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Valence electrons are those in the outermost shell and primarily determine an element's chemical reactivity. Elements in the same group have the same number of valence electrons and thus similar chemical properties.
From electron configurations, you know that electrons fill orbitals in a specific order and that each element has a characteristic arrangement of electrons across its energy levels. Valence electrons are the subset that occupy the outermost (highest principal energy level) shell — and they are the electrons that do almost all the chemical work. Core electrons, buried deep inside the atom and tightly bound to the nucleus, are shielded from neighboring atoms and rarely participate in bonding. It is the valence electrons, sitting on the atom's surface so to speak, that interact with other atoms to form bonds, get transferred, or get shared.
The periodic table, which you already know how to navigate, encodes valence electron count directly. Every element in Group 1 has one valence electron; every element in Group 17 has seven. This is why elements in the same group behave so similarly: sodium and potassium are both soft, reactive metals that lose one electron easily, because they both have a single valence electron. Chlorine and bromine are both reactive nonmetals that gain one electron readily, because they both need just one more to complete their valence shell. The group number (for main-group elements) essentially tells you the valence electron count, making the periodic table a map of chemical behavior.
Reactivity patterns follow directly from how close an atom is to achieving a filled valence shell — the stable configuration of a noble gas. Atoms with one or two valence electrons (like sodium or magnesium) find it energetically favorable to lose those electrons entirely, forming positive ions and exposing the already-complete shell underneath. Atoms with six or seven valence electrons (like oxygen or fluorine) find it favorable to gain one or two electrons to complete their shell. Atoms in the middle — with three, four, or five valence electrons — tend to share electrons through covalent bonding rather than fully transferring them, because neither gaining nor losing several electrons is energetically practical.
This framework explains why noble gases (Group 18) are famously unreactive: their valence shells are already full, so they have no driving force to gain, lose, or share electrons. It also explains trends within groups — for instance, reactivity increases going down Group 1 because the valence electron is farther from the nucleus and easier to remove. Understanding valence electrons transforms the periodic table from a wall of symbols into a predictive tool: given any main-group element's position, you can anticipate how many bonds it will form, what ions it will produce, and which other elements it will react with most vigorously.