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Lesson 6 of 914 minutes

Energy Transfer, Calorimetry, and Enthalpy

Start with the lesson question, connect the representations, and test the model with evidence.

system surroundingsheattemperaturecalorimetryenthalpyhess lawbond energy

Learning objectives

  • Track heat flow with explicit system and surroundings boundaries.
  • Apply calorimetry with units, signs, assumptions, and uncertainty.
  • Determine reaction enthalpy using calorimetry, Hess's law, and formation enthalpies.
Lesson flowHook, model, explanationShow guidance

Inspect the opening phenomenon

Predict what changes, then name the evidence.

Apply in the lab

Name the evidence before reading the answer.

Read only what helps

Then use the lab and recall check.

More when needed

Transcript and resources stay available below.

Course progress

AP Chemistry · Thermochemistry · Lesson 6

Energy Transfer, Calorimetry, and Enthalpy

In progress

Decision challenge

Observe the phenomenon. Then connect the representations.

Use the opening example to make a prediction, identify evidence, and explain which model supports it.

Does a Warmer Cup Mean the Reaction Absorbed Heat?

Predict the sign of reaction heat when the measured solution warms.

Define the boundary before interpreting a calorimetry temperature change.

Before

Predict the sign of reaction heat when the measured solution warms.

During

Track the boundary, temperature sign, surroundings heat, reaction heat, and molar conversion.

After

Explain why a positive solution heat implies a negative reaction enthalpy under the stated assumptions.

Reference drawerTranscript, source notes, scripts, and package status stay tucked away until you need them.7 files

Lesson reading

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14 min

Video script

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Transcript fallback

available

courses/ap-chemistry/modules/06-thermochemistry/lessons/01-energy-transfer-calorimetry-and-enthalpy/video-transcript.md

Calorimeter Energy Audit

approved

1 hr 30 min

Mastery check

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6 questions / 15 min

Book section:courses/ap-chemistry/modules/06-thermochemistry/lessons/01-energy-transfer-calorimetry-and-enthalpy/book-section.md
Transcript for accessibility and fallback

A hot cup does not prove the reaction gained heat. First draw the boundary. The solution is the measured surroundings. If it warms, delta T is positive, so q solution equals m c delta T and is positive. Energy conservation makes q reaction negative: the reaction is exothermic. For one hundred grams warming six point five degrees, q solution is about plus two point seven kilojoules. If zero point zero five zero zero mole reacted, delta H is about minus fifty-four kilojoules per mole. Retrieval pause: does breaking a bond release energy? No. Bond breaking requires energy; forming bonds releases it. Keep the boundary, sign, units, and assumptions together. Continue the free AP Chemistry lesson on EduQuest AI.

Reading lab

Core explanation

Connect the lesson's words, diagrams, graphs, evidence, and equations.

Driving question

How can a temperature change become defensible evidence about energy transfer in a chemical process?

Define the boundary before the sign

The system is the reaction or process being studied; everything else is the surroundings. Energy conservation requires qsystem+qsurroundings=0.q_{\mathrm{system}}+q_{\mathrm{surroundings}}=0. If the solution warms, the solution gained heat: qsurr>0q_{\mathrm{surr}}>0, so the reaction released heat: qrxn<0q_{\mathrm{rxn}}<0. Temperature measures average particle kinetic energy, not total thermal energy.

System and surroundings energy map

Coffee-cup calorimetry

For a solution treated as the surroundings, qsolution=mcΔT,ΔT=Tf−Ti.q_{\mathrm{solution}}=mc\Delta T,\qquad \Delta T=T_f-T_i. Suppose 100.0 g100.0\ \mathrm g of solution warms from 22.0∘C22.0^\circ\mathrm C to 28.5∘C28.5^\circ\mathrm C, and assume c=4.184 J g−1 K−1c=4.184\ \mathrm{J\,g^{-1}\,K^{-1}}: qsolution=(100.0)(4.184)(6.5)=2.7×103 J.q_{\mathrm{solution}}=(100.0)(4.184)(6.5)=2.7\times10^3\ \mathrm J. With negligible calorimeter heat loss, qrxn=−2.7 kJq_{\mathrm{rxn}}=-2.7\ \mathrm{kJ}. If 0.0500 mol0.0500\ \mathrm{mol} reacted, ΔHrxn≈−2.7 kJ0.0500 mol=−54 kJ mol−1.\Delta H_{\mathrm{rxn}}\approx\frac{-2.7\ \mathrm{kJ}}{0.0500\ \mathrm{mol}}=-54\ \mathrm{kJ\,mol^{-1}}. The two-significant-figure result reflects the measured temperature change. A better model includes qcal=CcalΔTq_{\mathrm{cal}}=C_{\mathrm{cal}}\Delta T and all relevant surroundings terms.

