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Covalent bonds form when atoms share electrons to fill their valence shells. Bonds can be nonpolar (equal sharing between identical atoms) or polar (unequal sharing based on electronegativity difference). Multiple bonds (double, triple) occur when atoms share more than one pair of electrons. Bond strength depends on both bond type and atomic size.
From your study of periodic trends, you know that atoms on the right side of the periodic table have high electronegativities and need only a few electrons to complete their valence shells. These atoms — carbon, nitrogen, oxygen, fluorine, and their neighbors — are unlikely to give up electrons entirely to form cations. Instead, when two such atoms come together, they reach stability by sharing electron pairs rather than transferring them. This mutual sharing is a covalent bond, and it is the dominant bonding mode in molecular compounds, from water to DNA.
Consider the simplest case: two hydrogen atoms approaching each other. Each has one electron and needs two for a filled 1s shell. When they share their electrons, both atoms simultaneously "see" two electrons in the space between the nuclei. This shared pair is attracted to both positive nuclei at once, pulling the atoms together. The distance where the attractive and repulsive forces balance is the bond length, and the energy you would need to pull the atoms apart is the bond energy. A single shared pair makes a single bond. When atoms need to share more electrons — as in O₂ or N₂ — they form double bonds (two shared pairs) or triple bonds (three shared pairs), which are progressively shorter and stronger.
Not all sharing is equal. When two identical atoms bond — H₂, Cl₂, O₂ — each atom pulls on the shared electrons with equal force, producing a nonpolar covalent bond with electron density distributed symmetrically. But when atoms with different electronegativities bond, the more electronegative atom pulls the electron density toward itself. In H–Cl, chlorine's higher electronegativity draws the shared pair closer, creating a polar covalent bond with partial charges: δ+ on hydrogen, δ− on chlorine. The degree of polarity depends on the electronegativity difference — a small difference gives a slightly polar bond, while a very large difference approaches ionic character. This continuum from nonpolar covalent to polar covalent to ionic is not three separate categories but a smooth spectrum determined by the periodic properties you already understand.
Bond strength follows predictable patterns rooted in the periodic table. Bonds between small atoms are stronger than bonds between large atoms because the shared electrons are closer to both nuclei and held more tightly. A C–C single bond (~348 kJ/mol) is stronger than a Si–Si bond (~226 kJ/mol) for exactly this reason. Multiple bonds between the same pair of atoms are stronger than single bonds — the C≡C triple bond (~837 kJ/mol) is much stronger than C=C (~614 kJ/mol) or C–C — though not simply three times as strong, because the second and third pairs occupy less favorable bonding regions. These bond energies matter because they determine which reactions are energetically favorable: breaking strong bonds requires energy input, and forming them releases energy.