1.4 Energetics Flashcards
(15 cards)
KEY DEFINITION: enthalpy change
The heat energy change at a constant pressure
Give 2 examples of important exothermic reactions:
- combustion of fuels
- respiration
Give 2 examples of important endothermic reactions:
- thermal decomposition of calcium carbonate
- photosynthesis
KEY DEFINITION: activation energy
The minimum energy required to start a reaction by the breaking of bonds
What is bond enthalpy?
The heat energy required to break one mole of a given covalent bond in the molecules in the gaseous state
Why are bond enthalpies always positive?
The bond breaking requires energy (endothermic)
KEY DEFINITION: mean bond enthalpy
The heat energy required to break one mole of a covalent bond, averaged for that type of bond in a range of different compounds
Why do bond energy calculations give less accurate values than Hess’s law?
- we assume all bonds are in the gaseous state
- values used are only averages taken across a range of compounds (not specific to the given molecule)
KEY DEFINITION: standard enthalpy of formation
The enthalpy change that occurs when one mole of a compound is formed from its constituent elements with all reactants and products in their standard states
What are the conditions for standard enthalpy?
- temperature: 298K
- pressure: 1 atmosphere (100KPa)
KEY DEFINTION: standard enthalpy of combustion
The enthalpy change that occurs when one mole of a compound reacts completely in oxygen with all reactants and products in their standard states
What is the equation for heat energy change?
q = mc(deltaT)
Why would an experimental enthalpy value be different from the data book value?
- heat loss to the surroundings
- incomplete combustion
KEY DEFINITION: Hess’s Law
The enthalpy change for a chemical reaction is independent of the route taken
Why isn’t it always possible to measure the enthalpy change directly?
- the reaction is very slow/fast, difficult to measure
- various products
