Download Thermochemistry PPT

Thermochemistry
Definitions
• Energy – capacity for doing work or
supplying heat.
• Thermochemistry – study of energy
changes that occur during phase changes
and chem. rxns.
• Chem. Potential Energy – energy stored in
chemical bonds.
Example
Lots of energy
stored in bonds!
Energy difference.
5473 kJ/mol
Little energy stored
in bonds.
Heat
• Represented by q.
• Energy that transfers from one object to
another because of a Temp. difference
between them.
• Heat flows from warm  cool until the
two objects are at the same Temp.
Exothermic vs. Endothermic
• In exothermic processes, the system loses
heat as its surroundings warm up.
– q has a negative value b/c the system is losing
heat.
• In endothermic processes, the system gains
heat as its surroundings cool down.
– q has a positive value b/c the system is gaining
heat.
Potential Energy 
Potential Energy Diagram of Ice
Melting at 0ºC.
Water
Ice
Time 
Is the melting of ice
an endothermic or
an exothermic
process? How can
you tell?
Measuring Heat Flow
• SI Unit of heat flow: Joule (J)
• Common unit used in chemistry: calorie
(cal)
– Amt. of heat needed to raise 1 gram of water
by 1ºC.
– 1 cal = 4.184 J
• Food Calorie (capital “C”) = 1000 cal, or 1
kilocalorie = 4184 J
Heat Capacity
• Amount of heat needed to raise an object’s
temperature by 1°C.
– Depends on the chemical composition and the
mass of the object.
• EXAMPLE: 1 gram of water requires 1
cal to raise its temperature by 1°C.
– 100. g of water require 100. cal to raise the
temp. by 1°C.
Heat Capacity
Same temperature change
10 g H2O
1 g H2O
Specific Heat (c)
• Amt. of heat needed to raise 1 gram of a
substance’s temperature by 1ºC.
– Expressed in J/g ºC, or cal/g ºC
• The higher a substance’s specific heat, the
more energy it takes to heat it.
• Substance’s with low specific heats heat up
and cool down quickly (most metals, e.g.)
Some Specific Heats
Substance
Water
Grain alcohol
Ice
Steam
Chloroform
Aluminum
Iron
Silver
Mercury
Specific Heat
J/gºC
cal/g ºC
4.18
1.00
2.4
0.58
2.1
0.50
1.7
0.40
0.96
0.23
0.90
0.21
0.46
0.11
0.24
0.057
0.14
0.033
Specific Heat (c)
• c = heat / (mass x change in Temp.)
• c = q / (m x ΔT)
• q = m x c x ΔT
Example Problem
• The temperature of a 95.4-g piece of Cu
increases from 25.0ºC to 48.0ºC when the
Cu absorbs 849 J of heat. What is the
specific heat of Cu?
– SOLUTION: q = m x c x ΔT
•
•
•
•
849 J = (95.4 g) c (48.0ºC – 25.0ºC)
849 J = (95.4 g) c (23.0ºC)
849 J = (2190 gºC) c
c = 0.388 J/gºC
• Based on what you know about metals,
does this answer make sense?
Example Problem
• When 435 J of heat is added to 3.4 g of
olive oil at 21ºC, the temperature increases
to 85ºC. What is the specific heat of olive
oil?
– SOLUTION: q = m x c x ΔT
•
•
•
•
435 J = (3.4 g) c (85ºC – 21ºC)
435 J = (3.4 g) c (64ºC)
435 J = (220 gºC) c
c = 2.0 J/gºC
Example Problem
• How much heat is required to raise the
temperature of 250.0 g of mercury by
52ºC? The specific heat of mercury is 0.14
J/gºC.
– SOLUTION: q = m x c x ΔT
– q = (250.0 g)(0.14 J/gºC)(52ºC)
– q = 1800 J = 1.8 kJ
Enthalpy Changes
• Enthalpy (H) – the heat content of a system
at constant pressure.
• Enthalpy change (ΔH) – the heat that
enters or leaves a system at constant
pressure.
• q = ΔH
• Neg. ΔH = exothermic process
• Pos. ΔH = endothermic process
Thermochemical Equations
• Enthalpy change can be written as a
reactant or a product.
– Reactant  endothermic
– Product  exothermic
• Example: The reaction of calcium oxide
with water is exothermic.
– It produces 65.2 kJ of heat per mole of CaO
reacted.
– CaO(s) + H2O(l)  Ca(OH)2(s) + 62.5 kJ