Calorimetry experiment

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Enthalpy measurement

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Measuring Enthalpy Changes

On this page:

Introduction What's Energy? Why Measuring Enthalpy is Important Measuring Enthalpy Changes 1) Calorimetry 2) Hess's Law 3) Calculating Bond Enthalpy Accuracy and Precision in Measuring Enthalpy Practical Considerations Units and Conversions Example Calculations Why This Matters in Chemistry Summary

Measuring Enthalpy Changes

Some reactions are hot, and some are cold. Some mix chemicals and release energy, while some mix and absorb energy.

When performing chemistry reactions, it is important to understand enthalpy changes to measure the energy involved in the reaction. This energy is important to know when designing a reaction that fuels the reaction (i.e., when designing a cooking fuel).

What's Energy?

Enthalpy (H) is the energy in the system of chemistry at constant pressure. Enthalpy will tell the chemists how much energy is released and how much energy is absorbed.

  • Exothermic reactions: release energy to the surroundings; therefore, the temperature of the reaction will increase.
  • Endothermic reactions: absorb energy from the surroundings, therefore temperature of the reaction will decrease.

Example:

  1. Exothermic: combustion of methane (there is heat released)
  2. Endothermic: photosynthesis (there is sunlight absorbed)

Enthalpy change (ΔH) is given by the formula,

ΔH = Hproducts - Hreactants

If ΔH is negative, it is exothermic; if ΔH is positive, it is endothermic.

Why Measuring Enthalpy is Important

Measuring enthalpy changes is important for several reasons.

  • Predicting the spontaneity of a reaction.
  • Helps chemists design safer reactions.
  • Determining the energy efficiency of fuels and batteries.

Measuring Enthalpy Changes

There are several different methods for measuring enthalpy changes, depending on the interpretation of the reaction.

Calorimetry

Direct measurement of heat changes using a calorimeter.

Hess's Law

Indirect calculation using known enthalpy changes of related reactions.

Bond Enthalpy

Calculation using average bond energies.

1) Calorimetry

Calorimetry is the measurement of heat changes in a chemical reaction. When heat is exchanged in a reaction, a device known as a calorimeter is used.

Simple calorimetry setup:

We can create an experiment to determine the thermal energy of a neutralization reaction rather simply. To determine the thermal energy of a reaction, we simply need to determine the temperature change caused by the reaction and then plug this value into a formula.

Equipment:

  • Styrofoam cup (acts as a calorimeter)
  • Thermometer
  • Acid + alkali (reactants)
  • Stirrer

We use the following to calculate the energy change (thermal energy) of the reaction (q) to find the change in heat (ΔH):

Q = mcΔT

m = mass of solution (g)

c = specific heat capacity (4.18 J/g/K for water)

ΔT = temperature change (K or °C)

There is a relationship between q and ΔH where: ΔH = -q / (moles of limiting reactant), where the negative sign states that the system lost heat to the surroundings.

An example of this reaction is the Neutralization of an acid with a base. Measure the temperature change (ΔT) of the water and use this data to calculate ΔH of the neutralization reaction.

Important Points to Remember

  • Make sure you have done a good enough insulating job to minimize heat lost to the surroundings.
  • It is important to stir the solution to make sure that the heat is distributed equally in the solution.

2) Hess's Law

Hess's Law states that we can find the change in heat (ΔH) for a reaction that cannot be directly measured by finding the ΔH for related (preceding) reactions, and then using their ΔH values to calculate the ΔH of the reaction in question.

Principles

The total enthalpy change (ΔH) for all the steps in a process in the same direction is equal to the ΔH for the total process. This is true regardless of the nature of the individual steps. ΔH is independent of the reaction pathway.

Example

The reaction of carbon with oxygen (C + O) to yield carbon dioxide (CO₂) is probably the hardest to measure directly. However, the enthalpy change for the reaction of carbon monoxide (CO) with oxygen (O₂) to yield carbon dioxide (CO₂) can be calculated using Hess's law.

Key facts

  • The chemical reaction must be written correctly.
  • The states of matter must be taken into account (ΔH is different for solid, liquid, or gas).

3) Calculating bond enthalpy

Another indirect method involves the use of average bond enthalpies.

