Chemistry for petroleum engineering
(PNGE3111; 3 credit hours)
Fall 2025
By
Dr. Khashayar Nasrifar
Department of Petroleum and Chemical
Engineering, SQU, Oman
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Introduction to organic chemistry
 Organic chemistry is the branch of chemistry that deals with the study of
carbon-containing compounds, particularly those involving carbonhydrogen (C-H) and carbon-carbon (C-C) bonds. While carbon is the
defining element in organic compounds, these compounds may also
include other elements such as oxygen, nitrogen, sulfur, and
halogens.
 Organic chemistry focuses on the structure, properties, composition,
reactions, and synthesis of these compounds.
 Due to carbon's unique ability to form stable covalent bonds with itself (a
property known as catenation) and with other elements, it can create a
vast variety of molecular structures, leading to the existence of
millions of organic compounds.
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Presence of organic compounds around us
 In Living Organisms: Proteins, carbohydrates, lipids/fats,
nucleic acids (DNA and RNA), hormones, enzymes, etc. in living
organism.
 Fuels: Fuels such as gasoline, diesel, fuel oils, natural gas, etc.
 Materials: Wood, paper, perfumes, carpets, clothes, etc.
 Plastics: such as polyethylene, polystyrene, polypropylene,
adhesives, foams, etc.
 Fruits, vegetables, wheat, rice, maize, etc.
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Presence of organic compounds around us
 Catenation is a chemical phenomenon where atoms of the same element
bond together to form long chains or rings through covalent bonds.
 This ability to form extended chains or rings is a unique property of
certain elements, particularly carbon. Carbon, due to its versatility in
bonding with other carbon atoms and various elements, exhibits strong
catenation properties. This results in the formation of diverse organic
compounds, including hydrocarbons, polymers, and complex organic
molecules, which play a fundamental role in chemistry and the structure of
organic matter.
 Other elements that can also form extended chains or rings: Si, S, B, P, O
and N.
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Why Carbon?
 Four covalent bonds allow for extensive bonding possibilities.
 Strong and stable C-C bonds support the formation of long chains and
rings.
 Small atomic size ensures strong bonds and efficient bonding with itself
and other elements.
 Versatility of bonding allows carbon to form single, double, and triple
bonds, increasing molecular diversity.
 Stable bonding with other elements enables a wide range of organic
molecules.
 No destabilizing tendency with multiple bonds compared to other
elements.
Together, these factors make carbon uniquely suited for catenation, forming the
backbone of organic chemistry and leading to the immense diversity of organic
compounds.
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Atomic Structure of Atoms (General
Concepts)
 Basic Atomic Model: Atom consists of protons, neutrons, and
electrons
Electrons (-ve
Charge)
Fundamental particles
1/1836th of a proton’s
mass
Chemical Bonding and
Reactivity
Electricity and
Conductivity
Neutrons (no
charge)
Contributes to Mass
Isotopes and Stability
Stabilizes the Nucleus
Involved in nuclear
reactions
Protons (+ve
Charged)
Define the element (atomic
number (Z))
Contribute to mass (1
amu)
Balance charge
Involved in nuclear
reactions
Bohr’s model
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Electron Configuration
Electron configuration describes the distribution of electrons among the
various orbitals around an atom's nucleus. This distribution follows
specific rules based on quantum mechanics, and understanding these rules
helps explain the atom’s chemical properties, such as bonding and
reactivity.
Energy Levels (Shells): Electrons in an atom are arranged in energy levels
or shells, denoted by the principal quantum number (n). These shells can
be thought of as layers surrounding the nucleus, with higher numbers (n = 1,
2, 3, 4, etc.) representing energy levels farther from the nucleus.
 The first shell (n = 1) is closest to the nucleus and has the lowest energy.
 The second shell (n = 2) is farther out and has higher energy, and so on.
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Electron Configuration
Subshells (Orbitals)
Each energy level (shell) is further divided into subshells, denoted as s, p, d,
f, which represent different types of atomic orbitals.
