Tag: gas laws

If you went to school and learned the “Gas Laws” early on and it seemed boring a LAME but then later things like mass flow rates and air flow conversions and compression ratios seemed HARD.

Well…

It’s because HVAC is hard if you don’t understand the concepts behind the gas laws, even if you never learn to do the math.

First, we need to make sure and define some words so we are all on the same page. Don’t skim over this part unless you are really sure you understand what these words mean.

Matter – Matter is “stuff”, anything that has weight and takes up space

Mass – You can think of it as weight for most purposes, it is a measurement of how much “stuff” there is.

Volume – How much space the stuff takes up.

I always think of volume as “boxes” and mass as “the stuff in the box”. The box defines the space but you can’t tell how much stuff is in the box until you try to pick it up or weigh it.

Density – A measure of “compactness” of matter. The more dense, the more mass it will have by volume. If you let a sponge take its natural shape it is less dense then when you ball it up in your fist and it becomes denser.

Temperature – The average intensity of heat energy at that point. It is literally the average molecular velocity of the stuff you are measuring.

Pressure –  The force exerted on, in or by matter.

In HVAC/R we are constantly dealing with concepts contained in the “Ideal Gas Law” which informs us of how matter and contained environments respond to changes in temperature, volume, pressure, and mass. There are several gas law(s) that are often taught but they are all combined in the ideal gas law to bring it all together. We will get back to the ideal gas law shortly, but first we need to understand a bit about systems.

Whenever you are thinking about matter and energy it is important to consider whether the situation you are observing or testing is an open, closed or isolated system.

 

Open System

An open system allows both matter (mass) and energy to move in and out of a system. You can think of this like an open pot of water where water molecules can leave or enter through the top and energy can also enter and leave through the top and through the walls of the pot.

Closed System

A closed system can allow energy to enter and leave but the mass of the matter remains fixed. This is like a pot of water with a lid tightly sealed on the top or for our purposed the refrigerant sealed inside an air conditioning system.

Isolated System 

An isolated system means a system that can neither transfer matter (mass) or energy in and out. This would be like a cooler full of cold beer with a top tightly sealed and with perfect insulation that would never let any heat in or out. In practice this is impossible but if it were the beer would never change temperature once the temperature stabilized inside.

In an open system like an open pot of water or a lake and we heat it, the water molecules on the surface begin to break free and move away as they evaporate. This leaves the surface below a little cooler as that high energy molecules leave. The lake or pot don’t pressurize when the molecules leave becasue the system is open to the atmosphere.

Now, if we add heat in a closed system like a pressure cooker the entire system begins to pressurize as the water evaporates and then boils. Energy is being added but the mass can’t change, this leads to a pressure and temperature increase inside the pot. If we keep adding more and more heat without a pressure relief the pot will either explode or melt.

This quick clip from Mythbusters shows what can happen in a closed system when energy is added to a water heater until it fails


What is an ideal gas? 

An ideal gas is a gas that obeys the ideal gas law, it’s ideal because it’s good at following rules. These ideal gasses walk in a straight line, they don’t run on the playground and they never fish without a proper permit. More like an ideal gas behaves in a predictable way with changes in volume, pressure, temperature, and mass.

The problem is, a truly “ideal” gas really doesn’t exist.

While many gasses behave close to ideal at normal temperatures there is no gas that obeys the ideal gas laws in all conditions.

The ideal gas law is –

P=  Absolute Pressure (gauge pressure + atmospheric pressure)
V = Volume (How much space the gas occupies)
n = Mass measured in “moles” (the number of molecules)
R = The universal gas constant (varies depending on the units of measure being used Example: [lbf ft/(lb mol oR)]= 8.3145 )
T = Absolute Temperature (temperature in a scale that starts at absolute zero like Kelvin or Rankine)

The result is that many gasses that we work with behave in about the same way with changes in mass, volume, temperature & pressure. This is the case because the primary force at play in a nearly ideal gas like nitrogen or CO2 is simply the velocity of the molecules bouncing around in the container and against one another like tiny little ping pong balls.

If the molecules react, or interact with one another through attraction or repulsion due to their intermolecular forces then they can cease to behave as an ideal gas. A perfect example is when a gas is in contact with its liquid form (saturation) it no longer obeys the gas laws. This is why most gasses behave more and more like an ideal gas the hotter they get (within a range) because the hotter they are the greater the force of molecular velocity (temperature) will be relative to the intermolecular interaction of the molecules.

From a practical standpoint understanding the relationships between temperature, pressure, volume and mass in contained (closed) environments like an HVAC refrigerant circuit and uncontained (open) environments like the air outside are really helpful.

