There are three somewhat different types of interactions
that can hold atoms together, all called bonds: there are ionic
bonds, and hydrogen
IONIC BONDS exist between, not too
surprisingly, ions of opposite charges: the attraction between
positive and negative ions, usually from atoms in Columns 1 and 2 in the
Periodic Table (which have positive charges from discarding electrons in the outer shell)
and atoms from Column 7 and sometimes 6 (negative ions from stealing
to fill their shells). This bond can hold ions together extremely
tightly, like two powerful magnets, and is commonly the bond found in salts,
including sodium (+1) chloride (-1), NaCl, table salt. Ionic bonds are often disrupted by water,
which itself is somewhat charged (see the next section on water
properties) and which tends to push apart and surround the ions;
this makes ionic compounds somewhat rare in
watery biological systems, while
dissolved ions are
common (when you "taste salt," you are not tasting
the compound sodium chloride, but detecting the ions
separated in your saliva). The ion form, with a full outer electron shell, is the
stable, natural form here, which makes the sharing of electrons we'll
see later in covalent bonds below not work. When
compounds separate into ions (or when radicals become ions), the act is called ionization
(water can spontaneously separate into a positive Hydrogen ion, H+,
and a negative Hydroxide ion, OH-). Larger molecules can
ionize when parts are lost that either take electrons with them or leave
them behind, and small parts of very large molecules might be ionized.
Ionic Bonding - Works even after ions have already formed.
Image of dry ionic compound atomic structure.
Video on ionization of water.
...And it all might work differently under non-Earth-surface
COVALENT BONDS are produced when atoms can fill
their outer electron shells more easily by borrowing electrons from
another atom, while sharing some of theirs. Atoms from the middle
Columns of the Periodic Table, from 3 to 6, often form these kinds of
bonds, even when coming together as ions. The number of bonds an
atom can form is the number of electrons it needs to fill its outer shell.
The vast majority of atoms in biological molecules come from just a set of
four: Hydrogen from Column 1 (but whose tiny
outer shell only needs one
more to be full), Oxygen from Column 6 (needing 2 electrons to fill its
bigger shell to its capacity of 8), Nitrogen from Column 5 (needs 3
electrons), and Carbon from Column 4 (needs 4 electrons). You might
remember them as HONC, which puts their bonding properties
in order, 1-2-3-4.
HONC molecules, showing bonds as lines.
more on bond types.
Although bonds often hold together atoms on a one-bond:
one-atom pairing, it is possible for more than one electron to be shared
between two atoms. These are multiple
bonds. There are double
bonds, as is found in the common atmospheric form of oxygen, O2,
or triple bonds, found in the atmospheric
form of nitrogen, N2.
Atmospheric nitrogen is held so tightly together by those bonds that it is
very stable and difficult to break; getting nitrogen for use in
biological molecules, where it is sometimes a critical ingredient, can only be done
by a handful of organisms on which the rest all depend. Quadruple
bonds probably do not appear much in nature, since the ones made in
labs tend to explode. Some bonds resonate,
shifting between single and double bonds in complex molecules, especially
rings of mostly carbon.
Imaging multiple bonds.
More on multiple bonds than you really want to know.
Why there aren't quadruple bonds.
Life is often talked about as Carbon-based
because, with its ability to share electrons and bond to 4 other atoms
simultaneously, carbon makes possible the
complicated atomic arrangements
needed in large molecules. Science fiction has toyed with the idea
of silicon-based life: silicon is also in Column 4 and is fairly
common on rocky planets like ours. There are reasons to accept that
this is theoretically possible but reasons to believe it is also extremely
unlikely. The term organic originally was applied to any
materials with carbon in them, on the assumption that any such compounds
were somehow related to living things. The term has been updated for
molecules, and now requires carbon and hydrogen, a combination that
is rare outside of living systems. The term should
require oxygen in the compounds as well, but it was never updated
Many carbon-based molecules.
Nitrogen, with its important ability to hold
three atoms together, is also important in such molecules as proteins and
DNA (breakdown products from these are called nitrogenous
wastes). Oxygen is commonly a "bridge" inside
molecules and sometimes, held on with a double bond, a potentially
reactive outer atom. Hydrogen can be thought of as a
"capper," attaching to those bond areas that face outward in
molecules and must be attached to something.
Bonds exist in three-dimensional space: depending
upon the connections, the angles at which atoms are arranged vary.
This is partly why strings of amino acids in proteins curl and bend and
kink and can produce molecules in a vast array of shapes, including small
rings and large specimens with very complex outer "surfaces."
The Nitrogen Cycle.
Bonding angles. (Video)
Covalent bonds may result from atoms sharing electrons,
but the sharing may not be completely equal: atoms with more protons may be able to hold
their own and other atom's borrowed electrons more strongly than their partner
atoms can, so that electrons spend more time around the larger atom in a
covalent bond. This can create slight charge differences across whole molecules or parts of
molecules - areas that electrons spend more time in are slightly more
negative, while areas where they spend less time are slightly positive.
Molecules with such partial charges are called
molecules. Water, with a big oxygen toward one side and little
hydrogens toward the other, tends to be somewhat negative on the oxygen
end and somewhat positive on the hydrogen end. This makes it
but not every polar arrangement is so neatly 2-ended. In water,
however, this is perhaps the most important feature of the molecule,
responsible for much of makes water a unique substance.
More on polarity.
Water, showing bipolar nature.
Like full ionic charges, these partial charges can
attract one another, which is the third type of bond: HYDROGEN
BONDS, called this because little, weak hydrogen is a common
participant. These can hold separate molecules loosely together, as
happens in water, or can hold different parts of one molecule together,
creating the spiral pattern of DNA or the complex forms of proteins.
Hydrogen bonds, based on a wide variety of partial charges, also can be
widely different in strength, approaching ionic bond strength in rare
cases. In molecules whose activity depends upon their shapes,
factors that disrupt hydrogen bonds, such as heat or loose ions, can
seriously denature them (shut down their activity).
Animation showing water and hydrogen bonds.
How hydrogen bonds hold proteins in particular shapes.
There is another class of bond,
bonds, solids where the electrons are shared among all of the atoms,
but they have no real role in biological systems. They do,
apparently, produce transitory examples of things like quadruple, even