How Are Metals, Non-Metals, and Metalloids Arranged on the Periodic Table?

Most periodic tables feature a bold staircase line descending from boron toward astatine. This boundary separates the metals on the left — the majority of all elements — from the non-metals in the upper right. Straddling that boundary is a small set of roughly six elements called metalloids (or semimetals): boron, silicon, germanium, arsenic, antimony, and tellurium. Depending on the textbook you encounter, one or two of these may be listed differently, but there is broad consensus on this core group.

3 Key Properties of Metallic Substances

Whether you are dealing with a pure elemental metal such as iron, or an alloy such as steel or brass, the material shares a characteristic set of physical properties. The umbrella term metallic substance covers both pure metals and alloys.

Electrical and Thermal Conductivity in Metals

Metals are excellent conductors of both heat and electricity. The key to their electrical conductivity lies in how their valence electrons behave: rather than being tightly bound to individual nuclei, these electrons are delocalized — shared across the entire bulk of the material. This sea of electrons model captures the idea that loosely held, mobile electrons can flow freely through the metal lattice, which is precisely what constitutes an electric current.

Non-metals, by contrast, hold their electrons tightly, making them poor conductors — insulators. Rubber gloves and rubber-soled shoes protect electrical workers because rubber is a hydrocarbon; carbon and hydrogen are both non-metals and resist electron flow. Metalloids occupy the middle ground: they are semiconductors, which explains why silicon — one of the most common metalloids — is the foundation of the entire semiconductor industry.

Luster, Malleability, and Ductility

Beyond conductivity, three additional properties define metallic substances:

• Luster — metals have a characteristic shine or reflective quality.
• Malleability — metals can be hammered or rolled into sheets without fracturing. A malleable metal deforms under stress rather than breaking.
• Ductility — metals can be drawn out into thin wires. Every metal wire you encounter exists because of this property.

What Are the 4 Types of Compounds and Their Properties?

Knowing where metals and non-metals sit on the periodic table lets us predict what kind of compound two elements will form when they combine. There are four major categories to understand: ionic compounds, molecular compounds, network covalent solids, and metallic substances (already covered above). Understanding these compound types is a core skill in general chemistry.

Ionic Compounds: High Melting Points and Brittleness Explained

When a metal combines with a non-metal, the result is an ionic compound — a crystal lattice of alternating cations (positive ions) and anions (negative ions). Ionic compounds have several characteristic properties that follow directly from the strength of their electrostatic bonds.

Why Do Ionic Compounds Have High Melting Points?

Ionic compounds are held together by strong electrostatic attractions between oppositely charged ions. Melting or boiling an ionic compound requires breaking a significant fraction of those ionic bonds — an energy-intensive process. Table salt (NaCl), for example, melts at approximately 800 °C and boils at around 1,400 °C. These figures are representative of ionic compounds as a class.

Why Are Ionic Compounds Brittle?

Despite their impressive thermal stability, ionic compounds are brittle. This seems paradoxical at first — if the bonds are so strong, why does a salt crystal shatter so easily? The answer lies in the geometry of the crystal lattice.

In an ionic lattice, every cation is surrounded exclusively by anions, and every anion is surrounded exclusively by cations. This alternating arrangement is what creates the strong cohesive forces. When a mechanical stress is applied — imagine striking a large salt block with a sledgehammer — entire planes of ions shift relative to one another. That shift brings cations into contact with neighboring cations and anions into contact with neighboring anions, replacing attractive forces with repulsive forces all along that plane. The crystal fractures instantly.

Lattice Energy: 2 Factors That Determine Ionic Bond Strength

The strength of the ionic bonds in a crystal is quantified by lattice energy — the energy required to completely separate one mole of an ionic solid into its constituent gaseous ions. Lattice energy is governed by two factors:

1. Ion charges: Lattice energy increases with the magnitude of the charges on the ions. A 2+ cation creates a stronger attraction than a 1+ cation; a 2− anion creates a stronger attraction than a 1− anion.
2. Ion size (distance): Lattice energy is inversely proportional to the distance between the ion centers (approximated by the sum of their radii). Smaller ions sit closer together, producing stronger attractions and higher lattice energies.

Molecular Compounds (Non-Metal + Non-Metal): Why Are Melting Points Low?

When two non-metals combine, the most common outcome is a molecular compound — a substance composed of discrete, finite molecules held together by covalent bonds within each molecule. Water (H₂O), carbon dioxide (CO₂), and methane (CH₄) are familiar examples of molecular compounds.

In the solid state, molecular compounds form crystals, but what holds those crystals together is not covalent bonds — it is a set of relatively weak intermolecular forces (to be covered in the next chapter). Because those forces are weak compared to ionic or covalent bonds, molecular compounds generally have lower melting and boiling points than ionic compounds or network covalent solids.

What Are Network Covalent Solids? (Diamond and Quartz)

A second, far less common possibility when non-metals combine is the formation of a network covalent solid. Instead of forming discrete molecules, the atoms bond covalently throughout the entire macroscopic sample, creating a continuous three-dimensional lattice held together entirely by covalent bonds.

The most familiar examples are diamond (pure carbon) and quartz (SiO₂, silicon dioxide). A single diamond crystal is, in a sense, one giant molecule — there is no such thing as an individual “diamond molecule.” Similarly, there is no discrete SiO₂ molecule; silicon and oxygen are always present in a 1:2 ratio across the entire crystal.

Because melting a network covalent solid requires breaking covalent bonds throughout the lattice, these materials have extremely high melting and boiling points — typically surpassing even those of ionic compounds.

