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Period 3 element
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Period 3 in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson

A period 3 element is one of the chemical elements in the third row (or period) of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behavior of the elements as their atomic number increases: a new row is begun when chemical behavior begins to repeat, meaning that elements with similar behavior fall into the same vertical columns. The third period contains eight elements: sodium, magnesium, aluminium, silicon, phosphorus, sulfur, chlorine and argon. The first two, sodium and magnesium, are members of the s-block of the periodic table, while the others are members of the p-block. All of the period 3 elements occur in nature and have at least one stable isotope.[1]

Atomic structure

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In a quantum mechanical description of atomic structure, this period corresponds to the buildup of electrons in the third (n = 3) shell, more specifically filling its 3s and 3p subshells. There is a 3d subshell, but—in compliance with the Aufbau principle—it is not filled until period 4. This makes all eight elements analogs of the period 2 elements in the same exact sequence. The octet rule generally applies to period 3 in the same way as to period 2 elements, because the 3d subshell is normally non-acting.

Elements

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Elements by number
Element # Symbol Block Electron configuration
Sodium 11 Na s-block [Ne] 3s1
Magnesium 12 Mg s-block [Ne] 3s2
Aluminium 13 Al p-block [Ne] 3s2 3p1
Silicon 14 Si p-block [Ne] 3s2 3p2
Phosphorus 15 P p-block [Ne] 3s2 3p3
Sulfur 16 S p-block [Ne] 3s2 3p4
Chlorine 17 Cl p-block [Ne] 3s2 3p5
Argon 18 Ar p-block [Ne] 3s2 3p6

Sodium

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Main article: Sodium

Sodium (symbol Na) is a soft, silvery-white, highly reactive metal and is a member of the alkali metals; its only stable isotope is 23Na. It is an abundant element that exists in numerous minerals such as feldspars, sodalite and rock salt. Many salts of sodium are highly soluble in water and are thus present in significant quantities in the Earth's bodies of water, most abundantly in the oceans as sodium chloride.

Many sodium compounds are useful, such as sodium hydroxide (lye) for soapmaking, and sodium chloride for use as a deicing agent and a nutrient. The same ion is also a component of many minerals, such as sodium nitrate.

The free metal, elemental sodium, does not occur in nature but must be prepared from sodium compounds. Elemental sodium was first isolated by Humphry Davy in 1807 by the electrolysis of sodium hydroxide.

Magnesium

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Main article: Magnesium

Magnesium (symbol Mg) is an alkaline earth metal and has common oxidation number +2. It is the eighth most abundant element in the Earth's crust[2] and the ninth in the known universe as a whole.[3][4] Magnesium is the fourth most common element in the Earth as a whole (behind iron, oxygen and silicon), making up 13% of the planet's mass and a large fraction of the planet's mantle. It is relatively abundant because it is easily built up in supernova stars by sequential additions of three helium nuclei to carbon (which in turn is made from three helium nuclei). Due to the magnesium ion's high solubility in water, it is the third most abundant element dissolved in seawater.[5]

The free element (metal) is not found naturally on Earth, as it is highly reactive (though once produced, it is coated in a thin layer of oxide [see passivation], which partly masks this reactivity). The free metal burns with a characteristic brilliant white light, making it a useful ingredient in flares. The metal is now mainly obtained by electrolysis of magnesium salts obtained from brine. Commercially, the chief use for the metal is as an alloying agent to make aluminium-magnesium alloys, sometimes called "magnalium" or "magnelium". Since magnesium is less dense than aluminium, these alloys are prized for their relative lightness and strength.

Magnesium ions are sour to the taste, and in low concentrations help to impart a natural tartness to fresh mineral waters.

Aluminium

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Main article: Aluminium

Aluminium (symbol Al) or aluminum (American English) is a silvery white member of the boron group of chemical elements and a p-block metal classified by some chemists as a post-transition metal.[6] It is not soluble in water under normal circumstances. Aluminium is the third most abundant element (after oxygen and silicon), and the most abundant metal, in the Earth's crust. It makes up about 8% by weight of the Earth's solid surface. Aluminium metal is too reactive chemically to occur natively. Instead, it is found combined in over 270 different minerals.[7] The chief ore of aluminium is bauxite.