These orbitals have different shapes and energy, and each can hold a
specific number of electrons.
 s orbitals: Spherical in shape, can hold 2 electrons.
 p orbitals: Dumbbell-shaped, can hold 6 electrons (2 electrons per
orbital, with 3 orbitals: px, py, pz).
 d orbitals: More complex shapes, can hold 10 electrons (5 orbitals, each
holding 2 electrons).
 f orbitals: Even more complex shapes, can hold 14 electrons (7 orbitals,
each holding 2 electrons).
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Electron Configuration
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Electron Configuration
Aufbau Principle
The Aufbau Principle dictates the order in which orbitals are filled:
Electrons occupy orbitals from lower energy to higher energy.
Lower energy subshells (like 1s) are filled before higher ones (like 2s, 2p).
Orbital Filling Order:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s
< 5f < 6d < 7p
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Electron Configuration
Hund's Rule
Hund's Rule states that, in a subshell, each orbital is singly occupied
before it is doubly occupied. The electrons present in singly occupied
orbitals possess identical spin.
.
 Carbon (C)
Atomic number = 6, so it has 6 electrons.
Electron configuration: 1s² 2s² 2p²
 Oxygen (O)
Atomic number = 8, so it has 8 electrons.
Electron configuration: 1s² 2s² 2p4
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Quantum Numbers
Quantum Numbers: Defining Electron Behavior within an Atom
Quantum numbers are essential in determining the position and behavior
of electrons in an atom. They provide a detailed "address" for each
electron, specifying its location and characteristics in terms of energy,
shape, orientation, and spin.
There are four quantum numbers that describe every electron:
1. Principal Quantum Number (n)
The principal quantum number defines the main energy level (shell)
where an electron resides. It is denoted by n and can take any positive
integer value: n = 1, 2, 3, 4,....The value of n determines the distance of the
electron from the nucleus. Electrons with lower n values are closer to the
nucleus, while higher n values mean the electron is farther away, thus
having more energy.
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Quantum Numbers
Angular Momentum Quantum Number (l)
The angular momentum quantum number (denoted by l) determines
the shape of the orbital where the electron is found.
The value of l depends on the principal quantum number n, and it
can take integer values from 0 to n-1. For example, if n = 3, then l
can be 0, 1, or 2.
Orbital Shape (s, p, d, f): The value of l corresponds to the type of
orbital:
l = 0: s orbital (spherical shape).
l = 1: p orbital (dumbbell shape).
l = 2: d orbital (cloverleaf shape).
l = 3: f orbital (complex shape).
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Quantum Numbers
Magnetic Quantum Number (ml)
The magnetic quantum number (denoted by ml) describes the orientation
of the orbital in space relative to an external magnetic field. The value of
ml depends on the value of l, and it can take integer values from - l to + l. For
example, if l = 1 (p orbital), then ml can be -1, 0, or +1, which corresponds to
the three different orientations of p orbitals (px, py, pz).
Orbital Orientation:
 For s orbitals (l = 0), there is only one orientation (ml = 0) because they are
spherical.
 For p orbitals (l = 1), there are three possible orientations: px, py, and pz,
which represent the three axes along which the p orbitals align.
 d and f orbitals have more orientations due to their more complex shapes
(d orbitals have five orientations, and f orbitals have seven).
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Quantum Numbers
Spin Quantum Number (ms)
 The spin quantum number (denoted by ms) describes the intrinsic
angular momentum (spin) of the electron.
 The spin quantum number can take one of two values: +½ or -½,
which are often referred to as "spin up" and "spin down.
 No two electrons in the same atom can have the same set of
quantum numbers (Pauli Exclusion Principle), meaning that if two
electrons occupy the same orbital, they must have opposite
spins. Electrons with opposite spins can pair up within an orbital,
while unpaired electrons in an orbital contribute to the atom’s
magnetic properties.
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Quantum Numbers
 Principal Quantum Number (n): Determines the energy level or
shell in which the electron resides. It tells us how far the electron is
from the nucleus.
 Angular Momentum Quantum Number (l): Describes the shape
of the orbital (s, p, d, f). This helps to define the space around the
nucleus where the electron is most likely to be found.
