1. The study of heat released or required by chemical reactions
THERMOCHEMISTRY
The study of heat released or required by chemical reactions
Fuel is burnt to produce energy - combustion (e.g. when fossil fuels are burnt)
CH4(g) O2(g) CO2(g) + 2H2O(l) + energy

2. Energy due to position (stored energy)
What is Energy?
Energy
Kinetic energy (EK)
Potential energy (EP)
Energy due to motion
Energy due to position (stored energy)

3. Total Energy = Kinetic Energy + Potential Energy E = EK + EP
Kinetic energy & potential energy are interchangeable
Ball thrown upwards slows & loses kinetic energy but gains potential energy
The reverse happens as it falls back to the ground

4. i.e. it is the total energy of all the atoms and molecules in a sample
Law of Conservation of Energy: the total energy of the universe is constant and can neither be created nor destroyed; it can only be transformed.
The internal energy, U, of a sample is the sum of all the kinetic and potential energies of all the atoms and molecules in a sample
i.e. it is the total energy of all the atoms and molecules in a sample

5. Systems & Surroundings
In thermodynamics, the world is divided into a system and its surroundings
A system is the part of the world we want to study (e.g. a reaction mixture in a flask)
The surroundings consist of everything else outside the system
SYSTEM
CLOSED
OPEN
ISOLATED

6. OPEN SYSTEM: can exchange both matter and energy with the surroundings (e.g. open reaction flask, rocket engine)
CLOSED SYSTEM: can exchange only energy with the surroundings (matter remains fixed) e.g. a sealed reaction flask
ISOLATED SYSTEM: can exchange neither energy nor matter with its surroundings (e.g. a thermos flask)

7. e.g. kettle heats on a gas flame
HEAT is the energy that transfers from one object to another when the two things are at different temperatures and in some kind of contact
e.g. kettle heats on a gas flame
cup of tea cools down (loses energy as heat)
Thermal motion (random molecular motion) is increased by heat energy
i.e. heat stimulates thermal motion

8. UNITS OF ENERGY S.I. unit of energy is the joule (J)
Heat and work ( energy in transit) also measured in joules
1 kJ (kilojoule) = 103 J
Calorie (cal): 1 cal is the energy needed to raise the temperature of 1g of water by 1oC
1 cal = J

9. First Law of Thermodynamics:
the internal energy of an isolated system is constant
Signs (+/-) will tell you if energy is entering or leaving a system
+ indicates energy enters a system
- indicates energy leaves a system

10. A change in internal energy can be identified with the heat supplied at constant volume
ENTHALPY (H)
(comes from Greek for “heat inside”)
the change in internal energy is not equal to the heat supplied when the system is free to change its volume
some of the energy can return to the surroundings as expansion work
 U < q

11. The heat supplied is equal to the change in another thermodynamic property called enthalpy (H)
i.e. H = q
this relation is only valid at constant pressure
As most reactions in chemistry take place at constant pressure we can say that:
A change in enthalpy = heat supplied

12. Energy & Chemical Reactions
Energy Changes
Exothermic Reactions
Endothermic Reactions

13. During a chemical reaction…
energy is used to break bonds
energy is released when new bonds are formed
breaking
bonds
making
bonds

14. EXOTHERMIC & ENDOTHERMIC REACTIONS
Exothermic process: a change (e.g. a chemical reaction) that releases heat.
A release of heat corresponds to a decrease in enthalpy
Exothermic process: H < 0 (at constant pressure)
Burning fossil fuels is an exothermic reaction

15. H2(l) + O2(l)  H2O(g) + energy
Exothermic Reaction
reaction that releases energy
energy released by making new bonds outweighs energy req’d to break old bonds
H2(l) + O2(l)  H2O(g) + energy
reaction that powers the space shuttle lift-off

16. An input of heat corresponds to an increase in enthalpy
Endothermic process: a change (e.g. a chemical reaction) that requires (or absorbs) heat.
An input of heat corresponds to an increase in enthalpy
Endothermic process: H > 0 (at constant pressure)
Forming Na+ and Cl- ions from NaCl is an endothermic process
Photosynthesis is an endothermic reaction (requires energy input from sun)

