Changes

*''Chemical thermodynamics'', which is heat and energy transfers involving chemical reactions within chemical systems.

*''Chemical thermodynamics'', which is heat and energy transfers involving chemical reactions within chemical systems.

Thermodynamics has an emphasis on a beginning or initial state of a system, and an end or final state of a system, with the system being all of the interacting components on this energy path.  Measurements in thermodynamics are usually reported on the [[International System of Units#Kelvin|Kelvin scale]] <ref>David Halliday, "Fundamentals of Physics Extended", John Wiley & Sons, New York, 1997 </ref>

Thermodynamics has an emphasis on a beginning or initial state of a system, and an end or final state of a system, with the system being all of the interacting components on this energy path.  Measurements in thermodynamics are usually reported on the [[International System of Units#Kelvin|Kelvin scale]] <ref>David Halliday, "Fundamentals of Physics Extended", John Wiley & Sons, New York, 1997</ref>

==History==

==History==

The downfall of Lavoisier's caloric theory happened at the arsenal in Munich, Germany.  The Bavarian minister of war was a British expatriate, Sir Benjamin Thompson, and he observed that work was being converted into heat by observing the boring of a cannon.  If caloric theory was correct, he reasoned, no more heat would be made once all of the caloric was removed from the cannon at the atomic level, yet his observations on this procedure - including a cannon bored while under water - demonstrated that work can be converted into heat, like the steam engines of his time converting heat into work.   

The downfall of Lavoisier's caloric theory happened at the arsenal in Munich, Germany.  The Bavarian minister of war was a British expatriate, Sir Benjamin Thompson, and he observed that work was being converted into heat by observing the boring of a cannon.  If caloric theory was correct, he reasoned, no more heat would be made once all of the caloric was removed from the cannon at the atomic level, yet his observations on this procedure - including a cannon bored while under water - demonstrated that work can be converted into heat, like the steam engines of his time converting heat into work.   

James Joule in 1849 made a precise determination of the mechanical equivalent of heat into work.  His stirring of water in a pot (work input with a mechanical stirring rod, caused by the motion of a 1-kg weight falling 42.4 cm) caused a temperature increase (heat output); his homemade, yet very-precise thermometers recorded the conversion factor.  The unit of energy is now called the ''[[International_System_of_Units#Joule|joule]]''.  4.2 joules of energy can raise one gram of water 1 degree Celsius, an amount called one calorie.  (The large-C "Calorie" used in nutritional measure is 1000 small-c calories.)

James Joule in 1849 made a precise determination of the mechanical equivalent of heat into work.  His stirring of water in a pot (work input with a mechanical stirring rod, caused by the motion of a 1-kg weight falling 42.4&nbsp;cm) caused a temperature increase (heat output); his homemade, yet very-precise thermometers recorded the conversion factor.  The unit of energy is now called the ''[[International System of Units#Joule|joule]]''.  4.2 joules of energy can raise one gram of water 1 degree Celsius, an amount called one calorie.  (The large-C "Calorie" used in nutritional measure is 1000 small-c calories.)

==Zeroth Law of Thermodynamics==

==Zeroth Law of Thermodynamics==

where the constant depends on the amount and type of the gas sample

where the constant depends on the amount and type of the gas sample

[[Charles' Law]], formulated in the 1780's, stated that, for a given sample of gas at a fixed pressure, the volume was directly proportional the the "absolute" temperature.  That is,

[[Charles' Law]], formulated in the 1780s, stated that, for a given sample of gas at a fixed pressure, the volume was directly proportional the "absolute" temperature.  That is,

:<math>V/T = constant</math>

:<math>V/T = constant</math>

Where X is some constant that is characteristic of the gas sample.  It's easy to see that X is proportional to the amount of gas (2 grams of gas will have twice the volume of 1 gram).  So X is actually the amount of gas, measured in some convenient units (grams, moles, molecules), times some number that is characteristic of the gas.   

Where X is some constant that is characteristic of the gas sample.  It's easy to see that X is proportional to the amount of gas (2 grams of gas will have twice the volume of 1 gram).  So X is actually the amount of gas, measured in some convenient units (grams, moles, molecules), times some number that is characteristic of the gas.   

In the 1810's, Gay-Lussac and Avogadro made an amazing discovery:  The mysterious constant is just the molecular weight of the gas, if the amount of gas is measured in the right units.  The right unit to use is the ''mole'', which is the mass, in grams, that matches the molecular weight of the gas.  (By then atomic weights and molecular weights were beginning to be understood.)  So, for example, a mole of Chlorine is 71 grams, because a Chlorine molecule has two atoms.  This led to the [[Ideal Gas Law|Universal Gas Law]] or Ideal Gas Law:

In the 1810s, Gay-Lussac and Avogadro made an amazing discovery:  The mysterious constant is just the molecular weight of the gas, if the amount of gas is measured in the right units.  The right unit to use is the ''mole'', which is the mass, in grams, that matches the molecular weight of the gas.  (By then atomic weights and molecular weights were beginning to be understood.)  So, for example, a mole of Chlorine is 71 grams, because a Chlorine molecule has two atoms.  This led to the [[Ideal Gas Law|Universal Gas Law]] or Ideal Gas Law:

:<math>PV = nRT\,</math>

:<math>PV = nRT\,</math>

where <math>n\,</math> is the amount of gas, measured in moles, and <math>R\,</math> is the ''Universal Gas Constant'' of 8.314 joules per Kelvin.  A mole has to be defined as that amount, in grams, equal to the molecular weight of the gas.  This required that the molecular weight of diatomic gases, like Hydrogen, Nitrogen, Oxygen, and Chlorine, be twice the atomic weight, because the molecules have two covalently bound atoms.  Inert gases, like Helium and Neon, have only one atom per molecule.  For something like Ammonia vapor (NH₃), the molecular weight is 17, the sum of the atomic weights of the atoms.

where <math>n\,</math> is the amount of gas, measured in moles, and <math>R\,</math> is the ''Universal Gas Constant'' of 8.314 joules per Kelvin.  A mole has to be defined as that amount, in grams, equal to the molecular weight of the gas.  This required that the molecular weight of diatomic gases, like Hydrogen, Nitrogen, Oxygen, and Chlorine, be twice the atomic weight, because the molecules have two covalently bound atoms.  Inert gases, like Helium and Neon, have only one atom per molecule.  For something like Ammonia vapor (NH₃), the molecular weight is 17, the sum of the atomic weights of the atoms.