Lab.9. Calorimetry
Key words:
Heat, energy, exothermic & endothermic reaction, calorimeter, calorimetry, enthalpy of
reaction, specific heat, chemical & physical change, enthalpy of neutralization, law of
conservation of energy, final temperature, initial temperature, lattice energy, hydratation
energy, enthalpy of solution
Literature:
J. A. Beran; Laboratory Manual for Principles of General Chemistry, pp. 245-256.
J.E. Brady, F. Senese: Chemistry – Matter and its Changes, 4th ed. Wiley 2003, Chapters 7,
20, .
M. Hein and S. Arena: Introduction to Chemistry, 13th ed. Wiley 2011; pp. 157 - 161
J. Crowe. T. Bradshaw, P. Monk, Chemistry for the Biosciences. The essential concepts.,
Oxford University Press, 2006; pp. 416 - 450.
Theoretical background
Accompanying all chemical and physical changes is a transfer of heat (energy); heat may be
either evolved (exothermic) or absorbed (endothermic). A calorimeter (Fig. 1) is the
laboratory apparatus that is used to measure the quantity and direction of heat flow
accompanying a chemical or physical change. The calorimeter is well-insulated so that,
ideally, no heat enters or leaves the calorimeter from the surroundings. For this reason, any
heat liberated by the reaction or process being studied must be picked up by the calorimeter
and other substances in the calorimeter. The heat change in chemical reactions is
quantitatively expressed as the enthalpy (or heat) of reaction, H, at constant pressure. H
values are negative for exothermic reactions and positive for endothermic reactions.
Lab.9. Calorimetry
Fig. 1. A set nested coffee cups is a good constant pressure calorimeter.
Two quantitative measurements of heat are detailed in his experiment: measurements of the
heat accompanying an acid-base reaction, and the heat associated with dissolution of a salt in
water.
Specific Heat
The energy (heat, expressed in joules, J) required to change the temperature of one gram of a
substance by 1⁰C is the specific heat of that substance:
 J 
energy J 
 
specific heat 
 g  C  mass g   T C 
(1)
or, rearranging for energy,
 J 
  mass g   T C 
energy J   specific heat 
 g  C 
(2)
T is the temperature change of the substance. Although the specific heat of a substance
changes slightly with temperature, for our purposes, we assume it is constant over the
temperature changes of this experiment.
Enthalpy (Heat) of neutralization of an Acid-Base Reaction
The reaction of a strong acid with a strong base is an exothermic reaction that produces water
and heat as products.
Lab.9. Calorimetry
Definition: Enthalpy of neutralization: energy released per mole of water formed in an acidbase reaction-an exothermic quantity.
H 3O  aq   OH  aq   2 H 2 O l   heat
The enthalpy (heat) of neutralization, Hn , is determined by (1) assuming the density and
specific heats of the acid and base solutions are equal to that of water and (2) measuring the
temperature change, T , when the two are mixed.
H n   specific heat H 2O  combined massesacid base  T
(3)
H n is generally expressed in units of kJ/mol of acid (or base) reacted. The mass (grams) of
the solution equals the combined masses of the acid and base solutions.
Enthalpy (Heat) of Solution for the Dissolution of a Salt
Lattice energy: energy required to vaporize one mole of salt into its gaseous ions-an
endothermic quantity.
Hydratation energy: energy released when one mole of a gaseous ion is attracted to and
surrounded by water molecules forming one mole of hydrated ion in aqueous solution-an
exothermic quantity.
When a salt dissolves in water, energy is either absorbed or evolved depending upon the
magnitude of the salt’s lattice energy and the hydration energy of its ions. For the dissolution
of KI:
2O
KI s  H

 K  aq  I  aq
Hs =+13 kJ/mol
The lattice energy (an endothermic quantity) of a salt, HLE, and the hydration energy (an
exothermic quantity), HHyd, of its composite ions account for the amount of heat evolved or
absorbed when one mole of the salt dissolves in water. The enthalpy (heat) of solution, Hs,
is the sum of two terms (Fig. 2):
H s  H LE  H Hyd
(4)
Whereas HLE and HHyd are difficult to measure in the laboratory, Hs is easily measured.
Lab.9. Calorimetry
Fig.2. Energy changes in the dissolving of solid KI in water.
The enthalpy of solution for the dissolution of a salt, Hs, is determined experimentally by
adding the heat changes of the salt and water when the two are mixed. Hs is expressed in
units of kilojoules per mole of salt.


H s   heat change H 2O  heat changesalt 
 specific heat H 2O  massH 2O  TH 2O
H S   
mole salt

(5)
  specific heat salt  masssalt  Tsalt 
   
 (6)

