Determination of Heat of solution of different solutes, comparing values and coming up with appropriate explanations

Project outline Heat of Solution.pdf
Final Report: Heat of Solution
(Draft deleted)

Project Objective (from Heat of Solution.pdf)
hen certain compounds are dissolved in water there is an enthalpy change (heat change) associated with this process.
In this experiment you will determine the heat of solution of compounds like potassium nitrate and ammonium chloride.
Compare their results and come up with logical arguments as to the reason for the observation.
This experiment can be performed using improvised calorimeters (Styrofoam cups)

  • Research and come up with a possible lab procedure for setting up your equipments and performing the experiment.
  • Include safety concerns that should be addressed
  • Make a list of equipments that may be used to perform the experiment.
  • Submit a draft of the procedure adopted for the lab to your teacher; and
    discuss if the materials required are available and the lab activity can be performed.
  • Choose a convenient time for your experiment, record your data and present your findings as instructed.

For the sake of organization, please follow the "Entry #, Topic Title, and Posted By" format to avoid any mix-up!
ENTRY 1: Basic Information and Links
The process of dissolving is a process which involves the breaking and making of bonds, and that involves energy.
- The breaking of bonds requires or absorbs energy. Using energy like that is called endothermic.
- The formation of bonds releases energy. This is called exothermic.
- Dissolution overall can be either endothermic or exothermic, depending on whether more energy was used to break the bonds,
or more energy was released when new bonds or more energy was released when new bonds were formed. If more energy is released in making bonds
than is used in breaking bonds, the process is exothermic. If more energy is used than is released, the process is endothermic.
^ That's just basic information for reference and I also typed up the project objective so it won't be a hassle to keep referring to the pdf :)
I didn't manage to find too much on the topic, but the links should provide a little bit of background info.
enthalpy change of a solution:
potassium nitrate:
ammonium chloride:$file/ST-AmmoniumChloride.pdf

  • will refer to textbook later to find additional info, 加油!
30/09/08 Posted by Yiing Hu

ENTRY 2: Basic Info Continued
- Expressed in kJ/mol at constant pressure
- the heat of solution of a substance is defined as the sum of the energy absorbed (positive kJ/mol) and the energy released (negative kJ/mol)
- According to, when using the term "heat of solution", the word "solution" in this case doesn't refer to the actual physical mixture that is formed, but rather the process of dissolving

Dissolution takes place in three steps:
1. Breaking solute-solute attraction (endothermic)
2. Breaking solvent-solvent attractions (endothermic)
3. Forming solvent-solute attractions (exothermic) through solvation (read more here)

04/10/08 Posted by Nicole Pun

ENTRY 3: Measuring Heat of Solution
- a calorimeter is used to measure the heat of chemical reactions, physical changes, and heat capacity
- to find heat change of a substance made of two liquids, add the liquids to the calorimeter and take note of the initial temperature, as well as the final temperature when the reaction is finished; multiply temperature change by mass and specific heat capacity of the two reactants to find the energy given off during the reaction (if the reaction is exothermic); divide energy change by how many moles of the substance was present [this method doesn't take into account heat loss through container or the heat capacity of the thermometer and container itself; we can use this as SOURCES OF ERROR]

Basically, we need to create a simple calorimeter to roughly measure the heat of solution of potassium nitrate or ammonium chloride (or similar compounds). I suppose we'll need a relatively good thermometer, Styrofoam cup(s), the compounds (in their separate form) and specific heat capacities, depending on what we end up using.

From Wikipedia:

Constant-pressure calorimeter
A constant-pressure calorimeter measures the change in enthalpy of a reaction occurring in solution during which the atmospheric pressure remains constant.
An example is a coffee-cup calorimeter, which is constructed from two nested Styrofoam cups and holes through which a thermometer and a stirring rod can be inserted. The inner cup holds the solution in which of the reaction occurs, and the outer cup provides insulation. Then
Cp = (W * DH / (M * DT))

DH = Enthalpy of solution
DT = change of temperature
W = weight of solute
M = molecular weight of solute

A. Finding the Heat of Solution from Calorimetry (taken from****
Review calorimetry from Chapter 11. This chapter tends to use specific heat instead of heat capacity. Using specific heat, Equation 11.8 becomes

heat energy (q) = (mass)(specific heat)(change of temperature)

Because there is normally much more solvent than solute, it is usually assumed that the specific heat is that of the solvent. For water (the most common solvent),the specific heat is 4.184 J/g•°C. Remember that the mass will refer to the mass of the entire solution, since it is the entire solution undergoing the temperature change.
The heat of solution refers to the dissolution of a solute. For an ionic compound, he reaction is the salt forming its composite ions. Although water is required for the reaction, it is not written as part of the reaction. For example,


So heat of solution is calculated from


Since the heat (q) comes from the chemical reaction of the solute, the external image delta_12_black.gifH value is per moles solute, not per mol solution.
The value of external image delta_12_black.gifH is positive for endothermic dissolutions(temperature decreases) and negative for exothermic dissolutions (temperature increase).

