Chemical bonds and polarity: covalent, ionic and hydrogen bonds

A chemical bond holds atoms together. In a covalent bond two atoms share electrons; in an ionic bond one atom transfers electrons to another and the oppositely charged ions attract. A hydrogen bond is weaker again: an attraction between a hydrogen on one molecule and an oxygen or nitrogen on another. Which kind forms comes down to electronegativity, and the answer decides whether a powder dissolves in water, how long it stays on an HPLC column and whether a chain folds into a helix.

Terms are defined in the glossary.

Why atoms bond

Atoms bond because the arrangement is more stable, at lower energy, than the separate atoms. Most atoms are most stable with a full outer shell of electrons, usually eight, or two for hydrogen. They get there in one of two ways: by sharing electrons with another atom, or by passing electrons from one atom to another.

Covalent bonds

A covalent bond forms when two atoms share a pair of electrons. Each atom counts the shared pair towards its own outer shell. Covalent bonds join non-metal atoms, and they are what hold molecules together: the bonds in water, in acetic acid and along the backbone of every peptide.

• Single bonds share one pair of electrons, such as C–C or C–H.
• Double bonds share two pairs, such as the C=O in a carbonyl group.
• Triple bonds share three pairs, such as the C≡N in acetonitrile.

More shared pairs make a bond shorter and stronger.

Bond energy is the energy needed to break the bond, and the difference between the bonds broken and the bonds made is what decides whether a reaction gives out heat or takes it in; see energy in reactions. These are averages (LibreTexts, data from Huheey et al.); the exact figure depends on the rest of the molecule.

The peptide bond is a covalent bond between the carbon of one amino acid's carboxyl group and the nitrogen of the next. It has partial double-bond character, which holds the atoms around it flat and limits how the chain can twist.

Ionic bonds

An ionic bond forms when one atom gives up electrons completely and another takes them. The atom that loses electrons becomes a positive ion, the one that gains them a negative ion, and the opposite charges attract.

Sodium chloride is the standard example. Sodium gives up one electron to become Na⁺, chlorine takes it to become Cl⁻, and in the solid the ions pack into a lattice held together by their charges. That lattice is strong, which is why salt melts at about 801 °C (PubChem).

In water the picture changes. Water molecules surround each ion and pull the lattice apart, so the salt dissolves into separate ions. The attraction between oppositely charged groups in solution, such as a positive lysine side chain and a negative aspartate, is much weaker than in a dry crystal. In peptides and proteins it is called a salt bridge. See isoelectric point and net charge for which groups carry charges, and trifluoroacetic acid counter-ions for the ionic pairing in peptide salts.

Ionic vs covalent bonds

There is a third kind. In a metallic bond, the atoms of a metal release their outer electrons into a shared pool that moves freely between the fixed positive ions. That pool is why metals conduct heat and electricity and can be bent into shape rather than shattering. Covalent, ionic and metallic are the three bonds between atoms; hydrogen bonds and the other forces below act between molecules, not inside them.

Electronegativity

Electronegativity is how strongly an atom pulls the electrons in a bond towards itself. The most widely used scale, devised by Linus Pauling, runs from about 0.8 to 4.0. The values below are Pauling values (PubChem).

Electronegativity rises across the periodic table from left to right and falls going down. Fluorine is the most electronegative element.

From covalent to ionic

The difference in electronegativity between two atoms decides what kind of bond they form. It is a sliding scale, not three separate boxes, but common rules of thumb are (LibreTexts):

Polarity

In a polar bond, the electrons sit closer to the more electronegative atom. That end gains a small negative charge, written δ−, and the other end a small positive charge, δ+.

A polar molecule has an uneven spread of charge overall. That depends on both its bonds and its shape:

• Water is polar. Its two O–H bonds are polar, and because the molecule is bent, their effects add up, leaving the oxygen side negative and the hydrogen side positive.
• Carbon dioxide is non-polar, even though each C=O bond is polar. The molecule is straight, so the two bonds pull in opposite directions and cancel out.
• Methane and other hydrocarbons are non-polar, because C–H bonds are barely polar at all.

