Transcript Thermochemistry - Parkway C-2

Thermochemistry
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
Energy is the capacity to do work
•
Thermal energy is the energy associated
with the random motion of atoms and
molecules
•
Chemical energy is the energy stored
within the bonds of chemical substances
•
Nuclear energy is the energy stored within
the collection of neutrons and protons in
the atom
•
Electrical energy is the energy associated
with the flow of electrons
•
Potential energy is the energy available
by virtue of an object’s position
6.1
Two main general forms of
energy
Kinetic energy
(EK) = ½ mv2
• Energy is measured in the
standard unit of Joules
• 1 J = 1 kg ∙ m2/s2
• mass must be in kg
• velocity must be in m/s
Energy due
to motion
• height must be in meters
• g = acceleration due to
gravity must be in m/s2
Potential
energy
(EP) = mgh
Energy due to
position (stored
energy)
First Law of Thermodynamics: the total energy
of the universe is constant and can neither be
created nor destroyed; it can only be
transformed.
The internal energy, U or E, 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
Total Internal Energy = Kinetic Energy
Potential Energy
Ut (Et) =
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
Energy Changes in Chemical Reactions
Heat is the transfer of thermal energy between two bodies that
are at different temperatures.
Temperature is an indirect measurement of the thermal
or heat energy.
Temperature is NOT Thermal
Energy
greater temperature
900C
400C
greater thermal energy 6.2
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 1 K
1 cal = 4.184 J
1000 cal = 1 Calorie (food calorie)
The specific heat (C) of a substance is the amount of heat (q)
required to raise the temperature of one gram of the
substance by one degree Celsius.
Heat (q) absorbed or released:
q = m C Dt
M = mass of substance
C = specific heat of substance
Dt = tfinal - tinitial
How much heat is given off when an 869 g iron bar cools
from 940C to 50C?
C of Fe = 0.444 J/g • 0C
Dt = tfinal – tinitial = 50C – 940C = -890C
q = mCDt
= 869 g x 0.444 J/g • 0C x –890C = -34,000 J
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
SURROUNDINGS
open
closed
isolated
6.2
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)
Exothermic process is any process that gives off heat –
transfers thermal energy from the system to the surroundings.
2H2 (g) + O2 (g)
H2O (g)
2H2O (l) + energy
H2O (l) + energy
Endothermic process is any process in which heat has to be
supplied to the system from the surroundings.
energy + 2HgO (s)
energy + H2O (s)
2Hg (l) + O2 (g)
H2O (l)
Burning fossil
fuels is an
exothermic
reaction
Photosynthesis is an
endothermic reaction
(requires energy input
from sun)
Forming Na+
and Cl- ions
from NaCl is an
endothermic
process
Enthalpy (H) is used to quantify the heat flow into or out of a system in a process
that occurs at constant pressure. (comes from Greek for “heat inside”)
DH = H (products) – H (reactants)
DH = heat given off or absorbed during a reaction at constant pressure
Hproducts < Hreactants
DH < 0
Hproducts > Hreactants
DH > 0
Thermochemical Equations
Is DH negative or positive?
System absorbs heat
Endothermic
DH > 0
6.01 kJ are absorbed for every 1 mole of ice that
melts at 00C and 1 atm.
H2O (s)
H2O (l)
DH = 6.01 kJ
Thermochemical Equations
Is DH negative or positive?
System gives off heat
Exothermic
DH < 0
890.4 kJ are released for every 1 mole of methane
that is combusted at 250C and 1 atm.
CH4 (g) + 2O2 (g)
CO2 (g) + 2H2O (l) DH = -890.4 kJ
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, like the pressure of air
Car engine: gasoline burns &
produces gases which push out
pistons in the engine and transfer
energy to the wheels of car
Heat and work are 2 equivalent ways of changing the
internal energy of a system
Change in
internal
energy
=
Heat either
gained or lost
by the system
+
Work done
either by or
on the system
DU (E) = q (heat) + w (work)
q
w
q
U
w
U (E) like reserves
of a bank: bank
accepts deposits or
withdrawals in two
currencies (q & w)
but stores them as
common fund, U (E).
First Law of Thermodynamics can be reworded….
- The internal energy of an isolated system is
constant – or, DU = q + w
•An important form of work
is EXPANSION WORK i.e.
the work done when a system
changes size and pushes
against an external force like
air pressure
e.g. the work done by hot
gases in an engine as they
push back the pistons
In a system that can’t expand, no work is done (w = 0)
DU (E) = q + w
when w = 0, DU (E) = q (at constant volume)
MOST CHEMICAL REACTIONS WE OBSERVE SIMPLY
WASTE HEAT AND DO NO WORK – THEY ARE NOT HOOKED
UP TO MACHINES IN THE LAB!
• What if work is done by
a chemical reaction?