04/10/08 Posted by Nicole Pun

ENTRY 4: Enthalpy Change of Solution
Summarized from Wikipedia's Enthalpy Change of Solution

  • enthalpy of solution = enthalpy of dissolution = enthalpy change associated with the dissolution of a substance in a solvent at constant pressure
  • enthalpy: heat content (denoted by H, h, or rarely as χ)
  • dissolution: the process by which a solid, gas, or liquid is dispersed homogeneously in a gas, solid, or, esp., a liquid.
  • one of the three dimensions of solubility analysis
  • kJ/mol, at constant temperature
  • heat of solution of a substance is defined as the sum of the energy absorbed, or endothermic energy, and energy released, or exothermic energy
    • endothermic = absorption of energy (i.e. heat) = positive kJ/mol
    • exothermic = liberation of energy (i.e. heat) = negative kJ/mol
  • heating decreases solubility of gases, exothermic
    • gas dissolves in liquid solvent, temperature decreases, solution releases E
    • an effect of the increase in heat (heat = E needed to attract solute and solvent molecules; outweighs the E needed to separate solvent molecules)
  • see Nicole's post for three steps of dissolution
  • value of overall enthalpy change = sum of individual enthalpy changes of each step
  • solutions with negative heats of solution form stronger bonds and have lower vapor pressure
Heat of solution for some selected compounds

hydrochloric acid



ammonium nitrate






potassium hydroxide



caesium hydroxide



sodium chloride



potassium chlorate



acetic acid



-Charissa 2008/10/04

ENTRY 5: Compounds
Background info on the compounds we'll be experimenting with

Potassium nitrate is a chemical compound with the chemical formula KNO3. A naturally occurring mineral source of Nitrogen, KNO3constitutes a critical oxidizing component of black powder/ gunpowder. In the past it was also used for several kinds of burning fuses, including slow matches. Since potassium nitrate readily precipitates, urine was a significant source, through various malodorous means, from the Late Middle Ages and Early Modern era through the 19th century.[citation needed]

Ammonium chloride(NH4Cl) (also Sal Ammoniac, salmiac, nushadir salt, zalmiak, sal armagnac, sal armoniac, salmiakki, salmiakand salt armoniack) is, in its pure form, a clear white water-soluble crystalline salt of ammonia. The aqueous ammonium chloride solution is mildly acidic. Sal ammoniac is a name of natural, mineralogical form of ammonium chloride. The mineral is especially common on burning coal dumps (formed by condensation of coal-derived gases), but also on some volcanoes.
Sorry for the huge font, the visual editor is very troublesome.
04/10/08 Posted by Yiing Hu

ENTRY 6: Calorimeter
Okay, this is some information on making an "improvised calorimeter", since we'll be using styrofoam cups instead of the real thing.
Calorimetry is used to determine the heat released or absorbed in a chemical reaction. The calorimeters shown here can determine the heat of a solution reaction at constant (atmospheric) pressure. The calorimeter is a double styrofoam cup fitted with a plastic top in which there is a hole for a thermometer. (It's crude, but very effective!) Key techniques for obtaining accurate results are starting with a dry calorimeter, measuring solution volumes precisely, and determining change in temperature accurately.

Calorimetric measurements led to the discovery that every substance requires a characteristic amount of heat to change its temperature over a temperature interval. The amount of heat required is approximately, but not exactly, uniform over any reasonable range of temperatures. The amount of heat required per unit of mass came to be known as the specific heat of the substance. The specific heat of a substance is an intensive property characteristic of the substance.

The ice calorimeter is simply a large insulated container of ice and water with a basket which can be used to remove the ice for weighing. The amount of heat evolved in whatever reaction takes place within the calorimeter is equal to the mass of ice melted multiplied by the heat of fusion of ice, 333.51 kJ/kg.

The water calorimeter is often used in undergraduate student laboratories in the form shown as the Figure below, the "coffee cup" calorimeter.

On Page 133
- Read more on the website, for more information on making a coffee cup calorimeter

Most reactions we perform in chemistry 30 are done in solution. The cheapest and most effective calorimeter to use to study solution chemists is a simple Styrofoam coffee cup. Water and the Styrofoam are the surroundings and the Styrofoam provides adequate insulation since most reactions in solution occur relatively rapidly. A thermometer is used to measure the temperature change in the water solution. Water is typically used because of it's high specific heat capacity, which means it can store large amounts of heat without large temperature changes occurring. (It can also release large amounts of heat for the same reason). Other substances , like oil can be used in place of the water.
The coffee cups may be nested (placed one inside another) to provide addition insulation if needed.