Aluminium is remarkable for the metal's low density and for its ability to resist corrosion due to the phenomenon of passivation. Structural components made from aluminium and its alloys are vital to the aerospace industry and are important in other areas of transportation and structural materials. The most useful compounds of aluminium, at least on a weight basis, are the oxides and sulfates.

Silicon

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Main article: Silicon

Silicon (symbol Si) is a group 14 metalloid. It is less reactive than its chemical analog carbon, the nonmetal directly above it in the periodic table, but more reactive than germanium, the metalloid directly below it in the table. Controversy about silicon's character dates from its discovery: silicon was first prepared and characterized in pure form in 1824, and given the name silicium (from Latin: silicis, flints), with an -ium word-ending to suggest a metal. However, its final name, suggested in 1831, reflects the more chemically similar elements carbon and boron.

Silicon is the eighth most common element in the universe by mass, but very rarely occurs as the pure free element in nature. It is most widely distributed in dusts, sands, planetoids and planets as various forms of silicon dioxide (silica) or silicates. Over 90% of the Earth's crust is composed of silicate minerals, making silicon the second most abundant element in the Earth's crust (about 28% by mass) after oxygen.[8]

Most silicon is used commercially without being separated, and indeed often with little processing of compounds from nature. These include direct industrial building use of clays, silica sand and stone. Silica is used in ceramic brick. Silicate goes into Portland cement for mortar and stucco, and combined with silica sand and gravel, to make concrete. Silicates are also in whiteware ceramics such as porcelain, and in traditional quartz-based soda–lime glass. More modern silicon compounds such as silicon carbide form abrasives and high-strength ceramics. Silicon is the basis of the ubiquitous synthetic silicon-based polymers called silicones.

Elemental silicon also has a large impact on the modern world economy. Although most free silicon is used in the steel refining, aluminum-casting, and fine chemical industries (often to make fumed silica), the relatively small portion of very highly purified silicon that is used in semiconductor electronics (< 10%) is perhaps even more critical. Because of wide use of silicon in integrated circuits, the basis of most computers, a great deal of modern technology depends on it.

Phosphorus

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Main article: Phosphorus

Phosphorus (symbol P) is a multivalent nonmetal of the nitrogen group, phosphorus as a mineral is almost always present in its maximally oxidized (pentavalent) state, as inorganic phosphate rocks. Elemental phosphorus exists in two major forms—white phosphorus and red phosphorus—but due to its high reactivity, phosphorus is never found as a free element on Earth.

The first form of elemental phosphorus to be produced (white phosphorus, in 1669) emits a faint glow upon exposure to oxygen – hence its name given from Greek mythology, Φωσφόρος meaning "light-bearer" (Latin: Lucifer), referring to the "Morning Star", the planet Venus. Although the term "phosphorescence", meaning glow after illumination, derives from this property of phosphorus, the glow of phosphorus originates from oxidation of the white (but not red) phosphorus and should be called chemiluminescence. It is also the lightest element to easily produce stable exceptions to the octet rule.

The vast majority of phosphorus compounds are consumed as fertilizers. Other applications include the role of organophosphorus compounds in detergents, pesticides and nerve agents and matches.[9]

Sulfur

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Main article: Sulfur

Sulfur (symbol S) is an abundant multivalent nonmetal, one of chalcogens. Under normal conditions, sulfur atoms form cyclic octatomic molecules with chemical formula S8. Elemental sulfur is a bright yellow crystalline solid when at room temperature. Chemically, sulfur can react as either an oxidant or a reducing agent. It oxidizes most metals and several nonmetals, including carbon, which leads to its negative charge in most organosulfur compounds, but it reduces several strong oxidants, such as oxygen and fluorine.

