Calorimetry

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[[fr:Calorim�trie]]

Calorimetry is the science of measuring the heat of chemical reactions or physical changes. Calorimetry involves the use of a calorimeter.

Contents 1 Heat 2 Work 3 Internal energy 4 Enthalpy 5 Constant-volume 6 Constant-pressure 6.1 References

Heat

Main article: Heat.

Heat is an amount of energy which is usually linked with a change in temperature or in a change in phase of matter. The SI unit for heat is the joule.

The equation for measuring heat is:

<math>q = C \Delta t</math>

• q is heat. When heat transfers energy from the system to the surroundings, the symbol is -q. When heat transfers energy from the surroundings to the system, the symbol is +q.
• C is heat capacity.
• Δt is change in temperature.

Work

Main article: Work.

Work is the energy transferred in applying force over a distance. In calorimetry, the force is generally pressure and instead of distance, volume is used. Work is given by the formula:

<math>w = -P \Delta V</math>

• w is work. When work transfers energy from the system to the surroundings, the symbol is -w. When work transfers energy from the surroundings to the system, the symbol is +w.
• P is the pressure of the system.
• ΔV is volume change.

Internal energy

Main article: Internal energy.

Internal energy is the kinetic energy associated with the motion of molecules, and the potential energy associated with the rotational, vibrational, and electric energy of atoms within molecules. Internal energy is a quantifiable state function of a system.

Internal energy can not be measured directly; it is only measured as a change (ΔU). The equation for change in internal energy is:

<math>\Delta U = q + w</math>

Enthalpy

Main article: Enthalpy.

Enthalpy is the sum of the internal energy of matter and the product of its volume multiplied by the pressure.

Enthalpy is defined by the following equation:

<math>H = U + PV</math>

• U is the internal energy.
• V is the volume.

The total enthalpy of a system cannot be measured directly; the enthalpy change of a system is measured instead. Enthalpy change is defined by the following equation:

<math>\Delta H = H_{final} - H_{initial}</math>

• ΔH is enthalpy change.
• H_final is the final enthalpy of the system. In a chemical reaction, H_final is the enthalpy of the products.
• H_initial is the initial enthalpy of the system. In a chemical reaction, H_initial is the enthalpy of the reactants.

Constant-volume

Constant-volume calorimetry is calorimetry performed at a constant volume. This involves the use of a constant-volume calorimeter.

No work is performed in constant-volume calorimetry, so the heat measured equals the change in internal energy of the system. The equation for constant-volume calorimetry is:

<math>q = C \Delta t = \Delta U</math>

Since in constant-volume calorimetry pressure is not kept constant, the heat measured does not represent the enthalpy change.

Constant-pressure

Constant-pressure calorimetry is calorimetry performed at a constant pressure. This involves the use of a constant-pressure calorimeter.

The heat measured equals the change in internal energy of the system minus the work performed:

<math> q = \Delta U - w </math>

Since in constant-pressure calorimetry, pressure is kept constant, the heat measured represents the enthalpy change:

<math>q = \Delta H = H_{final} - H_{initial}</math>

References

• Adapted from the Wikipedia article, "Calorimetry" http://en.wikipedia.org/wiki/Calorimetry, used under the GNU Free Documentation License