Protons have a positive charge, neutrons have no electrical charge, and electrons have a negative charge.

An atom contains protons, neutrons and electrons. Protons have a positive electrical charge and neutrons have no electrical charge. Both are located in the nucleus. Electrons have a negative electrical charge and occupy the electron cloud surrounding the nucleus.

8 protons and 8 electrons.

The atomic number tells us the number of protons in the nucleus. Oxygen has atomic number 8, so it has 8 protons. A neutral atom has the same number of electrons as protons, so a neutral oxygen atom also has 8 electrons.

Positively charged protons repel each other, yet the nucleus remains together. This indicates that a much stronger attractive force acts between particles in the nucleus. Nuclear reactions involve changes in the nucleus and much larger energy changes than chemical reactions, which involve electrons.

Protons all have positive electrical charge and therefore repel one another. Nevertheless, protons and neutrons can remain bound together in a very small nucleus. This provides evidence that another attractive interaction – the strong nuclear force – must overcome the electrical repulsion at these very short distances.

Chemical reactions involve changes in how electrons are arranged and shared between atoms, while the atomic nuclei remain unchanged. Nuclear reactions change the nuclei themselves. The energies associated with nuclear binding are much larger than those associated with chemical bonds, which is why nuclear reactions can release far more energy than chemical reactions.

Atomic number is the number of protons in an atom's nucleus. It uniquely identifies each element—no two elements have the same atomic number, and it never changes for a given element.

The atomic number is the total count of protons in the nucleus. This uniquely defines each element. Atomic number is more fundamental than atomic mass because atomic mass can vary between isotopes (different numbers of neutrons), but the atomic number remains constant. The periodic table is organized by increasing atomic number, which creates the repeating patterns that make chemistry predictable.

Mendeleev predicted gallium (eka-aluminum) would have a density of 5.9 g/cm³. When gallium was discovered in 1875, its actual density was 5.91 g/cm³—almost exactly as predicted.

In 1869, Mendeleev recognized gaps in his periodic table and predicted that undiscovered elements would fill them, along with their properties. For gallium, he predicted an atomic mass of about 68 (actual: 69.7), density 5.9 g/cm³ (actual: 5.91 g/cm³), and that it would be a soft metal that melts at a low temperature. The spectacular confirmation of these predictions convinced the scientific community that the periodic table revealed a fundamental truth about the organization of matter.

Mendeleev arranged known elements by increasing atomic mass and grouped elements with similar chemical properties. When patterns suggested that an element was missing, he left a gap and predicted the properties of the unknown element. The later discovery of elements such as gallium, scandium, and germanium showed that the periodic table revealed real patterns in matter.

Mendeleev’s periodic table was not simply a list of elements. He arranged the known elements by increasing atomic mass while also paying close attention to similarities in chemical properties. This arrangement revealed repeated, or periodic, patterns. When the pattern indicated that an element should exist but no known element fitted the position, Mendeleev left a gap instead of forcing the table to fit the known elements.

He then used the surrounding elements and the repeating pattern to predict properties of the missing elements, including approximate atomic mass, density, and chemical behaviour. These predictions were later confirmed when elements such as gallium, scandium, and germanium were discovered. This was important because it showed that the periodic table was more than a convenient classification system: it reflected a real underlying order among the elements. In modern terms, that order is based on atomic number, which is the number of protons in the nucleus.

The formula is 2n², where n is the shell number. Shell 1: 2 electrons. Shell 2: 8 electrons. Shell 3: 18 electrons.

Each electron shell has a capacity determined by quantum mechanics. The formula 2n² gives: Shell 1 (K-shell): 2(1)² = 2 electrons; Shell 2 (L-shell): 2(2)² = 8 electrons; Shell 3 (M-shell): 2(3)² = 18 electrons; Shell 4 (N-shell): 2(4)² = 32 electrons. These numbers arise from fundamental quantum mechanical principles about how electrons can be arranged around an atom.

Valence electrons are electrons in the outermost shell. Carbon has 4 valence electrons. This matters because valence electrons determine how an atom bonds with other atoms.