The most practical of them is –

A decrease in temperature causes a decrease in pressure or a decrease in pressure causes a drop in temperature. An increase in either results in the opposite effect.


Why Does it “Matter” (Pun Intended) 

In HVAC/R it is really important that we understand the impacts changes in temperature, pressure, volume and mass will have inside the system.

When we add or remove refrigerant we are changing the refrigerant mass which impacts the pressures and temperatures. When one system has a smaller condenser coil than another the decrease in internal volume will increase the pressure and temperature if the refrigerant mass is the same.

When the air passing over an evaporator coil is colder the pressure in the coil will be lower because the gas laws teach us that a decrease in temperature equals a decrease in pressure.

All of these factors are interconnected and if you don’t understand the gas laws you won’t understand how they impact one another.

Here’s one for the nerds!

— Bryan

Breaking the Sound Barrier

So what do you think of when you hear an “ideal gas”? R22, R12 maybe… Natural? Take a look at the F-18 above… It is breaking the sound barrier and that cloud is a shockwave… This has nothing to do with this article but I think it’s pretty darn cool!

An ideal gas is a gas that obeys the ideal gas law, it’s ideal because it’s good at following rules. These ideal gasses walk in a straight line, they don’t run on the playground and they never fish without a proper permit. More like an ideal gas behaves in a predictable way with changes in volume, pressure, temperature, and mass.

The problem is, a truly “ideal” gas really doesn’t exist.

While many gasses behave close to ideal at normal temperatures there is no gas that obeys the ideal gas laws in all conditions.

The ideal gas law is –

P=  Absolute Pressure (gauge pressure + atmospheric pressure)
V = Volume (How much space the gas occupies)
n = Mass measured in “moles” (the number of molecules)
R = The universal gas constant (varies depending on the units of measure being used Example: [lbf ft/(lb mol oR)]= 8.3145 )
T = Absolute Temperature (temperature in a scale that starts at absolute zero like Kelvin or Rankine)

The ideal gas law is really a combination of several different laws into one.

The result is that many gasses that we work with behave in about the same way with changes in mass, volume, temperature & pressure. This is the case because the primary force at play in a nearly ideal gas like nitrogen or CO2 is simply the velocity of the molecules bouncing around in the container and against one another like tiny little ping pong balls.

If the molecules react, or interact with one another through attraction or repulsion due to their intermolecular forces then they can cease to behave as an ideal gas. A perfect example is when a gas is in contact with its liquid form (saturation) it no longer obeys the gas laws. This is why most gasses behave more and more like an ideal gas the hotter they get (within a range) because the hotter they are the greater the force of molecular velocity (temperature) will be relative to the intermolecular interaction of the molecules.

Once the gas gets to the “supercritical” state all bets are off once again. So like most good kids, even the most ideal gasses have their limits where if pushed they become little molecular rebels.

— Bryan

 

The gas laws. We all learned about them in school and promptly forgot all about them. I really think that we need to dig our books out, dust that information off and work to understand and apply it.

Many will say that nitrogen pressure doesn’t change with pressure like other gasses. This is false but read on.

Let’s start by looking at the pressure a little differently. Pressure is a measure of the force exerted by a gas within a container. It exerts pressure because the individual molecules of the gas are colliding with the walls of the container. Those collisions are happening because each molecule has a specific amount of energy. So, in this way, we can view pressure as a measure of the amount of energy contained within our container of gas. That might sound complicated, so let’s kind of unpack it and see if we can understand it better.

We have a container that has a fixed volume, for example, 1 cubic foot. So at 0 psig, there is a certain number of gas molecules contained within that container and a certain number of collisions with the container walls occurring.

Now, let’s take that container and we’re going to double the number of molecules inside that container without changing its size at all.. We know that the pressure increased, but what did it take to do this? Energy.

Adding those additional gas molecules required that we add energy to force that extra gas into the container. The addition of energy to force additional molecules into the container resulting in an increase in pressure. The thing to remember now is the law of conservation of energy. Energy isn’t created nor destroyed, it simply changes form.

Since heat energy is simply another form of energy so it stands to reason that adding or removing heat energy from our system will affect the energy level of the gas molecules and ultimately the pressure exerted by them. Let’s return to our sample container of 1 cubic foot internal volume. We’re going to expend enough energy to put enough molecules into this container to raise the pressure to 100 psig at a temperature of 70°F. If we add more energy not in the form of compressing more gas but in the form of heat energy, what will happen to the pressure in the container?

