IB Chemistry HL Topic 5

Standard enthalpy change is the heat energy transferred under standard conditions—pressure 101.3 kPa, temperature 298 K. Only ∆H can be measured, not H for the initial or final state of a system.
Exothermic: A process that looses heat to the surroundings.
Endothermic: A process that absorbs heat from the surroundings.
Enthalpy of reaction: The amount of heat taken in or given out when one mol of substance is reacted in standard conditions.

State.

In an exothermic reaction, heat is given out to the environment and the temperature of the system increases. Endothermic reaction is one in which heat is taken in from environment and the temperature of the system decreases. The reaction mixture is called the system and anything around the system is called the surrounding or environment.

• Exothermic -> a reaction which produces heat (ΔH has a negative value by convention, -ve)
• Endothermic -> a reaction which absorbs heat (ΔH has a positive value by convention, +ve)

See notes.
The most stable state is where all energy has been released. When going to a more stable state, energy will be released, and when going to a less stable state, energy will be gained (from the surroundings). On an enthalpy level diagram, higher positions will be less stable (with more internal energy) therefore, if the product is lower, heat is released (more stable, ΔH is -ve) but if it is higher, heat is gained (less stable, ΔH is +ve).

• Formation of bonds causes an energy release (exothermic).
• Breaking of bonds requires energy (endothermic).

Students should be able to calculate the heat energy change for a substance given the mass, specific heat capacity and temperature change using q = mcΔT.

Students should consider reactions in aqueous solution and combustion reactions. Use of the bomb calorimeter and calibration of calorimeters will not be assessed.
Note that if a calorimeter is used that also absorbs energy, then its mass, m and specific heat, c (or heat capacity, m x c) must also be known: ΔHsol = heat absorbed by solution + heat absorbed by calorimeter = (msolution × csol × ΔTsol) + (mflask × cflask × ΔTflask)

A known mass of solution should be placed in a container, as insulated as possible, to prevent as much heat as possible from escaping. The temperature is measured continuously, the value used in the equation is the maximum change in temperature from the initial reading.The result will be a change in temperature. This can be converted into a change in heat (or energy) by using the above equation E = m x c x ΔT.Δ-H may then be calculated for the amount of reactants present, and then this can be used to calculate for a given number of mols.

Students should be aware of the assumptions made and errors due to heat loss.

Students should be able to use simple enthalpy cycles and enthalpy level diagrams and to manipulate equations. Students will not be required to state Hess’s law.
Hess law: If you add two or more thermochemical equations to give a final equation, then you can also add the heat changes to give the final heat change. The enthalpy of reaction will be the same independent of path.
According to Hess's Law, ΔH of a reaction is independent of any intermediate steps. Hess’s law is a special case of the law of conservation of energy.

Bond enthalpy: The energy required to break a bond in standard conditions.
Bond Enthalpy is the energy required when one mole of bonds is broken in the gaseous state (bond enthalpies are valid only in the gaseous state). Bond breaking needs energy and is an endothermic process. Bond making releases energy and it is an exothermic process. Average bond energy is the average energy required to break a mole of the same type of bonds in the gaseous state in a variety of similar compounds. If bond enthalpies are given then:
ΔHorxn = Σ BEbonds broken - ΣBEbonds made (not products – reactants)

Bond breaking needs energy and is an endothermic process. Bond making releases energy and it is an exothermic process.