How Does Enthalpy Change When Copper Sulphate Dissolves?
Molar heat of a solution or enthalpy of solution is defined as the amount of heat taken in or thrown out while per mole of a solution is being dissolved in any solvent, mostly water. In popular terms and academics, this molar heat is denoted by ΔH and measured in kJ/mol.
The reaction ensuing while being dissolved may be exothermic or endothermic. If heat is generated while the solute dissolves, then reaction is endothermic. On the other hand, if heat is absorbed while the solute dissolves, the reaction is called endothermic. The exchange of heat is determined by the amount of energy required to break down intermolecular bonds and also, heat released when new solute-solvent bonds are being made.
This process mentioned above is also called as heat of dissolution.
Aim of Experiment
To calculate value of enthalpy of dissolution of copper sulphate or potassium nitrate with help of water as a solvent and potassium nitrate (KNO3) or copper sulphate(CuSO4) as solute.
Theory behind Experiment
For all experiments such as this, aqueous solutions are used as a solvent. And among all aqueous solutions, water is mostly used due to the ease with which compounds are dissolved in water. Also, water shows much accurate changes in temperature when solutes are being dissolved in it.
Going by law of conservation of energy proposed in thermodynamics, sum of all enthalpy exchanges must amount to zero. Therefore, the below equation is followed while any reaction, which involves heat, takes place.
ΔH1 + ΔH2 + … + ΔHn = 0
While this heat exchange is taking place, the solution is also being formed simultaneously. Usually, a solution is declared as final when mixing any more solute inside the solvent does not result in any more heat exchange.
Amount of exchange of heat is measured with the help of a calorimeter.
Equipment Needed for Enthalpy of Dissolution of Copper Sulphate or Potassium Nitrate
Glass rod
Thermometer
Suitable beakers
Weight box
Copper sulphate in water
Measuring cylinder
Stirrer
Potassium sulphate water
Potassium nitrate
Water as a solvent
Block of wood
Cotton wool
Block of cardboard
Filler
How to Setup the Apparatus?
A crucial aspect while conducting this experiment, follow this image below to understand how to setup the equipments.
(image will be uploaded soon)
Procedure of Whole Experiment
This experiment, when carried on inside laboratories, is divided into three parts. First, we measure calorimeter constant to be used further. Then we first dissolve CuSO4 in water and find its heat of dissolution using the calorimeter constant value. Separately, we again dissolve potassium nitrate in water and find its heat of dissolution value.
Calculation of Calorimeter Constant
Take a polythene bottle and fix the thermometer and stirrer inside it as shown in the figure below.
Pour 100 ml of water (distilled) in the bottle.
Take down the temperature at which the water stands. Let this temperature be t1°C.
Take some water in a beaker. Place this beaker on a heater and heat it not more than 30°C above room temperature.
Keep aside 100 ml of this heated water and keep it aside in another beaker.
Take down the temperature at which this water stands. Let this temperature be t2°C.
Instantly, add this warm water into the polythene bottle. Do not waste any time. Otherwise, the temperature reading will be invalid.
Stir the mixed contents vigorously.
After mixing, take down the temperature at which the mixture stands and name it t3°C.
Pop Quiz 1
What material should be used to fill in the space between the two beakers?
Cotton wool (Answer)
Iron
Water
Should be kept empty
Calculation of Enthalpy of Dissolution of Copper Sulphate
Take a small beaker with a fixed calorimeter and put 100 ml of water inside it.
Place this smaller beaker inside a larger one of volume, say 500 ml.
Place the whole system on a block of wood.
Put the smaller beaker inside the larger beaker and fill the space with cotton wool.
Also, cover the entire system with a piece of card board.
Take down the temperature at which the water stands, say T1°C.
Take a fixed amount of copper sulphate powder and put it inside the water.
Let the copper sulphate dissolve and remove the excess solution.
Make sure your stirrer and thermometer are in place.
After the whole copper sulphate is dissolved, take down the temperature of the solution, say T2°C.
Calculation of Potassium Nitrate Enthalpy of Dissolution
Take a small beaker with a fixed calorimeter and put 100 ml of water inside it.
Place this smaller beaker inside a larger one of volume, say 500 ml.
Place the whole system on a block of wood.
Put the smaller beaker inside the larger beaker and stuff the empty space with cotton wool.
Also, cover the entire system with a piece of cardboard.
Take down the temperature at which the water stands, say T1°C.
Take a fixed amount of potassium nitrate powder and put it inside the water.
Let the potassium nitrate dissolve and remove the excess solution.
Make sure the thermometer and stirrer are in place.
After the whole potassium nitrate is dissolved, take down the temperature of the solution, say T2°C.
Inference and Observation
Calculations
If we assume water density to be constant and the specific heat of the solution equal to that of water, then the amount of heat absorbed or released can be given by the equation
Q = (W + 200) x (t1 – t2) cals
The above equation, when transformed to give result in joules, converts to
Q = (W + 200) x (t1 – t2) x 4.2 J
So, for w/M moles of solute in solvent (here, water),
Q = (W + 200) x (t1 – t2) x 4.2 J
Hence, for 1 mole of solute dissolved in solvent (here, water),
Q = (W + 200) x (t1 – t2) x 4.2 x M/w joules
So, the final formula for enthalpy of dissolution,
ΔH = (W + 200) x (t1 – t2) x 4.2 x M/w joules
The symbols in the above formula can be expressed as follows.
M = Formula mass of solute,
w = Mass of solute,
W = Equivalent weight of water calorimeter.
Did You Know?
ΔH is positive if the reaction is exothermic and heat is released during solution formation, and negative if the reaction is endothermic and heat is absorbed during solution formation.
