1. UNIT 11: CHEMICAL REACTIONS
A. THERMOCHEMISTRY
B. CHEMICAL KINETICS

2. ENERGY OF CHEMICAL REACTIONS
THERMOCHEMISTRY=
ENERGY OF CHEMICAL REACTIONS

3. ENERGY AND REACTIONS
Exothermic reactions
Release heat energy
The surroundings become warmer
E.g. Combustion
Endothermic reactions
Absorb heat energy
The surroundings become cooler
E.g. Activating cold packs
***Energy is measured in joules (J) or calories (cal)

4. surroundings

5. Bond Energy in chemical reactions
Breaking bonds is endothermic---energy is required
Making bonds is exothermic---energy is released

6. For example: Exothermic: heat is a product A + B → C + energy/heat
A + B → C kJ
CH4(g)+2O2(g)  CO2(g) + 2H2O(l) kJ
CH4(g)+2O2(g)  CO2(g) + 2H2O(l) ΔH = kJ
Endothermic: heat is a reactant
A + B + energy/heat → C
A + B kJ → C
27 kJ + NH4NO3(s)  NH4+(aq) + NO3-(aq)
NH4NO3(s)  NH4+ (aq) + NO3- (aq) ΔH = +27 kJ

7. ENTHALPY (H) Enthalpy--a measure of heat content of a chemical system
Depends on temperature and mass
H = change in heat content that accompanies a process
Exothermic: loss of heat by system
ΔH products have less energy
Endothermic: gain of heat by system
+ΔH products have more energy

8. POTENTIAL ENERGY DIAGRAMS
Show relationship between time and energy during a chemical reaction.
Terms associated with P. E. diagrams:
activation energy
activated complex
endothermic/exothermic reactions
forward/reverse reactions

9. Activation Energy
The activation energy (Ea) is the minimum energy needed for a reaction to take place upon proper collision of reactants.
activated
complex
The activated complex: exists while old bonds are breaking and new bonds are being formed.

10. - low Ea = fast reaction rate - takes less energy for the
Activation Energy:
- depends on reactants
- is always positive
- low Ea = fast reaction rate
- takes less energy for the
reaction to start. Ea

11. Potential Energy Diagrams
Show relationship between time and energy during the course of a chemical reaction.
Activated Complex (transition state) Ea
Ea(reverse)
Energy (E) in kJ/mol
Reactants
- - DH
Products
Forward Rxn
(exothermic)
Reverse Rxn
(endothermic)
Course of Reaction (time)

12. Practice #1
For the energy diagrams provided, label the reactants, products, DH, and Ea. Also, determine the values of DH for the forward reaction and the value of Ea.
DHforward = Eproducts – Ereactants 80 60 40 20
-20
Energy (kJ/mol)
products Ea DH
reactants
Forward Reverse

13. Practice #1 Continued
* DHforward = Eproducts – Ereactants = 55 kJ/mol – (-20 kJ/mol)
= 75 kJ/mol
What are 2 ways to determine from a potential energy diagram if a reaction is endothermic or exothermic? 80 60 40 20
-20
Energy (kJ/mol)
products Ea DH
reactants
Forward Reverse

14. Endothermic Reaction A reaction in which heat is absorbed
Products have higher potential energy than the reactants.
The pink curve shows the uncatalyzed reaction. The blue curve shows what happens when a catalyst is present. A catalyst lowers the activation energy and the reaction proceeds at a faster rate.
The energies and amounts of the products and reactants stays the same, and the DE stays the same. The catalyst just allows the reaction to reach equilibrium faster.
Energy
Course of Reaction

15. Exothermic Reactions Reactions in which heat is released.
Products have lower potential energy than the reactants.
The blue curve shows the uncatalyzed reaction. The red curve shows what happens when a catalyst is present. A catalyst lowers the activation energy and the reaction proceeds at a faster rate.
Again, nothing changes but the amount of time it takes for the reaction to reach equilibrium.
Exothermic rxns are referred to as “spontaneous” because they can proceed to products without outside intervention.
Energy
Course of Reaction

