What Is a Vacuum?

This chapter provides the reader with the basic details of a vacuum. This gives an appreciation of why as the pressure falls the evaporation rate of materials will increase. Also, as the amount of gas is reduced (low pressure or increasing vacuum) the number of gas collisions is reduced and so the distance vapor will travel before being scattered is increased. This line of sight deposition helps to maximize the material collection on the substrate as fewer atoms are scattered away before they reach the substrate.

Keywords

Vacuum; pressure; partial pressure; vapor pressure; saturated vapor pressure; mean free path; monolayer; line of sight; gas collisions

1.1 Introduction

The word vacuum is derived from the Latin word ‘vacua’ meaning empty. If we empty the chamber of gas we produce a vacuum.

A vacuum could be described as where within an enclosed volume there is less gas per unit volume than is present in a similar volume in the atmosphere surrounding the enclosed volume. This is something that can be used to our advantage.

Those of us who drink tea have all heard tales of not being able to brew a good cup of tea when on the higher slopes of Mount Everest because of the lower pressure and the problem of boiling water at a lower temperature. This effect of reducing the boiling point of materials when under vacuum is one advantage that can be used to good effect.

Many materials, particularly when raised in temperature to the boiling point, are prone to oxidation. Thus another advantage of operating in a vacuum is that materials that would normally be excessively affected by oxidation can have oxygen and water vapor kept away. This is achieved by being within a volume where there are few gas molecules, that is, a vacuum.

1.2 What Is a Gas?

If we look at materials in general they can be in the form of a solid, liquid, or gas. The structure changes with each form. Solids have atoms closely spaced and in fixed positions. Heating the material, the form changes to a liquid where the atoms are disordered and the distance between atoms is greater. With further heating the disorder is still greater and the spacing also much greater. The speed of motion of the atoms also increases with temperature. So let us look at a few facts and figures about gases.

A gas is where atoms or molecules are free to move in any direction and are in constant motion. Typically, these particles are traveling at speeds of approximately 1650 kph (1000 mph).

In air, gas molecules occupy approximately 0.01% of the space as compared to a solid where the molecules occupy approximately 74% of the space. The particles collide with each other or surfaces at a rate of 10,000,000,000 per second. These collisions mean that the gas particles have random motion and will rapidly expand to fill the whole volume. If the number of collisions looks to be large bear in mind that there will be around 20,000,000,000,000,000,000 particles per cubic centimeter and the mean free path (mfp) (the average distance a particle has to travel before it hits another particle) is only 100 nm.

1.3 Pressure

All atoms or molecules have mass and when they hit and bounce off a surface they exert a force on that surface. This force per unit area is known as pressure.

=ForceArea

All atoms or molecules in the atmosphere with their mass are subjected to the gravitational pull of the earth and are attracted to the earth. Thus at high altitudes the pressure is lower because the density of gas is lower.

At the top of Everest, the pressure is less than one-third than that found at sea level. A simple rule of thumb is that the pressure is halved every 5 km away from the earth surface (sea level).

Atmospheric pressure taken at sea level and 45°N latitude is 14.69 pounds per square inch (psi) or 1 kilogram force per square centimeter. If the pressure is taken at a temperature of 0°C the pressure is said to be 1 standard atmosphere (1 std atm)

Atmospheric pressure as a value is often rounded up for convenience.

psi~15psi15psi~760Torr~1stdatm~101325pascal(N/m2)~1.01325bar101.325kPa~1013.25mbar

There is an issue regarding the units used to designate pressure that can be problematic. As long ago as 1978 the Pascal replaced Torr as the acceptable measure of pressure. However, many vacuum systems that were built before this time are still in regular use and also old habits die slowly and so it is still common to find systems using Torr as the measure of pressure.

The Italian, Torricelli, in 1644 made a measurement of pressure using a mercury manometer. His measurement of 760 mm Hg for atmospheric pressure is the basis of the Torr scale used today.

mmHg=1Torr(anabbreviationofTorricelliinhonorofhiswork)1micronHg=1/1000mmHg=0.001mmHg=0.001Torr=1mTorr(mT)

The European preferred unit of pressure is the Pascal or the Newton per square meter. As ever, the Europeans have adopted a unit of measure of pressure that does not quite conform because of not being related to the SI unit by a factor of 103. The unit is the bar where the base unit of 1 bar equals 100000 Pa.

bar=1000mbar=750Torr

Vacuums are often categorized as one of four main types low, medium, high, or ultrahigh vacuum. There are other descriptors that are used that may not always be recognized.

=soft=poor=roughvacuum=1−1013mbarMedium=moderatevacuum=10−3−1mbarHigh=hard=goodvacuum=10−7−10−3mbarUltrahigh=below10−7mbar

It is worth noting that another confusion comes from the changing between describing a vacuum as a high vacuum and then changing to talking about a low-pressure system. Both are correct, a system having a low pressure does have a high vacuum, but it helps to be consistent in the terminology used.

1.4 Partial Pressure

The total pressure of a system is the sum of all the individual gas pressures of the gases present in the system. Each of the gases exerts a pressure and individually they are known as the partial pressures.

A common example is of a chamber open to atmosphere. The air around us is composed of a mixture of gases. Table 1.1 lists the gases present, their relative volumes, and the partial pressure of each. The partial pressure of each gas is relative to its percentage of the total volume; hence the partial pressure is the volume times the percentage present. Thus for oxygen it is 20.95% × 101325 pascal = 21227 pascal.

The most common gas pumped to produce a vacuum is air and Table 1.1 shows the most common constituents of air. It is important to note there is one omission from the table that has a big impact on vacuum systems and that is water vapor. Depending upon the local weather and temperature the water vapor content of the atmosphere can be between 0.6 and 6.0 weight percent. Hence in the table the figures are for a basic dry air and then the water vapor, as it is a variable, is given a range of values. If the water vapor present is at the upper end value, the other gases will have a slightly lower percentage volume and hence a slightly lower partial pressure than shown.

1.5 Vapor Pressure

A vapor is a gas that has a tendency to turn back into liquid. In the same way that gases hitting a surface produce a pressure, a vapor likewise bombards the surfaces and exerts a pressure. This is referred to as the vapor pressure.

If we take a pool of liquid such as water and leave it for some time it can be seen to have lost some water by evaporation. If we raised the temperature of the water this would be seen to happen more quickly. If we took the temperature up to the boiling point of the water, the water vapor would be seen as a cloud above the water and the water level would be seen to fall rapidly. So the vapor pressure gives an indication of the rate of evaporation of a material. Solids also have a vapor pressure but, as you would expect, they are tiny by comparison with liquids.

The vapor pressure of all gases is the same at the boiling point in atmosphere, 760 Torr, although the temperature at which they boil is different.

This becomes of practical interest when working with liquid precursors for chemical vapor deposition processes and also for physical...