The heat energy is going to ‘excite’ the molecules in the gas, increasing the number and force of the collisions that are occurring that are the basis of pressure existing. Since we’re adding energy, the pressure will rise and it will rise in a predictable and consistent way. The reverse is also true if we remove energy, the pressure will drop in the same consistent and predictable
way.

This is why we need to understand the gas laws as technicians. They allow us to predict and understand the pressure change caused by adding or removing heat energy from a sealed, pressurized system.

Practical application
Now that we understand how heat energy affects the pressure within a sealed system, we can apply this knowledge to pressure testing. A large number of factors are making proper leak testing at installation more important than ever and manufacturers are demanding more detailed leak testing procedures. Add to that the fact that our tools are more refined than old-school analog gauges and a leak of even 0.5 psi over a several hour period of time is easily something a technician can spot.

Let’s take a look at an imaginary but fairly realistic scenario to see how this works and what it means on the ground in the field.

New construction split system. Tonnage isn’t super important to this, but we just made the last brazed joint, it’s the end of a long day in the 90° heat and a nasty thunderstorm is brewing. Let’s get this thing pressurized and get home. Run the pressure up to 350 psig of nitrogen and get out of here. When we show up in the morning when it’s 65°F and find that the pressure has dropped almost 16 psig, that might make us a little nervous. We checked all of our joints with a mirror and with soap bubbles but we don’t see any leaks… where did the pressure go?

Before we get excited, let’s look at how the temperature change affected the pressure within this sealed system. We pressurized to 350 psig at 90°F and it’s now 65°F. With the gas law equations, we can know what the pressure in the system should be and eliminate time wasted looking for leaks that aren’t actually there. This is an expression of the gas laws known as Gay Lussac’s Law. In this, the system volume is a constant and can be disregarded. For our purposes, the copper piping we use to build systems is unchangeable, so we’ll use this equation.

The first step is for change the equation around to isolate the answer we wish to get.
P2= T2 (P1/T1)

Now, we have a simple equation we can plug our numbers into and get the answer, right? Not quite yet. We have one more step before we get the calculators out. We need to convert the pressure and temperature valves that we have to absolute pressure and temperature readings, so add 14.7 to the pressure and 459.76° (Rankine scale) to the temperature to get to absolute scales

Now, our numbers look like this:
T1 = 549.67°R (Rankine)
P1 = 364.7 psia
T2 = 524.67°R
NOW, let’s solve.
P2 = 524.67 (364.7/ 549.67)
P2 = 524.67 (0.6635)
P2 = 348.11

But wait, our system dropped to 334 psig, so we have a leak…
We forgot one VITAL step. We need to convert our P2 reading back to gauge pressure.
349.03 – 14.7
333.41 psig

This says that the pressure loss within the system was due ONLY to the temperature change and was not due to a leak.
Time to get the vacuum pump out and finish this job up.

In summary, every gas responds to the gas laws in the same way. We use nitrogen because it is readily available (the air is mostly made of nitrogen), dry and it doesn’t readily combine with other molecules under normal circumstances.

It does change pressure with temperature and all you need to do to find out how much it will change is by changing both the before and after temperatures to absolute scales (Rankine for Fahrenheit or Kelvin for Celcius) and convert your before and after pressure readings from gauge pressure (PSIG) to absolute pressure (PSIA). Once you have your solution you can convert back to Celcius or Fahrenheit

— Jeremy Smith CMS

P.S. – I made a little before and after calculator HERE

Most of the laws we refer to in air conditioning and refrigeration are pretty obvious and practical and Dalton’s law of partial pressures is no exception. John Dalton simply observed that the pressure of air was equal to the added pressures of each gas that make up air. This means that the pressure and density of air can vary based on the exact makeup of the gases contained in the air.

The law of partial pressures states that –

 In a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases.

This simply means that if you take two gases and you place them together in a contained space, you simply add the pressures together to get the total pressure. The only case when this does not apply is when the gasses “react” with one another to create new molecular structures.

In practical terms, this is why nitrogen or air in a refrigerant circuit increases the pressures. The pressure of the nitrogen is added to the pressure of the refrigerant resulting in higher pressure.

It is also one reason that refrigerant manufacturers blend refrigerants to create ideal boiling and condensing temperatures based on the percentage of one refrigerant over another. A common example of this is R407C vs. R407A, they both are made up of  R-32, R134a and R-125 but the % of each in the mixture dictates the pressure/temperature properties.

Now clearly, this law applies only to gas (vapor), not matter (refrigerant) in the saturated state like refrigerant in a tank, but when the refrigerant is in the vapor state it obeys Dalton’s law perfectly.

— Bryan

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