Transcript File - Chemistry 30

Unit B
Review
Thermochemical Changes
The study of energy changes by a chemical system during a chemical
reaction is called thermochemistry.
Calorimetry is the technological process
of measuring energy changes of an
isolated system called a calorimeter.
Recall that an isolated system does not
exchange matter or energy with its outside
environment.
No calorimeter is 100% sealed and insulated,
so they only approximate an isolated system.
Analyzing Energy Changes
Heat refers to the form of energy that is transferred
from an object at a higher temperature to an object at
a lower temperature.
Thermal energy is the total kinetic energy of the
entities of a substance.
Q = quantity of thermal energy (J)
Q = mcΔt
m = mass (g)
c = specific heat capacity
(J/g·°C)
Δt = temperature change
(°C)
The S.I. unit for energy is the joule (J).
The specific heat capacity of a
substance is the quantity of energy
required to raise one gram of a
substance by one degree Celcius.
The change in temperature of the water is
used to determine the quantity of heat energy
released or absorbed by the chemical system.
Heat Transfer and Enthalpy Change
Kinetic energy (energy of motion) of a chemical system includes:
• moving electrons within atoms
• the vibration of atoms connected by chemical bonds
• the rotation of molecules
• the translation of molecules (moving from one place to another)
The temperature of a chemical system is a measure of the average
kinetic energy of the entities that make it up.
So a change in temperature means a change in kinetic energy.
Potential energy (stored energy in chemical bonds) includes:
• covalent and/or ionic bonds between the entities (intramolecular)
• intermolecular forces between entities
hydrogen bonding
covalent bonding
The enthalpy (H) of a system is the sum of the kinetic and potential
energy within it.
An enthalpy change, ΔH, is the difference between the enthalpy of the
products and the enthalpy of the reactants for a system under constant
pressure.
ΔH = Hproducts – Hreactants
WE CANNOT MEASURE
ENTHALPY DIRECTLY!
We can calculate the quantity of heat that is
released or absorbed by the surroundings
of a chemical system by measuring a
change in temperature of the surroundings.
ΔH = Q
(system)
(calorimeter)
Zn(s) + 2 HCl(aq) → H2(g) + ZnCl2(aq)
The change in potential energy of the chemical system equals
the change in kinetic energy of the surroundings.
For an exothermic reaction (chemical system releases
heat), ΔH (the enthalpy change) is negative.
For an endothermic reaction (chemical system
absorbs heat), ΔH (the enthalpy change) is positive.
Molar Enthalpies and Calorimetry
Enthalpy of reaction (or enthalpy change of reaction) refers to the
energy change for a whole chemical system when reactants change to
products.
Δr H = enthalpy of reaction (kJ)
Δ r H = n Δ r Hm
n = chemical amount (mol)
Δr Hm = molar enthalpy of reaction
(kJ/mol)
Molar enthalpy of reaction is the enthalpy change in a chemical
system per mole of a specific chemical in a system at constant pressure.
In a calorimeter, the change in enthalpy of the chemical system is equal
to the change in thermal energy of the calorimeter.
ΔH = Q
n Δr Hm = mcΔt
Method 1: Molar Enthalpies of Reaction, ΔrHm
When reactants and products are in their
standard state, they are at a pressure of
100 kPa, an aqueous concentration of
1.0 mol/L. and liquids and solids are in
their pure state.
To communicate a molar enthalpy,
both the substance and the reaction
must be specified.
Formation Reaction
Δf Hm° = –239.2
CH3OH
kJ/mol
C(s) + 2 H 2 (g) + 12 O 2 (g)  CH 3OH(l)
When 1 mol of methanol is formed from its elements when they are in
their standard states at SATP, 239.2 kJ of energy is released.
Combustion Reaction
Δc Hm° = –725.9
CH3OH
kJ/mol
CH 3OH(l) + 32 O 2 (g)  CO 2 (g) + 2 H 2O(l)
The complete combustion of 1 mol of methanol releases 725.9 kJ of
energy.
Note that the above reactions are balanced for
one mole of the compound.
Method 2: Enthalpy Changes, ΔrH
Write an enthalpy change (Δr H) beside the chemical equation.
CO(g) + 2 H2(g) → CH3OH(l)
Δr H = –725.9 kJ
The enthalpy change is not a molar value, so does not require the “m”
subscript and is not in kJ/mol.
SO 2 (g) +
1
2
O 2 (g)  SO3 (g)
2 SO 2 (g) + O 2 (g)  2 SO3 (g)
Δc H° = –98.9 kJ
Δc H° = –197.8 kJ
When 2 moles of sulfur dioxide are burned, twice as much heat energy
is released as when 1 mole of sulfur dioxide is burned.
Method 3: Energy Terms in Balanced Equations
For endothermic reactions, the energy is listed
along with the reactants.
reactants + energy → products
For exothermic reactions, the energy is listed
along with the products.
reactants → products + energy
Method 4: Chemical Potential Energy Diagrams
During an exothermic reaction, the enthalpy of the system decreases.
Heat flows out of the system and into the surroundings and we observe
a temperature increase.
Method 4: Chemical Potential Energy Diagrams
During an endothermic reaction, the enthalpy of the system increases.
Heat flows into the system from the surroundings and we observe a
temperature decrease.
Not all chemical energy changes can be studied conveniently using
simple calorimetry.
