# Introduction to Chemical Engineering Processes/The most important point

## Component Mass Balance

[edit | edit source]Most processes, of course, involve more than one input and/or output, and therefore it must be learned how to perform mass balances on . The basic idea remains the same though. We can write a mass balance in the same form as the overall balance for each component:

For **steady state** processes, this becomes:

The **overall** mass balance at steady state, recall, is:

The mass of each component can be described by a similar balance.

The biggest difference between these two equations is that ** The total generation of mass is zero due to conservation of mass, but since individual species can be consumed in a reaction, for a reacting system**

## Concentration Measurements

[edit | edit source]You may recall from general chemistry that a **concentration** is a measure of the amount of some species in a mixture relative to the total amount of material, or relative to the amount of another species. Several different measurements of concentration come up over and over, so they were given special names.

### Molarity

[edit | edit source]The first major concentration unit is the **molarity** which relates the moles of one particular species to the total volume of the solution.

where

A more useful definition for flow systems that is equally valid is:

where

Molarity is a useful measure of concentration because it takes into account the volumetric changes that can occur when one creates a mixture from pure substances. Thus it is a very practical unit of concentration. However, since it involves volume, it can change with temperature so *molarity should always be given at a specific temperature*. Molarity of a gaseous mixture can also change with pressure, so it is not usually used for gasses.

### Mole Fraction

[edit | edit source]The **mole fraction** is one of the most useful units of concentration, since it allows one to directly determine the molar flow rate of any component from the total flow rate. It also conveniently is *always* between 0 and 1, which is a good check on your work as well as an additional equation that you can always use to help you solve problems.

The mole fraction of a component A in a mixture is defined as:

where signifies moles of A. Like molarity, a definition in terms of flowrates is also possible:

**Mole Fraction Definition**

If you add up all mole fractions in a mixture, you should always obtain 1 (within calculation and measurement error), because sum of individual component flow rates equals the total flow rate:

Note that **each stream has its own independent set of concentrations**. This fact will become important when you are performing mass balances.

### Mass Fraction

[edit | edit source]Since mass is a more practical property to measure than moles, flowrates are often given as *mass* flowrates rather than *molar* flowrates. When this occurs, it is convenient to express concentrations in terms of **mass fractions** defined similarly to mole fractions.

In most texts mass fraction is given the same notation as mole fraction, and which one is meant is explicitly stated in the equations that are used or the data given.

The definition of a mass fraction is similar to that of moles:

for batch systems

**Mass fraction of Continuous Systems**

where is the mass of A. It doesn't matter what the units of the mass are as long as they are the same as the units of the total mass of solution.

Like the mole fraction, the total mass fraction in any stream should always add up to 1.

## Calculations on Multi-component streams

[edit | edit source]Various conversions must be done with multiple-component streams just as they must for single-component streams. This section shows some methods to combine the properties of single-component streams into something usable for multiple-component streams(with some assumptions).

### Average Molecular Weight

[edit | edit source]The *average molecular weight* of a mixture (gas or liquid) is the multicomponent equivalent to the molecular weight of a pure species. It allows you to convert between the mass of a mixture and the number of moles, which is important for reacting systems especially because balances must usually be done in moles, but measurements are generally in grams.

To find the value of , we split the solution up into its components as follows, for k components:

where is the **mole fraction** of component i in the mixture. Therefore, we have the following formula:

where is the mole fraction of component i in the mixture.

This derivation only assumes that **mass is additive**, which it is, so this equation is valid for *any* mixture.

### Density of Liquid Mixtures

[edit | edit source]Let us attempt to calculate the density of a liquid mixture from the density of its components, similar to how we calculated the average molecular weight. This time, however, we will notice one critical difference in the assumptions we have to make. We'll also notice that there are **two** different equations we could come up with, depending on the assumptions we make.

#### First Equation

[edit | edit source]By definition, the density of a single component i is: The corresponding definition for a solution is . Following a similar derivation to the above for average molecular weight:

Now we make the assumption that **The volume of the solution is proportional to the mass**. This is true for any pure substance (the proportionality constant is the density), but it is further assumed that **the proportionality constant is the same for both pure k and the solution**. This equation is therefore useful for two substances with similar pure densities. If this is true then:

, where is the **mass fraction** of component i. Thus:

where is the **mass fraction** (not the mole fraction) of component i in the mixture.

#### Second Equation

[edit | edit source]This equation is easier to derive if we assume the equation will have a form similar to that of average molar mass. Since density is given in terms of mass, it makes sense to start by using the definition of **mass fractions**:

To get this in terms of only *solution* properties (and not *component* properties), we need to get rid of . We do this first by dividing by the density:

Now if we add all of these up we obtain:

Now we have to make an assumption, and it's different from that in the first case. This time we assume that the **Volumes are additive**. This is true in two cases:

1. In an **ideal solution**. The idea of an ideal solution will be explained more later, but for now you need to know that ideal solutions:

- Tend to involve similar compounds in solution with each other, or when one component is so dilute that it doesn't effect the solution properties much.
- Include Ideal Gas mixtures at constant temperature and pressure.

2 In a **Completely immiscible nonreacting mixture**. In other words, if two substances don't mix at all (like oil and water, or if you throw a rock into a puddle), the total volume will not change when you mix them. And the total volume in this case will be sum of volume of individual components.

If the solution is ideal, then we can write:

Hence, for an ideal solution,

where is the **mass fraction** of component i in the mixture.

Note that this is significantly different from the previous equation! This equation is more accurate for most cases. In all cases, however, it is most accurate to look up the value in a handbook such as Perry's Chemical Engineers Handbook if data is available on the solution of interest.