বিক্রিয়ার এনথালপি পরিবর্তন পরিমাপ

I fondly remember my first chemistry class. It was my sophomore year of high school. It was the first time that a science class made sense. I also remember my favorite experiment that year. Our teacher told us beforehand to bring food to our next class. When we walked into class, there were matches and soda cans on every lab bench.

We were doing calorimetry. So, we spent the whole class period burning different foods. The classroom quickly smelled of burning food. It was putrid! My group made a huge mess because we combusted marshmallows. We didn't think about the gases in the marshmallow expanding. We just wanted to light it on fire. Soon, we had sticky marshmallow all over the lab table.

In professional chemistry labs, the calorimeters are much more complex. But, we managed with just a soda can filled with water. We combusted the food directly underneath the can. It contained water, and we measured the water's temperature change. We were able to calculate the reaction's enthalpy change.

Calorimetry is the experimental method for measuring energy change. It's the basis for all thermochemical data. We've looked at how to calculate enthalpy change. We saw that there were a few methods. For example, we can use Hess's law or enthalpies of formation. But, where did the data come from? Chemists used calorimetry to find the reaction enthalpies of various chemical reactions. There are two calorimetry methods: constant pressure and constant volume. We'll explore each one.

Constant Pressure Calorimetry

Constant pressure calorimetry is the most common. So many chemical reactions happen under atmospheric pressure. It works well when the products and reactants are solids, liquids, or solutions. These phases are not very expandable or compressible. So, they do very little work against the atmosphere. So, our energy change equals the enthalpy. Remember, we defined enthalpy to equal heat at constant pressure.

We can do constant pressure calorimetry with reactions that produce gases. But, we need a container that has a moveable piston. The gas pushes against the piston. So, our vessel's volume changes. We know from Boyles law that the change in volume offsets the pressure. We calculate the work based on the piston's displacement. Unfortunately, moveable pistons can be challenging to use. You have to make sure they are correctly sealed. Also, they can only displace a certain distance. So, reactions that produce a lot of gas need constant volume calorimeters.

How do we calculate enthalpy at constant pressure? We use the formula $Q=-mc_{p}\Delta T$.

• The negative sign determines if the reaction is exothermic or endothermic.
• $Q$ is heat. At constant pressure $\Delta H=Q$
• $m$ is the mass.
• $c_{p}$ is the specific heat capacity.
• $\Delta T$ is the temperature change.

Constant Pressure Calorimetry Calculation Example

Let's look at an example. Let's imagine we did an acid-base reaction in a calorimeter. We mix 25.0 milliliters of 0.5 molar hydrobromic acid ($HBr$) and sodium hydroxide ($NaOH$). Our balanced chemical equation is $HBr_{(aq)} + NaOH_{(aq)} \rightarrow NaBr_{(aq)} + H_{2}O_{(l)}$. Let's say our temperature increased by 5.24 celsius. The solutions are dilute. So, we can assume their density equals water's density. That's $1.0\frac{g}{ml}$. Let's calculate the enthalpy change.

First, we need to find the mass. Our total volume is 50.0 ml. We know the density, so we can calculate mass.

• $50.0 ml\times 1.0\frac{g}{ml}= 50.0g$

We also need to know the specific heat capacity. For dilute solutions, it's an excellent approximation to use water's. It's $4.18\frac{J}{g ^{o}C}$.

Now we can use $Q=-mc_{p}\Delta T$.

• $\Delta H=Q=-mc_{p}\Delta T=-50.0g\times 4.18\frac{J}{g ^{o}C}\times 5.24^{o}C=-1095.16J$

Constant Volume Calorimetry

Constant volume calorimetry uses a bomb calorimeter. The bomb allows us to do reactions that produce a lot of gases. For example, combustion reactions. Keeping a constant volume eliminates pressure-volume work. So, our internal energy change equals heat. But, constant volume calorimeters have a trick. You have to determine the whole calorimeter's heat capacity. To do that, we use a standard. A common standard is benzoic acid ($C_{6}H_{5}COOH$). It releases $26.38\frac{KJ}{g}$. So, we could do a benzoic acid combustion. Then, we would know the heat ($Q$) and the temperature change ($\Delta T$). We can then solve for the calorimeter's heat capacity. We use it in later experiments.

Constant Volume Calorimetry Calculation Example

Let's imagine that we did a bomb calorimeter experiment. We used 1.3 grams of a benzoic acid standard. It raised the temperature by 6.4 Celsius. Next, we did a second reaction that raised the temperature by 2.1 Celsius. Let's find the heat value for our second reaction.

First, we need to find the calorimeter's heat capacity. We'll need the energy of our standard and the temperature change.

• $c_{cal}=\frac{Q_{standard}}{\Delta T}=\frac{26.38\frac{KJ}{g}\times 1.3 g}{6.4}=5.4\frac{KJ}{^{o}C}$.

Now we can find the heat from our second reaction. We just multiply the calorimeter's heat capacity by the temperature change.

• $\Delta U=Q_{v}=-c_{calorimeter}\times \Delta T=5.4\frac{KJ}{^{o}C}\times 2.1^{o}C= -11.3 KJ$

(Video) Coffee Cup Calorimeter - Calculate Enthalpy Change, Constant Pressure Calorimetry

by The Organic Chemistry Tutor

(Video) Bomb Calorimeter vs Coffee Cup Calorimeter Problem - Constant Pressure vs Constant Volume Calorimeter

by The Organic Chemistry Tutor

Between these two videos, the Organic Chemistry Tutor demonstrates some example problems for measuring enthalpy change. The first video uses a coffee cup calorimeter. So, that's constant pressure calorimetry. The second video is a bomb calorimetry example. So, that's constant volume calorimetry. These two videos will help you understand how we measure enthalpy changes.