Download Thermochemistry - Valdosta State University

Chapter 6- Thermodynamics (Thermochemistry)
Study of energy changes in chemical processes.
From Physics recall these definitions:
Force(F) - a kind of push or pull on an object
Energy – the capacity to do work.
Work(w) – force applied over a distance(d) w = F  d
Heat – form of energy transferred from a warmer
object to a cooler object.
Kinetic Energy (Thermal Energy) – enerEgy12 dmvue to motion.
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k
Potential Energy (Stored Energy) – the energy an object
possesses due to its position in a field (example: gravitation
field) or chemical energy stored in bonds.
- Potential energy can be converted into kinetic energy.
Example: a ball rolling down a hill
At rest. High potential energy(PE),
no kinetic energy (KE)
Moving. Increasing KE
Decreasing PE
PE due to position in
the field of gravity.
At rest. No KE.
Lower PE than at top of hill.
SI Unit for energy is the joule, J:
1 J = 1kg(m2)/s2
A more traditional unit is the calorie
calorie (cal) –energy required to raise 1.0 g of water 1oC.
1cal = 4.184J
On food packages it has Calories (note the capital "C").
These are really kilocalories.
1000 calories = 1 kilocalorie = 1 Calorie
Usually we need to divide the universe this way to study
something:
Temperature- measure of heat content (measures average
kinetic energy of the particles)Used to predict which way
heat will flow (hot→cold)
Endothermic(adj.) – a process that absorbs heat from the
surroundings. (Example: melting ice)
Exothermic(adj.) – a process that gives off heat to the
surroundings. (Example: combustion)
System – portion of the universe we wish to study.
Surroundings – everything else.
Universe = System + Surroundings
Energy can flow between a system and the surroundings,
but the __________________________________.
“The total amount of energy in the universe is fixed.”
This is the ________________________________
Also called the ________________________________
For an object (like a particular chemical) we can describe
something called the ________________________
Internal Energy – _____ of all ______________________
energy in an object.
Both ____ energy and _____ can _____________ a
system’s ____________________________
• It is very hard to determine an object’s internal energy,
but it is possible to determine the change in energy
(_ _ _ ).
• Change in internal energy, ΔE = __________________
– A ___________ ΔE means Efinal______ Einitial
or the _______________________ from the
surroundings (__________________)
– A ______________ ΔE means Efinal ____ Einitial
– or the system ______________ to the surroundings
(___________________)
_________________ – a process that is determined by its
initial and final conditions.
• “A process that is ______________________.”
• ______ (w) and _____ (q) _______ state functions.
• Energy change (____) ______ a state function.
Another State function
__________________) - Heat transferred between the
system and surroundings carried out under constant
pressure.
Most reactions occur under constant pressure, so
If volume is also constant,
So, Energy change is due to heat transfer,
Enthalpy Change (____) – The heat evolved or absorbed
in a reaction at constant pressure
H and ΔH are state functions, depending only on the initial
and final states.
There are many specific H’s, depending on what you
want to know
H
– enthalpy of vaporization (liquid  gas)
H
– enthalpy of fusion (solid  liquid)
H
– enthalpy of combustion
(energy from burning a substance)
H
- enthalpy for a neutralization reaction
H
- enthalpy of dissolving a substance
Hreaction - enthalpy of a general reaction
A fundamental H:__________________________________
For a reaction
1. Enthalpy is an _____________property (magnitude
ΔH is directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
ΔH = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) ΔH = _______
2. When we reverse a reaction, the sign of ΔH changes:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g)
ΔH = +802 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
ΔH =________
3. Change in enthalpy depends on state:
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
ΔH = -802 kJ
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l)
ΔH = -890 kJ
2 Mg(s) + O2(g)  2 MgO(s)
H = -1204 kJ
a) Is this reaction endothermic or exothermic?
b) Calculate the amount of heat transferred when 2.4 g of
Mg reacts at constant pressure.
c) How many grams of MgO are produced when the
enthalpy change is -96.0 kJ?
d) How many kilojoules of heat are absorbed when 7.50g
of MgO is decomposed into Mg and O2 at constant
pressure?
Calorimetry
Method of measuring heat flow and quantity of heat.
_______________ – The amount of heat required to raise
an object’s temperature by ________________
- As the heat capacity of a body increases, the amount of
heat required to produce a given temperature change is
increased.
__________Heat Capacity – The amount of heat required
to _______________________ of a substance by 1 Kelvin.
______________(Specific Heat Capacity, c) – The amount
of heat required to ______________________________
How many calories are needed to raise the temperature of
2.5 g of aluminum by 10oC?
(cAl= 0.897 J/g-K)
Changes of state: Temperature remains constant during
melting, freezing, boiling, condensation.
Use Hvaporization for boiling/condensation
Hfusion for freezing/melting
Pay attention to sign. ____for melting and boiling
____for freezing and condensation
The heat of vaporization of methanol is 2 kJ/g.
How much heat is required to vaporize 65.0 g of methanol?
How much energy is required to melt 5.0 g of pure ice and
bring it to room temperature (20oC)?
Hfusion =333 J/g
cwater = 4.184 J/g-K