William Kovarik
Radford University
and
Matthew E. Hermes
Kennesaw State University

 

Fuels and Society: b. Octane and Cracking

Back to 7. Poor Quality Gasoline

Back to a. Fuel Thermochemistry

Ahead to c. Fuel Thermodynamics

Back to Concept Map

Chemical Principles

4. Free Energy is the energy available to do work
5. Exothermic processes with decreasing entropy will be spontaneous at low temperatures, non-spontaneous at higher temperatures
6. Endothermic processes with increasing entropy will be non spontaneous at low temperature and spontaneous at high temperatures
7. A spontaneous reaction that CAN occur will occur only if the activation energy is present to take it through the transition state.

In chemistry we relate macroscopic, observable change to transformations at the molecular level that go unseen through the shorthand of symbolic representation. For the most part, we are comforted by the fact that we can at least observe the macroscopic.

Thermodynamics is troubling and demands a new level of trust and belief. In introducing the concept of Entropy (S), defined as a thermodynamic function that measures randomness or disorder, we ask the student to believe in a concept that can NOT be directly observed.

Nevertheless, Entropy is a state function, depending only upon the state of the system, not its history. And an understanding of entropy, or the degree of randomness is essential in understanding why a high compression engine is more efficient that a low compression engine.

Q. Think of a small syringe full of air. Cap it off so that none of the air can escape. Now push the plunger so that the volume is reduced by one half. The Gas Laws will tell us that if we decrease the volume by one half the pressure will double. But think about the disorder or randomness of the molecules inside the syringe. Have we increased or decreased the random nature, the freedom of movement as we compressed the air in the syringe?

 

Chemistry of Petroleum Cracking

To understand this section fully, you need to be familiar with the principles of thermochemistry and thermodynamics, kinetics and catalysis and the discussion of isomers described in the organic chemistry section of your chemistry text.

In the late 19th century, crude oil, as recovered from the ground, could be distilled to separate the liquid materials by their boiling points. Remember that crude oil is a mixture of thousands of individual chemical compounds, ranging from very low molecular weight hydrocarbons such as propane (C3H8) to individual hydrocarbons with 30 or more carbon atoms.

These early refiners discarded the low boiling fractions, the part we now call gasoline, and recovered the more valuable , higher boiling, kerosene for lighting and heating. But the internal combustion engine required a lower boiling fuel than kerosene because the fuel had to vaporize in the engine cylinder, and electrical service replaced much of the demand for kerosene.

So turn of the last century chemists developed the refinery practice of "cracking"; heating the crude to very high temperatures in the absence of air. They found they could produce additional amounts of the gasolene fractions by what seemed to be a process of rupturing chemical bonds to get smaller molecules.

But they also found that the hydrocarbons produced in the cracking process had no better quality; they knocked, as much as the fractions taken from the uncracked crude. They had learned that branched chain hydrocarbons improved fuel performance and the early fuel engineers developed an experimental measurement of fuel quality based on knocking called octane number.

Analysis of the cracking results showed that these cracked fuels consisted of about the same percentage of "branched-chain" hydrocarbons as the uncracked material. For example, the hexanes present (C6H14), were no more likely to have branches in the links among the six carbons than be a straight, linear chain.

We first want to show how the principles of chemistry explain how these high temperature cracking operations did not give branched chain substances.

Why Hexane Does Not Form Branched Chains in Cracking: Among the most common components of gasoline are the hexanes, (C6H14). A six carbon chain can be constructed in a number of configurations and many structural isomers are possible. The National Institute of Science and Technology lists for us thermochemical and other data on eleven of these hexane isomers.

We know from our thermochemistry and thermodynamics that we can determine a standard enthalpy and entropy of formation for any substance. These standards represents the properties of formation of the substance, from the elements. And the higher the standard enthalpy, the less stable, more energetic the species is. Standard enthalpy and entropy are state property, unaffected by the pathway by which it was generated.

Consider one potential reaction in a high temperature cracking process - the conversion of a straight-chain octane, n-hexane to its branched isomer, 2,2-dimethylbutane. The question is, if we had n-hexane in active transformation of one to the other and back, as shown below, what could we convert one to the other? Would this reaction be spontaneous? (Recognize, MANY other chemical reactions are possible; let us just choose this one for the moment:

(C6H14)n-hexane <=> (C6H14)2,2-dimethylbutane

We can get an idea on this by looking at standard states of 25 degrees Celsius and one atmosphere pressure operation in the gas phase. Remember that the cracking operation is FAR from the standard state.

