5.2 Calorimetry

Heat Capacity

  • Specific heat capacity (c): the quantity of thermal energy required to raise the temperature of 1g of a substance by 1 degrees Celsius.
    • SI unit: J/(g*C)

Heat

  • Thermal energy of a system is measured by the heat transferred to or from its surrounding as heat (q).
  • We can measure the change in thermal energy indirectly by measuring the temperature change of the surroundings (a direct measure of the kinetic motion of the molecules).

Calorimetry

  • <<Process of measuring energy changes during a physical or chemical change.<<
  • Calorimeter: device used to measure thermal energy changes in a physical or chemical change. Designed to mimic an isolated system.

 

The 3 Assumptions of Calorimetric Calculations

  • often sources of error in lab
    • inflated/underreported
  • Any thermal energy lost to the outside is negligible.
  • Any energy absorbed by the calorimeter is negligible.
  • All dilute solutions have the same density of water (1.0g/mL) and the same heat capacity as water (4.18 J/(g*C))

Calculating Heat Flow (q)

 Equation of Heat Flow

  • Where,
    • m = mass of the substance (mass of the water)
    • c = specific heat capacity of the substance (heat capacity of the water)
    • ΔT = temperature change experienced by the substance (temp of the water)
    • ∆T = Tfinal − Tinitial

q has to parts

  • Magnitude: the size of q, tells you how much energy is transferred
  • Direction: the sing of q, tells you if energy is. absorbed (+) or released (-)
  • In a system:
    • If q is negative, the reaction is exothermic.
    • If q is positive, the reaction is endothermic.
  • In surroundings:
    • If q is negative, the reaction is endothermic.
    • If q is positive, the reaction exothermic.
Example
  • How much heat is given off by a 50.00 g sample of copper when it cools from 80.00oC to 50.00oC. The specific heat of copper is 0.382 J/goC.