 Magnetic Quantum Number (ml): Describes the orientation of the
orbital in space (i.e., along which axis the electron’s orbital is
aligned).
 Spin Quantum Number (ms): Describes the intrinsic spin of the
electron, which can be either +½ (spin-up) or -½ (spin-down).
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Quantum Numbers
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Hybridization
The four bonds around C are of equal length and Energy
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Hybridization
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Hybridization
How to generate four equal orbitals?
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Hybridization
The concept of mixing atomic orbitals to form new hybrid
orbitals with different energies, shapes, etc., than the
component atomic orbitals) suitable for the pairing of
electrons to form chemical bonds in valence bond
theory.
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Hybridization
Hybridization in Methane (SP3 Hybridization)
For a tetrahedrally coordinated carbon e.g., methane CH4), the carbon
should have 4 orbitals with the correct symmetry to bond to the 4
hydrogen atoms
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Hybridization
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Hybridization
When an s and two p orbitals are mixed to form a set of
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three sp orbitals, one p orbital remains unchanged and
is perpendicular to the plane of the hybrid orbitals.
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Hybridization
(a) The orbitals used to form the bonds in
ethylene. (b) The Lewis structure for
ethylene.
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Hybridization
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Hybridization
Hybridization in acetylene
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Hybridization
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Hybridization
Property
sp Hybridization
sp² Hybridization
sp³ Hybridization
Number of Hybrid Orbitals
2
3
4
Bond Angle
180° (Linear)
120° (Trigonal Planar)
109.5° (Tetrahedral)
Geometry
Linear
Trigonal Planar
Tetrahedral
Unhybridized Orbitals
2 p orbitals (available for π
bonds)
1 p orbital (available for π
bonds)
No unhybridized orbitals
Types of Bonds Formed
1 σ (sigma) bond and 2 π (pi)
bonds (triple bonds)
1 σ bond and 1 π bond
(double bonds)
Only σ bonds (single bonds)
Example Molecule
C₂H₂ (Ethyne/Acetylene)
C₂H₄ (Ethene/Ethylene)
CH₄ (Methane)
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Bonding in Organic Molecules
 In organic molecules, bonding is critical for understanding the structure,
stability, and reactivity of compounds. Organic molecules are primarily
composed of carbon, hydrogen, oxygen, nitrogen, sulfur, and sometimes
other elements like phosphorus, and halogens. The bonding in these
molecules involves covalent bonds formed by the sharing of electrons
between atoms. Here’s a detailed discussion of the types of bonding
commonly found in organic molecules:
1. Covalent Bonding
Covalent bonding is the primary type of bonding in organic molecules, where
two atoms share one or more pairs of electrons to achieve a stable
electron configuration (usually an octet).
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Bonding in Organic Molecules
a. Single Bonds (Sigma Bonds, σ)
Definition: A single bond is a σ bond formed by the head-on overlap of
atomic orbitals. Each atom contributes one electron to the shared pair,
and the bond is symmetric around the bond axis.
Example: In methane (CH₄), carbon forms four single bonds with hydrogen
atoms. Each C-H bond is a σ bond formed by the overlap of the sp³ hybrid
orbital of carbon with the 1s orbital of hydrogen.
Characteristics: Single bonds are the strongest type of covalent bonds and
allow for free rotation around the bond axis.
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Bonding in Organic Molecules
b. Double Bonds (One Sigma and One Pi Bond, σ + π)
Definition: A double bond consists of one σ bond and one π bond. The σ bond is formed by
the head-on overlap of orbitals, while the π bond is formed by the side-by-side overlap of
unhybridized p orbitals.
Example: In ethylene (C₂H₄), each carbon atom is sp² hybridized and forms a σ bond with
the other carbon atom. The remaining unhybridized 2pz orbital on each carbon forms a π bond
through side-by-side overlap, resulting in a C=C double bond.
Characteristics: Double bonds are shorter and stronger than single bonds, but they
restrict rotation due to the rigidity of the π bond. This rigidity gives rise to cis-trans
isomerism in molecules with double bonds.
