1. Enthalpy, Entropy and Gibbs Law of Free Energy
Dr Nadeem Asad

2. 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

4. 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

5. 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

6. Delta H represents the transfer of heat
Enthalpy
Delta H = Hproducts - Hreactants
Delta H represents the transfer of heat

7. 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.

8. 2H2 (g) + O2 (g) 2H2O (l) + energy
Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings.
2H2 (g) + O2 (g) H2O (l) + energy
H2O (g) H2O (l) + energy
Endothermic process is any process in which heat has to be supplied to the system from the surroundings.
energy + H2O (s) H2O (l)
energy + 2HgO (s) Hg (l) + O2 (g)
6.2

9. Combustion and neutralization processes
Exothermic reaction
General Combustion Reaction Formula: Compound (usually hydrocarbon) + O2  CO2 + H2O + energy
CH O2  CO H2O + 890kJ
∆H = -890kJ
Neutralization
Acid + Base  Salt + Water + energy
HCl + NaOH  NaCl + H2O kJ
∆H = -57.3kJ

10. 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)
Since the products have more energy than the reactants, the ΔH value is positive. 10

11. DH = H (products) – H (reactants)
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
Hproducts > Hreactants
DH < 0
DH > 0
6.4

12. 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 = 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. 12

13. Terminology of rate of reaction (Δho)
Pseudonyms (other names) for H
Heat of Reaction: Hrxn heat produced in a chemical reaction
Heat of Combustion: Hcomb 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: Hsol heat produced by when something dissolves
Heat of Fusion: Hfus heat produced when something melts
Heat of Vaporization: Hvap heat produced when something evaporates
Heat of Sublimation: Hsub heat produced when something sublimes
Heat of formation: Hf change in enthalpy that accompanies the formation of 1 mole of compound from it’s elements (this has special uses in chemistry…)

14. 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

15. 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.

16. 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.

17. 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. 17

18. H = ∑ (energy required – ∑(energy released
to break bonds) when bonds are formed) OR
∑ (bond enthalpy – ∑(bond enthalpy
of reactants) of products)

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

20. 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 20

21. 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

22. 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

23. 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

24. Example reactions CO2 (s) --> CO2 (g) S > 0
H2O (g) --> H2O (s) S<0
2NH3(g) --> N2(g) + 2 H2 (g) S>0

25. Suniverse = Sreaction + Ssurroundings
Law of Thermodynamics
States: in any spontaneous process, the overal entropy of the universe always increases.
Suniverse = Sreaction + Ssurroundings

26. 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.

27. 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