Metals and non-metals: atomic structure differences

Atomic structure and the periodic table • The periodic table

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Why do metals tend to be dense?

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Close-packed lattice structures and strong metallic bonding place atoms tightly together, producing high density.

Key concepts

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Definition and general trend

Metals form a class of elements that typically conduct electricity, are malleable and form positive ions. Non-metals form a class that typically do not conduct electricity in solid form, are brittle when solid and form negative ions or share electrons. Periodic trends place most metals on the left and centre of the periodic table and most non-metals on the right. Metalloids lie along a dividing line and show mixed properties.

Physical properties of metals and atomic causes

Metals conduct electricity because delocalised electrons move freely through a lattice of positive ions; stronger delocalisation produces greater conductivity. Metals are malleable and ductile because metallic bonding is non-directional so metal ions slide past one another without breaking bonds. Typical metallic properties-high density and high melting points-arise when many electrons contribute to a strong bond between ions. Limited exceptions occur for some metals with weak metallic bonding, such as mercury, which is liquid at room temperature.

Chemical properties of metals and atomic causes

Metals react chemically by losing valence electrons to form positive ions (cations). Low ionisation energy and relatively low electronegativity cause metals to oxidise (react with oxygen) and to react with acids to produce hydrogen gas. Metals in group 1 and group 2 show increasing reactivity down a group because additional electron shells increase shielding and reduce the energy required to remove the outer electron(s).

Physical properties of non-metals and atomic causes

Non-metals show a range of physical properties because bonding varies: many non-metals form small covalent molecules (low melting/boiling points, poor electrical conductivity) while some form giant covalent lattices (very high melting points, poor electrical conductivity in solid form). Poor electrical conductivity in most non-metals results from the absence of delocalised electrons. Brittleness in solid non-metals arises from directional covalent bonds that break rather than allowing layers to slide.

Chemical properties of non-metals and atomic causes

Non-metals tend to gain electrons or share electrons in bonds. High ionisation energies and higher electronegativities produce a tendency to form negative ions (anions) or covalent bonds. Group trends show increased non-metallic character across a period because rising nuclear charge pulls electrons more strongly, making electron gain or sharing more favourable. Oxides of non-metals are often acidic because they form covalent bonds with oxygen that react with water to form acids.

Relation of atomic structure to periodic position

Position in the periodic table reflects electron configuration. Elements on the left have few valence electrons and lower ionisation energies, producing metallic character through electron loss and metallic bonding. Elements on the right have more valence electrons, higher ionisation energies and higher electronegativity, producing non-metallic behaviour through electron gain or sharing. Down a group, additional electron shells increase atomic radius and shielding, increasing metallic character; across a period, increased nuclear charge with similar shielding decreases atomic radius and increases non-metallic character.

Key notes

Important points to keep in mind

Metals conduct electricity because delocalised electrons carry charge through a lattice.

Metals are malleable because metallic bonding is non-directional and allows layers of atoms to slide.

Non-metals are often poor conductors because they lack freely moving electrons or ions in the solid state.

Metals form positive ions by losing electrons; non-metals form negative ions by gaining electrons or share electrons in covalent bonds.

Metallic character decreases across a period and increases down a group due to changes in nuclear charge and shielding.

Group 1 and 2 elements show predictable metallic reactivity because of low ionisation energies for their valence electrons.

Non-metal oxides are frequently acidic; metal oxides are frequently basic, reflecting electron transfer behaviour.

Metalloids sit on the dividing line and show mixed properties because of intermediate electron configurations.

Atomic radius, ionisation energy and electronegativity together explain most trends in physical and chemical behaviour.

Exceptions occur: hydrogen behaves as a non-metal, and some metals (for example, mercury) have anomalous physical states.