Valence electrons are located in the outermost electron shell and are the electrons that participate in chemical bonding. Inner shell electrons are tightly held and rarely participate in reactions. Carbon has 6 total electrons: 2 in shell 1 and 4 in shell 2. Since shell 2 is the outermost, carbon has 4 valence electrons. By sharing 4 electron pairs with other atoms, carbon achieves a full outer shell. This 4-bond capability makes carbon extraordinarily versatile—it is precisely why carbon forms the basis of organic chemistry and why hydrocarbons show such enormous structural diversity.

Hydrogen has 1 valence electron and 1 vacant position—it needs 1 more electron to reach a full shell of 2, so it forms 1 bond. Carbon has 4 valence electrons and 4 vacant positions—it needs 4 more electrons to reach a full shell of 8, so it forms 4 bonds.

The number of bonds an atom forms equals the number of vacant positions in its outermost shell. Hydrogen has 1 electron in shell 1, which has capacity 2. So hydrogen has 1 vacant position—it forms exactly 1 bond. Carbon has 4 electrons in shell 2, which has capacity 8. So carbon has 4 vacant positions—it forms exactly 4 bonds. Once all 4 bonds are formed, carbon's outer shell contains 8 electrons and carbon is stable. This simple rule—atoms form as many bonds as they have vacant positions—is the key to predicting molecular structures throughout this course.

A covalent bond is a shared pair of electrons between two atoms. Atoms bond because sharing electrons allows both atoms to fill vacant positions in their outermost shell, giving both a more stable, full outer shell.

A covalent bond forms when two atoms each contribute one electron to a shared pair. Both atoms count that shared pair as part of their own outer shell, allowing both to get closer to having a full outer shell without either atom giving up an electron permanently. The bond holds the atoms together because both benefit from the arrangement. The number of bonds an atom forms equals the number of vacant positions in its outer shell. Hydrogen, with 1 vacant position, forms 1 bond. Carbon, with 4 vacant positions, forms 4 bonds.

A single bond is one shared electron pair. A double bond is two shared electron pairs. Double bonds are more reactive because the second shared pair is more loosely held and more easily disrupted by other molecules.

A single bond (C—C) involves one shared electron pair and is stable—atoms can also rotate freely around it. A double bond (C=C) involves two shared electron pairs between the same two carbons. The bond is stronger overall, but the second electron pair is more loosely held and more exposed to interaction with other molecules, making double bonds chemically reactive. This is why alkenes react more readily than alkanes, and why double bonds are important reactive sites in petroleum refining chemistry.

Hydrogen has 1 valence electron and 1 vacant position—it needs 1 more electron to reach a full shell of 2, so it forms 1 bond. Carbon has 4 valence electrons and 4 vacant positions—it needs 4 more electrons to reach a full shell of 8, so it forms 4 bonds.

The number of bonds an atom forms equals the number of vacant positions in its outermost shell. Hydrogen has 1 electron in shell 1, which has capacity 2. So hydrogen has 1 vacant position—it forms exactly 1 bond. Carbon has 4 electrons in shell 2, which has capacity 8. So carbon has 4 vacant positions—it forms exactly 4 bonds. Once all 4 bonds are formed, carbon's outer shell contains 8 electrons and carbon is stable. This simple rule—atoms form as many bonds as they have vacant positions—is the key to predicting molecular structures throughout this course.

Hydrocarbons are classified mainly into aliphatic and aromatic classes. Aliphatic hydrocarbons include alkanes, alkenes, alkynes, and naphthenes (cycloalkanes). Aromatic hydrocarbons contain benzene-type rings with delocalized electrons.

The broadest classification used in this module divides hydrocarbons into two main classes: aliphatic hydrocarbons and aromatic hydrocarbons. Aliphatic hydrocarbons are non-aromatic and may be open-chain or cyclic. This class includes alkanes, saturated open-chain hydrocarbons; alkenes, unsaturated hydrocarbons with at least one carbon–carbon double bond; alkynes, unsaturated hydrocarbons with at least one carbon–carbon triple bond; and naphthenes, also called cycloalkanes, which are saturated cyclic hydrocarbons. Aromatic hydrocarbons differ because they contain benzene-type ring systems with delocalized electrons. This distinction is important because it shows that naphthenes are part of the aliphatic class rather than a separate top-level class.

