Lab Skills

Thermometer Reading

Interpolate between scale markings on a randomized analog thermometer. Practice °C only, °F only, or both with scored conversion problems.

Interpolation°C & °FAnalog ScalePractice
Topic L.3

Thermometers and Calorimetry Technique

Make observations or collect data from representations of laboratory setups or results, while attending to precision where appropriate.

Reading a thermometer: record to one estimated digit past the smallest division. A thermometer marked in 1 °C increments is read to ±0.1 °C. Read at eye level, with the bulb fully immersed and not touching the container wall.

Calorimetry technique:

  • Insulate. A nested pair of foam cups with a lid minimizes energy exchange with the room. Without a lid, an exothermic reaction reads a temperature rise that is too small, so the calculated |ΔH| is too small.
  • Stir. Uneven temperature means the thermometer reads a local value, not the system average.
  • Extrapolate. Because some energy always leaks, the standard practice is to record temperature versus time and extrapolate the cooling line back to the moment of mixing. The extrapolated maximum is closer to the true value than the observed maximum.
  • Use the solution mass, not the solute mass, in q = mcΔT.

Standard assumptions you may be asked to state: the calorimeter absorbs negligible energy, no energy is exchanged with the surroundings, and the dilute aqueous solution has the density (1.00 g/mL) and specific heat capacity (4.18 J·g⁻¹·°C⁻¹) of water.

Error direction: essentially every calorimetry error — no lid, slow mixing, delayed reading, a cold thermometer absorbing energy — reduces the measured ΔT and therefore underestimates the magnitude of ΔH.

Key points

  • Record one estimated digit past the smallest division.
  • Use the mass of the solution, not the solute, in q = mcΔT.
  • Most calorimetry errors cause energy loss and underestimate |ΔH|.
  • ΔT in °C equals ΔT in K — no conversion needed for a difference.

Equations

  • on the exam sheet

Common mistakes

  • Recording 25 °C instead of 25.0 °C. Precision is part of the measurement.
  • Using the solute mass in q = mcΔT.
  • Claiming heat loss makes |ΔH| too large. It makes it too small.

Worked example

A student measures ΔH for a neutralization and obtains −48 kJ/mol; the accepted value is −57.3 kJ/mol. Give two plausible experimental explanations and state why each lowers the magnitude.

1. Energy loss to the surroundings. If the calorimeter was not covered or was poorly insulated, some of the energy released escaped to the room instead of warming the solution. The measured ΔT is therefore smaller than the true value, so qsolution = mcΔT is too small and the magnitude of the calculated ΔH is too small.

2. Energy absorbed by the calorimeter and thermometer. The cups, stirrer, and thermometer all warm along with the solution, absorbing energy that is not counted in q = msolution × c × ΔT. Again the measured ΔT understates the energy released, and |ΔH| comes out low.

Also acceptable: delayed temperature reading, so the maximum was missed as the system cooled; or failure to stir, so the thermometer read a cooler region of the solution.

Note the pattern: nearly every calorimetry error reduces the observed temperature change, which is why student values for exothermic reactions are almost always low in magnitude.

Full notes for topic L.3 →