This lab puts you in the role of a product designer. You're testing different chemical or physical processes to figure out which one releases the most heat, then using calorimetry data to justify your choice. The core skill is connecting a temperature change you can measure to an energy change you can calculate, and then explaining what that means at the molecular level.

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Why This Lab Matters for the AP Exam

Thermochemistry questions show up consistently on the AP Chemistry exam, and they almost always ask you to do one of three things: calculate heat using $q = mc\Delta T$ , interpret an energy diagram, or explain why a process is exothermic or endothermic in terms of energy flow. This lab gives you hands-on practice with all three.

The design challenge framing also mirrors how the AP exam asks free-response questions. You won't just be asked to plug numbers into a formula. You'll need to make a claim, support it with calculated values, and explain the reasoning. That's exactly what you practice here.

CED Connections

This lab directly supports Unit 6: Thermochemistry across four topics.

Topic 6.1 - Endothermic and Exothermic Processes (LO 6.1.A)

The whole lab is built around 6.1.A. When you observe a temperature increase in the water inside your calorimeter, that's your experimental evidence that the reaction is exothermic (6.1.A.1, 6.1.A.2). You're watching energy leave the system (the dissolving salt or reacting chemicals) and enter the surroundings (the water). For dissolution reactions specifically, whether the process is exothermic or endothermic depends on the relative strengths of the interactions being broken versus the new ones being formed (6.1.A.4).

Topic 6.2 - Energy Diagrams (LO 6.2.A)

After collecting data, you should be able to sketch an energy diagram for each candidate reaction. The diagram shows reactants, products, and the activation energy barrier. An exothermic process has products sitting lower on the potential energy axis than reactants (6.2.A.1).

Topic 6.4 - Heat Capacity and Calorimetry (LO 6.4.A)

This is the calculation backbone of the lab. You use $q = mc\Delta T$ to calculate how much heat was transferred to the water in your calorimeter (6.4.A.1). The first law of thermodynamics tells you that the heat released by the reaction equals the heat absorbed by the water, so energy is conserved across the system and surroundings (6.4.A.2). Different substances have different specific heat capacities, which affects how much their temperature changes for the same amount of heat input (6.4.A.3).

Topic 6.5 - Energy of Phase Changes (LO 6.5.A)

Some hand warmer designs use phase changes rather than chemical reactions. If your design involves a substance freezing or crystallizing, you're dealing with the enthalpy of fusion concept. Freezing releases energy (the system loses energy as it transitions from liquid to solid), which is why sodium acetate hand warmers work the way they do (6.5.A.1, 6.5.A.2).

What You Need to Be Able to Do

Experimental Design

Identify the independent variable (the chemical or physical process being tested), the dependent variable (temperature change of the water), and the controlled variables (mass of water, starting temperature, calorimeter type)
Explain why a calorimeter is used and what it's trying to minimize (heat loss to the environment)
Recognize that your calorimeter is acting as a simplified closed system for the purpose of the experiment

Data Collection and Calculation

Record temperature vs. time data and identify the maximum temperature change
Use $q = mc\Delta T$ to calculate the heat absorbed by the water
Apply the assumption that $q_{water} = -q_{reaction}$ to find the heat released by the reaction
Calculate molar enthalpy change by dividing by moles of limiting reagent

Graphing and Interpretation

Plot temperature vs. time and identify the point of maximum heat transfer
Sketch energy diagrams that correctly show reactants, products, activation energy, and the sign of $\Delta H$

Claim-Evidence-Reasoning

Make a claim about which candidate material works best as a hand warmer
Support it with calculated $q$ values and molar enthalpy comparisons
Explain the reasoning using energy flow between system and surroundings

Core Concepts

Exothermic vs. Endothermic

An exothermic reaction releases energy to the surroundings. The products have less chemical potential energy than the reactants, so the difference gets released as heat. In your calorimeter, this shows up as a temperature increase in the water. The enthalpy change ( $\Delta H$ ) for an exothermic process is negative.

An endothermic reaction absorbs energy from the surroundings. The products have more potential energy than the reactants. Temperature of the surroundings drops. $\Delta H$ is positive.