Calorimetry calculation chain

Enthalpy is path independent

At constant pressure, heat transferred for the process is represented by ΔH\Delta H. Reversing a thermochemical equation changes the sign of ΔH\Delta H; multiplying the equation multiplies ΔH\Delta H. Adding equations adds their enthalpy changes—Hess's law—because enthalpy is a state function. For standard formation enthalpies, ΔHrxn∘=∑νΔHf∘(products)−∑νΔHf∘(reactants).\Delta H^\circ_{\mathrm{rxn}}=\sum \nu\Delta H_f^\circ(\mathrm{products})-\sum \nu\Delta H_f^\circ(\mathrm{reactants}). Coefficients ν\nu and physical states matter. Elements in their standard states have ΔHf∘=0\Delta H_f^\circ=0 by convention.

Hess law route comparison

Bonds and particulate reasoning

Breaking bonds requires energy; forming bonds releases energy. Average bond enthalpies estimate ΔHrxn≈∑D(bonds broken)−∑D(bonds formed).\Delta H_{\mathrm{rxn}}\approx\sum D(\text{bonds broken})-\sum D(\text{bonds formed}). This gas-phase average-bond model is approximate and is not interchangeable with measured formation-enthalpy data.

Retrieval challenge

  1. A reaction warms its solution. What are the signs of qsolutionq_{\mathrm{solution}} and qrxnq_{\mathrm{rxn}}?
  2. Why must the system boundary be named before assigning a sign?
  3. What changes when a thermochemical equation is reversed?
  4. Why is “breaking bonds releases energy” incorrect?

Sources

Practice labCalorimeter Energy AuditOpen this when you are ready to apply the model, collect evidence, and check your explanation.1 hr 30 min

Objective

Determine how including calorimeter heat capacity changes a measured molar enthalpy while maintaining a defensible system boundary and uncertainty record.

Materials

  • Nested insulated cups with lid and clamped temperature probe
  • Balance, graduated cylinders or volumetric pipets, and timer
  • Warm and cool water for calibration
  • Instructor-approved dilute acid/base pair or lower-risk dissolution system
  • Splash goggles, lab coat or apron, closed shoes, and SDS-selected gloves

Safety and supervision

Instructor supervision is required. Review every SDS and the local emergency and waste plan before work. Conduct the activity in a normally ventilated instructional laboratory; use local exhaust or a fume hood when the selected reagent SDS requires it. Never use concentrated reagents. Dilute acids and bases can still irritate or damage eyes and skin, and some dissolutions can make the cup noticeably hot or cold. Avoid skin/eye contact, splashing, inhalation of powders, and direct handling of a hot cup. Flush eye or skin exposures with water for at least 15 minutes, notify the instructor, and follow institutional eyewash, spill, and medical-response procedures. Stop the trial if the vessel leaks, the probe or cup becomes unstable, unexpected gas or odor appears, or the temperature approaches a teacher-set safe limit. Collect mixtures in the instructor-designated labeled aqueous-waste container; never use the drain without authorization.

The instructor must select the actual chemical system, concentration, microscale quantity, ventilation control, glove material, and waste stream from current SDS information and institutional rules. Warm/cool water calibration is the lower-risk option. A teacher-provided dataset or simulation is appropriate when chemical handling is unsuitable.

Steps

  1. Review the SDS, waste, spill, and emergency plan; put on required PPE.
  2. Calibrate the nested-cup calorimeter using measured masses of warm and cool water.
  3. Investigate an instructor-approved process at microscale and record reagent identity, concentration, amount, total solution mass, and temperature versus time.
  4. Extrapolate the temperature curve to mixing time where appropriate.
  5. Calculate qsolutionq_{\mathrm{solution}}, qcalq_{\mathrm{cal}}, qrxnq_{\mathrm{rxn}}, and molar enthalpy; compare results with and without the calorimeter term.
  6. Place all mixtures in the designated waste stream and clean the station.

Expected Result

Including a calibrated calorimeter term changes the magnitude of inferred reaction heat and generally reduces systematic underestimation caused by treating the cup as thermally invisible.

Data and energy audit

Record units and uncertainty for every measured quantity. For each trial, report ΔT\Delta T, qsolution=mcΔTq_{\mathrm{solution}}=mc\Delta T, qcal=CcalΔTq_{\mathrm{cal}}=C_{\mathrm{cal}}\Delta T, and qrxn=−(qsolution+qcal).q_{\mathrm{rxn}}=-(q_{\mathrm{solution}}+q_{\mathrm{cal}}). Divide by moles of the stated limiting reagent only after checking reaction stoichiometry. Include a signed conservation statement and compare the corrected value with the cup-ignored model.

Quality and accessibility

Use a probe with large display or screen-reader output, clamp vessels, provide seated access, and assign data or analysis roles when handling is unsuitable. A teacher-provided dataset or simulation is an accessibility and pre-lab alternative, not a replacement for required supervised hands-on laboratory time. Discuss heat loss, mixing, probe lag, density, specific-heat assumptions, uncertainty, and significant figures.

Reflection Questions

  1. Which objects belong to the system and surroundings?
  2. Which measurement most strongly affects uncertainty in molar enthalpy?
  3. How would heat lost to the room bias an exothermic result?

Extension Challenge

Fit the trend before and after mixing, extrapolate both fits to mixing time, and compare the corrected temperature change with the observed maximum.