In nature, sulfur can be found as the pure element and as sulfide and sulfate minerals. Elemental sulfur crystals are commonly sought after by mineral collectors for their brightly colored polyhedron shapes. Being abundant in native form, sulfur was known in ancient times, mentioned for its uses in ancient Greece, China and Egypt. Sulfur fumes were used as fumigants, and sulfur-containing medicinal mixtures were used as balms and antiparasitics. Sulfur is referenced in the Bible as brimstone in English, with this name still used in several nonscientific terms.[10] Sulfur was considered important enough to receive its own alchemical symbol. It was needed to make the best quality of black gunpowder, and the bright yellow powder was hypothesized by alchemists to contain some of the properties of gold, which they sought to synthesize from it. In 1777, Antoine Lavoisier helped convince the scientific community that sulfur was a basic element, rather than a compound.

Elemental sulfur was once extracted from salt domes, where it sometimes occurs in nearly pure form, but this method has been obsolete since the late 20th century. Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum. The element's commercial uses are primarily in fertilizers, because of the relatively high requirement of plants for it, and in the manufacture of sulfuric acid, a primary industrial chemical. Other well-known uses for the element are in matches, insecticides and fungicides. Many sulfur compounds are odiferous, and the smell of odorized natural gas, skunk scent, grapefruit, and garlic is due to sulfur compounds. Hydrogen sulfide produced by living organisms imparts the characteristic odor to rotting eggs and other biological processes.

Chlorine

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Main article: Chlorine

Chlorine (symbol Cl) is the second-lightest halogen. The element forms diatomic molecules under standard conditions, called dichlorine. It has the highest electron affinity and the one of highest electronegativity of all the elements; thus chlorine is a strong oxidizing agent.

The most common compound of chlorine, sodium chloride (table salt), has been known since ancient times; however, around 1630, chlorine gas was obtained by the Belgian chemist and physician Jan Baptist van Helmont. The synthesis and characterization of elemental chlorine occurred in 1774 by Swedish chemist Carl Wilhelm Scheele, who called it "dephlogisticated muriatic acid air", as he thought he synthesized the oxide obtained from the hydrochloric acid, because acids were thought at the time to necessarily contain oxygen. A number of chemists, including Claude Berthollet, suggested that Scheele's "dephlogisticated muriatic acid air" must be a combination of oxygen and the yet undiscovered element, and Scheele named the supposed new element within this oxide as muriaticum. The suggestion that this newly discovered gas was a simple element was made in 1809 by Joseph Louis Gay-Lussac and Louis-Jacques. This was confirmed in 1810 by Sir Humphry Davy, who named it chlorine, from the Greek word χλωρός (chlōros), meaning "green-yellow".

Chlorine is a component of many other compounds. It is the second most abundant halogen and 21st most abundant element in Earth's crust. The great oxidizing power of chlorine led it to its bleaching and disinfectant uses, as well as being an essential reagent in the chemical industry. As a common disinfectant, chlorine compounds are used in swimming pools to keep them clean and sanitary. In the upper atmosphere, chlorine-containing molecules such as chlorofluorocarbons have been implicated in ozone depletion.

Argon

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Main article: Argon

Argon (symbol Ar) is the third element in group 18, the noble gases. Argon is the third most common gas in the Earth's atmosphere, at 0.93%, making it more common than carbon dioxide. Nearly all of this argon is radiogenic argon-40 derived from the decay of potassium-40 in the Earth's crust. In the universe, argon-36 is by far the most common argon isotope, being the preferred argon isotope produced by stellar nucleosynthesis.

The name "argon" is derived from the Greek neuter adjective ἀργόν, meaning "lazy" or "the inactive one", as the element undergoes almost no chemical reactions. The complete octet (eight electrons) in the outer atomic shell makes argon stable and resistant to bonding with other elements. Its triple point temperature of 83.8058 K is a defining fixed point in the International Temperature Scale of 1990.

Argon is produced industrially by the fractional distillation of liquid air. Argon is mostly used as an inert shielding gas in welding and other high-temperature industrial processes where ordinarily non-reactive substances become reactive: for example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning. Argon gas also has uses in incandescent and fluorescent lighting, and other types of gas discharge tubes. Argon makes a distinctive blue–green gas laser.

