Bond polarity and molecular shape (VSEPR)

Not all covalent bonds share electrons equally. When two identical atoms bond (like H-H), they share the electron pair perfectly evenly -- a nonpolar covalent bond. When two different nonmetals bond (like H-Cl), the more electronegative atom (chlorine) pulls the shared pair closer to itself, giving it a slightly negative charge while hydrogen becomes slightly positive -- a polar covalent bond. As the electronegativity difference grows even larger (as with a metal and nonmetal), the pull becomes so lopsided that one atom essentially takes the electron entirely, and the bond becomes ionic rather than covalent at all.

A useful, commonly taught scale: an electronegativity difference below 0.4 is nonpolar covalent, between 0.4 and 1.7 is polar covalent, and above 1.7 is considered ionic. These are not hard walls in nature so much as a practical way to classify bonds for now -- but they work well and match real compounds consistently.

VSEPR theory (Valence Shell Electron Pair Repulsion) predicts a molecule's 3D shape from one simple idea: electron groups around a central atom (whether bonding pairs or lone pairs) repel each other and spread out as far apart as possible. Two electron groups with no lone pairs spread to a straight line (linear, 180 degrees, like CO2). Three groups with no lone pairs spread into a flat triangle (trigonal planar, 120 degrees). Four groups with no lone pairs spread into a tetrahedron (109.5 degrees, like methane). If some of those four groups are lone pairs instead of bonds, the shape is named for where the ATOMS end up, even though the repulsion math still uses all the electron groups -- four groups with one lone pair gives a trigonal pyramidal shape (like ammonia, NH3), and four groups with two lone pairs gives a bent shape (like water, H2O).