Transcript thermodynamics

Thermodynamics
Antoine Lavoisier
[1743-94]
Julius Robert Meyer
James Joule : 1818~1889
Hermann von Helmholtz (1821-1894)
Rudolf Clausius (1822--1888),
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) + 2O2(g)
CO2(g) + 2H2O(l) + energy
What is Energy?
Energy
Kinetic
energy
(EK)
Energy due
to motion
Potential
energy
(EP)
Energy due to
position (stored
energy)
Total Energy =
Kinetic Energy
E
=
EK
+
+
Potential Energy
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
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
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
OPEN
ISOLATED
CLOSED
Thermochemistry is the study of heat change in chemical
reactions.
The system is the specific part of the universe that is of
interest in the study.
SURROUNDINGS
SYSTEM
open
Exchange: mass & energy
closed
isolated
energy
nothing
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)
HEAT and WORK
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
Work is the transfer of energy that takes place when an
object is moved against an opposing force
i.e. a system does work when it expands against an
external pressure
Car engine: petrol burns &
produces gases which push out
pistons in the engine and transfer
energy to the wheels of car
•Work stimulates uniform motion
• Heat and work can be considered as energy in transit
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 = 4.184 J
INTERNAL ENERGY (U)
Internal energy changes when energy enters or leaves a
system
U = Ufinal - Uinitial
U
change in the internal energy
Heat and work are 2 equivalent ways of changing the
internal energy of a system
Change in
internal
energy
=
Energy
supplied to
system as
heat
+
Energy
supplied to
system as
work
U = q (heat) + w (work)
q
w
q
U
w
U like reserves of a
bank: bank accepts
deposits or
withdrawals in two
currencies (q & w)
but stores them as
common fund, U.
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
WORK
•An important form of work is EXPANSION WORK
i.e. the work done when a system changes size and
pushes against an external force
e.g. the work done by hot gases in an engine as they
push back the pistons
HEAT
In a system that can’t expand, no work is done (w = 0)
U = q + w
when w = 0, U = q
(at constant volume)
•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
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
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
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)
Photosynthesis is an
endothermic reaction
(requires energy input
from sun)
Forming Na+
and Cl- ions
from NaCl is an
endothermic
process
Measuring Heat
reaction
reaction
Exothermic reaction, heat
given off & temperature of
water rises
Endothermic reaction, heat
taken in & temperature of
water drops
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
b. Specific heat is the amount of
heat required to raise the temperature
of 1 kg of a material by one degree (C
or K).
1) C water = 4184 J / kg C
2) C sand = 664 J / kg C
This is why land heats up quickly
during the day and cools quickly at
night and why water takes longer.
Why does water have such a
high specific heat?
Water molecules form strong bonds
with each other; therefore it takes
more heat energy to break them.
Metals have weak bonds and do not
need as much energy to break
them.
How to calculate changes
in thermal energy
Q = m x T x Cp
Q = change in thermal energy
m = mass of substance
T = change in temperature (Tf – Ti)
Cp = specific heat of substance
c. A calorimeter is
used to help measure
the specific heat of a
substance.
First, mass and
Knowing
its Q value,
temperature
of
its mass,
and its
water
are measured
Then
heated
This
gives
the
T, its Cp can be
T is measured
sample
is
put
heat
lost
by
the
calculated
for water to help
inside
and heat
substance
get its heat gain
flows into water
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
Vaporisation
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)
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
e.g. photosynthesis
products (H +ve)
Bond Strengths
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
Enthalpy of hydration Hhyd
• the enthalpy change accompanying the hydration of gasphase 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
Breaking solid
into ions
Lattice
Enthalpy
+
+
Ions
associating
with water
Enthalpy of
Hydration
=
Dissolving
=
Enthalpy of
Solution
Substances dissolve because energy and matter tend to
disperse (spread out in disorder)
2nd law of Thermodynamics
USING ENTHALPY
Making H2O from H2 involves two steps.
H2(g) + 1/2 O2(g) ---> H2O(g) + 242 kJ
H2O(g) ---> H2O(liq) + 44 kJ
----------------------------------------------------------------------H2(g) + 1/2 O2(g) --> H2O(liq) + 286 kJ
Example of HESS’S LAW—
If a rxn. is the sum of 2 or more others, the net ∆H
is the sum of the ∆H’s of the other rxns.
Hess’s Law
& Energy Level Diagrams
Forming H2O can
occur in a single
step or in a two
steps.
∆Htotal is the same
no matter which
path is followed.
Hess’s Law
& Energy Level Diagrams
Forming CO2 can
occur in a single
step or in a two
steps.
∆Htotal is the same
∆H along one path =
∆H along another path
• This equation is valid because ∆H is
a STATE FUNCTION
• These depend only on the state of
the system and not on how the
system got there.
• V, T, P, energy — and your bank
account!
• Unlike V, T, and P, one cannot
measure absolute H. Can only
measure ∆H.
