Thermochemistry studies heat effects during chemical reactions and classifies systems as open, closed, or isolated based on their interaction with surroundings. It defines energy types, including kinetic and potential energy, and introduces key equations for calculating energy changes in systems, such as internal energy and heat capacity. The document also explains thermodynamic laws, particularly the conservation of energy, and details concepts of enthalpy related to endothermic and exothermic processes.

1. Thermochemistry
COLT JAMES D. ARAPOC, RPH

2. Thermochemistry is the branch of chemistry concerned with the
heat effects that accompany chemical reactions
A system is the part of the universe chosen for study, and it can
be as large as all the oceans on Earth or as small as the contents
of a beaker.
A surrounding is that part of the universe outside the system
with which the system interacts

3. An open system freely exchanges energy and matter with its surroundings
A closed system can exchange energy, but not matter, with its surroundings
An isolated system does not interact with its surroundings

4. ENERGY is the capacity to do work or transfer heat.
Represented by E.
HEAT is the form of energy that flows between 2
objects because of their difference in temperature.
WORK is done when a force acts through a distance.

5. Energy can be classified as kinetic or potential
Kinetic energy is energy that something has because it is moving
• Energy of motion of a macroscale object, such as a moving
baseball or automobile; this is often called mechanical energy.
• Energy of motion of nanoscale objects such as atoms,
molecules, or ions; this isoften called thermal energy.
• Energy of motion of electrons through an electrical conductor;
this is often called electrical energy.

6. Kinetic energy, Ek , can be calculated as where m
represents the mass and v represents the velocity of a
moving object.
Ek= 1/2 mv2

7. Potential energy is energy that something has as a result of its position and
some force that is capable of changing that position
• Energy that a ball held in your hand has because the force of gravity
attracts it toward the floor; this is often called gravitational energy.
• Energy that charged particles have because they attract or repel each
other; this is often called electrostatic energy. An example is the
potential energy of positive and negative ions close together.
• Energy resulting from attractions and repulsions among electrons and
atomic nuclei in molecules; this is often called chemical potential energy
and is the kind of energy stored in foods and fuels.

8. Potential energy can be calculated in different ways
gravitational potential energy, Ep, can be calculated as
Ep = mgh
where m is mass, g is the gravitational constant (g 9.8
m/s2 ), and h is the height above the surface

9. Energy Units
The SI unit of energy is the joule . symbol J.
The joule is a derived unit, which means that it can be expressed as
a combination of other more fundamental units:
1 J = 1 kg m2 /s2 .
1 kilojoule = 1 kJ = 1000 J
A joule is approximately the quantity of energy required for one
human heartbeat.

10. The internal energy ( E ) of a system is the sum of
the individual energies (kinetic and potential) of all
nanoscale particles (atoms, molecules, or ions) in that
system.
A positive value of E results when Efinal > Einitial , indicating that the system
has gained energy from its surroundings.
A negative value of E results when Efinal < Einitial, indicating that the system
has lost energy to its surroundings.

11. Calorie, cal. By definition 1 cal = 4.184 J is the
quantity of energy required to raise the
temperature of one gram of water by one degree
Celsius.
1 Cal=1kcal=1000cal
kilocalorie (kcal) in nutrition = 1000cal

13. Thermodynamics is the science of heat (energy)
transfer.
Heat transfers until thermal equilibrium is established.
Heat always transfer from hotter object to cooler one.
“Energy transfers from a sample with Hot temp to a
sample with Cold temp until the two reach the same
temperature”- thermal equilibrium

14. Directionality of Heat Transfer
EXOTHERMIC: heat transfers from SYSTEM to
SURROUNDINGS.
T(system) goes down
T(surr) goes up

15. Directionality of Heat Transfer
ENDOTHERMIC: heat transfers from SURROUNDINGS
to the SYSTEM.
T(system) goes up
T (surr) goes down

16. Heat Capacity
The heat capacity of a sample of matter is the
quantity of energy required to increase the
temperature of that sample by one degree( 1 )

17. Specific Heat Capacity
How much energy is transferred due to T difference?
The heat (q) “lost” or “gained” is related to
a) sample mass
b) change in T and
c) specific heat capacity

18. Specific heat capacity =
heat lost or gained by substance (J)
(mass, g)(T change, K)

19. EX. 25.0-g sample of ethylene glycol, 90.7 J is required to change the
temperature from 22.4 °C to 23.9 °C. (Ethylene glycol is used as a
coolant in automobile engines.)
Example 2: If 100.0 g water is cooled from 25.3 °C to 16.9
°C, what quantity of energy has been transferred from
the water?
Answer 3.5 kJ transferred from the water

20. Specific Heat Capacity
If 25.0 g of Al cool from 310 o
C to 37 o
C, how many joules
of heat energy are lost by the Al?

21. Specific Heat Capacity
heat gain/lose = q = (sp. ht.)(mass)(∆T)
where ∆T = Tfinal - Tinitial
q = (0.897 J/g•K)(25.0 g)(37 - 310)K
q = - 6120 J
Notice that the negative sign on q signals heat
“lost by” or transferred OUT of Al.

22. Molar Heat Capacity
It is often useful to know the heat capacity of a sample in
terms of the same number of particles instead of the same
mass.
molar heat capacity, symbol Cm.
This is the quantity of energy that must be transferred to
increase the temperature of one mole of a substance by 1 °C.
The molar heat capacity is easily calculated from the specific
heat capacity by using the molar mass of the substance.

24. Law of Thermodynamics
• First law of thermodynamics
Law of conservation of energy, which states that
energy can neither be created nor destroyed—the total
energy of the universe is constant.

25. FIRST LAW OF
THERMODYNAMICS
∆E = q + w
heat energy transferred
energy
change
work done
by the
system
Energy is conserved!

26. heat transfer out
(exothermic), -q
heat transfer in
(endothermic), +q
SYSTEM
∆E = q + w
w transfer in
(+w)
w transfer out
(-w)

27. ENTHALPY- heat content of a system
Most chemical reactions occur at constant P, so
and so ∆E = ∆H + w (and w is usually small)
∆H = heat transferred at constant P ≈ ∆E
∆H = change in heat content of the system
∆H = Hfinal - Hinitial
Heat transferred at constant P = qp
qp = ∆H where H = enthalpy

28. If Hfinal < Hinitial then ∆H is negative
Process is EXOTHERMIC
If Hfinal > Hinitial then ∆H is positive
Process is ENDOTHERMIC
ENTHALPY
∆H = Hfinal - Hinitial

29. When H is positive, the system has gained
heat from the surroundings which means the
process is endothermic.
When H is negative, the system has released
heat to the surroundings, which means the
process is exothermic.