Transcript Thermo Chemistry Ch 9 Notes

Unit 1 – Thermochemistry
General Outcomes:
1. Determine and interpret energy
changes in chemical reactions
2. Explain and communicate energy
changes in chemical reactions
Enthalpy of formation
 Enthalpy of reaction
 ΔH notation
 Hess’ Law
 Molar enthalpy
 Energy diagrams
 Activation Energy
 Catalysts
 Calorimetry
 Fuels and Energy Efficiency



the potential to do work.
examples of energy changes:
- chemical energy in food to motion
- solar energy to power
- geothermal energy to heat
- nuclear energy to power, heat,
motion


Thermodynamics: study of energy
changes in systems
Thermochemistry –study of energy
changes involved in physical and
chemical changes.
First Law - Energy can be neither
created nor destroyed, it simply
changes forms.
Esystem = -Esurroundings
Second Law – In any energy
conversion, energy is always lost in
the form of heat.



We can observe the effect it has on
our surroundings.
Energy systems either release energy
to the surroundings or absorb energy
from the surroundings.
Universe = System + Surroundings

Open systems – matter and energy
move freely in / out.

Closed systems – only energy moves
in / out.

Isolated system – neither matter nor
energy moves in / out.
Calculating Thermal Energy
1.
2.
3.
Q =quantity of heat flowing in / out
of that substance, measured as
energy in units of joules (J) or
kilojoules (kJ)
m-mass of substance undergoing
the change measured in units of
grams (g) or kilograms (kg)
c=heat capacity -heat required to
change the temp of 1 g of substance
by 1oC
1.
Many hot water heaters use the
combustion of natural gas to heat
the water in the tank. When 150 L
of water at 10oC is heated to 65oC,
how much energy flows into the
water?
2.
A solid substance has a mass of
250.0 g. It is cooled by 25.0oC, and
loses 4.937 kJ of heat. What is the
specific heat capacity and the
identity of the substance?
3.
If a copper pot weighing 0.500 kg is
heated with 2.500 kJ of energy, what
is the expected temperature change?

Energy from the sun is stored in the bonds
of chemical substances such as fossil fuels
and hydrocarbons.



Solar energy is captured by plants
during photosynthesis.
Living tissues from plants and
animals become buried. Instead of
decaying, these compounds form
hydrocarbons or fossil fuels.
The energy from fossil fuels is
released during combustion
reactions and converted into heat.
Photosynthesis is endothermic:
6CO2(g) + 6H2O(l) + energy  C6H12O6(s) + 6O2(g)

Cellular Respiration is exothermic:
C6H12O6(s) + 6O2(g)  6CO2(g) + 6H2O(l) + energy

Combustion is exothermic:
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) + energy

 There
are two types of energy
that matter may possess:
1. Kinetic energy – the energy of
motion
2. Potential energy – the energy of
position

faster the particle is moving, the
greater the kinetic energy.

Can be found in three types:
◦ Translational
◦ Rotational
◦ Vibrational


Ek measured indirectly by measuring
the temperature of a system.
is proportional to the amount of heat a
system holds.


Stored energy
found in the bonds
and inter-molecular
forces of molecules
Changes in
potential energy
include phase
changes, chemical
reactions and
nuclear reactions



Are changes in potential energy
energy is required to break bonds
and the energy is released when
new bonds are formed.
chemical reactions:micro/macro



enthalpy describes the total amount
of energy a system has.
Enthalpy Change (ΔH = kJ) refers to
how much the total energy of a
system increases or decreases
during a chemical reaction.
Molar enthalpy (ΔH = kJ/mol)
describes the amount of enthalpy
change per mole of substance.



Exothermic
reaction – more
energy is released
in making bonds
than is required to
break them.
(-ΔH)
Surrounding
temperature
increases



Endothermic
reaction – less
energy is released
in forming bonds
than was required
to break bonds.
(+ΔH)
Surrounding
temperature
decreases
H = change in enthalpy (kJ)
H = molar enthalpy kJ/mol
n = moles
n=m/M
1.
Hydrogen and oxygen gas react to
produce 0.534 g of water vapor. If
the molar enthalpy of reaction is
241.8 kJ/mol for water, what is the
enthalpy change?
2.
The molar enthalpy of combustion
for ethanol is -725.9 kJ/mol. What
mass of methanol must be burned to
generate 2.34 x 104 kJ of energy?
3.
What is the molar enthalpy of
combustion if a 10.0 g sample of
pure acetic acid is burned in oxygen
and produces 144.7 kJ of energy?
1. Molar enthalpy of a specific reaction relative to
one species in the reaction.
2H2(g) + O2(g)  2H2O(g) H: -241.8 kJ/mol H2O
2. Enthalpy change for a balanced reaction
equation.
2H2(g) + O2(g)  2H2O(g)
∆H: -483.6 kJ
3. Including the energy value as part of the
balanced equation.
2H2(g) + O2(g)  2H2O(g) + 483.6 kJ
2C(g) + 2H2(g) + 52.4 kJ  C2H4(g)
4. Drawing a potential energy diagram for the
reaction.


used to determine energy changes
during physical or chemical processes
calorimeter : ideally an isolated system
Energy in Chemical Reactions - Learning Activit - Flash
Player Installation


All components of the calorimeter
are included in the total energy
change.
Mathematically, calorimetry may be
described as:
Heat lost = heat gained
(* remember heat lost is expressed as a negative
value)



used for reactions in aqueous
solution
can calculate the heat change in the
water and use to determine
enthalpy change
P 353 lists assumptions made when
using a simple calorimeter
1.
50.00mL of 0.300mol/L CuSO4(aq) is
mixed with 50.00mL of NaOH(aq). The
initial temperature of both solutions is
21.40oC and the highest temperature
recorded after mixing is 24.60oC.
a. Determine the enthalpy change for the
system
b. Determine the molar enthalpy change for
CuSO4
c. Write the thermochemical equation.
2.
Barium chloride and sodium sulfate
react to produce a precipitate of
barium sulfate. When 0.100 mol of
precipitate is formed the
temperature of 3.00 kg of water is
raised from 20.123 to 20.316oC.
Calculate the molar enthalpy of
reaction for the production of barium
sulphate.


used for combustion reactions
total energy change of the calorimeter
includes changes in the container itself,
the water, the thermometer and the
stirrer.


bomb calorimeters may report a heat
capacity for the entire calorimeter
system. (C measured in J/oC)
Q = CT
1.
A bomb calorimeter has a heat
capacity of 40.00kJ/oC. Complete
combustion of 1.00g of hydrogen
causes the temperature of the
calorimeter to increase by 3.54oC.
What is the molar enthalpy of
combustion from this evidence?
2.
When 1.00 g of propane is burned,
about 2.36 J of heat is given off.
What mass of water at 50.0oC can
be heated to 80oC when 4.00 mol
of propane is burned?
3.
A strip of magnesium metal having a
mass of 1.22 g is placed in 100 mL of
1.00 mol/L HCl(aq) in a metal can with
a heat capacity of 562 J/oC. After the
reaction, the temperature increased
from 23.0oC to 45.5oC. Assuming the
solution in the calorimeter has the
same heat capacity as water,
determine the heat of reaction per
mole of magnesium.