• In this example, the
chemical reaction is
producing hydrogen gas
• The hydrogen gas is
doing work by pushing
the piston against the air
pressure
• Work is being done by
the system
•W=Fxd
• How can we translate
that here?
6.7
•w=Fxd
• The force the reaction is working
against is from the air
• Pressure = Force/Area
• So, w = Pressure x Area x distance
• The distance here is the height that
the cylinder has to move up against
the air
• Work is in Joules
• Pressure must be in the
standard unit of Pascals
• Volume must be in the
standard unit of m3
• The area would be the area of the
circle, or piston
• Area x height of a piston = volume
• So, w = P x DV
• Since the work is done by the system,
and should be negative, we say:
h
• w = - P x DV
6.7
• Mathematically, enthalpy is defined as:
•
H = U + PV
• We also know that DU = q + w, and
• w = - PDV
• So, DU = q – PDV
•Rearranging, we can also say that:
• q = DU + PDV
• Also, DH = DU + PDV
• If no work is done by a chemical reaction, then there is no volume
change at a constant pressure, and therefore:
• q = DU and DH = DU, so:
• DH = q
• Telling me that all of the heat lost by the reaction, since no work is done,
is equal to the change in enthalpy –
• This is common sense, since enthalpy is supposed to measure the heat
change of a reaction!
A state function is a property whose value does not depend on
the path taken to reach that specific value.
Change in enthalpy, pressure, volume, temperature,
potential energy are all state functions
For example: Density is a state
function because a substance's
density is not affected by how the
substance is obtained. Lets say we
have a certain amount H2O, it does
not matter whether that H2O is
obtained from one's tap, from a
well, or from a bottle, because as
long as all three have the same
states, they will have the same
density.
Potential energy of hiker 1 and
hiker 2 is the same even though
they took different paths.
When deciding whether a certain
property is a state function or not, keep
this rule in mind: is this property or
value affected by the path or way taken
to establish it? If the answer is no, then
there is a state function, but if the
answer is yes, then there is no state
function.
How do we measure the heat (DH) of a reaction…?
• There are four ways to measure the heat of a chemical reaction:
– Calorimetry – using a calorimeter to experimentally measure the
heat released or absorbed from a reaction
– Hess’ Law – using other chemical reactions to algebraically solve
for the heat of a reaction
– Heats of formation – using a table of potential energies to measure
the potential energies of the reactants and products of a reaction,
and thereby calculating the DH of a reaction:
DH = Σ H prod - Σ H reac
– Add the energy it takes to break the reactant’s bonds, and subtract
the energy released when the new bonds are made:
DH = Energy in – Energy out = Bonds broken – Bonds made
Using calorimetry…..
• A calorimeter is used….
• A calorimeter is an insulated
device used to capture all of
the heat either absorbed or
released by a reaction!
• The reaction is usually
surrounded by water….why?
• Water is stable, and has a
high specific heat
• It changes temperature
slowly!
• q reaction = - q surroundings
• By measuring the heat that
the water absorbs or releases,
we can calculate the heat of
the reaction!
No heat enters or leaves!
A 0.1964-g sample of solid quinone (C6H4O2) is burned in a bomb
calorimeter that contains 373 grams of water. The temperature of the
calorimeter increases by 3.2°C. Calculate the energy of combustion
of quinone per mole.
• First – write a balanced chemical equation!
• 1 C6H4O2 (s) + 6 O2 (g) → 6 CO2 (g) + 2 H2O (g)
• The heat released by the reaction is absorbed by the
calorimeter:
• q = mcDT
• q reaction = - q calorimeter
• q = (373 g H2O)(4.184 J/g0C)(3.2°C) = 4994.02 J gained by
calorimeter
• q reaction = -4994.02 J
This is not the energy of combustion, though!
• 1 C6H4O2 (s) + 6 O2 (g) → 6 CO2 (g) + 2 H2O (g)
• The energy of combustion, or DH, is the energy
released for the reaction the way it was written!
• We only used .1964 grams of the chemical!
• The reaction calls for one mole of the chemical!
• 1 mole C6H4O2 (s)
• So I set up a ratio:
• -4994.02 J/.1964 g = X/108 g
• X = -2,746,203.764 J
• So, DHcombustion = -2,700,000 J (2 significant
figures
 DHcombustion = -2,700 KJ
•The other two methods, Hess’ Law and
Heats of Formation, will be discussed in
class next time!
Just to clarify….
• Molar heat of combustion is the heat change for one mole
of a substance burning: DHcombust
• The standard heat of formation is the heat change for a
substance being formed from the elements: DHform
• The molar enthalpy is the DH or change in heat for one
mole of a substance – not necessarily for how the reaction
is written!
• Standard enthalpy change is DHo, which means the heat
change for a reaction the way it is written under standard
conditions.
• Enthalpy change or DH is the heat change for the reaction
the way it is written at any conditions.