Activity
Perform the above experiment in your school laboratory and mention what is the enthalpy of dissolution of both these compounds.
Precautions to be Taken
In the first step of the experiment, transfer the hot water immediately into the cold water such that the temperature reading does not change.
Mix the solute inside the solvent well enough. Do not stir too fast, or else the heat of the solution will increase due to friction.
Copper sulphate, being hygroscopic, should be measured well to take the initial reading of the powder.
For potassium nitrate, a small amount, lesser than 3 gm, has to be dissolved in 100 ml of water.
The space between the small beaker and the large beaker should be stuffed with cotton wool. Cotton wool is an insulator and does not let heat pass through it.
Use a well-calibrated thermometer to measure the solution temperature. The thermometer should be accurate till 0.1°C such that minor changes in temperature can be noticed.
This was all about the enthalpy of copper sulphate or potassium nitrate, when dissolved in water. If you are curious to read more about enthalpy of solution and other physio-chemical processes, check out our reference notes, sample papers and free study material, available on our Vedantu app.
FAQs on Enthalpy of Dissolution of Copper Sulphate: Key Concepts and Steps
1. What is meant by the enthalpy of dissolution for copper sulphate?
The enthalpy of dissolution (ΔH_sol) of copper sulphate refers to the total heat energy that is absorbed or released when one mole of the substance dissolves completely in a large amount of a solvent, typically water, at constant pressure. This value differs significantly depending on whether the copper sulphate is anhydrous (CuSO₄) or hydrated (CuSO₄·5H₂O).
2. Why is dissolving anhydrous copper sulphate (CuSO₄) exothermic, but dissolving hydrated copper sulphate (CuSO₄·5H₂O) is endothermic?
This difference arises from the energy changes involved in breaking and forming bonds:
- Anhydrous Copper Sulphate (Exothermic): When anhydrous CuSO₄ dissolves, a large amount of energy, known as hydration enthalpy, is released as water molecules surround the Cu²⁺ and SO₄²⁻ ions. This energy release is much greater than the energy required to break the crystal lattice, resulting in a net release of heat (negative ΔH).
- Hydrated Copper Sulphate (Endothermic): For CuSO₄·5H₂O, energy must first be used to break apart the crystal lattice, which already contains water molecules. The energy absorbed to overcome these forces is greater than the energy released when the ions are further hydrated by the solvent. This results in a net absorption of heat from the surroundings (positive ΔH), making the solution feel cold.
3. How is the enthalpy of dissolution of copper sulphate typically determined in a lab?
The enthalpy of dissolution is determined experimentally using a technique called calorimetry. The basic steps are:
- A known mass of copper sulphate is added to a known volume of water inside an insulated container (calorimeter).
- The initial temperature of the water (T₁) is recorded.
- The mixture is stirred until the salt is fully dissolved, and the final temperature (T₂) is recorded.
- The heat absorbed or released by the solution (q) is calculated using the formula q = mcΔT, where 'm' is the mass of the solution, 'c' is its specific heat capacity, and 'ΔT' is the temperature change (T₂ - T₁).
- This heat change is then used to calculate the molar enthalpy of dissolution in kJ/mol.
4. What is the significance of a positive or negative sign for the enthalpy of dissolution (ΔH_sol)?
The sign of the enthalpy of dissolution (ΔH_sol) tells you whether heat is released or absorbed during the process:
- A negative ΔH_sol indicates an exothermic process. Heat is released into the surroundings, causing the temperature of the solution to increase. This is observed when dissolving anhydrous copper sulphate.
- A positive ΔH_sol indicates an endothermic process. Heat is absorbed from the surroundings, causing the temperature of the solution to decrease. This is what happens when dissolving hydrated copper sulphate.
5. How does Hess's Law relate to the enthalpies of dissolution for anhydrous and hydrated copper sulphate?
Hess's Law is crucial for indirectly determining the enthalpy of hydration of anhydrous copper sulphate. Since it's practically impossible to measure the energy change of turning gaseous Cu²⁺ and SO₄²⁻ ions into a solid crystal, we use a thermochemical cycle. By experimentally measuring the enthalpy of dissolution for both anhydrous (ΔH_sol(anhydrous)) and hydrated (ΔH_sol(hydrated)) copper sulphate, we can calculate the enthalpy of hydration (ΔH_hydration) using the following relationship derived from Hess's Law:
ΔH_hydration = ΔH_sol(anhydrous) - ΔH_sol(hydrated).
6. What is the difference between enthalpy of solution and enthalpy of hydration?
While related, these two terms describe different parts of the dissolution process:
- Enthalpy of Solution (or Dissolution): This is the overall energy change when a solute dissolves in a solvent. It is the net result of two competing energies: the energy required to break the solute's lattice (lattice enthalpy) and the energy released when ions are solvated (hydration enthalpy).
- Enthalpy of Hydration: This refers specifically to the energy released when one mole of gaseous ions is surrounded by water molecules to become hydrated ions. It is just one component that contributes to the overall enthalpy of solution.
7. What are some real-world applications of understanding the enthalpy of dissolution?
Understanding the enthalpy of dissolution is important in various fields:
- Chemical Manufacturing: Engineers must account for heat changes when dissolving reactants on a large scale to prevent overheating or to provide necessary energy.
- Hot and Cold Packs: Instant cold packs often use the endothermic dissolution of salts like ammonium nitrate, while hot packs use the exothermic dissolution of salts like calcium chloride.
- Pharmaceuticals: The dissolution rate and energy of a drug can affect its bioavailability and formulation.
- Environmental Science: It helps in understanding the thermal effects of mineral dissolution in natural water bodies.
