Structure and Polarity of the Carbonyl Group

The carbonyl carbon in both aldehydes and ketones is $sp^2$ hybridised: it forms three $\sigma$ bonds (to oxygen and to the other two attached groups) that lie in a single plane at roughly $120^\circ$ to one another, and it carries one unhybridised $p$ orbital perpendicular to that plane. This $p$ orbital overlaps side-on with a $p$ orbital on the oxygen to form the $\pi$ bond of the $\text{C=O}$ double bond, so overall the carbonyl group consists of one $\sigma$ bond and one $\pi$ bond between carbon and oxygen, exactly analogous to the $\text{C=C}$ double bond of an alkene in its geometry, but very different in its chemistry because one of the two atoms (oxygen) is far more electronegative than the other.

Why the bond is polarised. Oxygen's higher electronegativity means it attracts the shared bonding electrons -- both the $\sigma$ and especially the more loosely held $\pi$ electrons -- more strongly than carbon does. The $\pi$ electron cloud is therefore skewed towards oxygen, leaving oxygen with a partial negative charge ($\delta^-$) and the carbonyl carbon with a matching partial positive charge ($\delta^+$). This can be represented by two resonance structures: a major, non-charged structure with a normal $\text{C=O}$ double bond, and a minor but chemically decisive structure with a full positive charge on carbon and a full negative charge on oxygen, $\text{C}^+\text{-O}^-$. The true electronic structure is a hybrid of the two, but it is the contribution from the charged form that makes the carbonyl carbon behave, for practical purposes, as an electrophilic centre. …