17. process used to obtain aluminum from aluminum ore
Endothermic Reaction
reaction that absorbs energy
energy req’d to break old bonds outweighs energy released by making new bonds
2Al2O3 + energy  4Al + 3O2
process used to obtain aluminum from aluminum ore

18. Measuring Heat
reaction
Exothermic reaction, heat given off & temperature of water rises
reaction
Endothermic reaction, heat taken in & temperature of water drops

19. How do we relate change in temp. to the energy transferred?
Heat capacity (J/oC) = heat supplied (J)
temperature (oC)
Heat Capacity = heat required to raise temp. of an object by 1oC
more heat is required to raise the temp. of a large sample of a substance by 1oC than is needed for a smaller sample

20. Specific Heat Capacity (Cs)
Specific heat capacity is the quantity of energy required to change the temperature of a 1g sample of something by 1oC
Specific Heat Capacity (Cs)
Heat capacity
Mass =
J / oC / g
J / oC g =

21. Vaporisation Melting Freezing
Energy has to be supplied to a liquid to enable it to overcome forces that hold molecules together
endothermic process (H positive)
Melting
Energy is supplied to a solid to enable it to vibrate more vigorously until molecules can move past each other and flow as a liquid
endothermic process (H positive)
Freezing
Liquid releases energy and allows molecules to settle into a lower energy state and form a solid
exothermic process (H negative)
(we remove heat from water when making ice in freezer)

22. All chemical reactions either release or absorb heat
Reaction Enthalpies
All chemical reactions either release or absorb heat
Exothermic reactions:
Reactants products energy as heat (H -ve)
e.g. burning fossil fuels
Endothermic reactions:
Reactants + energy as heat products (H +ve)
e.g. photosynthesis

23. Bond Strengths Lattice Enthalpy
Bond strengths measured by bond enthalpy HB (+ve values)
bond breaking requires energy (+ve H)
bond making releases energy (-ve H)
Lattice Enthalpy
A measure of the attraction between ions (the enthalpy change when a solid is broken up into a gas of its ions)
all lattice enthalpies are positive
I.e. energy is required o break up solids

24. Enthalpy of hydration Hhyd
the enthalpy change accompanying the hydration of gas-phase ions
Na+ (g) + Cl- (g) Na+ (aq) + Cl- (aq)
-ve H values (favourable interaction)
WHY DO THINGS DISSOLVE?
If dissolves and solution heats up : exothermic
If dissolves and solution cools down: endothermic

25. + = + = Ions associating with water Breaking solid into ions
Dissolving
Lattice Enthalpy +
Enthalpy of Hydration =
Enthalpy of Solution
Substances dissolve because energy and matter tend to disperse (spread out in disorder)
2nd law of Thermodynamics

26. Second Law of Thermodynamics:
the disorder (or entropy) of a system tends to increase
ENTROPY (S)
Entropy is a measure of disorder
Low entropy (S) = low disorder
High entropy (S) = greater disorder
hot metal block tends to cool
gas spreads out as much as possible

27. = + Total entropy change entropy change of system
entropy change of surroundings = +
Dissolving
disorder of solution
disorder of surroundings
must be an overall increase in disorder for dissolving to occur

28. 1. If we freeze water, disorder of the water molecules decreases , entropy decreases
( -ve S , -ve H)
2. If we boil water, disorder of the water molecules increases , entropy increases (vapour is highly disordered state)
( +ve S , +ve H)

29. A non-spontaneous change is a change that occurs only when driven
A spontaneous change is a change that has a tendency to occur without been driven by an external influence
e.g. the cooling of a hot metal block to the temperature of its surroundings
A non-spontaneous change is a change that occurs only when driven
e.g. forcing electric current through a metal block to heat it

30. A chemical reaction is spontaneous if it is accompanied by an increase in the total entropy of the system and the surroundings
Spontaneous exothermic reactions are common (e.g. hot metal block spontaneously cooling) because they release heat that increases the entropy of the surroundings.
Endothermic reactions are spontaneous only when the entropy of the system increases enough to overcome the decrease in entropy of the surroundings

31. System in Dynamic Equilibrium  S  0 (zero entropy change)
A B C D
Dynamic (coming and going), equilibrium (no net change)
no overall change in disorder
 S  0 (zero entropy change)