What is a chemical bond? Covalent, ionic and metallic bonds

A chemical bond is a lasting attraction between atoms that holds them together in a molecule or a solid. It forms because the atoms' electrons, shared between them or handed from one to another, end up in an arrangement with lower energy than the atoms had apart. Atoms share electrons in a covalent bond, transfer them in an ionic bond, and pool them in a metallic bond. Every substance more complicated than a lone atom is held together this way.

In short

• Bonds form because the bonded arrangement has less energy than the separate atoms, so pulling them apart takes energy.
• Covalent bonds share electrons, ionic bonds come from electrons moving from one atom to another, and metallic bonds share electrons across a whole lump of metal. Real bonds sit on a spectrum between these.
• Breaking a bond never releases energy. A reaction gives off heat when the bonds it forms are stronger than the bonds it breaks.

Why do atoms bond at all?

Atoms bond for the same reason a ball rolls downhill: the bonded state has lower energy. An atom's electrons are attracted to its own nucleus, but when two atoms come close, electrons can also be attracted to the other nucleus. At the right distance, those attractions outweigh the repulsion between the two nuclei and between the electrons, and the total energy dips. The atoms settle at the bottom of that dip, the bond length, and it takes energy to pull them out of it.

Which arrangement wins depends on the atoms' outer electrons, the valence electrons described in the electron note. Atoms tend to bond in ways that leave each with a filled outer shell, often eight electrons, which chemists call the octet rule. It is a useful pattern for the lighter elements, not a law, and many molecules break it.

What are the main kinds of chemical bond?

Kind	What the electrons do	Example	Typical strength	What it gives the substance
Covalent	A pair is shared between two atoms	Water (H₂O), methane (CH₄), diamond	About 150 to 950 kJ/mol per bond: H–H 436, N≡N 946	Distinct molecules; very hard solids when the bonds form a network, as in diamond
Ionic	One atom gives electrons to another, and the charged ions attract	Table salt (NaCl)	Pulling 1 mole of solid NaCl apart into separate ions takes about 770 to 790 kJ, depending on how it is worked out	Hard, brittle crystals with high melting points (NaCl melts at 801 °C); conduct electricity when melted or dissolved
Metallic	Outer electrons spread through the whole metal, shared by all its atoms	Copper, iron	About 100 kJ/mol for sodium to over 400 for iron	Conducts electricity and heat, bends without breaking, shiny
Hydrogen bond	A hydrogen on one molecule is drawn to an oxygen, nitrogen or fluorine on another	Between water molecules; between DNA's two strands	About 10 to 40 kJ/mol, around 20 in water	Keeps water liquid at room temperature; holds DNA's strands together but lets them unzip

The unit, kJ/mol, is kilojoules of energy per mole, about 6.02 × 10²³ bonds. The covalent values are averages from the OpenStax Chemistry textbook. Hydrogen bonds are attractions between molecules rather than true bonds inside one, and they are ten to fifty times weaker than a covalent bond, but there are so many of them that they shape the behavior of water and the folding of proteins.

Is a bond ionic or covalent? How to tell

Most bonds are neither purely one nor purely the other. What decides it is electronegativity, how strongly an atom pulls on shared electrons, measured on a scale devised by Linus Pauling that runs from about 0.8 to 4.0. The bigger the difference between two atoms, the more lopsided the sharing.

• Little or no difference (under about 0.4): electrons shared evenly, a nonpolar covalent bond. Example: the two hydrogens in H₂, difference 0.
• Moderate difference (about 0.4 to 1.7): shared unevenly, a polar covalent bond. Example: oxygen (3.44) and hydrogen (2.20) in water, difference 1.24, so the oxygen end is slightly negative.
• Large difference (above about 1.7 to 2.0): the electron effectively moves over, an ionic bond. Example: chlorine (3.16) and sodium (0.93), difference 2.23.

