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Lecture

# CHEM 112 Lecture Notes - Isothermal Process, Vapor Pressure, Joule

Department
Chemistry
Course Code
CHEM 112
Professor
John Carran

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Systems:
Open system – open to transfer of heat and matter
Closed – allows transfer of energy, but not matter
Isolated system – closed to transfer of both energy and matter
Units:
Energy, U (state function)
Work, w (not state function; half dependent)
Kinetic Energy, eK
Potential Energy, V
Heat, q (not state function)
Calorie, cal – quantity of heat required to change temperature of 1 gram of water by 1
degree C
Joule, J – SI unit for heat (1 cal = 4.184J)
***Pay attention to capital or lowercase symbol/abbreviation***
Thermo Energy
kinetic energy associated with random molecular motion
Generally proportional to temperature
Intensive property
Internal energy is transferred between a system and its surroundings as a result of temperature
difference
Heat flows from hotter to cooler
Heat Capacity
quantity of heat required to change temperature of a system by 1 degree C
Specific heat, c (q=mcT)  always positive
Molar heat capacity (q= CT)
In the law of conservation of energy, total energy remains constant

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- qsystem + qsurroundings = 0
- qsystem = -qsurroundings (surroundings is absorbing energy)
***look at calculation example in journal***
Bomb Calorimetry – energy transferred only by heat, no work involved
The “Coffee-Cup” Calorimeter – assumed to be completely insulated
– open to the atmosphere/open to expansion `
*** pay attention to how many moles you are solving for, so that the balanced equation
reflects the same values***
Pressure-volume work (involves volume change)
F x d
(m x g) x h
Compression = negative h
Expansion = positive h
First Law of Thermodynamics
a system contains only internal energy
a system does not contain heat or work
only occur during a change in the system
delta U = q+ w
*** know sign conventions used: when is q and w positive/negative***
Hard to measure individual state functions, easiest to find change in the state function
In a reversible, you get more work out of the system than a regular system
Constant volume = no work performed
*** Know figure 7-15 ***
Hess’s Law:
Delta H is an extensive property, therefore how much substance is present. In other
words, make sure the coefficients in the reactions properly reflect the molar value

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State function – half independent
Enthalpy change for the overall process is the sum of the enthalpy changes for the
individual steps (look at “if you’re wondering” box in 7.1)
reference/standard state – most common state that the element occupies at standard
temperature
for ionic reactions, you need to make a reference coefficient. (In the power point, pg 53,
the hydrogen is assigned an arbitrary value of 0)
Enthalpies of formation
can be positive or negative
you must be able to work out a complete reaction, given partial information
THERMO 2 – W11L1
Spontaneous process
occur without intervention
directional (entropy, S, provides basis for predicting direction)
equilibrium is established when driving force or tendency of the spontaneous change is
expended
Direction of a spontaneous change is often from higher to a lower energy, but there are
exceptions
Non-spontaneous reactions
require the system to by acted on by an external agent
examples include: ball not rolling up a hill, gases do not collected at one end of a
container, heat does not flow from cold body to hot body, water doesn’t freeze at
temperature above 0 C.
For chemical systems, the internal energy, U, is equivalent to potential energy