Thermochemistry is the study of heat changes that take place in a change of state or a chemical reaction  -  heat energy is either absorbed or released.  If a process releases energy in the form of heat, the process is called exothermic.  A process that absorbs heat is called endothermic. 

 

Heat is defined as the energy transferred from one object to another due to a difference in temperature.  We do not observe or measure heat directly  -  we measure the temperature change that accompanies heat transfer.  In a chemical reaction it is often not possible to measure the temperature of the reactants or products themselves.  Instead, we measure the temperature change of  the surroundings.

 

The difference between the system and the surroundings is a key concept in thermochemistry.  The system consists of the reactants and the products of the reaction.  The solvent, the container, the atmosphere about the reaction (in other words, the rest of the universe) are considered the surroundings.  Heat may be transferred from the system to the surroundings (the temperature of the surroundings will increase) or from the surroundings to the system (the temperature of the surroundings will decrease).

 

When a system releases heat to the surroundings during a reaction, the temperature of  the surroundings increases and the reaction container feels warm to the touch.  This is an exothermic reaction.  Heat flows out of the system.  An example of an exothermic reaction is the combustion of propane (C₃H₈) in a barbecue grill to produce carbon dioxide, water, and heat.  Note that heat appear in the product side of the equation:

 

          C₃H₈ (g)  +  5 O₂ (g)  à  3CO₂ (g)  +  4H₂O (g)  +  heat

 

When a system absorbs heat from the surroundings during a reaction, the temperature of the surroundings decreases and the reaction container feels cold to the touch.  This in an endothermic reaction.  Heat flows into the system.  A familiar example of an endothermic process is the melting of ice.  Solid water (ice) needs heat energy to break the forces holding the molecules together in the solid state.

 

The heat absorbed by water in turning it to steam has been used to break apart the forces (i.e., hydrogen bonding) between water molecules in the liquid phase.  The amount of heat that must be absorbed to vaporize a specific quantity of liquid (usually one gram or one mole) is called the heat of vaporization.  In a similar manner, heat is also required to melt ice.  The amount of heat that must be absorbed to melt a specific quantity of solid is called the heat of fusion.

 

The amount of heat transferred in these processes depends on the difference in the energy stored in each substance.  This stored energy is called the heat content or enthalpy, and is represented by the symbol H.. The enthalpy change ( ΔH) for a physical process or a chemical reaction is defined as the heat change that occurs at a constant pressure.  This is convenient, because most of the reactions that are carried out in the lab are in flasks or containers that are open to the atmosphere – that is, they take place at a constant pressure equal to the barometric pressure.

 

Equation 1 shows the equality between the change in enthalpy (ΔH) of a system and the amount of heat transferred, symbolized by q_p, for a reaction carried out at a constant pressure.

ΔH = q_p       Equation 1

 

The amount of heat (q_p) transferred to a substance or object depends on three factors:  the mass (m) of the object, its specific heat (c_p) and the resulting temperature change (Δt)

q_p  =  m • Δt • c_p     Equation 2

 

The specific heat (c_p) of a substance reflects its ability to absorb heat energy and is defined as the amount of heat needed to raise the temperature of one gram of a substance by one degree Celsius.  The specific heat of water is equal to 4.185 J/gºC.  In most laboratory situations, the temperature change is measured not for the system itself (the reactants and products) but for the surroundings (the solution and reaction vessel).  The amount of heat released by the system must be equal to the amount of heat absorbed by the surroundings.  The sign convention in Equation 3 reveals that the heat change occurs in the opposite direction.

q_(system)   =  - q_{(surroundings)}

 

For an exothermic reaction, the heat released by the system results in a temperature increase for the surroundings (Δt is positive) and the heat absorbed by the surroundings will be a positive quantity.  The heat released by the system must have the reverse sign – it must be a negative quantity.  According to this convention, the enthalpy change for an exothermic reaction is always a negative value.  For an endothermic reaction, in contrast, the heat absorbed by the system results in a temperature decrease for the surroundings (Δt) is negative) and the heat released by the soundings will be a negative quantity.  The heat absorbed by the system must have the opposite sign  -  it must be a positive quantity.  According to this convention, the enthalpy change for an endothermic reaction is always a positive value.

 

The energy or enthalpy change associated with the process of a solute dissolving in a solvent is called the heat of solution (ΔH_soln).  In the case of an ionic compound dissolving in water, the overall energy change is the net result of two processes  -  the energy required to break the attractive forces (ionic bonds) between the ions in the  crystal lattice, and the energy released when the dissociated (free) ions form ion-dipole attractive forces with the water molecules.

 

Heats of solution and other enthalpy changes are generally measured in an insulated vessel called a calorimeter that reduces or prevents heat loss to the atmosphere outside the reaction vessel.  The process of a solute dissolving in water may either release heat into the aqueous solution or absorb heat from the solution, the amount of heat exchange between the calorimeter and the outside surroundings should be minimal.  When using a calorimeter, the reagents being studied are mixed directly in the calorimeter and the temperature is recorded both before and after the reaction ahs occurred.  The amount of heat change occurring in the calorimeter may be calculated using the equation:  q_p  =  m • Δt • c_p     The specific heat of the solution is generally assumed to be the same as that of the water, namely  4.185 J/gºC.

 

The mass of grams of solute is usually the independent variable and will be varied in different trials.  The temperature change that is produced depends on the mass of the solute and is thus the dependent variable in a calorimetry experiment.