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Bonding in Organic Molecules
c. Triple Bonds (One Sigma and Two Pi Bonds, σ + 2π)
Definition: A triple bond consists of one σ bond and two π bonds. The σ bond is
formed by the head-on overlap of orbitals, while the two π bonds are formed by the
side-by-side overlap of two pairs of unhybridized p orbitals.
Example: In acetylene (C₂H₂), the carbon atoms are sp hybridized. They form a σ
bond with each other, and the unhybridized 2py and 2pz orbitals on each carbon atom
overlap to form two π bonds, resulting in a C≡C triple bond.
Characteristics: Triple bonds are even shorter and stronger than double bonds.
They also restrict rotation and make the molecule linear.
Ethyne/
Acetylene
Propyne
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Bonding in Organic Molecules
1. Nonpolar Covalent Bonds:
Definition: A nonpolar covalent bond occurs when two atoms share
electrons
equally
because
they
have
the
same
or
very
similar
electronegativity values. The electron distribution is uniform, meaning there
are no partial charges (δ+ or δ−) across the bond.
Conditions for Nonpolar Covalent Bonds: When two atoms have equal
or nearly equal electronegativities (typically between 0 and 0.4 difference
in electronegativity). Common in molecules made up of identical atoms
(like H₂, O₂, N₂) or in hydrocarbons (where carbon and hydrogen have
similar electronegativities).
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Bonding in Organic Molecules
Examples:
Hydrogen gas (H₂): Two hydrogen atoms share their electrons equally,
making it a nonpolar bond.
Methane (CH₄): The carbon-hydrogen bonds in methane are considered
nonpolar because the difference in electronegativity between carbon (2.55)
and hydrogen (2.20) is very small.
Effect in Organic Molecules: Molecules with nonpolar covalent bonds
tend to be hydrophobic (insoluble in water) because they do not form
strong interactions with water molecules. Examples include alkanes, oils,
and fats.
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Bonding in Organic Molecules
2. Polar Covalent Bonds:
Definition: A polar covalent bond occurs when two atoms share
electrons unequally because one atom is more electronegative than the
other. The more electronegative atom attracts the shared electrons more
strongly, creating partial charges (δ− and δ+).
Conditions for Polar Covalent Bonds: When there is a moderate
difference in electronegativity between the two bonded atoms (typically
between 0.5 and 1.7 difference in electronegativity).This unequal
sharing results in one atom having a partial negative charge (δ−)
and the other having a partial positive charge (δ+).
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Bonding in Organic Molecules
Effect in Organic Molecules:
Polar covalent bonds are common in organic compounds that contain
atoms like oxygen, nitrogen, and halogens. These molecules can interact
with water and other polar substances, making them hydrophilic (watersoluble). Examples include alcohols (e.g., ethanol, CH₃CH₂OH) and
carboxylic acids (e.g., acetic acid, CH₃COOH).
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Bonding in Organic Molecules
3. Coordinate Covalent Bonding (Dative Bond)
Definition: A coordinate covalent bond is a type of covalent bond where both
electrons in the shared pair come from the same atom. This often occurs when
one atom has a lone pair of electrons, which it donates to form a bond with
another atom that has an empty orbital.
Example: In the ammonium ion (NH₄⁺), nitrogen in ammonia (NH₃) donates a lone
pair of electrons to bond with a proton (H⁺), forming a coordinate covalent bond.
In Organic Molecules: Coordinate covalent bonding is often seen in the formation
of complexes.
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Bonding in Organic Molecules
4. Ionic Bonds
 An ionic bond forms between a metal and a nonmetal, where one atom
transfers electrons to another, resulting in the creation of ions:
 Metals lose electrons to form cations (positively charged ions).Nonmetals
gain electrons to form anions (negatively charged ions).
 These oppositely charged ions are held together by strong electrostatic
forces, forming a crystal lattice structure in ionic compounds.
Key Characteristics:
 High melting/boiling points due to strong ionic bonds.
 Conduct electricity when molten or dissolved in water, but not in the solid
state.