The cutoffs are conventions, not sharp lines. Pauling's work on these ideas, published in his 1939 book The Nature of the Chemical Bond, earned him the 1954 Nobel Prize in Chemistry.

A worked example: making table salt from sodium and chlorine

The usual story is that sodium "wants" to lose an electron and chlorine "wants" to gain one. The numbers tell a more interesting story. Follow 1 mole of sodium metal and half a mole of chlorine gas becoming 1 mole of sodium chloride, step by step, with energy put in counted as positive:

1. Turn solid sodium into separate gas atoms: +107 kJ.
2. Split half a mole of Cl₂ molecules into chlorine atoms (half of the 243 kJ/mol Cl–Cl bond energy): +122 kJ.
3. Pull one electron off each sodium atom, its ionization energy: +496 kJ.
4. Give that electron to each chlorine atom, its electron affinity: −349 kJ.
5. Let the Na⁺ and Cl⁻ ions pull together into a crystal, the lattice energy: −787 kJ.

Total: 107 + 122 + 496 − 349 − 787 = −411 kJ, the measured heat released when sodium chloride forms. A lattice energy can't be measured directly, so step 5 is the value that makes this cycle add up; estimates from the ions' charges and spacing land close to it, and OpenStax quotes 769 kJ.

The electron transfer on its own is uphill: step 3 costs 496 kJ and step 4 gives back only 349. Sodium does not "want" to lose its electron at all. What pays for the whole process is step 5, the strong attraction between oppositely charged ions packed in a lattice. An ionic bond is that attraction, not the transfer.

A second example: why burning hydrogen releases heat

Hydrogen burns in oxygen to make water: 2H₂ + O₂ → 2H₂O. Using average bond energies:

1. Break 2 H–H bonds: 2 × 436 = 872 kJ put in.
2. Break 1 O=O bond: 498 kJ put in. Total in: 1,370 kJ.
3. Form 4 O–H bonds (two in each water molecule): 4 × 464 = 1,856 kJ released.
4. Net: 1,370 − 1,856 = −486 kJ for 2 moles of water, or about −243 kJ per mole.

The measured value for making water vapor is −242 kJ/mol, so the averages land close. Every bond broken cost energy; the heat came from forming stronger bonds than the ones broken. That is how energy is released in every fire, battery and meal.

Do chemical bonds store energy?

This is the most common misconception about bonds, and it is worth getting right. People say food or fuel "stores energy in its bonds," which is released "when the bonds break." But breaking a bond always takes energy, as OpenStax puts it: separating any pair of bonded atoms requires energy. The energy in fuel is better described as belonging to the whole arrangement: fuel and oxygen together sit higher in energy than the carbon dioxide and water they become, because the products' bonds are stronger.

The same applies to ATP, often said to have a "high-energy bond." Splitting off its end phosphate releases energy only because the products, together with the water they bond to, end up more stable overall. The bond itself still costs energy to break.

Where do bonds lead next?

Bonds decide a substance's shape and its behavior. They explain why water, held together by hydrogen bonds between its molecules, is a liquid while heavier carbon dioxide is a gas; why diamond and graphite, both pure carbon, are so different; and why metals conduct. When bonds break and new ones form, the result is a chemical reaction. And the outer-electron patterns behind bonding are what line up the columns of the periodic table.

Check yourself

• Without looking back, explain why two atoms that bond end up harder to separate than they were to bring together.
• In the table salt example, which step pays for the whole process, and why does that change the "sodium wants to lose an electron" story?
• Why is it wrong to say energy is released when a bond breaks?

Sources

• Strengths of ionic and covalent bonds — OpenStax, Chemistry 2e
• Covalent bonding — OpenStax, Chemistry 2e
• Linus Pauling, Nobel Prize in Chemistry 1954 — Nobel Prize Outreach
• Electronegativity — Wikipedia
• Sodium chloride — Wikipedia
• Chemical bond — Wikipedia