Metallic vs Ionic Bonding: Same Attraction, Different Structures

Pick up a copper wire and a crystal of table salt. Both are solids at room temperature, both are held together by electrical attraction between positive and negative charges, and both form regular, repeating arrangements of particles. Yet you can bend the wire into a loop, while the salt crystal cracks if you press it with a spoon. The wire carries a current straight out of the reel; the salt refuses to conduct at all until you melt it or dissolve it.

Those differences are not accidents. They come directly from where the negative charge sits in each structure. In a metal, the negative charge is spread out as a mobile sea of electrons. In an ionic compound, the negative charge is locked onto specific ions. Once that single idea clicks, nearly every property on the comparison table below follows from it.

The two bonds in one sentence each

Metallic bonding is the attraction between a lattice of positive metal ions and a sea of delocalised electrons that are free to move through the whole structure.

Ionic bonding is the attraction between oppositely charged ions, formed when a metal atom transfers one or more electrons to a non-metal atom, arranged in a giant lattice where every ion is surrounded by ions of the opposite charge.

Notice what the definitions share: a giant structure, positive ions, and electrostatic attraction acting in every direction. Notice, too, what they do not share: in a metal there are no negative ions at all.

Side-by-side comparison

Feature	Metallic bonding	Ionic bonding
Elements involved	Metal atoms only (a pure metal or an alloy)	A metal and a non-metal
What happens to electrons	Outer electrons leave each atom and become shared by the whole structure	Electrons move from the metal atom to the non-metal atom
Particles in the solid	Positive metal ions + delocalised electrons	Positive ions + negative ions
Where the negative charge is	Spread out, mobile, belonging to no single atom	Fixed on each negative ion
Structure type	Giant metallic lattice	Giant ionic lattice
Direction of the bonding	Non-directional	Non-directional
Conducts as a solid?	Yes, very well	No
Conducts when molten?	Yes	Yes
Conducts in solution?	Metals generally don’t dissolve in water unchanged	Yes, if the compound is soluble
Response to a blow	Malleable and ductile: layers slide without breaking	Brittle: the crystal shatters along planes
Melting points	Very wide range (low for group 1, high for transition metals)	Generally high
Appearance	Shiny (lustrous) surface	Often colourless or white crystals; many transition-metal salts are coloured
Typical examples	Sodium, magnesium, copper, iron, brass	NaCl, MgO, CaF₂, Na₂O, Al₂O₃

Why metals conduct in the solid state and salts do not

Electrical conduction needs charged particles that can move. In a solid metal, the delocalised electrons are already free to drift through the lattice. Connect a cell and they flow towards the positive terminal straight away, which is why copper is the standard material for household wiring.

In a solid ionic compound the charged particles are ions, and ions in a lattice can only vibrate on the spot. The electrons are held tightly on the negative ions, so there is nothing mobile to carry a current. Melt the solid, or dissolve it in water, and the ions break free of their fixed positions. Now they can move, and the liquid or solution conducts. Note the difference in what moves: electrons in a metal, ions in a molten or dissolved salt. The article on bonding and conductivity takes this further, including graphite and the special case of electrolysis.

Why metals bend and salts shatter

This is the property difference that examiners love, and it is worth being able to explain with a quick sketch.

In a metal, all the ions are positive and identical (or nearly so, in an alloy). If a force pushes one layer of ions along past the next, the new arrangement looks exactly like the old one. The sea of electrons simply flows along with the ions and keeps holding them together. Nothing repels anything, so the metal changes shape without breaking. That is malleability (hammering into sheets) and ductility (drawing into wires).

In an ionic lattice, each layer is a strict alternation of positive and negative ions. Shift one layer by just one ion’s width and suddenly positive ions sit next to positive ions, and negative next to negative. Those like charges repel strongly, the layers push apart, and the crystal cleaves. That is brittleness. The bonding in the salt is strong, but it cannot cope with being rearranged.

A useful way to summarise it for revision: metallic bonding tolerates disorder, ionic bonding punishes it.

Melting points: overlapping, but for different reasons

Students are often told that both types of giant structure have “high melting points”. That is fair for ionic compounds and for most metals, but the metallic range is much wider.