Naphthenes are saturated cyclic aliphatic hydrocarbons with only single bonds. Aromatics contain benzene-type rings with delocalized electrons. Naphthenes behave more like saturated hydrocarbons, while aromatics have special stability due to aromaticity.

Naphthenes and aromatics are both ring-containing hydrocarbons, but they are fundamentally different. Naphthenes are cyclic saturated hydrocarbons whose rings contain only single carbon–carbon bonds. Because of this, their chemistry resembles that of other saturated hydrocarbons. Aromatics, by contrast, contain benzene-type rings with delocalized electrons spread over the ring system. This delocalization gives aromatic compounds unusual stability and strongly influences their reactivity.

In petroleum chemistry, naphthenes are important because they may make up a significant fraction of crude oil and can be converted in reforming processes. Aromatics are important because they contribute to octane rating and are valuable chemical feedstocks, but some, especially benzene, are also toxic and regulated.

Carbon has four valence electrons and usually forms four covalent bonds. This allows it to form chains, branches, rings, and multiple bonds. These structural differences change physical properties, chemical reactivity, refining behavior, and the value of hydrocarbon streams.

Carbon’s four valence electrons and its tendency to form four covalent bonds make it the natural foundation of organic and hydrocarbon chemistry. With four bonds available, carbon can link to other carbon atoms in long chains, branched structures, cyclic rings, and aromatic systems, and it can form single, double, and triple bonds. Because of this, molecules made only of carbon and hydrogen (hydrocarbons) can have many different shapes, stabilities, and reactivities.

In petroleum chemistry, these structural differences determine boiling point and combustion behavior, influence octane rating, and define how hydrocarbons can be changed. Refining processes such as cracking, reforming, and hydrogenation have been developed specifically to break and rebuild carbon–carbon and carbon–hydrogen bonds in ways that turn the complex mixture in crude oil into useful products. The diversity of petroleum molecules follows directly from these bonding properties of carbon.

IUPAC stands for International Union of Pure and Applied Chemistry. Its naming rules follow a clear hierarchy: first choose the parent structure (chain or ring), then decide the main family (alkane, alkene, alkyne, cycloalkane), number the parent, add substituents, and finally assemble the name.

IUPAC (International Union of Pure and Applied Chemistry) is the worldwide organization that standardizes chemical nomenclature, symbols, and terminology. Before IUPAC rules, the same compound could have many local or historical names, making communication between laboratories, companies, and countries confusing.

IUPAC nomenclature solves this by making the name a direct description of the structure. For hydrocarbons, the rules are hierarchical: first choose the parent structure (the longest chain or main ring), then decide the hydrocarbon family from its bond types (alkane, alkene, alkyne, cycloalkane). Next, number the parent so important features get clear position numbers, identify and name any branches as alkyl substituents, and finally assemble the complete name.

This system ensures that any chemist who knows the rules can go from structure to name and from name back to structure, which is essential in petroleum technology documents and safety data sheets.

(a) Pentane; (b) But-1-ene or 1-butene; (c) Cyclohexane; (d) Ethene.

(a) Five carbons in a straight chain means the parent has 5 carbons, so the prefix is pent-. The chain contains only single bonds, so the family is alkane and the suffix is -ane. There are no branches or multiple bonds to locate, so the name is pentane (C₅H₁₂).

(b) A 4-carbon chain gives the prefix but-. The presence of one double bond means the family is alkene and the suffix is -ene. The double bond starts at carbon 1, so the locant is 1-. Combining these, the full name is but-1-ene (or 1-butene), which means a four-carbon parent with a double bond beginning at carbon 1.

(c) A 6-carbon saturated ring is a cycloalkane. The ring has 6 carbons, so the prefix is hex-, and the cyclo- prefix indicates a ring. With only single bonds, the suffix is -ane. The name is cyclohexane (C₆H₁₂), a saturated six-membered ring that is particularly important and common in petroleum.

