Fuel combustion knowledge is fundamental to automotive engineering technology, as it directly affects engine performance, fuel efficiency, and emissions control. Understanding the chemical composition of fuels and the reactions they undergo during combustion enables engineers to optimize engine design and troubleshoot operational issues. In Kenya’s automotive sector, where diverse fuel types are used, mastery of combustion principles supports innovation in vehicle maintenance and adaptation to alternative fuels. This chapter explores the elements and compounds in fuels and the chemical equations governing their combustion.

4.1 Elements and compounds of fuel

Fuel composition decides how a fuel burns, how much energy it releases and which products leave the exhaust. Most fuels used in engines and boilers are made of a few elements, mainly carbon and hydrogen, with small amounts of oxygen, nitrogen and sulphur. Knowing the terms used to describe these substances, and being able to use atomic and molecular masses, is the first step towards writing and balancing combustion equations. This topic defines the key terms and gives the table of elements and symbols used in fuel calculations.

4.1.1 Define of terms

The terminology of fuel chemistry gives a common language for describing substances and their reactions. Precise definitions show how elements, compounds, atoms and molecules are related, and how their masses are used in calculations. Example: petrol contains octane, a compound whose molecules are made of carbon and hydrogen atoms in a fixed ratio.

4.1.1.1 Element

An element is a pure substance made of only one kind of atom, and it cannot be broken down into simpler substances by chemical means. Each element has a unique atomic number and a chemical symbol. Carbon (C) and hydrogen (H) are the principal elements in hydrocarbon fuels. Example: a piece of coal is mostly the element carbon, which burns to release heat.

• Identity: Every atom of an element has the same number of protons, called the atomic number.
• Symbol: A one- or two-letter code, such as C for carbon and O for oxygen, represents the element.
• Pure substance: An element contains only one kind of atom.
• Fuel elements: Carbon, hydrogen and sulphur are the combustible elements in fuels.
• Supporting elements: Oxygen supports combustion, and nitrogen is usually inert.

4.1.1.2 Compound

A compound is a substance formed when two or more elements combine chemically in a fixed ratio. Its properties differ from those of the elements in it, and it can be split into its elements only by a chemical reaction. Hydrocarbons such as octane (C₈H₁₈) are compounds of carbon and hydrogen, and they are the main constituents of petrol and diesel. Example: when methane burns, the compound reacts with oxygen to form carbon dioxide and water, which are also compounds.

• Fixed ratio: The elements are always present in the same proportion, as in water, H₂O.
• Chemical bonds: The atoms are joined by bonds, so they cannot be separated by physical means.
• Hydrocarbon fuels: They are compounds made of hydrogen and carbon, and their combustion releases energy and produces carbon dioxide and water.
• Formula: A compound is written with a formula that shows the elements and their numbers.
• Products of combustion: Carbon dioxide (CO₂), carbon monoxide (CO) and water (H₂O) are common compounds in exhaust.

4.1.1.3 Atoms

An atom is the smallest particle of an element that keeps the chemical properties of that element. It has a central nucleus of protons and neutrons with electrons moving around it. Example: a carbon atom has 6 protons, 6 neutrons and 6 electrons in its most common form.

• Nucleus: It contains protons, which carry positive charge, and neutrons, which carry no charge.
• Electrons: Negative charges move around the nucleus and take part in chemical bonds.
• Atomic number: The number of protons identifies the element.
• Neutral atom: The numbers of protons and electrons are equal.
• Combining: Atoms join in chemical reactions to form molecules and compounds.

4.1.1.4 Molecules

A molecule is a group of two or more atoms held together by chemical bonds. It is the smallest unit of a compound that keeps the properties of that compound, and some elements also exist as molecules. Example: oxygen in the air exists as molecules of two atoms, O₂, and each molecule of octane has 8 carbon atoms and 18 hydrogen atoms.

• Molecular formula: It shows the exact number of atoms of each element in one molecule, such as C₈H₁₈ for octane.
• Elemental molecules: Oxygen (O₂), hydrogen (H₂) and nitrogen (N₂) are molecules made of identical atoms.
• Compound molecules: Water (H₂O) and carbon dioxide (CO₂) are made of different atoms.
• Subscripts: A small number after a symbol shows how many atoms of that element are in the molecule.
• In equations: Combustion equations are written in molecules, so that the numbers of atoms can be balanced.