For a hand warmer, you obviously want exothermic. But understanding endothermic processes matters too, because the AP exam will ask you to compare both.

System and Surroundings

The system is the chemical reaction or physical process you're studying (the dissolving salt, the reacting chemicals). The surroundings is everything else, including the water in your calorimeter. Energy flows between them. When the system releases heat, the surroundings absorb it, and you see a temperature rise.

The First Law of Thermodynamics

The first law of thermodynamics says energy is conserved. It doesn't appear or disappear. In calorimetry, this means:

$q_{system} + q_{surroundings} = 0$

So if the water gains 500 J of heat, the reaction released 500 J. You'll use this relationship constantly.

Heat Transfer Equation

$q = mc\Delta T$

$q$ = heat transferred (in joules)
$m$ = mass of the substance being heated or cooled (in grams)
$c$ = specific heat capacity (in J/g·°C), which tells you how much energy it takes to raise 1 gram of a substance by 1°C
$\Delta T$ = change in temperature (final minus initial, in °C)

The specific heat of water is 4.184 J/g·°C. This value is high compared to most substances, which is why water is such a useful calorimeter medium.

Enthalpy Change

Enthalpy change ( $\Delta H$ ) is the heat released or absorbed at constant pressure. In a coffee-cup calorimeter (which is open to the atmosphere), pressure is essentially constant, so the heat you measure equals $\Delta H$ . To get a molar enthalpy value, divide the total heat by the number of moles of the substance that reacted.

Energy Diagrams

An energy diagram (also called a potential energy diagram or reaction coordinate diagram) plots potential energy on the y-axis against the reaction coordinate on the x-axis. The activation energy ( $E_a$ ) is the energy barrier the reactants have to overcome to get to the transition state. For exothermic reactions, the products sit lower than the reactants. The difference in height between reactants and products is $\Delta H$ .

Phase Change Enthalpy

When a substance changes phase, temperature stays constant even though heat is being transferred. The energy goes into breaking or forming intermolecular forces rather than speeding up particles.

Heat of fusion ( $\Delta H_{fus}$ ): energy needed to melt a solid. The reverse (freezing) releases the same amount.
Enthalpy of vaporization ( $\Delta H_{vap}$ ): energy needed to vaporize a liquid. Condensation releases this energy.

For water: $\Delta H_{fus} = 6.01 \text{ kJ/mol}$ and $\Delta H_{vap} = 40.7 \text{ kJ/mol}$ . Vaporization requires much more energy because you're completely separating molecules, not just loosening them.

Kinetic and Potential Energy

Kinetic energy is the energy of motion. Temperature is a measure of average kinetic energy of particles. Potential energy is stored energy, including the energy stored in chemical bonds and intermolecular forces. When a reaction releases heat, chemical potential energy is being converted into kinetic energy of the surroundings.

How the Lab Works

The investigation gives you several candidate materials that could potentially be used in a hand warmer. These might include ionic compounds that dissolve in water (like calcium chloride or sodium acetate), or other chemical processes. Your job is to test each one, measure the temperature change it produces in a known mass of water, and use that data to calculate how much heat each process releases.

You're essentially building a simple calorimeter out of a foam cup (or similar insulating container). The foam minimizes heat loss to the room, so you can assume most of the heat from the reaction stays in the water. This is the closed system assumption. It's not perfect, but it's good enough for the calculations you need to do.

For each candidate material, you dissolve or react a measured amount in a measured volume of water and track the temperature over time. The material that produces the largest temperature increase (for a comparable amount of substance) is releasing the most heat per gram or per mole.

The design challenge part means you also have to think about practical constraints. A material might release a lot of heat but be too expensive, too slow, or not safe for skin contact. The lab asks you to weigh the evidence and make a justified recommendation, which is exactly the kind of reasoning the AP exam tests.

Data and Analysis Moves

Calculating Heat Released

Once you have your temperature change, plug into $q = mc\Delta T$ :

Use the mass of the water (in grams) as $m$
Use 4.184 J/g·°C as $c$ for water
Use the maximum temperature change as $\Delta T$

This gives you the heat absorbed by the water. Because of conservation of energy, the heat released by the reaction has the same magnitude but opposite sign:

$q_{rxn} = -q_{water}$

Getting to Molar Enthalpy

To compare materials fairly, convert to a per-mole basis. Divide $q_{rxn}$ by the number of moles of the substance you used. This gives you $\Delta H$ in kJ/mol (convert from J first by dividing by 1000).