 Brittle and tend to break under stress.
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Bonding in Organic Molecules
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Interactions in Organic Molecules
 Types of Interactions in Organic Molecules
1. London Dispersion Forces
o
Present in all molecules (e.g., hydrocarbons).
2. Dipole-Dipole Interactions
o
Occur between polar molecules (e.g., ketones, aldehydes).
3. Hydrogen Bonding
o
Critical in organic compounds with O-H and N-H bonds (e.g.,
alcohols, amines).
4. π-π Stacking
o
Common in aromatic systems (e.g., benzene rings).
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Interactions in Organic Molecules
1. London Dispersion Forces
Description:
Van der Waals forces, specifically London dispersion forces, are weak, temporary
attractions between molecules. These forces arise from momentary dipoles
created by fluctuations in electron distribution around atoms or molecules. All
molecules, regardless of polarity, exhibit Van der Waals forces, but these are the
only interactions in nonpolar molecules.
Strength: Weakest among intermolecular forces, but their strength increases
with the size and surface area of the molecules.
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Interactions in Organic Molecules
Example:
Hydrocarbons: In alkanes, the Van der Waals forces are the only interactions. As the
chain length of the alkane increases, the boiling point increases due to stronger
London dispersion forces.
Key Point: In large nonpolar molecules, like long-chain hydrocarbons, these forces
become significant enough to affect physical properties such as melting and boiling
points.
Energy: 0.05–40 kJ/mol, depending on the size and surface area of the molecules
involved.
Small nonpolar molecules: 0.05–2 kJ/mol.
Larger nonpolar molecules: 2–40 kJ/mol (due to increased surface area for
interactions).
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Interactions in Organic Molecules
2. Dipole-Dipole Interactions
Description: Dipole-dipole interactions occur between polar molecules, where the
positive end of one molecule is attracted to the negative end of another. These
interactions arise from the permanent dipoles present in polar molecules.
Strength: Stronger than Van der Waals forces but weaker than hydrogen bonds.
Example:
Ketones and Aldehydes:
In molecules Carbonyl compounds, the C=O bond creates a permanent dipole. The
oxygen atom, being more electronegative, attracts electrons, creating a dipole, and
molecules align so that oppositely charged ends are near each other.
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Interactions in Organic Molecules
Hydrogen Chloride (HCl): HCl molecules experience dipole-dipole attractions due to
the permanent dipole from chlorine's higher electronegativity compared to hydrogen.
Key Point:
Dipole-dipole interactions significantly affect boiling points and solubility. Polar
molecules tend to have higher boiling points than nonpolar molecules of similar size
due to these interactions, and tend to have higher solubility in polar solvents.
Energy: 5–25 kJ/mol
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Interactions in Organic Molecules
3. Dipole-induced dipole forces
 Description: Dipole-induced dipole forces, also known as Debye forces, occur
between a polar molecule (which has a permanent dipole) and a nonpolar
molecule (which lacks a permanent dipole). The permanent dipole of the polar
molecule can distort the electron cloud of the nonpolar molecule, inducing a
temporary dipole. This results in a weak attraction between the polar molecule and
the induced dipole of the nonpolar molecule.
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Interactions in Organic Molecules
 Strength: Weaker than dipole-dipole interactions and hydrogen bonds, but stronger
than London dispersion forces.
 Bond Energy: Typically ranges from 2 to 10 kJ/mol, depending on the polarizability
of the nonpolar molecule and the strength of the permanent dipole.
 Examples:
 Iodine (I₂) in ethanol: I₂ is nonpolar, but when dissolved in ethanol (polar), the
permanent dipole in ethanol induces a dipole in I₂, enhancing its solubility.
 Key Points: The strength of the interaction increases with the polarizability of the
nonpolar molecule (i.e., larger, more electron-rich molecules induce stronger
dipoles).
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Interactions in Organic Molecules
4. Hydrogen Bonding
 Description: Hydrogen bonding is a special type of dipole-dipole interaction that
occurs when a hydrogen atom is covalently bonded to a highly electronegative
atom (oxygen, nitrogen, or fluorine) and is attracted to another electronegative
atom in a nearby molecule.
