Calorimetry

Calorimetry is the science of measuring the heat of chemical reactions or physical changes. Calorimetry is performed with a calorimeter. The word calorimetry is derived from the Latin word calor, meaning heat.  Calorimetric techniques are therefore based on the measurement of heat that may be generated (exothermic process), consumed (endothermic process) or simply dissipated by a sample.

To determine the energy change of a system, it is often easier to look at what is happening to the surrounding as it is easier to measure. Three factors will influence the amount of heat transferred  (q) between the system and surrounding during a physical, chemical, or nuclear change – mass (m), temperature change of the surrounding (∆T), and the specific heat capacity (c) of the surrounding.

     q = mc∆T

specific heat capacity (c): quantity of heat required to raise the temperature of a unit mass of a substance
1°C or 1K
Specific Heat Capacity Table
Substance	Specific Heat Capacity  
at 25oC in J/goC	Substance	Specific Heat Capacity  
at 25oC in J/goC	
H2 gas	14.267	calcium	0.650	
He gas	5.300	iron	0.444	
H2O(l)	4.184	nickel	0.440	
ice @ 0oC	2.010	zinc	0.39	
steam @ 100oC	2.010	copper	0.385	
vegetable oil	2.000	brass	0.380	
air	1.020	sand	0.290	
magnesium	1.020	silver	0.240	
aluminum	0.900	tin	0.21	
Concrete	0.880	lead	0.160	
glass	0.840	mercury	0.14	
potassium	0.75	gold	0.129	
sulphur	0.73			

Calorimeter
A calorimeter is a device used to measure heat of reaction. It can be sophisticated and expensive or simple and cheep. In our labs, a Styrofoam cup is used as a calorimeter, because it is a container with good insulated walls to prevent heat exchange with the environment. In order to measure heats of reactions, we often enclose reactants in a calorimeter, initiate the reaction, and measure the temperature difference before and after the reaction. The temperature difference enables us to evaluate the heat released in the reaction.

For our calculations around the calorimeter three assumptions will be used:
•	all reactions will be assumed to be conducted as an isolated system (all of the energy transfer will occur between the system and surrounding, with none of the energy lost to the universe)
•	no significant energy will be lost or gained from the calorimeter
•	for dilute solutions (concentration < 1 mol / L) it is assumed that they have the density and specific heat capacity of pure water (c = 4.184 J/goC and d = 1.000 g / ml)

Sample Calculations:
When 600 mL of water in an electric kettle is heated from 20°C to 85°C to make a cup of tea, how much heat flows into the water?

First, use the density formula to calculate the mass of water.
m = dV
     = 1.00 g/m_L x 600 _mL   = 600 g

Use the heat formula, q = mc∆T, to calculate the quantity of heat transferred.
q  = ?		m = 600 g	c = 4.18 J/(g•°C)      ∆T = 85°C - 20°C = 65°C

q = mc∆T
   = (600 g)(4.184 J/goC)(65oC)
   = 1.63 x 105 J or 163 kJ.

163 kJ of heat flows into the water

What would the final temperature be if 250.0 J of heat were transferred into 10.0 g of  methanol initially at 20.0°C?

m = 10.0 g
c = 2.918 J/(g•°C)
T1 = 20.0 °C
∆T= T2 – T1 = T2 – 20.0°C
q = 250 J
  q = mc∆T								T2 = ∆T + T1
∆T = q /mc								T2 = 20°C + 8.57°C
     =(250 J)/ (10.0 g x 2.918 J/(g•°C))					T2 = 28.6°C
     = 8.57°C

The final temperature of the methanol is 28.6°C.