We start with all n-hexane. We know that the spontaneity of a process comes from considering the Free Energy (G) and that for a process to be spontaneous, the value of delta G in the equation below must be negative:

delta G0 = delta H0 - T delta S0

We find at the NIST site, values for the standard enthalpies and entropies for the two substances in the gas phase:

  delta H0 kcal.mol S0 cal/mol*K
n-hexane -167 +389
2,2-dimethylbutane -186 +359

Enthalpy for the reaction:

(Change in enthalpy, products minus reactants)

delta H0 2,2-dimethylmethylbutane - delta H0 n-hexane = delta H

-186kcal/mol - (-167kcal/mol) = -19kcal/mol

Entropy for the reaction:

(Change in entropy, products minus reactants)

S0 2,2-dimethylmethylbutane - S0 n-haexane = delta S

359cal/mol*K - 389cal/mol*K = -30cal/mol*K

Free Energy:

delta G = delta H - T delta S

at standard temperature:

delta G0 = -19kcal - (298K)(-30cal/K) = -10kcal

So at room temperature, the reaction is spontaneous - that is it CAN occur. But please be warned. REmember your study of kinetics. A reaction that CAN occur, will NOT OCCUR if the activation energy is not available to bring the reactant to the "transition state". This process, just like many spontaneous process does not occur at modest temperatures.

Let us look higher temperatures, let us say 400 degrees Celsius, where we might expect the cracking to occur, we find:

delta G0 = -19kcal - (673K)(-30cal/K) = +1kcal

So the process of branching at 400 degress Celsius is not spontaneous.

A process can be spontaneous at a low temperature - and not occur because of kinetic reasons such as the rate is so low because a large activation energy is required.

The same process can reach a temperature at which it will not take place because it is no longer spontaneous. The negative effects of decreasing entropy overcome the effect of an exothermic process.

Why Cracking Works - Why Hexane Cracks to Lower Molecular Weight Fragments: Let us propose an alternate equation for hexane behavior on heating. Let us write a breakdown of the hexane again to give two products, propane and propene. The balanced equation is:

C6H14g(Hexane) --> C3H8g (Propane)+ C3H6g(Propene)

  delta H0 kcal.mol S0 cal/mol*K
n-hexane

-167

+389
propane -104 +270
propene +21 +267

Enthalpy for the reaction:

(Change in enthalpy, products minus reactants)

delta H0 (propane+propene) - delta H0 n-hexane = delta H

-83kcal/mol - (-167kcal/mol) = 84kcal/mol

Entropy for the reaction:

(Change in entropy, products minus reactants)

S0 (propane+propene) - S0 n-haexane = delta S

537cal/mol*K - 389cal/mol*K = 148cal/mol*K

Free Energy:

delta G = delta H - T delta S

at standard temperature:

delta G0 = 84kcal - (298K)(148cal/K) = 44kcal

So at room temperature, the reaction is not spontaneous.

Let us look higher temperatures, let us say 400 degrees Celsius, where we might expect the cracking to occur, we find:

delta G0 = 84kcal - (673K)(148cal/K) = -16kcal

So the process of splitting at 400 degress Celsius is spontaneous.

A process can be non spontaneous at a low temperature - because it requires heat.

The same process can reach a temperature at which it will be spontaneous. The endothermic heat properties are overcome by the increased entropy of the products. In this reaction, we are making two molecules and that factor of increased disorder allows the cracking process to occur.

1. From the principles of isomerism in organic compounds, recall that carbon atoms form long chains. And because a carbon atom can bond with as many as four other carbon atoms, we find extensive branching among the possible "structural isomers".

1I. Thermochemical and other data on elements and compounds is catalogued and reported by the National Institute of Science and Technology (NIST).

III. You can see stick-figure representations of the carbon chains of these two isomers at the NIST page.

College of Science and Mathematics
Kennesaw State University
1000 Chastain Rd.
Kennesaw, GA 30114
770-423-6160
 
t_logo.gif (12525 bytes) ChemCases.com is a National Science Foundation supported curriculum development project.