΔH)
CaO + H2O Ca(OH) kJ

16. G. Heat of Combustion
The heat of reaction for the complete burning of one mole of a substance
Written the same was as change in enthalpy

17. Write the thermochemical equation for the oxidation of Iron (III) if its ΔH= kJ Fe(s) + O2(g)→ Fe2O3(s) How much heat is evolved when 10.00g of Iron is reacted with excess oxygen? 2 4 3
10.00g Fe
1 mol
= mol Fe
55.85g Fe
mol Fe
1652 kJ
=73.97 kJ of heat
4 mol Fe