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Chapter 6

CHEM 1050 Chapter Notes - Chapter 6: Thermodynamic System, Kinetic Theory Of Gases, Calorie


Department
Chemistry
Course Code
CHEM 1050
Professor
Lori Jones
Chapter
6

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Chem 1050 Text Chapters 6.1-6.3, 6.6
Chapter 6.1- Energy and Its Units
Energy: the potential or capacity to move matter
- energy is a property of matter and exists in many forms that can be interconverted
- for example, can be interconverted from heat energy to electrical energy to light energy etc
Kinetic Energy; Units of Energy
Kinetic Energy: energy associated with an object by virtue of its motion
Ek = ½ mv2 m = mass І v = velocity І Ek = kinetic energy
- kinetic energy depends both on mass and speed, thus a heavier object could move more
slowly than a lighter object and still retain the same kinetic energy
- the SI unit of energy is kg · m2/s2, otherwise known as joules (J)
- one watt is 1 J/s (ie a 100 watt bulb uses 100 J of energy every second)
- a kilowatt hour is 3600 kilowatt seconds, or 3.6 million joules
Calorie (cal): a non SI unit of energy commonly used by chemists, originally defined as the
amount of energy required to raise the temperature of one gram of water by one degree Celsius
- 1 cal = 4.184 J
Potential Energy
Potential Energy: the energy an object has by virtue of its position in a field of force
Ep= mgh m = mass І g = gravity (9.807m/s2) І h = height of water
- only differences in potential energy are important in any physical situation
- for example, as something falls, its potential energy decreases from mgh at the top to 0 at
standard level
- as an object is falling, the potential energy is converted to kinetic energy
- as potential energy decreases, kinetic energy increases
Internal Energy
Internal Energy (U): the sum of the kinetic and potential energies of the particles making up a
substance
Etot = Ek + Ep + U
- when studying a substance in a lab, its kinetic energy is 0 (at rest), and its potential energy is a
constant (and thus can be taken as 0), and therefore the total energy of the substance is equal
to its internal energy, U
Law of Conservation of Energy
Energy may be converted from one form to another, but the total quantity of energy remains
constant
- energy is conserved
- the first law of thermodynamics is a specific statement of the conservation of energy

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Chapter 6.2- First Law of Thermodynamics; Work and Heat
- first law of thermodynamics relates the change in internal energy of a physical or chemical
change taking place in a container to the flows (transfers) of energy into our out of the container
- the energy transfers come in two kinds; heat and work
Thermodynamic System and Its Surroundings
System: the substance or mixture of substances that we single out for study (perhaps including
the vessel)
Surroundings: everything in the vicinity of the thermodynamic system (its environment)
Definition of Work
Work: an energy transfer (or energy flow) into or out of a thermodynamic system whose effect
on the surroundings is equivalent to moving an object through a field of force
- energy flows from the system to its surroundings, and thus does work on the surroundings
- can also have the surroundings to work on the system
- work (w) is;
--> + when work is done on the system (when you do work on the system, you add
energy to it)
--> - when work is done by the system (when the system does work, energy is
subtracted)
Definition of Heat
Heat: an energy transfer (energy flow) into or out of a system that results from a temperature
difference between the system and its surroundings
- as long as the system and its surroundings are in thermal contact (not insulated from one
another), energy (heat) flows between them to establish equal temperatures, or thermal
equilibrium
- heat flows from regions with higher temperatures to those with lower temperatures, and once
thermal equilibrium is reached, heat flow stops
Kinetic-theory Explanation of Heat: Imagine there are two vessels, one vessel contains
molecules at a higher temperature than the other vessel which contains oxygen molecules.
Molecules collide with the vessel walls, thereby losing or gaining energy. The faster molecules
tend to slow down, and the slower molecules tend to speed up. The net result is that energy is
transferred through the vessel walls from the hot gas to the cold gas, and this transfer is
referred to as heat
- heat is not the same as temperature
--> absolute temperature of a gas is directly proportional to the average kinetic energy of
the molecules
--> when you add heat to a gas, you increase its internal energy, and therefore its total
kinetic energy
--> as a result, the increase in average kinetic energy per molecule (and consequently
the increase in temperature) depends on the size of the gas sample
--> a given quantity of heat will raise the temperature of a sample more if the sample is
small
- heat (q) is;
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