04/10/08 Posted by Yiing Hu
Entry 7: Tentative Procedure and Materials

1. Trace the outline of the thermometer’s bulb onto the lid. Carefully cut it out with an Exacto knife. Ensure that the thermometer slides through the lid easily.
2. Measure 100mL of water with a graduated cylinder and pour it into the Styrofoam cup.
3. Measure 10mL of the chemical and pour it into the water gently. Stir.
4. Put the lid onto the cup.
5. Place the thermometer through the lid and into the cup. Ensure that the bulb is below the water surface.
6. For the next few minutes (est. 3?), observe and take note of the changes in temperature.
7. Repeat.
8. Calculate heat flows.
Note: Take note of the experimental errors that take place during the experiment, and any outside factors that might have influenced the results.

100mL x (# of compounds used) of water
10mL of each compound
2 Styrofoam cups (to put one in the other to reduce heat loss)
1 lid that fits the cup
1 thermometer
1 Exacto knife

-Charissa 2008/10/04

Entry 8

Determining Heat of Solution - Rough Outline
To determine (molar) heat of solution for Ammonium nitrate:

Part 1
In this part of the experiment, the calorimeter is filled with 60.0 g of water. A small sealed glass bulb containing 5.00 g of pure NH4NO3 is placed in the calorimeter. The reaction is initiated by breaking the glass bulb, allowing the NH4NO3 to dissolve in the water.
The heat capacity of the calorimeter (Ccal) is 153. J oC-1. The formula weight of NH4NO3 is 80.04.
Part 2
Repeat the procedure in Part 1 for determining the heat of solution of ammonium nitrate using different values for the mass of ammonium nitrate and/or the mass of water. You should obtain the same heat capacity as you did in Part 1.

We can adopt the method; it's here to give you an idea
For more info go to

The purpose of Calorimetry and Molar Enthalpy

  • The purpose of calorimetry = determine the enthalpy of a substance undergoing chemical change
  • bomb calorimeter = enthalpy of combustion measured
    • This is how the caloric content of foods is determined
  • Since the heat absorbed/released is proportional to the amount of reactant used, molar enthalpy = DH/n is a more meaningful and characteristic quantity.
Go to:

Calorimetry - Measuring Heats of Reactions
  • Any process that results in heat being generated and exchanged with the environment is a candidate for a calorimetric study
  • has a very broad range of applicability, with examples ranging from drug design in the pharmaceutical industry to the study of metabolic rates in biological (people included) systems
  • two types of calorimetry: measurements based on constant pressure and measurement based on constant volume.


  • device used to measure heat of reaction
  • A simple styrofoam cup is sufficient
    • because it is a container with good insulated walls to prevent heat exchange with the environment
  • In order to measure heats of reactions, we often enclose reactants in a calorimeter, initiate the reaction, and measure the temperature difference before and after the reaction
    • The temperature difference enables us to evaluate the heat released in the reaction.
  • Calorimeter may be operated under constant (atmosphere) pressure, or constant volume
  • Whichever kind to use, we must know its heat capacity
    • heat capacity = amount of heat required to raise the temperature of the entire calorimeter by 1 K
      • determined experimentally before or after the actual measurements of heat of reaction
  • heat capacity of the calorimeter is determined by transferring a known amount of heat into it and measuring its temperature increase
  • B/C temperature differences = VERY small, sensitive thermometers are needed
The following example will demonstrate how this can be done.
Example 1
The temperature of a calorimeter increases 0.10 K when 7.52 J of electric energy is used to heat it. What is the heat capacity of the calorimeter? Solution:

Dividing the amount of energy by the temperature increase yields the heat capacity, C,
C = 7.52 / 0.10 = 75.2 J/K.

For more information, go to:

-Thamy G. 05/10/08

Entry 9
Just a note: I added two more compounds from Wikipedia onto the draft because Mr. Vincent said we need at least three compounds in order to compare them.

-Charissa 2008/10/05

Entry 10

We need to update our materials list. Here are the things that we used in our lab:

- safety goggles
- scoopula
- 2 Styrofoam "bowls"
- 1 thermometer
- graduated cylinder
- water
- 10g each of NaOh, NH4Cl, NaCl, K2NO3

Please add more if you guys remember something I missed.

And just a quick reminder about what our objective is:

When certain compounds are dissolved in water there is an enthalpy change (heat change) associated with this process. In this experiment you will determine the heat of solution of compounds like potassium nitrate and ammonium chloride. Compare their results and come up with logical arguments as to the reason for the observation. This experiment can be performed using improvised calorimeters (Styrofoam cups)

Posted by Nicole Pun [2008/11/23]