Contributed by:
The highlights are:
1. Definitions and Concepts
2. Heat Transfer
3. Calorimetry
4. The first law keeps track of the heat transfer
1.
Heat Transfer
and
Calorimetry
Dr. Keith Baessler
2.
Definitions and Concepts
• Energy • Specific Heat
• Law of Conservation of • Heat Capacity
Energy • Heat Transfer
• System vs Surroundings • Sign Conventions of q
• Types of Systems
• Calorimetry
• Heat vs Temperature • Calorimeters
• Endothermic vs • Formula Applications
Exothermic
2
3.
Energy is the capacity to do work.
• Radiant energy comes from the sun and is
earth’s primary energy source
• Thermal energy is the energy associated with
the random motion of atoms and molecules
• Chemical energy is the energy stored within the
bonds of chemical substances
• Nuclear energy is the energy stored within the
collection of neutrons and protons in the atom
• Potential energy is the energy available by virtue
of an object’s position
3
4.
First law of thermodynamics – energy can be
converted from one form to another, but cannot be
created or destroyed.
C3H8 + 5O2 3CO2 + 4H2O
Exothermic chemical reaction!
Chemical energy lost by combustion = Energy gained by the surroundings
system surroundings
4
5.
Heat versus Temperature
Heat is the transfer of thermal energy between
two bodies that are at different temperatures.
Temperature is a measure of the
thermal energy.
Temperature = Thermal Energy
5
6.
Exothermic process is any process that gives off heat –
transfers thermal energy from the system to the surroundings.
2H2 (g) + O2 (g) 2H2O (l) + energy
H2O (g) H2O (l) + energy
Endothermic process is any process in which heat has to be
supplied to the system from the surroundings.
energy + 2HgO (s) 2Hg (l) + O2 (g)
energy + H2O (s) H2O (l)
6
7.
The system is the specific part of the universe that is of
interest in the study.
open closed isolated
Exchange: mass & energy energy nothing 7
8.
Heat Transfer, q
• Heat (q) = the transfer of energy which causes the temperature of
an object to change
• units: joules (j), calories (cal)
– A calorie is the amount of energy needed to raise the
temperature of 1.00 g water by 1°C.
1cal = 4.184 joules
• Heat spontaneously moves regions of high temperature to regions
of lower temperature.
• A metal spoon at 25°C is placed in boiling water. What happens?
8
9.
The Sign Convention of q
•Endothermic systems require the surroundings to add
energy to the system.
q is positive (+)
•Exothermic reactions release energy to the surroundings.
q is negative (-)
•Energy changes are measured from the point of view of the
9
10.
The specific heat (s) of a substance is the amount of heat (q)
required to raise the temperature of one gram of the
substance by one degree Celsius.
The heat capacity** (C) of a substance is the amount of heat
(q) required to raise the temperature of a given quantity (m)
of the substance by one degree Celsius.
C=mxs
Heat (q) absorbed or released:
q = m x s x t
q = C x t
t = tfinal - tinitial
10
**Heat Capacity (C) is sometimes called the calorimeter constant
11.
How much heat is given off when an 869 g iron bar cools
from 94oC to 5oC?
s of Fe = 0.444 J/g • oC
t = tfinal – tinitial = 5oC – 94oC = -89oC
q = mst = 869 g x 0.444 J/g • oC x –89oC = -34,000 J
11
12.
Calorimetry
• Calorimeter: a closed container used to measure temperature
changes in physical and chemical processes.
• From the temperature changes we can calculate the heat of
the reaction, q
– qv; heat measured under constant volume conditions
– qp: heat measured under constant pressure conditions
q = m x s x t
12
13.
Two of the most common types of calorimeters are the
coffee cup calorimeter and the bomb calorimeter.
13
14.
Coffee Cup Calorimeter
• The open system allows the
pressure to remain constant.
• Thus we measure qp = ΔH, the
enthalpy change.
• ΔH =change in heat at constant
pressure.
14
15.
The 1st Law of Thermodynamics
Keeps Track of Heat Transfer
• If we monitor the heat transfers (q) of all materials
involved, we can predict that their sum will be zero.
• By monitoring the surroundings, we can predict what is
happening to our system.
• Heat transfers until thermal equilibrium, thus the final
temperature is the same for all materials.
15
16.
Example of Using 1st Law for Heat Transfer
A 43.29 g sample of solid is transferred from boiling water (T=99.8°C)
to 152 g water at 22.5°C in a Styrofoam coffee cup calorimeter. The
Twater increased to 24.3°C. Calculate the specific heat of the solid
assuming no heat was lost or gained by the calorimeter cup.
qsample+ qwater + qcup= 0
qcup is neglected in problem = 0
Hence qsample = - qwater
16
17.
qsample = - qwater
qsample 43.29 g s 24.3 99.8 C
4.184 J
qwater 152 g 24.3 22.5 C
g C
4.184 J
43.29g s 24.3 99.8 C 152g 24.3 22.5 C
g C
s - 3.2684(10 ) g C 1.1447(103) J
3
0.45 J
s
g C
17
18.
Chemistry in Action:
18
19.
Chemistry in Action:
Fuel Values of Foods and Other Substances
1 cal = 4.184 J
1 Cal = 1000 cal = 4184 J
Substance Hcombustion (kJ/g)
Apple -2
Beef -8
Beer -1.5
Gasoline -34
19
20.
A metal alloy was heated for 2 minutes in a hot water bath at 101.20 C before
being transferred into an aluminum calorimeter cup containing water at 23.42 °C.
A rise in the temperature of the water in the calorimeter cup was observed. In fact
the maximum temperature reached was 28.25 °C. Using the data below, calculate
the specific heat of the metal object in cal/g°C. Note that the specific heats of
water and aluminum are 1.00 cal/gC and 0.22 cal/gC respectively.
Wt. Alloy 243.42 g
Wt. water in calorimeter cup 213.14 g
Wt. Calorimeter cup 45.37 g
Initial temp. of water 23.42 C
Initial temp. of calorimeter cup 23.42 °C
Initial temp. of metal alloy 101.20 °C
Final temp. of water 28.25 C
20
+1-315-636-4485
+91 98151-74050 
