Source: LibreTexts Organic Chemistry, Chapter 1.3 | Purdue University
Difficulty: Introductory
Prerequisites: Basic understanding of atomic structure (protons, neutrons, electrons, shells). If you are unsure what electron shells are or how the periodic table is organised, review Chapter 1.1 and 1.2 first.
This section explains why atoms react the way they do, and the answer comes down to valence electrons. Atoms "want" a full outer shell (a closed valence shell), and whether they gain, lose, or share electrons to get there determines whether they form ionic or covalent bonds. This is the foundation for understanding all of organic chemistry: every reaction mechanism you study later traces back to how electrons move between atoms.
Atoms with a full outer electron shell (s²p⁶) are stable and unreactive. Atoms that are one or two electrons away from a full shell are highly reactive: they will lose electrons (forming cations), gain electrons (forming anions), or share electrons (forming covalent bonds) to reach that stable configuration. Valence electrons can also absorb or release energy as photons, which drives atomic excitation and ionisation.
Valence electrons
The electrons in the outermost shell of an atom that participate in chemical bonding. Think of them as the electrons that are "available" for reactions.
Closed valence shell (noble gas configuration)
A full outer electron shell, corresponding to the electron configuration s²p⁶ (eight electrons in the outer shell, or two for helium). In simple terms, this is the stable arrangement that atoms are trying to achieve.
Open valence
An incomplete outer electron shell. An atom with open valences has room to gain, lose, or share electrons. This is what makes an atom reactive.
Cation
A positively charged ion formed when an atom loses one or more valence electrons. Think of it as: fewer electrons than protons means a net positive charge.
Anion
A negatively charged ion formed when an atom gains one or more valence electrons. More electrons than protons means a net negative charge.
Covalent bond
A bond formed when two atoms share one or more pairs of electrons, rather than transferring them outright. This is the dominant bonding type in organic chemistry.
Atomic excitation
When an electron absorbs energy (in the form of a photon) and jumps to a higher energy shell. The atom is now in an "excited state" rather than its ground state.
Ionisation
The process in which an electron absorbs enough energy to break free from its atom entirely, leaving behind a positive ion. This is excitation taken to its extreme.
Photon
A packet (quantum) of electromagnetic energy. Electrons absorb photons to gain energy and emit photons when they lose energy and drop to a lower shell.
An atom with a closed valence shell (s²p⁶ configuration) is chemically inert
The noble gases (He, Ne, Ar, Kr, Xe) all have closed shells
This is the most energetically stable arrangement for the outer electrons
All other atoms are trying to reach this closed-shell configuration through bonding
These atoms are highly reactive
They readily lose their "extra" valence electrons to form cations (positive ions)
Example: sodium (Na) has one valence electron beyond a closed shell. It easily loses that electron to become Na⁺
Example: magnesium (Mg) has two extra valence electrons. It loses both to become Mg²⁺
This is the basis of metallic behaviour in Groups 1 and 2 of the periodic table
These atoms are also highly reactive
They can gain electrons to form anions (negative ions)
Example: chlorine (Cl) needs one more electron to complete its shell. It gains one to become Cl⁻
Example: oxygen (O) needs two more electrons. It gains two to become O²⁻
Alternatively, they can share electrons with another atom, forming a covalent bond
This is the bonding type that dominates organic chemistry
Carbon, with four valence electrons, forms four covalent bonds to complete its octet
Valence electrons can absorb or release energy in the form of photons
When a valence electron absorbs energy:
It can jump to a higher (outer) energy shell, known as atomic excitation
If it absorbs enough energy, it can break free from the atom entirely, causing ionisation and forming a positive ion
When a valence electron loses energy:
It emits a photon
It drops to a lower (inner) shell that is not fully occupied
The number of valence electrons corresponds to an element's group (column) number:
Group 1: 1 valence electron
Group 2: 2 valence electrons
Groups 13–17: 3–7 valence electrons
Group 18: 8 valence electrons (closed shell, the 4s shell is complete and cannot hold more electrons)
This pattern repeats down the periodic table for each period
Neon signs work because of atomic excitation: an electric current excites the valence electrons of neon gas, and when those electrons fall back to lower energy levels, they emit photons of visible light. Different gases produce different colours because their electrons emit photons at different wavelengths.
