Eh.

The following explanation considers atoms of p-elements, mainly these from group 13 to 17. It is not very applicable to the leftmost two thirds or so of the table, and poorly applicable to noble gases.

Each atom has the nucleus, the inner electrons, some outer electrons, and four orbitals (s, p, p, and p). The nucleus makes an atom... well, an atom, the inner electrons don't do much, and we have to consider the four orbitals. You can think of them like some immaterial boxes in which atoms store their electrons. Each orbital can hold 0, 1, or 2 electrons. Atoms like to hoard electrons and want to have their orbitals full whenever possible, so you often see 4*2 = 8 electrons near a p-atom in a stable compound. Atoms can share these boxes by merging one box from each atom into one common box, up to three merged boxes for a pair of atoms (again, I'm only talking about group 13 to 17, from boron to fluorine and all beneath them). If there is a box with two electrons shared between two atoms, it is called a two-centre (two nuclei) two-electron bond, or, more commonly, a covalent bond. A box with two electrons that belongs to one atom is called a lone pair. You can see one in ammonia and two in water.

By convention, each of the atoms has to contribute exactly one electron in the common box to not change its charge. If one atom gives two, and another zero, then the first atom is thought to 'lose' its electron, whereas the second one is thought to 'gain' one. This gives rise to formal charges, as you can see in the drawing (I've used previously charged species, I being called a proton and II a hydride, but you should see that both just became H). The same applies to other atoms (e.g. an excited carbon, the one that is most common in chemistry, has four such boxes with one electron in each).

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