molesalt

 
A temperature rise for the dissolution of a salt, indicating an exothermic process, means that
the HHyd is greater than the HLE for the salt; conversely, a temperature decrease in the
dissolution of the salt indicates that HLE is greater than HHyd and Hs is positive.
Lab.9. Calorimetry
Table 1. Specific Heat of Some Salts
Salt
Formula
Specific Heat (J/g•⁰C)
NH4Cl
1.57
Ammonium nitrate
NH4NO3
1.74
Ammonium sulfate
(NH4)2SO4
1.409
Calcium chloride
CaCl2
0.657
Sodium carbonate
Na2CO3
1.06
Sodium hydroxide
NaOH
1.49
Sodium sulfate
Na2SO4
0.903
Ammonium chloride
Sodium thiosulfate pentahydrate NaS2O3•5H2O
1.45
Potassium bromide
KBr
0.439
KNO3
0.95
Potassium nitrate
Plot the Data.
Plot the temperature ( y axis) versus time ( x axis) on the top half of a sheet of linear graph
paper. The maximum temperature is the intersection point of two lines: (1) the best line drawn
through the data points on the cooling portion of the curve and (2) a line drawn perpendicular
to the time axis at the mixing time [when the acid is added to the base or when a salt is added
to the water]. (Fig.3).
Lab.9. Calorimetry
Fig. 3. Extrapolation of temperature vs. time data (not scale) for an exothermic process.
EXPERIMENTAL PART
PART 1
Enthalpy (Heat) of Neutralization for an Acid-Base Reaction
Calculate the enthalpy of neutralization reaction:
q = m∙ c ∙ΔT = m∙ c∙ ( Tf - Ti )
∆H = q /n,
where c - specific heat of mixture, n – moles of water created
Step 1
Measure 50.0 mL of 1.1 M HCl in a clean, graduated cylinder. Measure temperature with
temperature sensor and record it in Table 2.
Step 2
Using another clean, graduated cylinder transfer 50.0 mL of a standard 1.0 M NaOH solution
to the dry calorimeter. Record the temperature and exact molar concentration of the NaOH
solution in Table 2.
Lab.9. Calorimetry
Step 3
Carefully, but quickly add the acid to the base, replace the calorimeter lid, and swirl gently.
Read and record the temperature and time every 5 seconds for 1 minute and thereafter every
15 seconds for other 2 minutes.
Step 4
Plot the temperature (y axis) versus time (x axis) on the top of half of a sheet of linear graph
paper. Determine the maximum temperature following instruction presented on Fig. 3.
[THEORETICAL BACKGROUND].
Step 5
Repeat the acid-base experiment. Plot the data on the bottom half of the same sheet of graph
paper.
Step 6
Calculate the thermodynamic data and fill the proper places in Table 3.
Step 7
Repeat steps 1-6, for 1.1 M HNO3.
Table 2. Data for acid–base neutralization
HCl + NaOH
HNO3 + NaOH
Trial 1 Trial 2 Trial 1 Trial 2
1. Volume of acid (mL)
2. Temperature of acid ⁰C
3. Volume of NaOH (mL)
4. Temperature of NaOH (⁰C)
5. Exact molar concentration of NaOH (mol/L)
6. Maximum temperature from graph( ⁰C)
Table 3. Calculation for Enthalpy (Heat) of Neutralization for an Acid-Base Reaction
HCl + NaOH
Trial Trial
1
2
Average initial temperature of acid and base ( ⁰C )
Temperature change, T ( ⁰C )
HNO3 + NaOH
Trial 1 Trial 2
Lab.9. Calorimetry
Volume of final mixture (mL)
Mass of final mixture (g) (Assume the density of the
solution is 1.0 g/mL)
Specific heat of mixture
4.18J/g∙◦C
4.18J/g∙◦C
Heat evolved (J)
Amount of OH- reacted, the limiting reactant (mol)
Amount of H2O formed (mol)
Heat evolved per mole of H2O, ∆Hn (kJ/mol H2O)
Average ∆Hn (kJ/mol H2O)
Answer the question: Is the enthalpy different for both neutralization reactions or is the same
(in frame of experimental error)?
Comment on your answer.
Disposal: Rinse the calorimeter twice with deionized water before next step
PART 2.
Enthalpy (Heat) of Solution for the Dissolution of a Salt
Calculation for Enthalpy:
 specific heat H 2O  massH 2O  TH 2O
H S   
mole salt

  specific heat salt  masssalt  Tsalt 
   


molesalt

 
Step 8.
Prepare the salt. On weighing paper, measure about 5.0 g of the assigned salt. Record the
name of the salt in Table 4.
Lab.9. Calorimetry
Step 9
Using your clean, graduated cylinder, add 20 g (20.0 mL) of deionized water to the
calorimeter and record its temperature in Table 4. Secure the thermal sensor below the water
surface.
Step 10
Carefully add (do not spill) the salt to the calorimeter, replace the lid, and swirl gently. Read
and record the temperature and time at 5-second intervals for 1 min and thereafter every 15
seconds for 2 min. Record the data.
Step 11
Plot the temperature (y axis) versus time (x axis) on the top half of the sheet of linear graph
paper. Determine the maximum (for an exothermic process) or minimum (for an endothermic
process) temperature as done in first part under point 4.
Step 12
Calculate thermodynamic data for salt and fill the Table 4.
Step 13
Repeat steps 7-12, for other salt assigned by TA.
Disposal: Discard the salt solution into the “Waste Salts” container, followed by
additional tap water.
Table 3. Enthalpy (Heat) of solution for the Dissolution of a Salt
Name of the salt
Mass of salt (g)
Moles of salt
Mass of water (g)
Initial temperature of water ( ⁰C )
Final temperature of mixture from the graph ( ⁰C )
..................................
........................................
Lab.9. Calorimetry
Table 4.Calculation for Enthalpy (Heat of Solution for the Dissolution of a Salt
Name of salt
........................................... ......................................
Change in temperature of solution
Heat loss of water (J)
Heat loss of salt (J)
Total enthalpy change
∆Hs (J/mol salt)
Average ∆Hs (J/mol salt)
Additional information
Table 1. Specific Heat of Some Salts
Salt
Formula
Specific Heat (J/g•⁰C)
NH4Cl
1.57
Ammonium nitrate
NH4NO3
1.74
Ammonium sulfate
(NH4)2SO4
1.409
Calcium chloride
CaCl2
0.657
Sodium carbonate
Na2CO3
1.06
Sodium hydroxide
NaOH
1.49
Sodium sulfate
Na2SO4
0.903
Ammonium chloride
Sodium thiosulfate pentahydrate NaS2O3•5H2O
1.45
Potassium bromide
KBr
0.439
KNO3
0.95
Potassium nitrate