sigma and pi bonds
Think of two ways to join two boards. You can press their ends head-on so the join sits right along the line between them — strong and hard to twist. Or you can lay them side by side and clamp along their lengths — a join above and below the line. Atoms bond in these same two geometric styles, and we call them sigma (σ) bonds and pi (π) bonds.
Precisely, a sigma bond forms when two orbitals overlap directly along the axis joining the nuclei; its electron density is concentrated on that line, so it is strong and lets the atoms rotate freely. A pi bond forms when two p orbitals overlap sideways, above and below the bond axis; its density sits off the axis, it is generally weaker, and it locks the atoms against rotation. Every single bond is one sigma bond; a double bond is one sigma plus one pi; a triple bond is one sigma plus two pi.
The distinction matters because pi bonds explain rigidity and reactivity. Because pi bonds resist twisting, double bonds make molecules flat and give rise to cis/trans isomers; because their exposed electrons are easier to attack, double and triple bonds are where many reactions happen. A caveat: the first bond between two atoms is always the sigma bond — pi bonds only come on top of an existing sigma bond, never alone.
The carbon–carbon triple bond in acetylene (HC≡CH) is one sigma bond holding the carbons together end-to-end, wrapped by two pi bonds at right angles to each other. The pi electrons form a tube of charge around the axis, which is why acetylene is so reactive and burns so hot.
A triple bond = one sigma + two pi; the pi electrons are the reactive part.
Bond order and bond strength are not perfectly proportional. A double bond is stronger than a single bond, but not twice as strong — its pi component is weaker than its sigma component. That is why it is often energetically favorable for a double bond's pi part to break and form two new single bonds in a reaction.