Alkali Metals (Group 1): Reactivity, Water Reactions, and Periodic Trends

The alkali metals — lithium, sodium, potassium, rubidium, cesium, and francium — are defined by having the lowest ionization energies on the periodic table. Their single valence electron is loosely held and extremely easy to remove, which has sweeping consequences for their chemistry.

This low ionization energy means alkali metals are eager to give up electrons, making them highly reactive with non-metals to form ionic compounds. More dramatically, it makes them react vigorously — sometimes violently — with water.

How Do Alkali Metals React with Water?

When an alkali metal contacts water, the result is rapid formation of the metal hydroxide plus hydrogen gas, with a large release of heat. The balanced equation for sodium is:

Alkali metal reactivity with water increases down the group as ionization energies decrease. Lithium reacts steadily; sodium reacts vigorously with visible flames; potassium reacts explosively with a characteristic violet flame; cesium and rubidium react so violently that they must be stored under mineral oil and handled with extreme care.

Alkaline Earth Metals (Group 2): How Do They Compare to Alkali Metals?

The alkaline earth metals — beryllium, magnesium, calcium, strontium, barium, and radium — share the “alkaline” descriptor with the alkali metals for good reason: they too produce basic (alkaline) solutions when they react with water.

However, with two valence electrons and correspondingly higher ionization energies than their Group 1 neighbors, the alkaline earth metals are less reactive than the alkali metals. Beryllium barely reacts with water at all, and magnesium reacts only slowly with liquid water (though readily with steam). Calcium, strontium, and barium react noticeably with water, and the reactivity increases as you descend the group — barium is quite reactive, though still far less dramatic than sodium or potassium.

Noble Gases (Group 18): Why Are They Chemically Inert?

Moving to the far right of the periodic table, the noble gases — helium, neon, argon, krypton, xenon, and radon — are the chemical recluses of the periodic table. They are described as chemically inert, meaning they participate in virtually no chemical reactions under ordinary conditions.

The lightest noble gases, helium and neon, are completely unreactive under any known laboratory conditions. The heavier noble gases — especially xenon and krypton — can be coaxed into a very small number of reactions with highly electronegative elements, most notably fluorine. Compounds such as xenon difluoride (XeF₂) and xenon tetrafluoride (XeF₄) exist, but they are laboratory curiosities rather than everyday chemistry.

The reason for this extraordinary unreactivity will become clear in the bonding chapter: the noble gases possess a completely filled valence shell (a full octet, or in helium’s case a full duet), the most stable electron configuration possible. This full valence shell is also why noble gases have the highest ionization energies on the periodic table.

Halogens (Group 17): High Electron Affinity and Reactivity Trends

The halogens — fluorine, chlorine, bromine, iodine, and astatine — occupy the column immediately to the left of the noble gases, and their chemistry is essentially the mirror image of the alkali metals’. While alkali metals excel at losing electrons (low ionization energies), halogens excel at gaining electrons. They have the most negative electron affinities of any group on the periodic table — chlorine holds the record for the most negative electron affinity of all elements.

This powerful tendency to gain electrons makes halogens highly reactive with metals, producing ionic compounds (metal halides) with large releases of energy. The reaction can be surprisingly dramatic for something as familiar as the product it produces.

Chalcogens (Group 16): Sulfide Insolubility and Oxygen Allotropes

Group 16 is formally called the chalcogens, though this group name appears less frequently in introductory courses than the others. Two topics within this group are worth knowing for general chemistry: the insolubility of sulfides and the allotropes of oxygen.

Why Are Most Sulfide Compounds Insoluble?

Sulfide compounds are among the most insoluble ionic compounds known. When nearly any transition metal combines with sulfur to form a sulfide salt, the result is essentially insoluble in water. Manganese(II) sulfide (MnS), copper(I) sulfide (Cu₂S), and copper(II) sulfide (CuS) are all classic examples of highly insoluble compounds.

The main exceptions to this rule are alkali metal sulfides (e.g., Na₂S) and ammonium sulfide ((NH₄)₂S), which are soluble. For essentially every other metal, assume the sulfide salt is insoluble.

What Are the Allotropes of Oxygen?

Oxygen’s most stable elemental form is diatomic oxygen, O₂ — the standard state of oxygen and the form that makes up about 21% of Earth’s atmosphere. However, oxygen also exists in the atmosphere as ozone, O₃, particularly in the stratosphere.

When the same element exists in more than one distinct structural form, those forms are called allotropes. O₂ and O₃ are allotropes of oxygen, with O₂ being the more stable of the two. Carbon provides another well-known example: diamond and graphite are both pure carbon, yet they differ dramatically in structure and properties — they are allotropes of carbon.

Metal Oxides vs. Non-Metal Oxides: Which Produce Acidic or Basic Solutions?

One of the most useful predictive rules in descriptive chemistry concerns what happens when an oxide dissolves in water. The type of element bonded to oxygen determines whether the resulting solution is acidic or basic:

• Metal oxides are basic. Dissolve sodium oxide (Na₂O) in water and you get sodium hydroxide (NaOH), a strong base. The solution pH rises well above 7. This mirrors the behavior of sodium metal itself reacting with water to produce NaOH and hydrogen gas.
• Non-metal oxides are acidic. Dissolve carbon dioxide (CO₂) in water and it forms carbonic acid (H₂CO₃), lowering the pH below 7. The same principle applies to nitrogen dioxide (NO₂) forming nitric acid and sulfur trioxide (SO₃) forming sulfuric acid — the very reactions responsible for acid rain.