Download calorimetry

CALORIMETRY
Objective: To determine a metal’s specific heat and approximate atomic mass; to determine the
quantity of heat evolved in the neutralization of a strong acid with a strong base; and to determine the
quantity and direction of heat flow for dissolving a salt in water.
Discussion: In nearly all chemical and physical changes occurring in nature, energy is either
evolved or absorbed. In the laboratory a calorimeter is used to measure the heat flow for a chemical
reaction. Reactions which evolve heat are exothermic; their H values are negative. Reactions that
absorb heat are endothermic; their H values are positive.
We will make three calorimetric determinations in this experiment: a metal’s specific heat; the heat of
neutralization for an acid-base reaction; and the heat of solution for a salt dissolving in water.
SPECIFIC HEAT OF A METAL: THE LAW OF DULONG AND PETIT
The quantity of heat required to change the temperature of 1 gram of a substance by 1 oC is the
specific heat of that substance. The specific heat of a nonreactive metal with water is determined by
heating the metal to a known temperature, placing it into a calorimeter containing a measured amount
of water at a known temperature, and measuring the final equilibrium temperature. This process
assumes no heat is lost to the calorimeter. By assuming that the heat lost by the metal is equal the
heat absorbed by the water the specific heat of the metal can be determined.
Pierre Dulong and Alexis Petit in 1819 proposed that 1 mole of any pure element has the same
capacity for absorbing heat, approximately 25 joules per degree Celsius per mole. Expressed
mathematically,
Specific heat (J/goC) x atomic mass (g/mole) = to 25 J/moloC
Although the value of their “constant” varies 10% or more, a metal’s atomic mass can be estimated
from its specific heat.
HEAT OF NEUTRALIZATION FOR AN ACID-BASE REACTION
The reaction of a strong acid with a strong base produces water and heat:
H+ + OH-  H2O + heat
The heat of neutralization, Hn, is determined by: a) assuming the density and specific heats of the
acidic and basic solutions are the same as water’s, and b) measuring the temperature change when
the two are mixed.
HEAT OF SOLUTION FOR DISSOLVING A SALT
A salt’s lattice energy and the hydration energy of its composite ions govern the heat evolved or
absorbed when it dissolves in water. Lattice energy, an endothermic quantity, is the energy required
to vaporize 1 mole of a solid salt into its gaseous ions; hydration energy, an exothermic quantity, is
the energy released when the gaseous ions are spontaneously attracted by and surrounded by water
molecules as the solution forms. The heat of solution, DHs, is the difference between these two. See
the example for NaCl below:
NaCl(s)  Na+(g) + Cl-(g)
Na+(g) + Cl-(g)  Na+(aq) + Cl-(aq)
H = lattice energy (positive value)
H = hydration energy (negative value)
NaCl(s)  Na+(aq) + Cl-(aq)
Hs = sum of the above H values
The heat of solution, Hs, is determined by adding the heat changes for the solid and water as the
salt dissolves.
Qrxn = heat change (H2O) + heat change (solid) = mct(water) + mct(solid)
Then Qrxn can be converted to Hs by changing the units. Hs is generally expressed in units of kJ/g
salt. The specific heats of some common salts are listed below.
Salt
Specific heat (J/goC)
Ammonium chloride
0.506
Ammonium nitrate
1.28
Ammonium sulfate
1.18
Sodium sulfate
0.845
Sodium thiosulfate pentahydrate
1.45
Potassium bromide
0.435
Potassium nitrate
0.895
Procedure: You should complete two trials for each part of this experiment.
A. SPECIFIC HEAT OF A METAL
1. Mass an unknown metal sample in a dry, previously massed 200-mm test tube. Place the test
tube in a 400-mL beaker filled with water well above the level of the metal in the test tube.
Heat to boiling and maintain this temperature for at least 5 minutes so that the metal reaches
thermal equilibrium with the water. Measure the water’s temperature.
2. Obtain two polystyrene coffee cups and a thermometer. Place one cup inside the other. Using
a graduated cylinder, add 75.0 mL of water to the calorimeter. Allow the thermometer to sit in
this water until the temperature stops changing. Once thermal equilibrium is attained, read,
and record the temperature.
3. Once both temperatures have been recorded in steps 1 and 2 remove the test tube from the
boiling water and quickly transfer the metal to the calorimeter. Be careful not to break the
thermometer. Swirl gently. Do not use the thermometer as a stirring rod! Record the highest
temperature attained by the metal/water mixture
B. HEAT OF NEUTRALIZATION
4. Pour 50.0 mL of 1.0 M NaOH solution, measured with a graduated cylinder, into a clean, dry
calorimeter. Record the temperature and concentration of the solution.
5. Measure 50.0 mL of 1.1 M HCl in a clean, graduated cylinder. Rinse the thermometer to
remove any NaOH and then measure the acid’s temperature. Carefully and quickly, add the
acid to the base, and swirl gently. Do not use the thermometer as a stirring rod. Record the
highest temperature attained by the acid/base mixture.
C. HEAT OF SOLUTION
6. Using a graduated cylinder, add 25.0 mL of water to a clean, dry calorimeter. Record the
water’s temperature.
7. Mass about 5.0 grams of your assigned salt. Add the salt to the calorimeter and swirl gently
until the salt dissolves.
8. Measure the highest (or lowest) temperature attained by the solution.
Calculations:
A. SPECIFIC HEAT OF A METAL
1. Calculate the specific heat of the metal for each trial, and the average specific heat.
2. Determine the approximate atomic mass of the metal
B. HEAT OF NEUTRALIZATION
1. Calculate the heat evolved for each trial, and the average heat evolved.
2. Calculate Hn in kJ/mole H2O.
3. Obtain the true value for Hn from your textbook or another reference, and calculate the percent
error of your result.
C. HEAT OF SOLUTION
1. Calculate the Hs in J/g for each trial and the average Hs.
Pre-lab Questions:
1. A metal weighing 15.08 g at 81.2oC is placed in a calorimeter containing 25.0 mL of water at
24.8oC. The final temperature is 29.1oC. What is the metal’s specific heat and its approximate
atomic mass?
2. A 5.0 g sample of KBr dissolves in 30.0 mL of water at 25.0oC. The final temperature of the
solution is 18.1oC. What is the heat of solution per gram and per mole of KBr?
3. A metal’s specific heat is 0.665 J/goC. An unknown mass of the metal undergoes a 49.7oC
drop in temperature when added to 25.0 g H2O. IF the heat which the metal evolves causes a
9.7oC temperature rise in the water, how many grams of metal are there?
4. In part B, excess mole of HCl are added to NaOH. Why is this procedure used rather than
adding a 1:1 mole ratio?
Post-lab Questions:
1. In Parts A and B the calorimeter, although a good insulator, absorbs some heat when the
system is above room temperature. a) How does this affect the metal’s specific heat value?
b) How does this affect the reported Hn value in the acid/base reaction? c) How does this
affect the reported Hs value for the dissolving of a salt if the dissolution is endothermic?
2. The specific heat capacity for the calorimeter is 1.34 J/goC. If we assume that the entire inner
cup reaches thermal equilibrium with the solution, how much heat is lost to the calorimeter in
part B? Assume the cup’s mass is 2.35 grams.
3. The value for the heat of neutralization is essentially the same for all strong acid-base
reactions. Explain.
4. Suppose that in part A the thermometer is miscalibrated by +0.5 oC. Will this affect the reported
specific heat of the metal? If so, how?
5. If the HCl solution I Part B is mistakenly prepared as a 0.91 M solution rather than the labeled
1.1 M solution, how will the reported Hn be affected? Explain.