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Chemistry · 9. Chemical energetics

Exothermic and endothermic reactions and enthalpy

Track energy transfer and use simple calorimetry and Hess’s law to calculate enthalpy changes.

45 minutesContent version: 2.2
Start the interactive pathwayOpen the virtual experiment

Driving question

How can a measured temperature change reveal the energy change of a reaction?

Curriculum coverage

Vietnam Chemistry 10; Cambridge IGCSE 0620 and AS & A Level 9701: energetics, calorimetry and Hess cycles.

Model scope and limitations

Simple cup calorimetry often assumes solution density 1.00 g/cm³ and specific heat capacity 4.18 J g⁻¹ K⁻¹; heat loss and apparatus heat capacity shift results from standard values. The simulation illustrates direction of change and does not replace calibrated calorimetry data.

Learning objectives

  • Distinguish system from surroundings and determine the sign of ΔH.
  • Use q = mcΔT and ΔH = −q/n with consistent units.
  • Apply Hess’s law and bond energies with the correct break–make convention.

Prerequisite knowledge

  • Moles, balanced equations and J–kJ conversion.
  • Temperature, mass and specific heat capacity.

Core knowledge

Sign of ΔH

An exothermic reaction transfers energy from system to surroundings, so ΔHsystem < 0; an endothermic reaction absorbs energy, so ΔHsystem > 0.

Calorimetry

The solution gains qsolution = mcΔT. Neglecting losses, qreaction = −qsolution and ΔH = qreaction/n for the defined amount of reaction.

Hess’s law

Enthalpy is a state function, so overall ΔH is the algebraic sum of steps; reversing an equation changes the sign and scaling it scales ΔH.

Worked example

Mixing 50.0 cm³ of 1.00 mol/dm³ HCl and 50.0 cm³ of 1.00 mol/dm³ NaOH raises temperature by 6.80 K. With m = 100 g and c = 4.18 J g⁻¹ K⁻¹, qsolution = 2842 J. As 0.0500 mol H₂O forms, ΔH = −2.842/0.0500 = −56.8 kJ mol⁻¹.

Misconceptions to avoid

  • A temperature rise does not mean positive system ΔH; the surroundings warm because the reacting system loses energy.
  • Bond formation releases energy whereas bond breaking requires it; using mean bond enthalpies, ΔH ≈ energy to break − energy released on forming.
  • A temperature interval has the same numerical size in kelvin and Celsius, but absolute temperatures are not interchangeable without conversion.

Virtual experiment procedure: Zinc–acid reaction and thermal change

  1. In the Zn–HCl simulation, identify system and surroundings and predict the sign of ΔH before mixing.
  2. Record observations and the direction of temperature change; do not invent measurements if the simulation has no quantitative sensor.
  3. Use the lesson calorimetry data to calculate q, limiting moles and ΔH, then state a directional source of error.

Safety and cautions

  • Metal–acid reactions produce flammable H₂ and may splash acid; real work requires eye protection, ventilation and no ignition source near the reaction vessel.

Evidence to collect

  • An energy-transfer diagram showing system, surroundings and the sign of ΔH.
  • A unit-labelled q and ΔH calculation with reacting moles and at least one assumption or heat-loss source.
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<iframe src="https://visuallab.win/chem-lab/?experiment=exp_sgk_zn_hcl" width="100%" height="560" frameborder="0" allow="fullscreen" loading="lazy"></iframe>

Diagnostic check

1. A reaction that warms the surroundings normally has what system ΔH?

  1. Negative
  2. Positive
  3. Always zero
Show answer and explanation

Correct answer: Negative

Energy leaves the system, so system enthalpy decreases.

2. In q = mcΔT, what is a suitable unit for c?

  1. J g⁻¹ K⁻¹
  2. J mol⁻¹
  3. kg m s⁻¹
Show answer and explanation

Correct answer: J g⁻¹ K⁻¹

Specific heat capacity is energy per mass per temperature interval.

3. When a Hess equation is reversed, what happens to ΔH?

  1. Change its sign
  2. Keep it unchanged
  3. Square it
Show answer and explanation

Correct answer: Change its sign

The reverse path transfers the same energy in the opposite direction.

Application practice

1. A 100 g solution warms by 5.0 K with c = 4.18 J g⁻¹ K⁻¹. How much heat does it gain?

  1. 2.09 kJ
  2. 20.9 kJ
  3. 0.209 kJ
Show answer and explanation

Correct answer: 2.09 kJ

q = 100×4.18×5.0 = 2090 J = 2.09 kJ.

2. If 0.050 mol of reaction releases 2.85 kJ, what is the molar ΔH?

  1. −57 kJ mol⁻¹
  2. +57 kJ mol⁻¹
  3. −0.143 kJ mol⁻¹
Show answer and explanation

Correct answer: −57 kJ mol⁻¹

ΔH = −2.85/0.050 = −57 kJ mol⁻¹; negative because the reaction releases heat.

3. For an exothermic calorimetry experiment, heat loss to the room usually makes measured |ΔH| do what?

  1. Smaller than the true value
  2. Larger than the true value
  3. Unaffected
Show answer and explanation

Correct answer: Smaller than the true value

The observed ΔT is too small, so q and the magnitude of ΔH are underestimated.

Academic references

  1. Cambridge IGCSE Chemistry 0620 — official syllabus and resources
  2. Cambridge AS & A Level Chemistry 9701 — official syllabus
  3. OpenStax Chemistry 2e — peer-reviewed general chemistry
  4. IUPAC Color Books — authoritative chemical terminology and nomenclature
  5. BIPM — SI Brochure, 9th edition

This lesson is maintained against the cited sources below. Simulations are learning models and do not replace supervised physical-laboratory safety procedures.

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WebLab STEM
Technical maintainer
Đức Tiến — Control & Automation Engineer