The entire semiconductor industry depends on controlling valence electrons. Silicon has four valence electrons, and by "doping" it with atoms that have three or five, engineers create the p-type and n-type materials that make transistors and computer chips possible.
Students often think "inert" means an atom has no electrons. It does not. Inert atoms have a full set of valence electrons in a closed shell, which is exactly why they do not react.
Students frequently confuse cations and anions. Remember: a cat has paws (cat-ion is paws-itive). An anion is negative.
Students sometimes believe covalent bonds involve one atom "giving" electrons to another. They do not. In a covalent bond, electrons are shared between atoms. Transfer of electrons produces ionic bonds, not covalent ones.
Students often assume that atomic excitation and ionisation are the same thing. Excitation moves an electron to a higher shell within the atom. Ionisation removes the electron from the atom entirely. Excitation is reversible; ionisation requires significant energy input.
⚠️ You will almost certainly be asked to predict whether an atom will form a cation or anion based on its position in the periodic table. Know the rule: 1–2 electrons over a closed shell means it loses electrons (cation); 1–2 electrons short means it gains electrons (anion) or shares them (covalent bond).
⚠️ Expect questions that ask you to identify the number of valence electrons from an element's group number. This is foundational and comes up repeatedly in later chapters.
⚠️ The distinction between ionic bonding (electron transfer) and covalent bonding (electron sharing) is tested heavily in organic chemistry. Organic molecules are held together by covalent bonds, so understanding why atoms share electrons rather than transfer them is essential.
⚠️ Questions about atomic excitation and ionisation often appear as conceptual multiple-choice. Be clear on the difference: excitation is a jump within the atom, ionisation is escape from the atom.
True or False: An atom with a closed valence shell (s²p⁶) is highly reactive. (False. It is chemically inert.)
Fill in the blank: An atom that loses valence electrons forms a ____. (Cation)
True or False: A covalent bond involves the transfer of electrons from one atom to another. (False. Covalent bonds involve sharing electrons.)
Fill in the blank: When a valence electron absorbs a photon and jumps to a higher shell, this is called atomic ____. (Excitation)
True or False: Group 1 elements have 8 valence electrons. (False. Group 1 elements have 1 valence electron.)
Q: Sodium (Na) is in Group 1 of the periodic table. How many valence electrons does it have, and what type of ion does it most readily form?
A: Sodium has 1 valence electron. It readily loses that electron to form a cation (Na⁺) with a +1 charge, achieving a closed-shell configuration.
Q: Chlorine (Cl) is in Group 17. Explain two ways it can achieve a closed valence shell.
A: Chlorine has 7 valence electrons and needs one more to complete its octet. It can gain one electron to form the chloride anion (Cl⁻), or it can share an electron with another atom by forming a covalent bond (as in HCl or Cl₂).
Q: What is the difference between atomic excitation and ionisation?
A: In atomic excitation, an electron absorbs energy and jumps to a higher energy shell but remains bound to the atom. In ionisation, the electron absorbs enough energy to escape the atom entirely, leaving behind a positive ion. Both are triggered by energy absorption, but ionisation requires more energy.
Q: Why are noble gases chemically inert?
A: Noble gases have a closed valence shell (s²p⁶ configuration, or s² for helium). Because their outer shell is already full, they have no energetic incentive to gain, lose, or share electrons, so they do not form bonds under normal conditions.
Q: Carbon has 4 valence electrons. Why does it typically form covalent bonds rather than ionic bonds?
A: To form an ionic bond, carbon would need to either gain 4 electrons or lose 4 electrons, both of which require a large amount of energy. Sharing 4 electrons through covalent bonding is far more energetically favourable, which is why carbon forms four covalent bonds in most organic molecules.
This material connects directly to Lewis structures and the octet rule (Chapter 1.4), because drawing Lewis structures is just a visual way of tracking how valence electrons are shared or transferred between atoms.
Understanding cation and anion formation leads into electronegativity and bond polarity (Chapter 1.5+), where you will learn that electrons in a covalent bond are not always shared equally.
The concept of atomic excitation and photon emission connects to spectroscopy techniques (UV-Vis, IR) covered later in the course, which are used to identify organic compounds based on how their electrons interact with light.
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