4.1.1.5 Atomic Mass

Atomic mass is the mass of an atom of an element, expressed in atomic mass units (u), where 1 u is one twelfth of the mass of a carbon-12 atom. It is shown on the periodic table, and in calculations it is used in kilograms per kilomole. Example: carbon has an atomic mass of 12, so 12 kg of carbon is one kilomole.

Values commonly used in fuel calculations:

• Relative mass: It compares the mass of an atom with that of carbon-12.
• Round values: Fuel calculations usually use whole-number values such as those in the table.
• Link to the kilomole: One kilomole of an element has a mass in kilograms equal to its atomic mass.
• Use: It is needed to find molecular masses and the mass of fuel and air in combustion.
• Isotopes: The table gives the average value for the natural mixture of isotopes.

4.1.1.6 Molecular Mass

Molecular mass is the sum of the atomic masses of all the atoms in a molecule. It tells how heavy one kilomole of the substance is, and it is used to convert between mass and amount in combustion calculations. Example: the molecular mass of carbon dioxide is found by adding the atomic mass of one carbon atom and two oxygen atoms.

Formula: molecular mass of CO₂ = 1 × 12 + 2 × 16

Working: 12 + 32 = 44

Result: molecular mass of CO₂ = 44, so 1 kmol of CO₂ has a mass of 44 kg

Formula: molecular mass of octane C₈H₁₈ = 8 × 12 + 18 × 1

Working: 96 + 18 = 114

Result: molecular mass of octane = 114

Formula: molecular mass of water H₂O = 2 × 1 + 1 × 16

Working: 2 + 16 = 18

Result: molecular mass of water = 18

• Common values: O₂ = 32, N₂ = 28, CO = 28, H₂ = 2 and CH₄ = 16.
• Kilomole: The mass of 1 kmol in kilograms equals the molecular mass.
• Mass of a substance: Mass = number of kilomoles × molecular mass.
• Balanced equations: The total mass of the reactants equals the total mass of the products.
• Use in combustion: It is used to find the mass of oxygen needed and the mass of products formed.

4.1.2 Element and symbols table sketches

A table of elements and their symbols is a quick reference for writing fuel formulas and combustion equations. Below are the elements found in fuels, with their symbols, atomic numbers and roles in combustion. Example: an engineer writing the combustion equation for octane uses the symbols C, H and O from this table to write the formulas correctly.

Atomic symbols are used in combustion equations to represent the fuel and the oxygen. For instance, the formula C₈H₁₈ for octane uses these symbols to show its composition.

• Capital letters: The first letter of a symbol is a capital, and a second letter, if any, is small.
• Sketching the table: When sketching it from memory, list each element with its symbol, atomic number and atomic mass.
• Fuel elements: Carbon, hydrogen and sulphur burn, and oxygen supports burning.
• Air: It is mostly nitrogen and oxygen, which are written N₂ and O₂.
• Pollutants: Sulphur and nitrogen compounds in the exhaust cause pollution.

Chapter Summary

This chapter explored the fundamental elements and compounds that constitute fuels, providing clear definitions of terms such as element, compound, atoms, molecules, atomic mass, and molecular mass, alongside illustrative tables of element symbols. It then progressed to balancing combustion equations for both solid and gaseous fuels, emphasizing the importance of stoichiometric accuracy in combustion processes. The chapter detailed conversion analysis techniques, explaining how to determine the theoretical or minimum volume of air required for complete combustion and how to interconvert volumetric and gravimetric analyses of fuel-air mixtures. Calculations related to the mass of carbon in flue gases were covered, showing how to quantify carbon content per kilogram of exhaust gases. Further, methods to calculate the mass of dry flue gases per kilogram of fuel burnt were described by comparing carbon masses in fuel and flue gases. The chapter also examined how to determine the mass of excess air supplied using the mass of unused oxygen found in the flue gases and how to compute the total mass of air supplied to the combustion process. Finally, it introduced the Orsat apparatus for flue gas analysis, detailing its components, operation, and practical use supported by diagrammatic sketches.