ENTHALPY, ENTROPY
AND GIBBS LAW OF
FREE ENERGY
Dr Nadeem Asad
2-8-2014
Energy reactions
Review the Energy Diagram
Endothermic have a high startup energy
Exothermic have a low start up energy
3
Heat and Temperature
• Heat is energy that is transferred from one
object to another due to a difference in
temperature
• Temperature is a measure of the average
kinetic energy of a body
• Heat is always transferred from objects at a
higher temperature to those at a lower
temperature
3
Energy Curve
• Activation energy: energy required to get the reaction to
move forward
• Energy released or absorbed is noted at the end of the
curve
Factors that effect the Reaction Rate
• Temperature: Endothermic vs Exothermic reactions
• Concentration: Increase the reactants will increase the
products
• Surface Area: Smaller particles have large surface area
• Catalysts and inhibitors
Enthalpy
Delta H = Hproducts - Hreactants
Delta H represents the transfer of
heat
Exothermic vs endothermic reactions
Exothermic Reaction: A process that releases heat to its
surroundings. Products have less energy than the
reactants
Endothermic Reaction : A process that absorbs heat from
the surroundings. Products have more energy than the
reactants.
Exothermic process is any process that gives off heat –
transfers thermal energy from the system to the surroundings.
2H2 (g) + O2 (g)
H2O (g)
2H2O (l) + energy
H2O (l) + energy
Endothermic process is any process in which heat has to be
supplied to the system from the surroundings.
energy + 2HgO (s)
energy + H2O (s)
2Hg (l) + O2 (g)
H2O (l)
6.2
Combustion and neutralization processes
Combustion
Exothermic reaction
General Combustion Reaction Formula:
Compound (usually hydrocarbon) + O2
 CO2 + H2O + energy
CH4 + 2O2  CO2 + 2H2O + 890kJ
∆H = -890kJ
Neutralization
Exothermic reaction
Acid + Base  Salt + Water + energy
HCl + NaOH  NaCl + H2O + 57.3 kJ
∆H = -57.3kJ
Activation energy and enthalpy diagrams
Exothermic Reactions
Products more stable than reactants
(lower energy).
ΔH = Hproducts – Hreactants
Since the products have less energy
than the reactants, the ΔH value is
negative.
Endothermic Reactions
Products less stable than reactants
(higher energy)
ΔH = Hproducts – Hreactants
Since the products have more energy
than the reactants, the ΔH value is
positive.
Enthalpy (H) is used to quantify the heat flow into or out of a
system in a process that occurs at constant pressure.
DH = H (products) – H (reactants)
DH = heat given off or absorbed during a reaction at constant pressure
Hproducts < Hreactants
DH < 0
Hproducts > Hreactants
DH > 0
6.4
Standard enthalpy of reaction
Standard Enthalpy Change of Reaction (∆H): The heat
energy exchanged with the surroundings when a reaction
happens under standard conditions (NOT STP… see
below).
Since the enthalpy change for any given reaction will vary
with the conditions, esp. concentration of chemicals, ΔH
are measured under standard conditions:
•
•
•
•
pressure = 101.3 kPa
temperature = 25ºC = 298 K
Concentrations of 1 mol dm-3
The most thermodynamically stable allotrope (which in the case of
carbon is graphite)
Only ΔH can be measured, not H for the initial or final
state of a system.
Terminology of rate of reaction (Δho)
Pseudonyms (other names) for DH
 Heat of Reaction: DHrxn heat produced in a chemical reaction
 Heat of Combustion: DHcomb heat produced by a combustion reaction
 Heat of Neutralization: heat produced in a neutralization reaction (when an
acid and base are mixed to get water, pH = 7)
 Heat of solution: DHsol heat produced by when something dissolves
 Heat of Fusion: DHfus heat produced when something melts
 Heat of Vaporization: DHvap heat produced when something evaporates
 Heat of Sublimation: DHsub heat produced when something sublimes
 Heat of formation: DHf change in enthalpy that accompanies the formation
of 1 mole of compound from it’s elements (this has special uses in
chemistry…)
Spontaneous Reactions
• If the reaction moves forward as written without an
intervention then the reaction is said to be spontaneous.
• Remember Delta H (enthalpy)
• Negative signs represent exothermic
• Always spontaneous
• Positive sign represents endothermic
• Sometimes spontaneous
Examples
H2O (s) --> H2O (l) delta H = +6kJ
2Na(s) + Cl2 --> 2NaCl (s) delta H = -822kJ
The above reactions are all spontaneous. But most
endothermic rxn are Nonspontaneous.
Delta H
• Most enthalpy reactions that are endothermic rxn are not
spontaneous. This is because the reaction has a high
energy barrier.
• Only because of a large difference between entropy and
enthalpy can the reaction happen.
Average bond enthalpy
Enthalpy changes of reactions are the result of bonds breaking and
new bonds being formed. Remember…
•
Breaking bonds requires energy
•
Forming new bonds releases energy
Bond enthalpy is the energy required to break one mole of a certain
type of bond in the gaseous state averaged across a variety of
compounds.
FYI: Bond enthalpies for unlike atoms will be affected by surrounding
bonds and will be slightly different in different compounds so
average bond enthalpies are used.
• DH =
∑ (energy required
to break bonds)
OR
–
∑(energy released
when bonds are formed)
–
∑(bond enthalpy
of products)
• DH =
∑ (bond enthalpy
of reactants)
19
Energy Changes in endothermic and
exothermic processes
In an
endothermic
reaction there is
more energy
required to
break bonds
than is released
when bonds are
formed.
The opposite is
true in an
exothermic
reaction.
19
Exothermic and endothermic reactions in light of average bond enthalpies
If the amount of energy required to break the bonds in the reactants
is greater than the amount of energy released when bonds are
formed in the products, the reaction is endothermic.
average bond enthalpy reactants > average bond enthalpy products
If the amount of energy required to break the bonds in the reactants
is less than the amount of energy released when bonds are
formed in the products, the reaction is exothermic.
average bond enthalpy reactants < average bond enthalpy products
Calorimetry
• Calorimetry involves the measurement of heat
changes that occur in chemical processes or
reactions. Determines
the ΔH by measuring temp Δ's created from the rxn
• The heat change that occurs when a
substance absorbs or releases energy is
really a function of three quantities:
• The mass
• The temperature change
• The heat capacity of the material
21
Entropy
• Entropy is the measurement of disorder of particles in a
reaction.
Equation: Delta S = SProduct - Sreactant
*not easy to measure, but we can compare based on the
states of matter
Comparing Entropy
• Sproduct > Sreactants delta S is positive
• Sproduct< Sreactants delta S is negative
• Solids have a very low entropy
• Liquids have a slightly higher entropy
• Gases have a very high entropy
Example reactions
• CO2 (s) --> CO2 (g) S > 0
• H2O (g) --> H2O (s) S<0
• 2NH3(g) --> N2(g) + 2 H2 (g) S>0
Law of Thermodynamics
• States: in any spontaneous process, the overal entropy of
the universe always increases.
Suniverse = Sreaction + Ssurroundings
Law of Thermodynamics
• Sreaction - calculated using the states of matter
• Ssurround - calculated based on enthalpy
• If the reaction is endothermic it is taking energy away from the
surrounding making it negative
• If the reaction is exothermic it is placing energy into the system and
making it positive.
Gibbs Free Energy
• Delta G = Delta H - T deltaS
Compares entropy, enthalpy and temperature
* If delta G is negative the reaction is
spontaneous
* If delta G is positive the reaction is
nonspontaneous
* If delta G is 0 the reaction is at equilibrium