#"Electronegativity"# is defined as the ability of an atom in a covalent bond to polarize electron density towards itself. Non-metals towards the right of the Periodic Table as we face the table, e.g. #F#, #O#, #N#, tend to be HIGHLY electronegative. Why? Because nuclear charge has increased, and the unfilled valence shell of electrons does not effectively shield the nuclear charge. As evidence of this we can cite the well-known decrease of atomic radii across the Period from right to left.
On the other hand, in the bulk metal (of whatever kind), the valence electrons tend to be delocalized across the entire metallic lattice; i.e. each metal contributes one or two (or more) electrons to the bulk metallic lattice. This delocalization of electrons, typically described as #"positive ions in a sea of electrons"#, is conceived to give rise to typical metallic properties: malleability; ductility; conductivity to electricity and heat. Because the positively charged metal nuclei can move with respect to each other, WITHOUT disrupting the metallic bond which is retained between metal nuclei and the free electrons, metals tend to be (i) malleable, capable of being beaten out into a sheet, and (ii) ductile, capable of being drawn out into a wire.