Difficulty: Introductory / Foundational
Prerequisites: A-level or AP-level general chemistry (periodic table, electron shells, basic bonding). If you are shaky on electron configuration or the difference between ionic and covalent bonds, review your general chemistry notes first.
Organic chemistry is the chemistry of carbon compounds, and this opening chapter lays the groundwork you will rely on for the entire course. Before you can understand why molecules react the way they do, you need a solid grasp of how atoms are built, how their electrons are arranged, and how those electrons form bonds. Everything here, from orbital shapes to hybridisation to electronegativity, feeds directly into reaction mechanisms, stereochemistry, and spectroscopy in later chapters.
Atoms have electrons arranged in orbitals of increasing energy. These orbitals can overlap to form covalent bonds (sharing electrons) or atoms can transfer electrons to form ionic bonds. The geometry around a bonded atom depends on how many groups and lone pairs surround it, which in turn depends on how the atom's orbitals hybridise (sp, sp2, or sp3). Polar bonds arise when atoms of different electronegativity share electrons unevenly, producing dipole moments.
Atomic number
The number of protons in the nucleus of an atom. In a neutral atom, it also equals the number of electrons. Think of it as the atom's identity number on the periodic table.
Orbital
A mathematical function (wavefunction, ψ) that describes the wave motion of an electron. The square of the wavefunction, ψ², gives the probability of finding the electron at any point in space. In simple terms, an orbital is the region where an electron is most likely to be found.
Principal quantum number (n)
An integer (1, 2, 3, ...) that designates the energy level (shell) of an orbital. Higher n means higher energy and greater average distance from the nucleus. Think of it as the "floor number" in a building: floor 1 is closest to the ground.
Sublevel quantum number (s, p, d, f)
Designates the shape of an orbital within a given principal level. s orbitals are spherical, p orbitals have two lobes, d orbitals have more complex shapes. In simple terms, this tells you the orbital's shape.
Node
A region where the electron density (ψ²) is zero. An orbital with principal quantum number n has n − 1 nodes. Nodes can be planar (like cutting through a p orbital) or spherical (concentric shells of zero density). Think of it as a dead zone where the electron will never be found.
Valence electrons
The electrons in the highest principal energy level of an atom. These are the electrons involved in chemical bonding. The number of valence electrons equals the group number in the "A" groups of the periodic table. In simple terms, these are the outermost electrons that do the chemistry.
Core electrons
All the electrons that are not valence electrons. They sit in lower energy levels and do not normally participate in bonding.
Pauli exclusion principle
No more than two electrons can occupy the same orbital, and if two do, they must have opposite spin. Think of it as: each orbital is a two-seater, and the passengers must face opposite directions.
Hund's rule
When filling orbitals of equal energy (degenerate orbitals), electrons occupy them singly before pairing up, and all single electrons have the same spin. In simple terms, electrons spread out before doubling up.
Aufbau principle
Electrons fill orbitals from lowest energy to highest energy. "Aufbau" is German for "building up."
Octet rule
Atoms tend to gain, lose, or share electrons so that each is surrounded by eight valence electrons (two for hydrogen). This is the driving force behind most bonding in organic chemistry.
Wavefunction (ψ)
The mathematical function that describes the wave behaviour of an electron in an orbital. Electron density at any point is proportional to ψ². Think of it as the blueprint for where electrons live.
The atomic number of an element tells you the number of protons, which equals the number of electrons in a neutral atom.
Electrons occupy orbitals, which are arranged in principal energy levels (n = 1, 2, 3, ...).
Each principal level n contains n sublevel types. n = 1 has only 1s. n = 2 has 2s and three 2p orbitals. n = 3 has 3s, three 3p, and five 3d orbitals.
Orbital energies increase with n and, within a level, in the order s < p < d.
s orbitals are spherical. A 1s orbital has no nodes. A 2s orbital has one spherical node. A 3s orbital has two spherical nodes.
p orbitals have two lobes separated by a planar node at the nucleus. The three 2p orbitals are oriented along the x, y, and z axes. A 3p orbital has both a planar node and a spherical node.
The number of nodes in any orbital equals n − 1.
Nodes are regions of zero electron density. They can be planar (flat surfaces through the nucleus) or spherical (concentric shells around the nucleus).
The electron density at any point is given by ψ² (the square of the wavefunction).
The probability of finding an electron in a spherical shell at distance r from the nucleus is given by r²ψ². This is the radial probability distribution.
For a 1s orbital, the maximum in r²ψ² occurs at 52.9 pm, which is the Bohr radius.
For a 2s orbital, the maximum probability is at about 280 pm, with a smaller peak nearer the nucleus and a spherical node at 110 pm.
Aufbau principle: fill orbitals from lowest to highest energy.
Pauli principle: maximum two electrons per orbital, with opposite spin.
Hund's rule: within a sublevel, occupy each orbital singly (same spin) before pairing.
Example: nitrogen (7 electrons) has the configuration (1s)²(2s)²(2p_x)¹(2p_y)¹(2p_z)¹. The three 2p electrons are unpaired.
