Transcript CHAPTER 6 Thermodynamics

Thermochemistry
Lecture Notes
Thermodynamics
It is the science that studies the
relationship between heat and other forms
of energy.
First Law of thermodynamics
It is also known as the law of conservation
of energy: “the total energy of the Universe
is constant” or sometimes expressed as
"Energy can neither be created nor
destroyed".
Another approach is to say that the total
energy of an isolated system remains
constant.
Energy
• Energy (E or U) - is the capacity to do
work (W) or to produce heat (Q). Energy is
a state function, does not depend on a
pathway, it depends only on the present
state.
ΔE = Q
+ W
Internal Energy = Heat + work
Work & Heat
• Work is defined as force acting over a
distance.
• Heat is the energy that flows between
objects that become in contact if they are
at different temperatures. It is an extensive
property of matter, depends on the amount
of matter.
Energy Units
Name
Symbol
Unit conversion
1 Joule
J
1 kg.m2∕ s2
1 kilojoule
kJ
1000 J
1 calorie
cal
4.184 J
1 Calorie
Cal
1000 cal
Other forms of Energy
Kinetic
Potential
Thermal energy; all matter has
thermal energy at the submicroscopic
level due to motion of submicroscopic
particles.
Chemical potential energy;
contained in chemical bonds,
observed in chemical reactions.
Mechanical energy; all moving
objects have it.
Gravitational energy; caused by the
effect of gravity over matter.
Electric energy; generated by moving Electrostatic energy; attraction
electrons.
between particles of opposite charges
placed at a certain distance.
Sound; due to compression &
expansion of spaces between
molecules.
Heat and Temperature
• Temperature – T
It is the property of matter that determines whether heat
energy can be transferred from one body to another and
in what direction. It is an “intensive” property, it does not
depend on the amount of matter.
• Heat – q
It is thermal energy. It is directly proportional to the
motion of molecules in matter. It is an extensive
property. It depends on the amount of matter.
Second Law of
Thermodynamics
• The second law states that heat energy
cannot be transferred from a body at a
lower temperature to a body with a higher
one without the addition of energy.
Heat flows from matter at a higher
temperature to matter at a lower
temperature.
Third law of Thermodynamics
The Third Law of Thermodynamics states that
the entropy of any pure substance in thermodynamic equilibrium
approaches zero as the temperature approaches zero (Kelvin), or
conversely the temperature (Kelvin) of any pure substance in
thermodynamic equilibrium approaches zero when the entropy
approaches zero.
Entropy is the measure of disorder in a system.
System & Surroundings
“Thermo” Vocabulary
• System – designated part of the universe under study.
• Surroundings - Everything outside the system, the rest
of the universe. In practice one considers an “immediate
surroundings” and then surroundings.
• Heat flow - the direction at which thermal energy flows.
It always goes from the matter at higher temperature to
the one at a lower temperature.
• Thermal equilibrium – It is the point at which all heat
transfer stops.
• Law of conservation of energy – The quantity of heat
lost is equal to the quantity of heat gained.
Heat Transfer
ENDOTHERMIC
EXOTHERMIC
SURROUNDINGS
↓
SYSTEM
qf > q i
SURROUNDINGS
↑
SYSTEM
qf<qi
The system absorbs heat
from the surroundings
The system releases heat to
the surroundings
Specific Heat Capacity or Specific Heat
• It is the quantity of heat required to raise
the temperature of 1 gram of a substance
by one kelvin degree. Its units are joules
per gram per kelvin.
• Heat = joules (energy unit)
• Mass = grams
• Temperature = kelvin scale
Specific Heat Capacity, C
(J/g.K)
C = ___q____
m x (Tf-Ti)
q = total heat transferred in Joules
m = amount of matter in grams
Tf = Final temperature in Kelvin degrees
Ti = initial temperature in Kelvin degrees
Sign Conventions (page 209)
ΔTsys = Tf –Ti
Δqsys= qf - qi
Direction of
heat transfer
+
Tf > Ti
+
qf>qi
Endothermic
process
Tf<Ti
qf<qi
Exothermic
process
Example for calculating C
• If 1160 Joules of heat were used to raise the
temperature from 25 0C to 325 0C of 10.0 g of a metal
sample, what is its specific heat capacity? Identify this
metal using Table 6.1* page 210.
DATA
q = 1160 J
Tf = 325 0C + 273 = 598 K
Ti = 25 0C + 273 = 298 K
m = 10.0g
C=?
SOLUTION
C = ___ q____
= _____1160 J_____ = 0.385_ J_ (*copper)
m . (Tf-Ti)
10.0 g . (598-298)K
g.K
Heat Transfer and specific heat capacity
q sys + q surr = 0
This assumption is used to calculate experimentally the specific heat capacities of
unknown metals.
System = piece of metal , with known mass and initial high temperature.
Immediate surrounding = water, with known specific heat, known mass and initial
cool temperature.
For the metal
q sys = q metal= m metal. Cmetal. (Tf –Ti metal )
For the water
q surr = q water = m water . (4.184 J/g.K) . (Tf-Ti water)
q water + q metal = 0
Note that Tf is the same for both. Tf is the temperature at which thermal
equilibrium has been reached and no more heat transfer occurs.
Calorimetry
• Heat transferred at constant pressure is
called enthalpy, H.
q rxn + q sol = 0  q rxn = - q sol
• Heat transferred at constant volume is
called internal energy , ΔE
q rxn + q bomb + q sol = 0 
q rxn = -(q bomb + q sol)
Enthalpy Changes in Chemical
Reactions
• Enthalpies are measured taking the system (chemical
reaction) as its reference point.
• ΔH > 0 means the system has absorbed heat, it is an
endothermic reaction.
• ΔH < 0 means the system has released heat, it is an
exothermic reaction.
• ΔH of the reverse reaction is equal to - ΔH of the forward
reaction.
• ΔH is molar dependent.
Hess’s Law
“ If a reaction is the sum of two or more
other reactions, ΔH for the overall process
is the sum of the ΔH values of those
reactions”
ΔH3 = ΔH1 + ΔH2 + …
Enthalpy Changes in Chemical
Reactions
• Enthalpies are measured taking the system (chemical
reaction) as its reference point.
• ΔH > 0 means the system has absorbed heat, it is an
endothermic reaction.
• ΔH < 0 means the system has released heat, it is an
exothermic reaction.
• ΔH of the reverse reaction is equal to - ΔH of the forward
reaction.
• ΔH is molar dependent.
Standard Molar Enthalpy of
Formation
• Standard Molar Enthalpy of Formation,
ΔHf, is the enthalpy change for the
formation of 1 mol of a compound directly
from its component elements in their
standard states, usually measured at 298
K and 1 at.
• Example ΔHf (CO2, 1at, 298 K) is
• C(s) + O2(g)  CO2(g) ΔH f = -393.5 kJ
Enthalpy Change for a Chemical
Reaction
• Reactants  Products
• ΔH rxn = ΣΔH (products) – ΣΔH (reactants)
• Exothermic Reactions are product favored,
because the final enthalpy is lower, more
stable.
• Endothermic reactions are reactant
favored.