bond polarity
Picture a tug-of-war where one team is a little stronger: the rope's center never crosses to either side completely, but it settles closer to the stronger team. In a covalent bond, when one atom pulls the shared electrons harder than the other, the electron cloud shifts toward it. That uneven sharing is bond polarity.
Precisely, bond polarity is the separation of electric charge along a bond caused by a difference in the two atoms' electronegativity. The more electronegative atom gains a small partial negative charge (written δ−) and the other a small partial positive charge (δ+). The bigger the electronegativity difference, the more polar the bond; if the two atoms are identical, there is no difference and the bond is nonpolar.
Bond polarity matters because it gives molecules their 'sticky ends' — the partial charges that drive attraction between molecules, dissolving, and how molecules dock together in biology. The caveat is that a molecule with polar bonds is not always a polar molecule overall: if the bond polarities point in opposite directions and cancel by symmetry, the whole molecule can be nonpolar despite having polar bonds.
Carbon dioxide (O=C=O) has two strongly polar C=O bonds, yet the molecule is nonpolar. Because it is linear and symmetric, the two bond pulls point in exactly opposite directions and cancel. Water, with the same kind of polar bonds but a bent shape, cannot cancel them, so it is strongly polar.
Polar bonds can cancel by symmetry, leaving a nonpolar molecule.
Partial charges (δ+, δ−) are not full charges like those on ions. They mark a slight, fractional imbalance in shared electrons within an intact covalent bond — no electron has actually been handed over. Only when the imbalance becomes extreme does the bond count as ionic.