IB CHEMISTRY - Topic 5 Energetics Notes

IB CHEMISTRY:
TOPIC 5 ENERGETICS NOTES

5.1 Measuring energy changes


Heat and Temperature
 Energetics is the study of heat changes in chemical reactions
o Heat is a form of energy
 Law of conservation of energy: Energy can neither be created nor destroyed. Total
energy is always conserved and can only be transferred
 Basic law of thermodynamics: Heat energy always flows from a higher temperature
object to a lower temperature object
 The Kelvin scale is based on kinetic energy , so 0K means that there is no kinetic
movement at all


Definitions

Heat – A measure of the total kinetic energy of particles in a substance
Temperature – A measure of the average kinetic energy of particles in a substance




Enthalpy (H)
 Enthalpy (H): The amount of heat energy contained in a substance
 Enthalpy is stored in the chemical bonds as potential energy
 When substances react, the total enthalpy of a system cannot be measured (due to loss
of heat), but it is possible to measure the difference
in the enthalpy between the reactants and products
 Enthalpy is denoted as H, however heat change is
denoted as ΔH
 The enthalpy change for chemical reactions is
denoted kJ mol-1
 The reaction mixture is called the system (the
chemical reaction), which gives heat to or takes
heat from the surroundings (anything around the system)
 Chemical reactions that involve transfer of heat between system and the surroundings
are described as exothermic and endothermic


Exothermic: Heat is Released (forming bonds)
 In exothermic reactions heat is released to the surroundings
o This is because more heat energy is released than what is added

, o So, the overall heat energy is released from the system, causing the
surroundings to become hotter
 Exothermic reactions have negative ΔH values, because heat is released (thus enthalpy
decreases, -ΔH)
 In an exothermic reaction, the products are more stable than the reactants as they
have a lower enthalpy (as the heat has been released).
o Less heat means more stable
 This means the reaction is downhill in terms of heat energy (hence the negative deltaH)
 Exothermic reactions release energy (as heat)
 Examples include:
o Bond forming: Removing heat brings atoms closer together, forming bonds
o When chemical bonds are formed, heat is released (See 5.3)
o Gas -> Liquid -> Solid: Heat is removed, thus these are exothermic reactions
o Rain: The condensation of water vapor into rain releases energy in the form of
heat
o Combustion: The burning of carbon compounds uses oxygen from air, and
produces CO2, H2O and lots of heat

Endothermic: Heat is Absorbed – system gets cooler (bond breaking)
 In endothermic reactions heat is absorbed from the surroundings
o This is because more heat energy is added than what is released
o So, the overall heat energy is absorbed by the system, causing the surroundings
to become cooler.
 Endothermic reactions have positive ΔH values, because heat is absorbed (thus,
enthalpy increases +ΔH)
 In an endothermic reaction, the products are less stable than the reactants as they
have a higher enthalpy
o More heat means less stable
 This means the reaction is uphill in terms of heat energy
 Exothermic reactions require energy (through heat)
 Examples include:
o Bond breaking: Adding heat separates atoms, breaks bonds (See 5.3)
o Photosynthesis: Plants absorb heat energy from sunlight to convert CO2 and
water into glucose and oxygen
o Solid -> Liquid -> Gas: Heat is added, so the reactions are endothermic

Energy Diagrams: Endothermic and Exothermic Reaction

Enthalpy
Enthalpy



Reaction Pathway (EXO) Reaction Pathway (ENDO)




Standard Enthalpy change notation: ∆ H θx