16. Hrxn = Hfinal – Hinitial Hrxn = Hproducts – Hreactants
Calculating Hrxn
Hrxn = Hfinal – Hinitial
Hrxn = Hproducts – Hreactants

17. Practice Calculate Hrxn for the following reaction.
CaCO3(s) → CaO(s) CO2(g)
Hreactants = kJ
Hproducts = kJ
Hrxn = kJ
Is it endothermic or exothermic?
endothermic

18. Find Hrxn for the following reaction if the enthalpy of the products is -966 kJ and the enthalpy of reactants is -75 kJ.
CH4(g) + 2O2(g) → CO2(g) H2O(l)
Hrxn = kJ
It is exothermic

19. HEATS OF FORMATION
When compounds are formed, either from the elements that make them up or in a reaction, energy accompanies the formation.
This energy is known as the heat of formation
Some compounds require energy and the heat of formation is a positive value. It is an endothermic reaction.
Some compounds release energy and the heat of formation is a negative value. It is an exothermic reaction.
Only compounds have heats of formation. Pure elements are not made from anything and therefore have no heats of formation.

20. Find the ΔHrxn for the following reaction
2H2S(g) O2(g) → 2H2O(l) SO2(g)
ΔHf H2S(g) = kJ/mol
ΔHf H2O(l) = kJ/mol
ΔHf SO2(g) = ─296.8 kJ/mol
Hrxn = kJ

21. B. REACTION RATES (OTHERWISE KNOWN AS CHEMICAL KINETICS)

22. WHAT FACTORS DRIVE A CHEMICAL REACTION?
How does a chemical reaction take place? Collision theory is used to explain that in order for a reaction to occur, particles MUST collide (this is reasonable), but not all collisions result in the formation of a new product

23. Collision theory
States that atoms, molecules, ions must collide in order to react
Must collide with:
a. correct orientation
b. enough energy to form activated complex or transition state

24. A. Correct orientation

25. B. Enough energy to break bonds and form activated complex19

26. REMEMBER----- A reaction won’t happen if:
Insufficient energy to break bonds.
N O2
N O2
Molecules are not aligned correctly. 19

28. The study of reaction rates (speed)
Chemical kinetics
The study of reaction rates (speed) 17

29. Fast: Slow: Oxidation: Paper burning Oxidation: Nails rusting
REACTION RATE IS DEFINED AS:
Speed at which reactant is used up.
Speed at which product forms.
Fast:
Oxidation: Paper burning
Slow:
Oxidation: Nails rusting
Paper turning yellow

30. Factors that affect reaction rates
Concentration
Surface area
Temperature
Catalyst
Nature of the reactants

31. Concentration - high concentration = fast reaction rate
- more opportunities for collisions because there are more particles in the same volume that can react.
There are less red particles in the same volume so there is less chance of a collision
There are more red particles in the same volume so there is more chance of a collision so the reaction goes faster

32. Surface area high surface area = fast reaction rate
more opportunities for collisions
Increase surface area by…
using smaller particles – if we make the pieces of the reactants smaller, we increase the number of particles on the surface which can react.
dissolving in water – gases & dissolved particles can mix & collide freely. Reactions happen rapidly.

33. Temperature - high temperature = fast reaction rate - high KE
- when we increase the temperature, we give the
particles energy
- this makes the particles move faster
- so there are more opportunities for collision
- it is easier to reach activation energy

34. Catalyst
substance that increases reaction rate without being consumed in the reaction
lowers the activation energy

35. Nature of the reactants
reactant structure(polar vs. nonpolar)
physical state of reactants ---ionic compounds dissolve faster than covalent compounds in water (salt dissolves faster than sugar)
reactions occur faster in solutions or if reactants are gaseous
more active elements will react more vigorously. Cs is more reactive than Na so it will react faster and more vigorously than Na when placed in water. 25