Standard Enthalpy Values
Most ∆H values are labeled ∆Ho
Measured under standard conditions
P = 1 bar = 105 Pa = 1 atm /1.01325
Concentration = 1 mol/L
T = usually 25 oC
with all species in standard states
e.g., C = graphite and O2 = gas
Enthalpy Values
Depend on how the reaction is written and on
phases of reactants and products
H2(g) + 1/2 O2(g) --> H2O(g)
∆H˚ = -242 kJ
2 H2(g) + O2(g) --> 2 H2O(g)
∆H˚ = -484 kJ
H2O(g) ---> H2(g) + 1/2 O2(g)
∆H˚ = +242 kJ
H2(g) + 1/2 O2(g) --> H2O(liquid)
∆H˚ = -286 kJ
Standard Enthalpy Values
NIST (Nat’l Institute for Standards and Technology) gives
values of
∆Hfo = standard molar enthalpy of formation
— the enthalpy change when 1 mol of compound is formed
from elements under standard conditions.
See Table 6.2
∆Hfo, standard molar
enthalpy of formation
Enthalpy change when 1 mol of compound
is formed from the corresponding
elements under standard conditions
H2(g) + 1/2 O2(g) --> H2O(g)
∆Hfo (H2O, g)= -241.8 kJ/mol
By definition,
∆Hfo = 0 for elements in their standard
states.
Using Standard Enthalpy
Values
Use ∆H˚’s to calculate enthalpy change for
H2O(g) + C(graphite) --> H2(g) + CO(g)
(product is called “water gas”)
Using Standard Enthalpy
Values
H2O(g) + C(graphite) --> H2(g) + CO(g)
From reference books we find
• H2(g) + 1/2 O2(g) --> H2O(g) ∆Hf˚ = - 242
kJ/mol
• C(s) + 1/2 O2(g) --> CO(g)
kJ/mol
∆Hf˚ = - 111
Using Standard Enthalpy
Values
H2O(g) --> H2(g) + 1/2 O2(g) ∆Ho = +242 kJ
C(s) + 1/2 O2(g) --> CO(g)
∆Ho = -111 kJ
-------------------------------------------------------------------------------
H2O(g) + C(graphite) --> H2(g) + CO(g)
∆Honet = +131 kJ
To convert 1 mol of water to 1 mol each of
H2 and CO requires 131 kJ of energy.
The “water gas” reaction is ENDOthermic.
Using Standard Enthalpy
Values
Calculate ∆H of
reaction?
In general, when ALL
enthalpies of formation are
known:
∆Horxn =  ∆Hfo (products) -  ∆Hfo (reactants)
Remember that ∆ always = final – initial
Using Standard Enthalpy
Values
Calculate the heat of combustion of methanol, i.e.,
∆Horxn for
CH3OH(g) + 3/2 O2(g) --> CO2(g) + 2 H2O(g)
∆Horxn =  ∆Hfo (prod) -  ∆Hfo (react)
Using Standard Enthalpy
Values
CH3OH(g) + 3/2 O2(g) --> CO2(g) + 2 H2O(g)
∆Horxn =  ∆Hfo (prod) -  ∆Hfo (react)
∆Horxn = ∆Hfo (CO2) + 2 ∆Hfo (H2O)
- {3/2 ∆Hfo (O2) + ∆Hfo (CH3OH)}
= (-393.5 kJ) + 2 (-241.8 kJ)
- {0 + (-201.5 kJ)}
∆Horxn = -675.6 kJ per mol of methanol
CALORIMETRY
Measuring Heats of Reaction
Constant Volume
“Bomb”
Calorimeter
• Burn combustible
sample.
• Measure heat
evolved in a
reaction.
• Derive ∆E for
reaction.
Calorimetry
Some heat from reaction warms
water
qwater = (sp. ht.)(water mass)(∆T)
Some heat from reaction warms
“bomb”
qbomb = (heat capacity, J/K)(∆T)
Total heat evolved = qtotal = qwater + qbomb
Measuring Heats of Reaction
CALORIMETRY
Calculate heat of combustion of octane.
C8H18 + 25/2 O2 --> 8 CO2 + 9 H2O
• Burn 1.00 g of octane
• Temp rises from 25.00 to 33.20 oC
• Calorimeter contains 1200 g water
• Heat capacity of bomb = 837 J/K
Measuring Heats of Reaction
CALORIMETRY
Step 1 Calc. heat transferred from reaction to water.
q = (4.184 J/g•K)(1200 g)(8.20 K) = 41,170 J
Step 2 Calc. heat transferred from reaction to bomb.
q = (bomb heat capacity)(∆T)
= (837 J/K)(8.20 K) = 6860 J
Step 3 Total heat evolved
41,170 J + 6860 J = 48,030 J
Heat of combustion of 1.00 g of octane =
- 48.0 kJ
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
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
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)
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
•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
System in Dynamic Equilibrium
A
+
B
C
+
D
Dynamic (coming and going), equilibrium (no net change)
• no overall change in disorder
 S  0 (zero entropy change)
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