ΔH ≈ Σ(Broken bonds) - Σ(Formed bonds)

Example: ΔH for the combustion of CH₄ (methane)

  1. break all the C-H bonds;
  2. break all the O=O bonds, and
  3. form the bonds in CO₂ and in H₂O.
  4. Calculate ΔH.

Limitation: This method is the least accurate; it only works for approximate values. Preferred for gaseous states.

Accuracy and Precision in Measuring Enthalpy

Both in physics and in chemistry, measurement must be accurate and precise. This is also true in the measurement of enthalpy.

  • Accuracy refers to how close your measured ΔH is to the actual value of ΔH
  • Precision refers to how close your repeated measurements of ΔH are to each other, regardless of whether they are close to the actual value or not.

Sources of Error

  • Random Errors: Temperature readings fluctuate for a variety of reasons, including inconsistency in stirring or measuring. These errors can be reduced by taking several measurements and averaging the results.
  • Systematic Errors: Understood heat lost to the environment. A thermometer that is not calibrated correctly. Repeating the measurements will not help in these cases; the measuring equipment or setup must be improved.
  • Zero Errors: Thermometers that are not calibrated to the correct starting temperature will give false readings. These can be corrected by subtracting the false starting temperature.

Uncertainty: Every measurement has a built-in uncertainty, which is a possible range of values around the measured value, expressed as ±

For example, ΔH = −57.3 ± 0.5 kJ mol⁻¹

Practical Considerations

When measuring enthalpy in the laboratory, we must:

  • Use an insulated calorimeter to reduce the loss of heat to the environment.
  • Temperatures must be evenly mixed in the solution, and the mass and concentration of the solution must be measured exactly to give the correct values of moles.
  • To check the precision of the measurement, the experiment must be repeated.

Units and Conversions

  • ΔH values are expressed in kJ mol⁻¹
  • Mass must be measured in grams and temperature in degrees Celsius or Kelvin, and moles must be measured in mol
  • For calculations, you may need to convert joules to kilojoules, where 1000 J = 1 kJ

Example Calculations

Neutralization Reaction:

50.0 cm³ of 1.0 mol dm⁻³ HCl reacts with 50.0 cm³ of 1.0 mol dm⁻³ NaOH. The temperature increases from 20.0°C to 25.5°C.

Step 1: Calculate q

Mass of solution = m = 100g

Specific heat capacity: c = 4.18 J g⁻¹ K⁻¹

ΔT = 25.5 - 20.0 = 5.5 K

Formula: Q = mcΔT

Q = (100 g)(4.18 J g⁻¹ K⁻¹)(5.5 K)

Q = 2299 J ≈ 2300 J

Step 2: Convert to kJ

q = 2300 J ÷ 1000 = 2.30 kJ

Step 3: Calculate ΔH per mole

The moles of HCl: = 0.050 L × 1 mol/L = 0.050 mol

ΔH = -2.30 kJ / 0.050 mol = -46.0 kJ/mol

Negative → exothermic

Why This Matters in Chemistry

Measuring enthalpy changes shows:

  • How energy flows in chemical reactions
  • Why reactions feel hot or cold
  • How fuels, batteries, and industrial reactions are designed

All of this hinges on measuring:

  • Mass of substances
  • Volume and concentration of solutions
  • Temperature changes

For example, to accomplish this you need:

  • Accurate mass for q
  • Correct concentration for moles
  • Precise temperature for ΔT

ΔH involves all of these. Errors in any of these will give wrong ΔH and wrong conclusions.

Summary

  • Enthalpy is a measure of the heat content of a system at constant pressure.
  • Exothermic processes release heat and involve a negative ΔH, while endothermic processes absorb heat and have a positive ΔH.
  • Enthalpy changes can be measured directly through calorimetry or calculated indirectly through Hess's Law or bond enthalpy.
  • In experiments, accuracy, precision, and uncertainty are important.
  • When calculating ΔH, the mass and temperature of the system, along with the number of moles involved in the reaction, should always be taken into consideration. The units of ΔH are kJ mol⁻¹.

Energy is not visible in Chemistry. However, enthalpy changes provide a measure that allows us to appreciate it.