 Strength: Stronger than dipole-dipole interactions but weaker than covalent or ionic
bonds
 Bond Energy: 10–40 kJ/mol
 Example:
 Water (H₂O): The hydrogen atoms in one water molecule
are attracted to the oxygen atom in neighboring water
molecules, leading to hydrogen bonding.
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Interactions in Organic Molecules
 Methanol (CH₃OH): The O-H bond in methanol enables
hydrogen bonding with other methanol molecules or
with water, making methanol miscible in water.
 Ammonia (NH₃): Hydrogen bonds form between
nitrogen in one ammonia molecule and hydrogen in
another.
 Key Point: Hydrogen bonding significantly raises
the boiling point, melting point, and solubility of
substances in water. It is critical in biological
molecules like DNA (between nitrogenous bases) and
proteins (in their secondary structures).
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Interactions in Organic Molecules
5. π-π Stacking (Aromatic Interactions)
 Description: π-π stacking refers to the interactions between aromatic rings, where
the π-electrons in the delocalized pi system of one ring are attracted to the πelectrons in another ring. This interaction is common in aromatic compounds.
 Strength: Weaker than hydrogen bonds but stronger than Van der Waals forces.
 Bond Energy: 5–10 kJ/mol
 Example:
 Benzene (C₆H₆): In benzene molecules,
the π-systems of the benzene rings
interact
with
one
another,
forming
stacked structures.
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Interactions in Organic Molecules
 π-π stacking between graphene and pyrene is an important mechanism for noncovalent functionalization and immobilization of various species.
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Interactions in Organic Molecules
 Nucleic Acids (DNA): In DNA, the stacking of aromatic bases (like adenine, guanine,
cytosine, and thymine) contributes to the overall stability of the double helix
structure.
 Key Point: π-π stacking is crucial in biological systems (e.g., in DNA stability) and in
materials science (e.g., organic semiconductors). Aromatic rings prefer to align in
a stacked fashion to maximize interactions between their delocalized πelectrons.
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Writing structural formulas
Different
Types of
Formulas
Empirical
Formula
The simplest wholenumber ratio of atoms
of each element in a
compound.
For glucose the
formula is CH₂O.
empirical
Molecular
Formula
The actual number
of atoms of each
element
in
a
molecule
or
compound
For Glucose C₆H₁₂O₆
Structural
Formula
It provides a detailed
representation of how
atoms are bonded and
arranged
in
a
molecule.
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Writing structural formulas
Or without bond
representation
CH3CH2OH
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Abbreviated structural formulas
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Writing structural formulas
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Writing structural formulas
59
Writing structural formulas
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Functional Groups in Organic Compounds
 A functional group is a specific group of atoms within a molecule that is
responsible for the characteristic chemical reactions of that molecule.
 Each functional group behaves in a predictable way in chemical
reactions, regardless of the rest of the molecule's composition. Thus,
molecules with the same functional group generally have the quite
similar chemical and physical properties .
 Functional groups are the key to understanding the reactivity and
properties of organic molecules. They determine how a molecule will
react with others and can give a molecule unique chemical behaviors.
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Functional Groups
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Classification of organics: functional groups
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Classification of organics: functional groups
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Classification of organics: functional groups
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Isomerism in Organic Compounds
 Isomerism refers to the phenomenon where compounds have the
same molecular formula but differ in the arrangement of atoms or
spatial orientation, resulting in different properties.
 Isomerism is a fundamental concept in organic chemistry because the
arrangement of atoms within a molecule can drastically alter its physical
and chemical characteristics.
 Types of Isomerism: Isomerism is broadly divided into two categories:
 Structural (or Constitutional) Isomerism: Structural isomers have the
same molecular formula but differ in the connectivity or bonding
arrangement of atoms. There are several types
 Stereoisomerism: Stereoisomers have the same molecular formula and
connectivity but differ in the spatial arrangement of atoms.
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Isomerism in Organic Compounds
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