Americas
Europe
Explore our app and discover over 50 million learning materials for free.
Lerne mit deinen Freunden und bleibe auf dem richtigen Kurs mit deinen persönlichen Lernstatistiken
Jetzt kostenlos anmeldenIdeal gases are precisely what their name assumes them to be: ideal! And, just like most ideal things, they are non-existent! However, there are two gases that behave closely to what you expect from an ideal gas, although not in the most perfect way. Hydrogen (H2) and Helium (He) are examples of real gases that behave closely to the assumptions of ideal gas behavior. This is because they have weak Intermolecular Forces and occupy a very small volume. So, you might be thinking: when do we consider a gas ideal? Or, when does deviation from the Ideal Gas Law happen?
In this article, we will look at the definition and application of the Ideal Gas Law, and the deviation of real gases from this ideal behavior!
Before diving into the definitions and Ideal and Real Gases, let's review about the basic properties of gases and the Kinetic Molecular Theory.
Gases are one of the three most common States of Matter. Gases are commonly known for:
Assuming the shape and volume of a container.
Being compressible.
Having random, fast motion of molecules.
If you want to know more about the difference between the three most common States of Matter, take a look at the article "Solids, Liquids, and Gases"
If a gas molecule behaves according to the "Kinetic Molecular Theory", then we consider them an ideal gas. When you are dealing with an ideal gas, the ideal gas law can be used to relate its volume, pressure, the number of moles, and temperature.
Pressure is an intensive property, meaning that it does not depend on the amount of matter or substance present. The pressure of a gas is the force exerted by the gas on the surface of its container.
Pressure is the force exerted per unit area.
The SI unit of pressure is the pascal (1 Pa = 1 N/m2 or 1 kg m-s2) but it can be converted between different units. Some conversions that you might come across are:
1 atm = 760 torr.
1 atm = 760 mmHg = 1.013x105 Pa = 101.325 kPa.
1 atm = 101,325 Pa.
1 bar = 0.9869 atm.
1 bar = 105 Pa.
1 mmHg = 133.322 Pa.
1 psi = 6.895 kPa.
Let's look at an example where you need to convert between different units of pressure!
Convert 700 mmHg to Pa and torr.
We can use the conversions above to answer this question.
a) Convert mmHg to Pa:
$$700mmHg\cdot \frac{133.322Pa}{1mmHg}=93325.4Pa$$
b) Convert Pa to torr:
$$93325.4Pa\cdot \frac{1atm}{1.013\cdot 10^{5}Pa}\cdot \frac{760torr}{1atm}=700.17torr$$
Did you know that although a pressure of 1 atm is considered the standard atmospheric pressure of the Earth, it can change at certain points depending on the weather? For example, if there is low atmospheric pressure in a certain area (also known as depression) compared to the surrounding pressure, you might notice a lot of clouds, high winds, warm air, and even tropical storms. On the other hand, if an area has high pressure (also called anticyclone), then you will see an absence of clouds and tranquil weather!
Gases also have Partial Pressure, which is the pressure exerted by an individual gas within a mixture. Partial Pressure is a part of Dalton's Law of Partial pressures, which states that the sum of the partial pressures of each individual gas in a container is equal to the total pressure of the mixture of gases in the same container.
Partial pressure is defined as the pressure exerted by an individual gas within a mixture.
The equation Dalton's Law of Partial pressure is:
$$P_{total}=P_{A}+P_{B}+...$$
Fig. 1: Visualizing Partial Pressure, Isadora Santos - StudySmarter Originals.
If you are starting to feel confused, let's look at an example of how to calculate the total pressure of a container using Dalton's Law!
A sample of Nitrogen gas exerts a partial pressure of 1.25 atm inside a container. This container also has Neon gas exerting a partial pressure of 80 torr. Calculate the total pressure inside the container.