Biological roles

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Sodium is an essential element for all animals and some plants. In animals, sodium ions are used against potassium ions to build up charges on cell membranes, allowing transmission of nerve impulses when the charge is dissipated; it is therefore classified as a dietary inorganic macromineral.

Magnesium is the eleventh most abundant element by mass in the human body; its ions are essential to all living cells, where they play a major role in manipulating important biological polyphosphate compounds like ATP, DNA, and RNA. Hundreds of enzymes thus require magnesium ions to function. Magnesium is also the metallic ion at the center of chlorophyll, and is thus a common additive to fertilizers.[11] Magnesium compounds are used medicinally as common laxatives, antacids (e.g., milk of magnesia), and in a number of situations where stabilization of abnormal nerve excitation and blood vessel spasm is required (e.g., to treat eclampsia).

Despite its prevalence in the environment, aluminium salts are not known to be used by any form of life. In keeping with its pervasiveness, it is well tolerated by plants and animals.[12] Because of their prevalence, potential beneficial (or otherwise) biological roles of aluminium compounds are of continuing interest.

Silicon is an essential element in biology, although only tiny traces of it appear to be required by animals,[13] though various sea sponges need silicon in order to have structure. It is much more important to the metabolism of plants, particularly many grasses, and silicic acid (a type of silica) forms the basis of the striking array of protective shells of the microscopic diatoms.

Phosphorus is essential for life. As phosphate, it is a component of DNA, RNA, ATP, and also the phospholipids that form all cell membranes. Demonstrating the link between phosphorus and life, elemental phosphorus was historically first isolated from human urine, and bone ash was an important early phosphate source. Phosphate minerals are fossils. Low phosphate levels are an important limit to growth in some aquatic systems. Today, the most important commercial use of phosphorus-based chemicals is the production of fertilizers, to replace the phosphorus that plants remove from the soil.

Sulfur is an essential element for all life, and is widely used in biochemical processes. In metabolic reactions, sulfur compounds serve as both fuels and respiratory (oxygen-replacing) materials for simple organisms. Sulfur in organic form is present in the vitamins biotin and thiamine, the latter being named for the Greek word for sulfur. Sulfur is an important part of many enzymes and in antioxidant molecules like glutathione and thioredoxin. Organically bonded sulfur is a component of all proteins, as the amino acids cysteine and methionine. Disulfide bonds are largely responsible for the mechanical strength and insolubility of the protein keratin, found in outer skin, hair, and feathers, and the element contributes to their pungent odor when burned.

Elemental chlorine is extremely dangerous and poisonous for all lifeforms, and is used as a pulmonary agent in chemical warfare; however, chlorine is necessary to most forms of life, including humans, in the form of chloride ions.

Argon has no biological role. Like any gas besides oxygen, argon is an asphyxiant.

Table of elements

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Period 3 in the periodic table
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Group →
↓ Period
3 So­dium11Na​22.990 Magne­sium12Mg​24.305 Alumin­ium13Al​26.982 Sili­con14Si​28.085 Phos­phorus15P​30.974 Sulfur16S​32.06 Chlor­ine17Cl​35.45 Argon18Ar​39.95

Notes

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References

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Revisions and contributorsEdit on WikipediaRead on Wikipedia

