Explore the PPT 5.1 Measuring energy changes in the above slider and also the notes below. Thermodynamics considers the entire universe in two parts:
System and Surrounding
- System: Anything under observation during chemical reaction
Example: A cup of hot tea
- Surrounding: Everything apart from the system
Example: Entire space of the universe surrounding the cup of hot tea
Heat is exchanged between system and surrounding during a chemical reaction because bonds of reactants break and new bonds of products are formed.
According to the law of conservation of energy, the energy is conserved during any reaction.
Endothermic reactions: Heat is absorbed by the system. In endothermic reactions the energy required to break the bonds of reactants is greater than the energy released by the formation of new bonds in products.
Example: Thermal decomposition of calcium carbonate
Exothermic reactions: Heat is liberated by the system. In exothermic reactions the energy required to break the bonds of reactants is lesser than the energy released by the formation of new bonds in products.
Example: Burning of fuels
Heat is a form of energy measured in Joule or Kilo Joule and it is observed by change in temperature. Temperature is a measure of degree of hotness or coldness. It is measured in Kelvin or 0 Celsius.
Kelvin Scale = 0 Celsius + 273.15
Enthalpy(H): It is the stored internal energy in a chemical substance. It is not measurable however the change in enthalpy(∆H) is measurable at constant pressure.
∆H is positive when heat is absorbed by the system.
∆H is negative when heat is released from the system into the surroundings.
- Pressure 100kPa
- Concentration of 1 mol dm-3 for solutions
- All substances in their standard states
- Temp 298K
Standard Enthalpy of reaction(ΔHθ ) is the enthalpy change during a particular reaction occurred at standard conditions.
Example: Standard Enthalpy of Combustion(ΔHθC ), Standard Enthalpy of formation(ΔHθf ),
Standard Enthalpy of Combustion(ΔHθC ): It is the enthalpy change which takes place by the complete combustion of 1 mole of a substance under standard conditions in enough oxygen gas.
Example: CH4 (g) + 2O2 (g) —> CO2 (g) + 2H2O (l)
ΔHθC = 890kJ/mol
Standard Enthalpy of formation(ΔHθf ): It is the enthalpy change which takes place when 1 mole of a substance is formed in its standard state from its constituent elements in their standard states.
2C (s) + 3H2 (g) + ½O2 (g) —>C2H5OH (l)
Specific Heat Capacity: It is the amount of heat energy required to heat 1 gram of a substance by 10C or 1K.
q= m C ΔT where q= heat produced, m= mass of water or sample solution, ΔT is the difference in temperature of the sample heated.
Example: Water has Specific Heat Capacity of 4.18 J g– K– which means 1 g water needs 4.18 J g– K– energy to heat by 1K.
In all the calorimetry experiments, we assume that aqueous solutions have same density and specific Heat Capacity as that of water.
It is the technique in which a substance is burned and the produced heat energy is used to heat the sample of water. We record the initial and final temperature and calculate the difference.
ΔT = Tf -Ti where Tf is the final temperature of the sample of water or solution while Ti is the initial temperature of the sample of water or solution.
The following equation is used to calculate the heat produced by the combustion of a substance.
- q= m C ΔT where q= heat produced, m= mass of water or sample solution, ΔT is the difference in temperature of the sample heated and C is the specific heat capacity.
You can calculate the enthalpy of combustion by using the following equation.
ΔH=- where q= heat produced and n= moles of the substance combusted.
Example: In an experiment if 100 g water is heated by combustion of 5 g of ethanol to raise the temperature from 25 0C to 45 0C. To calculate the enthalpy of combustion, we use the following steps.
ΔT= Tf -Ti = 45 – 25 = 20 0C
q= m C ΔT
q= 100 X 4.18 X 25
q= 10450 J
q= 10450/1000 = 10.45 KJ
n= 5g/46gmol– = 0.109 mole
ΔH = 10.45/0.109
= 95.87 KJmol–
I hope you have read PPT 5.1 Measuring energy changes.