$\Delta H = \frac{q_{rxn}}{n}$

A more negative $\Delta H$ means more heat released per mole, which means a better hand warmer candidate.

Identifying Variables

Independent variable: the identity (and possibly the amount) of the candidate material
Dependent variable: the temperature change of the water, or the calculated $q$
Controlled variables: mass of water, starting temperature, calorimeter type, volume of solution

Graphing

Plot temperature (y-axis) vs. time (x-axis) for each trial. For an exothermic dissolution, you'll see a sharp rise followed by a gradual decrease as heat leaks out to the room. The maximum temperature is your best estimate of the temperature at the moment all the heat has transferred to the water.

Error Considerations

Your calorimeter isn't perfect. Some heat escapes to the foam cup itself and to the surrounding air. This means your measured $q_{water}$ is slightly less than the actual $q_{rxn}$ . This is a systematic error that makes your calculated $\Delta H$ less negative than the true value. You should acknowledge this in your analysis and explain the direction of the error.

Comparing Phase Change Processes

If one of your candidates involves a phase change (like a supersaturated solution crystallizing), the energy calculation shifts. Instead of $q = mc\Delta T$ , you'd use:

$q = n \cdot \Delta H_{fus}$

where $n$ is moles of substance and $\Delta H_{fus}$ is the molar enthalpy of fusion. Freezing releases energy equal in magnitude to melting, just with the opposite sign.

Common Mistakes

Mixing up system and surroundings in the sign convention. The water's temperature goes up, so $q_{water}$ is positive. That means $q_{rxn}$ is negative (exothermic). Students often flip this and report a positive $\Delta H$ for an exothermic reaction. If the hand warmer gets hot, $\Delta H$ is negative.

Using the wrong mass in $q = mc\Delta T$ . You use the mass of the water (the surroundings), not the mass of the salt or chemical you added. The formula is calculating how much heat the water absorbed.

Forgetting to convert to kJ/mol. If you calculate $q$ in joules and report $\Delta H$ in joules per mole, that's off by a factor of 1000 compared to the standard kJ/mol unit. Always check your units.

Assuming a bigger temperature change always means a better hand warmer. A large $\Delta T$ could just mean you used more material. The fair comparison is molar enthalpy, which accounts for how much substance you used.

Drawing energy diagrams with the wrong relative heights. For an exothermic reaction, products must be lower than reactants on the potential energy axis. The activation energy hump sits above the reactants, not the products. Students sometimes draw the hump on the wrong side or put products higher than reactants for an exothermic process.

Treating the calorimeter as a perfect closed system. It's not. Heat loss to the cup and air is real. Your calculated $\Delta H$ will be slightly less negative than the actual value. Acknowledging this and explaining the direction of error is expected in AP-level analysis.

Confusing specific heat with heat of fusion. Specific heat applies when temperature is changing. Heat of fusion applies during a phase change when temperature is constant. They're different quantities used in different situations.

Quick Review Checklist

You can explain why a temperature increase in the calorimeter water indicates an exothermic process in the system.
You can use $q = mc\Delta T$ to calculate heat absorbed by water and then find $q_{rxn}$ using conservation of energy.
You can convert $q_{rxn}$ to a molar enthalpy value (in kJ/mol) and use the sign correctly (negative for exothermic).
You can sketch an energy diagram for an exothermic or endothermic process, correctly labeling reactants, products, activation energy, and $\Delta H$ .
You can identify the controlled, independent, and dependent variables in a calorimetry experiment.
You can explain the difference between heat released during a chemical reaction and heat released during a phase change, and use the correct equation for each.
You can describe at least one source of error in a coffee-cup calorimeter setup and explain how it affects your calculated result.
You can make a claim about which material makes the best hand warmer and support it with calculated $\Delta H$ values and a reasoning statement about energy flow.