First, we need to convert the neon pressure from torr to atm. Then, we can just add up the partial pressures to get the total pressure inside the container.
$$80torr\cdot \frac{1atm}{760torr}=0.105atm$$
$$P_{total}=1.25atm+0.105atm=1.355atm$$
Feeling the need to review partial pressures? Check out the article "Ideal Gas Law"!
Volume is an extensive property that also affects the behavior of gases. Extensive properties are properties that depend on the amount of matter being measured. In gases, the volume is not fixed because gases take the shape of the container. Gases can also be compressed when pressure is applied, resulting in a change in volume.
Volume is the amount of space that matter occupies.
The SI base unit for volume is the liter (L). But, similar to pressure, it can be converted to other units!
So, if you understand how to use these different conversions, you can apply them to various examples like the one below!
A certain gas is found inside a 3.00 in3 container. Convert the volume of the container to liters (L).
By using the volume conversions, we can quickly find the volume of the container in liters.
$$3.00in^{3}\cdot \frac{16.39cm^{3}}{1in^{3}}\cdot \frac{1L}{10^{3}cm^{3}}=0.0492L$$
Temperature is an intensive property, so it does not depend on the Amount of Substance. Changing temperature causes Phase Changes.
Temperature is defined as the measurement of the average kinetic energy
You might be asked to convert between other temperature units in your exam, so these are the conversions you should know:
$$T_{K}=T_{Celsius}+273.15$$
Temperature from Kelvin to Celsius:
$$T_{Celsius}=T_{K}-273.15$$
Temperature from Fahrenheit to Celsius:
$$T_{Celsius}=\frac{T_{Fahrenheit}-32}{1.8}$$
Temperature from Celsius to Fahrenheit:
$$T_{Fahrenheit}=(1.8\cdot T_{Celsius})+32$$
Convert a temperature of 23 °C to Kelvin and Fahrenheit.
$$T_{K}=23^{o}C+273.15=296.15K$$
$$T_{Fahrenheit}=(1.8\cdot 23^{o}C)+32=73.4F$$
There are three laws that, when combined, create the Ideal Gas law Formula. These three laws are Boyle's Law, Charles's Law, and Avogadro's Law:
Fig. 2: Graphical representation of Boyle's law, Charles's Law, and Avogadro's Law, Isadora Santos - StudySmarter Originals.
If you want to know more about these three laws, check out the article "Ideal Gas Law"!
The formula for the Ideal Gas Law is:
$$P\cdot V=n\cdot R\cdot T$$
Where:
Now, let's work through an example and apply the ideal gas law equation!
You have a container with 2 moles of oxygen gas at 298 K. The container has a pressure of 1.40 atm. Find the volume (in L) of the container?
Now, we can rearrange the ideal gas law formula to solve for the volume!
$$V=\frac{n\cdot R\cdot T}{P}=\frac{(2.0mol\ O_{2})(0.08206)(298K)}{1.4atm}=34.9L$$
We can also use the ideal gas law to find the density of gases. Density is also an intensive property and can be calculated by dividing mass over volume. When dealing with density, we say that gases deviate from the ideal gas law at moderate and higher densities.
Density is defined as the amount of mass per volume.
The ideal gas law formula relating density is:
$$d=\frac{m}{V}=\frac{PM}{RT}$$
Where:
Let's utilize the ideal gas/density formula and solve an example!
Calculate the density of Carbon monoxide at 0.89 atm and a temperature of 335 K.
First, we need to find the molar mass of Carbon monoxide. So:
Now, we can plug with are the variables into the ideal gas law-density equation!
$$d=\frac{PM}{RT}=\frac{(0.89atm)\cdot (28.01g)}{(0.08206\frac{J}{mol\cdot K})\cdot (335K)}=0.907\frac{g}{L}$$
So far, we have only looked at ideal gases and how they behave. Now, we have to consider what happens when gases do not behave ideally! If a gas molecule deviates from ideal behavior (according to the kinetic molecular theory), then we call them real gases. Compared to ideal gases, real gases do have volume and forces of attraction/repulsion present, so particles do interact with one another.