(d) Ethylene is the common or commercial name for ethene. The molecule has two carbon atoms, so the prefix is eth-, and it contains a carbon–carbon double bond, so the suffix is -ene. Therefore, the corresponding IUPAC name is ethene (C₂H₄). It is an alkene, not an alkyne.

Common names can be useful but may not identify one exact structure. “Isooctane” is commonly associated with branched C₈H₁₈ hydrocarbons, but the IUPAC name 2,2,4-trimethylpentane identifies one specific molecule: a five-carbon parent chain with three methyl branches. This exact molecule is important because it is the reference compound used to define high octane rating.

Common or commercial names are widely used in petroleum and petrochemistry because they are short and familiar, but they do not always describe molecular structure precisely. The name “isooctane” is especially important in fuel discussions because it is associated with branched eight-carbon hydrocarbons that resist knocking in gasoline engines. However, as a common name, it can be less precise than a systematic IUPAC name because different C₈H₁₈ isomers can have different carbon skeletons.

The IUPAC name 2,2,4-trimethylpentane removes this ambiguity by describing the exact structure. The parent chain is pentane, meaning the longest selected chain has five carbon atoms. The name also shows that there are three methyl branches: two attached to carbon 2 and one attached to carbon 4. From the name alone, we can reconstruct the molecule and confirm that it has the molecular formula C₈H₁₈.

This exact molecule is important in octane rating because it is used as the high-octane reference compound. In octane rating, 2,2,4-trimethylpentane is used as the high-octane reference compound and is assigned an octane rating of 100 because it resists engine knocking very well.

The comparison shows why both naming systems matter: common names are useful in industry, but IUPAC names are needed when an exact molecular structure must be identified.

Full structural formulas show every atom and every bond. Condensed structural formulas write carbon atoms and attached hydrogens in a compact sequence. Skeletal formulas show the carbon framework as lines, with carbon atoms at line ends and corners and carbon-bound hydrogens implied.

A full structural formula shows every atom and every bond explicitly. It is very clear, but it becomes difficult to use for large molecules.

A condensed structural formula groups hydrogen atoms with the carbon atoms they are attached to, such as CH₃CH₂CH₃ for propane. This is faster to write while still showing the basic connectivity.

A skeletal formula, also called a line-angle formula, is the most compact representation. It is widely used in organic chemistry. In a skeletal formula, each line end and corner represents a carbon atom. Hydrogens bonded to carbon are not shown but are inferred from the usual four-bond pattern of carbon in stable hydrocarbons.

Atoms other than carbon and hydrogen, such as O, N, S, or Cl, are written explicitly, and hydrogens attached to such atoms are usually shown explicitly.

Two hydrogen atoms. The two lines represent two carbon-carbon bonds. In a stable hydrocarbon, carbon normally has four bonds in total, so the remaining two bonds are to hydrogen atoms.

In a skeletal formula, each line represents a bond. If two single lines meet at a vertex, the carbon atom at that vertex is bonded to two neighboring carbon atoms.

In stable hydrocarbons, carbon normally forms four bonds in total. Therefore, this carbon still needs two more bonds to complete its usual bonding pattern. Since hydrogen forms one bond, the carbon has two implied hydrogen atoms.

The same logic is used throughout skeletal formulas: count the bonds shown to a carbon atom and subtract from four to find the number of implied hydrogens. For example, an endpoint with one single line is usually a CH₃ group, and a vertex with three single lines is usually a CH group.

A molecular formula gives only the number of atoms, not how they are connected. C₆H₁₄ can represent five different structural isomers, including hexane and 2-methylpentane. Skeletal formulas show the connectivity directly. Benzene is a special skeletal-formula case: it is often drawn as a hexagon with a circle to show its delocalized aromatic electrons.

A molecular formula tells us only how many atoms of each element are present. It does not show which atoms are connected or how the carbon skeleton is arranged.

