When cooking have you ever used a thermometer to get your food just right? Whether it was to not overcook meat, burn caramel, or something else the thermometer helped us measure these temperature changes.
Similar to this in chemistry, we can measure temperature or heat changes to better understand the reactions taking place around us. As heat changes can tell scientists a lot about an element and its characteristics and therefore help us understand the world around us.
First, we’ll go over what constant pressure calorimetry is.
Next, we’ll look at the equation and its importance.
After, we’ll read about some examples of constant pressure calorimetry.
Lastly, we’ll compare constant pressure and constant volume calorimetry.
Constant-pressure Calorimetry Definition
Let's begin by looking at the definition of calorimetry.
Calorimetry is the technique used to measure the amount of heat transferred to or from a substance in a chemical reaction.
We use a calorimeter to do this as it’s a calibrated and insulated device.
The temperature we measure from the calorimeter is what’s used to find out the amount of heat transferred to or from a substance in a chemical reaction.
Important things to know to understand how a calorimeter works are:
What a system and surroundings are.
A system is a substance that's undergoing the change caused by a chemical reaction.
In contrast, the surroundings are all the other components around the system, which includes the calorimeter itself.
When an exothermic reaction occurs in a calorimeter, the heat created (Q) is transferred from the system to its surroundings making Q negative.
When an endothermic reaction occurs in a calorimeter, the heat created (Q) is transferred from the surroundings to the system making Q positive.
An exothermic reaction occurs when energy is released to the surroundings, as the bonds being made are stronger than the ones being broken in a chemical reaction.
In comparison, endothermic signifies that energy is absorbed from the surroundings due to the bonds being broken being stronger than the ones being created in a chemical reaction. The system absorbs energy from the surroundings to break the stronger bonds.
Constant-pressure Calorimetry Formula
In calorimetry, we assume that the energy released and absorbed by the surroundings is the same as the energy released and absorbed by the system.
For example, if we wanted to determine the heat change (enthalpy) or heat of fusion for ice melting within the calorimeter we can assume that the energy absorbed by ice while melting is equal to the energy in value but opposite in sign to the energy released by water inside the calorimeter at constant pressure (1 atm). Please see Figure 1, below:
Qice = -Qsurroundings
Another easier way to think about this is if you have a chocolate bar, and you give it to your crying sibling. You released or lost your chocolate bar to your younger sibling who has gained or absorbed your chocolate bar. In this case, you are the surroundings and the chocolate bar is the energy from the surroundings being transferred to the system, which is your younger sibling.
In giving your sibling your chocolate bar, you have performed an equal exchange of energy with opposite signs (as you lost the chocolate bar while your sibling gained it simultaneously).
Qsibling chocolate bar = -Qyour chocolate bar
Figure 1: Constant-pressure calorimetry, exothermic versus endothermic processes. Daniela Lin, Study Smarter Original
This means that the net heat change is zero as Qice + Qsurroundings = 0 or Qsibling chocolate bar + Qyou without chocolate bar = 0.
This can also be translated to Qreaction + Qsolution = 0.
This is due to the law of conservation of energy.
The law of conservation of energy states that energy is neither created nor destroyed.
This means that in this case the heat given off is equal to the heat absorbed.
Overall, exothermic reactions result in a negative Q as heat is transferred from the system (sodium hydroxide crystals in the first experiment depicted in the figure above) to the surroundings (water) leading to a temperature increase.
In contrast, endothermic reactions result in a positive Q as heat is transferred from the surroundings (water) to the system (ice cubes in the second experiment depicted in the figure above) leading to a temperature decrease.
Now that we understand the basics concerning calorimetry we can go over the equation.
Figure 2: Enthalpy change equation. Daniela Lin, Study Smarter Originals.
Enthalpy or the heat of change equation is the formula we can use to find the amount of heat transferred in a calorimeter.
The specific heat capacity of a substance is the number of joules required to raise 1g of X substance by 1 Celsius.
Constant-pressure Calorimetry Importance
Constant-pressure calorimetry is important because we use it to measure the amount of heat gained or lost during a chemical reaction.
Knowing this allows us to know whether a reaction is exothermic, releasing heat, or endothermic, absorbing heat.
In chemistry, it’s important for scientists to understand all of this because heat transfer plays a crucial role in machine efficiency, material sciences, etc.
Some real-world applications of calorimetry include controlling homeostasis in humans, counting the number of calories humans consume, and other thermal changes necessary to maintain our survival.
Constant-pressure Calorimetry Examples
Now that we understand how important constant-pressure calorimetry is, we understand the need to use the formula to calculate the heat changes associated with it.
We can now move on to reading about examples involving constant-pressure calorimetry:
A 100 g sample of coal is put into 200 g of water at room temperature (25 Celsius). Afterward, the water temperature is measured to be 50 Celsius. Find the initial temperature of the coal. Assume that there’s no heat exchange between the surroundings only between the coal and the water.