Example: oxygen (8 electrons) is (1s)²(2s)²(2p_x)²(2p_y)¹(2p_z)¹. One 2p orbital is now paired.
Valence electrons are those in the highest principal energy level.
The number of valence electrons equals the group number for the "A" groups. Boron (Group 3A) has 3, carbon (Group 4A) has 4, nitrogen (Group 5A) has 5, oxygen (Group 6A) has 6, halogens (Group 7A) have 7.
When atoms form ions, they tend to gain or lose electrons to reach the electron configuration of the nearest noble gas (the octet rule for atoms, the duet for hydrogen).
Students often think orbitals are fixed paths or orbits, like planets around the sun. They are not. An orbital is a probability distribution: it tells you the likelihood of finding an electron in a given region, not a track it follows.
Students sometimes confuse the number of nodes with the number of lobes. A 2p orbital has one node and two lobes, not two nodes.
Hund's rule is frequently misapplied: students pair electrons in one 2p orbital before filling the other two. The correct approach is to place one electron in each 2p orbital (all same spin) before pairing any of them.
The octet rule is a guideline, not a law. Elements in the third period and beyond (like sulfur and phosphorus) can have expanded octets because they have available d orbitals. Boron commonly has only six electrons around it.
⚠️ You will almost certainly be asked to write electron configurations and identify valence electrons for common elements (C, N, O, halogens, S, P).
⚠️ Expect questions asking how many nodes an orbital has and what type (planar vs spherical). Remember: nodes = n − 1.
⚠️ Predicting whether an atom follows the octet rule, or has an expanded octet, is a recurring exam question. Know which elements can exceed eight electrons and why.
⚠️ Drawing Lewis dot structures of neutral elements and common ions (Br⁺, S²⁻, Mg²⁺) is foundational. If you cannot do this quickly, everything downstream will be harder.
True or false: A 3p orbital has two nodes.
True. It has one planar node and one spherical node (n − 1 = 3 − 1 = 2).
Fill in the blank: The number of valence electrons in a neutral oxygen atom is ______.
Six.
True or false: According to Hund's rule, the two electrons in nitrogen's 2p sublevel should be placed in the same 2p orbital with opposite spins.
False. Nitrogen has three 2p electrons, and each goes into a separate 2p orbital before any pairing.
Fill in the blank: The Bohr radius of a 1s orbital is approximately ______ pm.
52.9 pm.
True or false: A 1s orbital has one spherical node.
False. A 1s orbital has zero nodes (n − 1 = 0).
Q: Write the full electron configuration of a bromine atom (Z = 35). Which electrons are the valence electrons?
A: (1s)²(2s)²(2p)⁶(3s)²(3p)⁶(4s)²(3d)¹⁰(4p)⁵. The valence electrons are the 4s and 4p electrons (seven total). Even though the 3d orbital has a lower energy than 4s, the 4s and 4p electrons are the ones involved in the chemistry of bromine.
Q: A 4s orbital should look like a series of concentric spheres. How many spherical nodes does it have, and how does its size compare to a 3s orbital?
A: A 4s orbital has three spherical nodes (n − 1 = 3). It is larger than a 3s orbital, which has only two spherical nodes.
Q: Why do nitrogen and phosphorus have the same number of valence electrons despite being in different periods?
A: Both are in Group 5A, so both have five valence electrons. The difference is that nitrogen's valence electrons are in the n = 2 level, while phosphorus's are in the n = 3 level.
Q: Sulfur has the electron configuration Ne²(3p)⁴. Draw the dot structure of a neutral sulfur atom and explain why two of its 3p electrons are paired.
A: Sulfur's 3s orbital is full (paired). Of its four 3p electrons, by Hund's rule two 3p orbitals each get one unpaired electron, and the third 3p orbital gets the remaining two electrons as a pair. The dot structure shows two paired dots and two single dots around the S symbol.
Q: What is the difference between a planar node and a spherical node? Give one example of each.
A: A planar node is a flat surface (plane) passing through the nucleus where electron density is zero; every p orbital has at least one planar node. A spherical node is a concentric shell around the nucleus where electron density is zero; the 2s orbital has one spherical node at about 110 pm from the nucleus.
This material connects directly to Chapter 1's later sections on chemical bonding (Part 2 of these notes) and molecular geometry and hybridisation (Part 3). Understanding electron configuration and orbital shapes is essential for grasping how atoms form covalent and ionic bonds, and for predicting molecular geometry with VSEPR and hybridisation models. It also underpins spectroscopy (IR and NMR) later in the course, where the electronic environment around atoms determines the signals you observe.
atomic number, electron configuration, orbital, wavefunction, psi, ψ², principal quantum number, sublevel, s orbital, p orbital, d orbital, node, planar node, spherical node, valence electrons, core electrons, Pauli exclusion principle, Hund's rule, aufbau principle, octet rule, Bohr radius, radial probability distribution, electron density, noble gas configuration, Lewis dot structure, periodic table groups