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
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