Given
C = 1.262 J/(g °C) for coal
C = 4.184 J/(g °C) for water
Q = m x C x \( \Delta T \)
The water temperature increase means that it absorbed heat. The heat came from the coal, which we can assume was at a higher temperature initially.
Since there's only heat transfer between coal and water we can say that:
Qcoal = -Qwater
m x C x \( \Delta T \)= m x C x \( \Delta T \)
\( \Delta T \) = Final temperature \( T_f \) - Initial temperature \(T_i\)
100 g x 1.262 J/(g °C) x (50 °C- \(T_i\)) = -200 g x 4.184 J/(g °C) x (50 °C-25 °C )
126.2 J/ °C x (50 °C- \(T_i\)) = -20,920 J
The initial temperature of the coal was \(T_i\)= 215.77 °C.
When 0.5 g of potassium hydroxide is dissolved in 100 g of water in a styrofoam cup calorimeter the temperature changes from 25 Celsius to 35 Celsius. What is the amount of heat created by this solution? Assume that there’s no heat exchange between the surroundings only between water and potassium hydroxide and that because the solution is aqueous the specific heat of the solution is the specific heat of water.
\(KOH (s) + H_2O (l) \longrightarrow K^+ (aq) + OH^- (aq) \)
Given
C = 4.184 J/(g °C) for water
C = 321.85 J/(g °C) for potassium hydroxide
Q = m x C x \( \Delta T \)
Since, we assume that there's only heat transfer between potassium hydroxide (reaction) and water (solution) we can say that:
Qreaction = -Qsolution
m x C x \( \Delta T \)= m x C x \( \Delta T \)
\( \Delta T \) = Final temperature \( T_f \) - Initial temperature \(T_i\)
However, since we are already given the initial and final temperatures, and because of our assumption is that heat only transfers from the potassium hydroxide to water and is not lost to the surroundings, we can state that in these conditions:
Qreaction = -Qsolution = -100 g x 4.184 J/(g °C) x (35 °C-25 °C )
Thus, The heat produced by the reaction is Qreaction = -4184 J
The negative sign indicates that this reaction is exothermic and that it produces -4,184 Joules of heat.
For more detailed information on solutions, please refer to our “Dilution” and “Concentration” articles.
Constant-pressure vs Constant-volume Calorimetry
Throughout this entire article, we’ve been addressing constant-pressure calorimetry. A coffee cup or a double styrofoam cup with a stopper on top is an example of constant-pressure calorimetry.
Figure 3: Constant-pressure calorimeter example. Daniela Lin, Study Smarter Originals.
A constant-pressure calorimeter keeps the system isolated from other surroundings like air, dust, etc. The two holes in the cover of the cup allow us to stir the mixture to make sure it's dissolved and to measure the temperature change when the pressure is constant.
But did you know that the constant-pressure calorimeter is not the only type of calorimeter?
Well, from the title above you probably guessed that the other type of calorimetry we’ll talk about is the constant-volume calorimeter.
A constant-volume calorimeter does the same thing in that it’s also used to find out the amount of heat transferred to or from a substance in a chemical reaction.
Except it keeps the volume constant and not the pressure.
A constant-volume calorimeter is also called a bomb calorimeter and this is because it’s a container that can withstand a large pressure change that can occur in chemical reactions.
Similar to a constant-pressure calorimeter, a bomb calorimeter also is a sealed system. But instead of a more simple setup it consists of a large enclosed steel vessel usually shaped like a bomb that contains the system (reactants), water in which the large steel vessel is submerged in, a thermometer, a stirrer, and lastly an ignition wire.
Examples of when we use a constant-volume calorimeter include in combustion reactions.
Combustion reactions involve an exothermic reaction between usually an oxidant, like oxygen, and a gas that results in smoke, high-temperatures, and flames.
You’ve reached the end of the article, and now you should understand the importance of a constant-pressure calorimeter, how to calculate enthalpy changes in a calorimeter and the differences between a constant-pressure and constant-volume calorimeter.
For more practice and examples click on our flashcards associated with this article.
Constant-pressure Calorimetry - Key takeaways
Calorimetry is the technique used to measure the amount of heat transferred to or from a substance in a chemical reaction. We use a calorimeter to do this as it’s a calibrated and insulated device.
Some real-world applications of calorimetry include controlling homeostasis in humans, counting the number of calories humans consume, and other thermal changes necessary to maintain our survival.
A constant-volume calorimeter does the same thing in that it’s also used to find out the amount of heat transferred to or from a substance in a chemical reaction. Except it keeps the volume constant and not the pressure.
A constant-volume calorimeter is also called a bomb calorimeter and this is because it’s a container that can withstand large pressure changes that can occur in chemical reactions.
References
- 1. Libretexts. (2021, June 30). 6.7: Constant pressure calorimetry- measuring ΔH for chemical reactions. Chemistry LibreTexts.
- 2. 5.2 calorimetry - chemistry 2E. OpenStax. (n.d.).
- 3. Libretexts. (2020, July 14). Constant volume calorimetry. Chemistry LibreTexts.
Explanations 
Exams 
Magazine 