Methods used to study these
reactions are based on the principle
that net changes in all properties of
a system are independent of the
way the system changes from the
initial state to the final state.
Predicting ΔrH: Hess’ Law
Hess’ law states that the addition of chemical equations yields a net
chemical equation whose enthalpy change is the sum of the individual
enthalpy changes.
ΔrH° = Δ1H° + Δ2H° + Δ3H° + … =
ΣΔrH°
Two things to remember:
• If a chemical equation is reversed, then the sign of ΔrH changes.
• If the coefficients of a chemical equation are altered by multiplying
or dividing by a constant factor, then the ΔrH is altered by the same
factor.
Use the following given equations and their standard enthalpy changes.
When comparing enthalpy changes for formation reactions of different
compounds, we must choose a reference energy state.
It is convenient to set the enthalpies of elements in their most stable
form at SATP to be zero.
As an arbitrary convention, for the sake of simplicity, all other
enthalpies of compounds are measured relative to that reference energy
state.
A formation reaction always begins with elements, so any standard
enthalpy of formation reactions are measured from the reference
energy state of zero.
Thermal stability is the tendency of a compound to resist
decomposition when heated.
The lower (i.e. more negative) the value of a compound’s standard
molar enthalpy of formation, the more stable it is.
Δf Hm° = – 280.7
kJ/mol
Δf Hm° = – 577.6
kJ/mol
SnO
SnO2
Tin(IV) oxide has a greater thermal stability than tin(II) oxide.
The standard enthalpy change of a reaction is the sum of the
standard enthaplies of formation of the products minus the sum
of the standard enthalpies of formation of the reactants.
ΔrH° = ΣnΔfPHm° – ΣnΔfRHm°
ΔrH° = – 64.5
kJ
– 985.2
= – 64.5 kJ
Reaction Progress
• Why do some chemicals react faster than others, when all other
variables are controlled, except for the type of chemicals?
e.g. Mg(s) reacts much faster with HCl(aq) than Zn(s).
• Why do some reactions need an initial input of external energy to
start?
e.g. A match is needed to start the combustion of a hydrocarbon.
Collision-Reaction Theory
I2
H2
HI
• A chemical sample consists of entities (atom, ions, or molecules) that
are in constant random motion at various speeds, rebounding
elastically from collisions with each other.
• A chemical reaction must involve collisions of reactant entities.
• An effective collision requires sufficient energy. Collisions with the
required energy have the potential to react.
• An effective collision also requires the correct orientation
(positioning) of the colliding entities so that bonds can be
broken and new bonds can be formed.
• Ineffective collisions involve entities that rebound elastically from
the collision.
Activation Energy of a Reaction
Activation energy is the
minimum energy that colliding
entities must have in order to
react.
This initial input energy may be
in the form of heat, light, or
electricity.
CO(g) + NO2(g) → CO2(g) + NO(g)
ΔrH° = – 224.9 kJ
CO(g) + NO2(g) → CO2(g) + NO(g)
The diagram on the right just tells
us about the “before and after.”
ΔrH° = – 224.9 kJ
H2(g) + I2(g) → 2 HI(g)
ΔrH° = + 53.0 kJ
Breaking bonds between atoms or ions requires energy, making it an
endothermic process.
bonded particles + energy → separated particles
Bond energy is the energy required to break a chemical bond. The
stronger the bond, the greater the energy needed to break it.
Forming bonds between atoms or ions releases energy, making it
an exothermic process.
bonded particles → separated particles + energy
Bond energy is also the energy released when a bond is formed.
The stronger the bond, the greater the energy released.
Endothermic Reactions
2 H2O(l) → 2 H2(g) + O2(g)
ΔrH° = +571.6
kJ
In any endothermic reaction, the energy required to break bonds
(2 O─H bonds in H2O) is greater than the energy released when
bonds are formed (O═O and H─H).
Exothermic Reactions
H2(g) + Cl2(g) → 2 HCl(g)
ΔrH° = –184.6
kJ
In any exothermic reaction, the energy required to break bonds (H─H
and Cl ─Cl) is less than the energy released when bonds are formed
(H─Cl).
Empirical Effect of Catalysis
Catalysis deals with the properties and development of catalysts, and
the effects of catalysts on chemical reactions.
A catalyst is a substance that increases the rate of a chemical reaction
without being consumed in the overall process.
Chlorophyll is a catalyst
for photosynthesis.
The inside of a catalytic converter in a car
exhaust system is coated with an alloy that acts
as a catalyst for the combustion of exhaust gases.
Theoretical Explanation of Catalysis
A catalyst lowers the activation energy for a reaction.
This results in a larger number of effective collisions between entities,
so the reaction rate increases.
The rate of the forward reaction and of the reverse reaction increases
by the same amount.
Biological catalysts (enzymes) have binding sites specific to the
reactants (substrate). The enzyme temporarily binds the substrate
molecules and aligns them in the proper orientation.
The uncatalyzed reaction proceeds slowly at room temperature.
The catalyzed reaction has a lower activation energy, thus having an
increased rate.
Uses of Catalysts
The Oil Industry
Catalysts are used in the cracking and reforming of crude oil which
reduces the temperature required for these processes, making them
more efficient.
Upgrading of Bitumen from Oil Sands
Emissions Control
Enzymes