Real gases are defined as gases that do not have ideal behavior and therefore are not considered ideal gases.
Deviations from the ideal gas law are less at low pressure and high temperature, and greater at high pressure and low temperature. Also, heavier gases tend to deviate most from the ideal gas behavior.
This deviation at high pressure and low temperature happens because, at low temperatures, the attractive forces can interact with the gas molecules because they have low speeds. When gases are at a low temperature, then they will have low kinetic energies.
Since real gases deviate from the ideal gas law, what can we do if we needed to calculate the pressure, volume, amount of gas, or temperature of real gases? Easy. We can use van der Waal's equation for real gases!
$$[P+a(\frac{n}{V})^{2}]\cdot [V-bn]=nRT$$
The van der Waal's equation was actually derived from the ideal gas law formula, and some corrections were made to consider the volume occupied by the gas and the Intermolecular Forces that are present. This is what the constants a and b are: correction factors! The values for these constants are specific for each gas. If you come across a question asking to use the Real Gas formula, you will be given these values, so you worry about memorizing them!
So, try not to feel overwhelmed when you see this equation! It's not as complicated as it looks and can be easily applied in calculations.
You have 0.990 moles of O2 in 1.00 L at 298 K. Using the Van der Waals formula, calculate the pressure for O2.
a = 1.36 atm·L2/mol2
b = 0.0318 atm/mol
R = 0.0821 atm/mol·K
$$P=\frac{nRT}{V-nb}-\frac{n^{2}a}{V^{2}}$$
$$P=\frac{(0.990)(0.0821)(298)}{1.00-(0.990\cdot 0.0318)}-\frac{(0.990)^{2}(1.36)}{(1.00)^{2}}$$
$$P=23.68atm$$
Let's talk about Boyle's law! This law states that, at a constant temperature and number of moles, the volume (V) of a confined gas is inversely proportional to the pressure (P) exerted on the gas In other words, when volume increases, pressure decreases (and vice versa).
When we say that real gases deviate from the ideal gas law, it means that real gases do not follow Boyle's law, Charles's law, or Avogadro's law perfectly!
To make things simple, let's summarize the conditions for deviation from the ideal gas law!:
Now, you should have an understanding of ideal and real gases, and also the different equations that are used to analyze their behavior!
Chemicals that deviate from ideal gas laws are usually those at low temperatures and high pressure.
Molecules that deviate most from the ideal gas law are heavier molecules like Xenon.
At low temperatures, gases will deviate from the ideal gas law.
Deviations from ideal gas law are observed when gases are considered real gases and these gases are subjected to high pressures and low temperatures.
Deviation from ideal gas behavior is caused by low temperature and high pressure.
YOUR SCORE
Your score:
Already have an account? Log in
SIGNUP SIGNUPFlashcards in Deviation From Ideal Gas Law17
Start learningGases are commonly known for:
Assuming the shape and volume of a container
True or false: The kinetic molecular theory is used to explain the behavior of gas molecules.
True
If a gas molecule behaves according to the kinetic molecular theory, they are called:
Ideal gases
Definition of intensive property:
A type of property that does not depend on the amount of matter or substance present.
Definition of Extensive property:
Properties that depend on the amount of matter being measured
What is partial pressure?
The pressure exerted by an individual gas within a mixture.
Already have an account? Log in
60%
of the users don't pass the Deviation From Ideal Gas Law quiz! Will you pass the quiz?
Start QuizHow would you like to learn this content?
How would you like to learn this content?
Free chemistry cheat sheet!
Everything you need to know on . A perfect summary so you can easily remember everything.
The first learning app that truly has everything you need to ace your exams in one place
Sign up to highlight and take notes. It’s 100% free.
Save explanations to your personalised space and access them anytime, anywhere!
Sign up with Email Sign up with AppleBy signing up, you agree to the Terms and Conditions and the Privacy Policy of StudySmarter.
Already have an account? Log in
Explanations 
Exams 
Magazine 