Polar vs non-polar: like dissolves like

Polar substances tend to dissolve in polar solvents, and non-polar substances in non-polar ones. Water surrounds polar and charged groups and holds them in solution. It can't do the same for non-polar groups, which cluster together instead. That clustering, called the hydrophobic effect, is why oil and water separate.

Peptides contain both. Charged and polar side chains make a sequence more water-soluble; long runs of non-polar side chains make it harder to dissolve. See peptide solubility for how this guides the choice of solvent, and hydrophobic and polar amino acids for which side chains fall where.

Hydrogen bonds

A hydrogen bond is an attraction between a hydrogen atom bonded to a strongly electronegative atom, usually oxygen or nitrogen, and another electronegative atom with a lone pair of electrons. The hydrogen carries a partial positive charge and is drawn towards the partial negative charge on the other atom. IUPAC's 2011 definition describes it as an attraction between a hydrogen on one group and an atom on another where there is evidence of bond formation (Arunan et al., Pure Appl. Chem. 2011).

• The group that supplies the hydrogen, such as O–H or N–H, is the donor.
• The atom with the lone pair, such as the O in C=O, is the acceptor.

A hydrogen bond isn't a full bond. It is much weaker than a covalent bond, and is constantly breaking and re-forming in water. But there are usually a great many of them, and together they have large effects (Alberts et al., Molecular Biology of the Cell).

What hydrogen bonds explain

• Water's high boiling point. Water boils at 100 °C, while hydrogen sulfide, a heavier molecule of similar shape, boils at about −60 °C (PubChem). Sulfur is barely more electronegative than hydrogen, so H₂S molecules hardly hydrogen-bond at all.
• Solubility. Polar groups that can hydrogen-bond with water, such as –OH, –NH₂ and C=O, help a molecule dissolve.
• Peptide and protein shape. The N–H and C=O groups along a peptide backbone hydrogen-bond to each other. Regular patterns of these bonds make the alpha helix and beta sheet. Proline lacks a backbone N–H, which is why it tends to break these patterns; see the 20 amino acids.

Intermolecular forces: hydrogen bonds, dipole–dipole and London dispersion

Bonds hold atoms together inside a molecule. Intermolecular forces are the much weaker attractions between separate molecules. They decide whether a substance is a gas, liquid or solid at room temperature, and what it dissolves in.

• London dispersion forces act between all molecules, even non-polar ones. Electrons shift for an instant, making a brief, weak dipole that attracts a neighbour (OpenStax). They grow stronger as molecules get larger.
• Dipole–dipole attractions act between polar molecules. The slightly positive end of one is drawn to the slightly negative end of the next (OpenStax).
• Hydrogen bonds are an especially strong kind of dipole–dipole attraction, involving hydrogen bonded to N, O or F. They are only about 5 to 10% as strong as a covalent bond, but stronger than the other two (OpenStax).

From strongest to weakest:

All of these are much weaker than the covalent bonds inside a molecule. That is why heating ice breaks hydrogen bonds and melts it, without breaking water molecules apart. For how these forces fold a chain into helices and sheets, see peptide and protein structure.

Bonds and chromatography

The same forces explain how HPLC separates peptides. In reversed-phase HPLC, the column is coated with non-polar C18 chains. Non-polar parts of a peptide stick to them through the hydrophobic effect and van der Waals forces. As the mobile phase becomes less polar during a run, these attractions weaken and the peptide is released. More non-polar peptides hold on longer. Ion-exchange chromatography instead uses ionic attraction to separate by charge; see how peptides are purified. The same polarity rules decide how far a spot travels in thin-layer chromatography.

Related guides

For the bond that links amino acids, see what is a peptide. For how water breaks that bond, see hydrolysis. For the groups that carry charge and form hydrogen bonds, see functional groups in peptides. For how a bond's stretching shows up in an infrared spectrum, see the IR spectrum table.