For example, the formula C₆H₁₄ can represent five different structural isomers. One is hexane, CH₃CH₂CH₂CH₂CH₂CH₃, which has a straight six-carbon chain. Another is 2-methylpentane, CH₃CH(CH₃)CH₂CH₂CH₃, which has a five-carbon chain with a methyl branch at carbon 2. Both molecules have the same molecular formula, but their structures and properties are different.

Skeletal formulas solve this problem by showing the carbon framework directly, while omitting carbon-bound hydrogens for clarity.

Benzene is a special case: it contains six carbon atoms in a ring and is often drawn as a hexagon with a circle inside. The hexagon shows the carbon skeleton, while the circle represents the delocalized aromatic electron system. This tells us that benzene is not simply an ordinary ring with three fixed double bonds, but an aromatic structure with equivalent carbon-carbon bonds and special stability.

Carbon has four valence electrons and usually forms four covalent bonds. This allows carbon atoms to connect to other carbon atoms and hydrogen in many different ways, forming straight chains, branched chains, rings, multiple bonds, and aromatic structures. This creates the large variety of hydrocarbons found in petroleum.

Carbon has six electrons, with four valence electrons in its outermost occupied shell. These valence electrons allow carbon to form four covalent bonds in stable organic compounds. Carbon can bond to hydrogen, but importantly it can also bond to other carbon atoms. Carbon atoms can therefore form straight chains, branched chains, and rings, and they can be connected by single, double, or triple bonds. Aromatic ring systems provide another important type of structure.

These different ways of connecting carbon and hydrogen atoms produce the different hydrocarbon families—such as alkanes, alkenes, alkynes, naphthenes, and aromatics—and also allow different isomers within the same molecular formula. Carbon’s electron structure and four-bond capability are therefore the starting point for understanding the enormous molecular diversity found in petroleum.

Covalent bonds hold atoms together within a hydrocarbon molecule, while weaker intermolecular forces act between separate molecules. Distillation overcomes intermolecular attractions and separates molecules without changing their chemical identity, while conversion processes such as cracking, reforming, and hydrogenation break, form, or rearrange chemical bonds.

Within a hydrocarbon molecule, carbon and hydrogen atoms are held together by strong covalent bonds. These bonds determine the molecular structure. Changing these bonds changes the molecule itself and therefore requires a chemical reaction.

Separate molecules also attract one another through much weaker intermolecular forces. These forces do not determine which atoms are connected, but they influence physical properties such as boiling point and viscosity.

This distinction is important in petroleum refining because different refining processes rely on these two types of interaction in different ways. In distillation, heating allows molecules to overcome intermolecular attractions and move into the vapour phase. The molecules themselves remain chemically unchanged, and different hydrocarbons can be separated because of differences in their physical behaviour.

In conversion processes such as cracking, reforming, and hydrogenation, chemical bonds are broken, formed, or rearranged, producing molecules with different structures and properties.

Molecules with the same molecular formula but different structures are isomers. Because their atoms are connected differently, they can have different physical and chemical properties. A molecular formula alone cannot distinguish them, so structural formulas, skeletal formulas, or systematic IUPAC names are needed. Their different properties can affect how they behave and how valuable they are in petroleum refining.

A molecular formula tells us only how many atoms of each element a molecule contains; it does not tell us how those atoms are connected. Two molecules can therefore have exactly the same molecular formula but different molecular structures. Such molecules are called isomers.

For example, butane and 2-methylpropane both have the molecular formula C₄H₁₀, but butane has a straight carbon chain while 2-methylpropane has a branched structure. Because molecular structure influences physical and chemical properties, isomers can differ in properties such as boiling point and in their behaviour as fuels or refinery components.

The molecular formula C₄H₁₀ cannot tell us which isomer is present. A structural or skeletal formula shows how the atoms are connected, while an IUPAC name such as 2-methylpropane systematically identifies the particular structure.

This illustrates an important principle that connects much of the course: electron structure determines bonding; bonding allows different molecular structures; molecular structure influences properties; and those properties influence how hydrocarbons behave in petroleum refining. This is why it is important to understand both molecular structure and the ways chemists name and represent it when studying refinery processes.