Substance	Bond type	Melting point
Sodium	Metallic	97.8 °C
Magnesium	Metallic	about 650 °C
Copper	Metallic	1,084.6 °C
Iron	Metallic	about 1,538 °C
Sodium chloride, NaCl	Ionic	801 °C
Magnesium oxide, MgO	Ionic	about 2,852 °C

Sodium melts below the boiling point of water because each atom contributes only one electron to the sea and the Na⁺ ion is relatively large, so the attraction per ion is modest. Magnesium contributes two electrons per atom and forms a smaller, doubly charged ion, so its metallic bonding is noticeably stronger.

The same two factors control ionic compounds: charge and size. Magnesium oxide contains Mg²⁺ and O²⁻, both doubly charged and both small, so the attraction between them is far stronger than between Na⁺ and Cl⁻. That is why MgO melts at a temperature more than three times higher than table salt on the Celsius scale, and why it is used to line furnaces. You can read more about this in the post on lattice energy.

So the rule of thumb is the same for both bonds: smaller ions with bigger charges bond more strongly. What differs is the spread. Ionic solids cluster at the high end; metals run from soft and low-melting (the alkali metals) to extremely high-melting (tungsten and the other heavy transition metals).

How the electrons end up where they do

Why does sodium bonded to sodium give a metal, but sodium bonded to chlorine give a salt? The answer lies in how strongly each atom attracts electrons.

Sodium atoms hold their single outer electron weakly. When only sodium atoms are present, none of them is able to grab electrons from the others, so the outer electrons are pooled and shared across the whole lattice. Pauling electronegativity for sodium is only 0.93.

Chlorine, with an electronegativity of 3.16, pulls on electrons strongly. When sodium meets chlorine, the difference of 2.23 is large enough that the electron does not get shared at all; it ends up on the chlorine. The chloride ion keeps it firmly, and it is no longer free to roam. Metallic bonding needs a shortage of atoms that want electrons; ionic bonding needs one atom that gives and one that takes.

When to use which model

Both descriptions are models, and each is best suited to a particular set of substances.

Use the metallic model when:

• the substance is a pure metal or an alloy, such as steel, brass or bronze;
• you need to explain electrical and thermal conductivity in a solid;
• you are explaining malleability, ductility or lustre;
• you are comparing melting points across a period of metals (Na, Mg, Al) in terms of electrons per atom and ionic charge.

Use the ionic model when:

• a metal has combined with a non-metal, especially from groups 1 and 2 with groups 16 and 17;
• you are working out formulas from ion charges, for example Mg²⁺ with N³⁻ giving Mg₃N₂, or Al³⁺ with O²⁻ giving Al₂O₃;
• you need to explain why a compound conducts only when molten or dissolved;
• you are explaining brittleness or high melting points in terms of lattice attraction.

Be cautious when: the metal is a transition metal or the electronegativity difference is modest. Many compounds, such as aluminium chloride, show significant covalent character, and the pure ionic picture starts to strain. The bonding continuum explains how real substances sit between the ideal types.

Alloys: a metallic twist

An alloy is a mixture of a metal with other elements, usually other metals. The bonding stays metallic because the electron sea is still there, but the different-sized atoms disrupt the neat rows of ions. Layers can no longer slide as easily, so alloys are typically harder and stronger than the pure metals they are made from. There is no equivalent in ionic compounds: swapping in a different ion simply gives a different compound, or a mixed crystal, and it stays brittle. See alloys and bonding for worked examples.

Common mix-ups to avoid

• “Metals contain negative ions.” They do not. The negative charge is carried by delocalised electrons, not by ions.
• “Ionic solids conduct because they have ions.” Ions must be free to move; in a solid lattice they are not.
• “Metallic bonds are weak because sodium is soft.” Sodium is an extreme case. Iron and copper have strong metallic bonding.
• “An ionic bond is between one pair of ions.” Each ion is attracted to all the oppositely charged ions around it; the bonding belongs to the whole lattice.
• “Salt is brittle because ionic bonds are weak.” The bonds are strong. Brittleness comes from like charges being forced together when layers shift.

Key takeaways

• Both are giant structures held together by non-directional electrostatic attraction involving positive ions.
• In metallic bonding, the negative charge is a mobile sea of delocalised electrons; in ionic bonding, it sits on fixed negative ions.
• Metals conduct as solids; ionic compounds conduct only when molten or in solution, because ions must be free to move.
• Metals are malleable because shifting layers changes nothing; ionic crystals are brittle because shifting layers brings like charges together.
• Melting point rises with ionic charge and falls with ion size for both types. Ionic compounds are generally high-melting; metals span a much wider range.
• Metal with metal gives metallic bonding; metal with non-metal (large electronegativity difference) gives ionic bonding.