Period 3 element

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The Period 3 elements comprise the eight chemical elements in the third row of the periodic table, spanning atomic numbers 11 to 18: sodium (Na), magnesium (Mg), aluminium (Al), , phosphorus (P), sulfur (S), chlorine (Cl), and argon (Ar). These elements fill the 3s and 3p subshells with electrons, marking a transition from highly reactive metals on the left to non-metals and an inert on the right, with silicon serving as a intermediary. Across Period 3, physical properties exhibit clear trends driven by increasing effective nuclear charge and decreasing atomic radius, which falls from 0.157 nm for sodium to 0.099 nm for chlorine. First ionization energies generally rise from a low of 496 kJ/mol for sodium to 1521 kJ/mol for argon, though slight dips occur at aluminium (due to the start of 3p orbital filling) and sulfur (due to electron pairing in the p subshell). Melting and boiling points increase from sodium (371 K) through the metals to a peak at silicon (1683 K, owing to its giant covalent structure), then drop sharply for the molecular non-metals phosphorus, sulfur, chlorine, and argon, where weak van der Waals forces dominate. Electrical conductivity is high in the metals sodium (2.1 × 10^7 S m⁻¹), magnesium, and aluminium (3.8 × 10^7 S m⁻¹), moderate in semiconducting silicon, and negligible in the non-metals. Chemically, reactivity shifts from strong reducing agents on the left to oxidizing agents on the right, with electronegativity increasing from 0.9 for sodium to 3.0 for . The elements react vigorously with oxygen to form oxides: basic oxides like Na₂O and MgO on the metallic side, amphoteric Al₂O₃ in the middle, and acidic oxides such as SiO₂, P₄O₁₀, SO₂, and Cl₂O₇ toward the non-metals. Reactions with highlight this progression—sodium reacts explosively to produce NaOH and H₂, magnesium requires to form MgO or Mg(OH)₂, while is passivated by its oxide layer, and the non-metals through generally do not react with cold to liberate hydrogen (unlike metals), though reacts with to form acidic solutions of HCl and HClO, and white can slowly hydrolyze; remains inert. Chlorides follow a similar pattern, hydrolyzing to produce HCl for covalent non-metal chlorides like SiCl₄ and PCl₅, but remaining neutral for ionic ones like NaCl. These trends underscore the periodic law, illustrating how influences bonding, reactivity, and applications—from sodium in streetlights and magnesium in alloys, to in semiconductors, in fertilizers, in production, in disinfectants, and as an inert atmosphere in .

Overview

Definition and periodic table position

Period 3 elements refer to the eight chemical elements positioned in the third horizontal row (period) of the periodic table, where the 3s and 3p atomic orbitals are successively filled with electrons. These elements span atomic numbers 11 through 18 and include sodium, magnesium, aluminum, , , , , and . The following table lists the period 3 elements with their symbols and atomic numbers:
ElementSymbolAtomic Number
SodiumNa11
MagnesiumMg12
AluminumAl13
Si14
P15
S16
Cl17
Ar18
In the overall structure of the periodic table, period 3 lies between period 2 (elements through , atomic numbers 3–10) and period 4 (elements through , atomic numbers 19–36), with elements arranged in order of increasing from left to right across each row. This positioning highlights period 3 as a key segment where elemental character shifts gradually from metallic to non-metallic properties with rising , beginning with highly reactive metals on the left and ending with the inert on the right. The electronic structure of period 3 elements involves the addition of electrons to the third principal , specifically filling the 3s subshell first (for sodium and magnesium) and then the 3p subshell (for aluminum through ), achieving a full octet of eight valence electrons in the case of .

Historical context

The period 3 elements encompass a diverse group whose discoveries spanned from antiquity to the late , laying foundational insights into chemical periodicity. , known since prehistoric times for its use in pigments, medicines, and early , was recognized as an element in 1809 by French chemists Louis-Joseph Gay-Lussac and Louis-Jacques Thénard through decomposition studies confirming its indivisibility. was first isolated in 1669 by German alchemist during experiments distilling urine, where it appeared as a glowing substance due to slow oxidation in air. was discovered in 1774 by Swedish chemist through the reaction of with , though its elemental status was firmly established in 1810 by via . In the early 19th century, advances in enabled the isolation of several metallic period 3 elements. was discovered in 1807 by English , who obtained the metal by electrolyzing molten caustic soda (). followed in 1808, also isolated by Davy from magnesia () through electrolysis, though pure samples were later refined in 1831 by Antoine-Alexandre Brutus Bussy using a reaction with . Aluminum was first isolated in impure form in 1825 by Danish by reducing aluminum with potassium amalgam, with purer metal achieved in 1827 by German via a similar reduction of aluminum . was prepared in amorphous form in 1824 by Swedish by heating potassium fluosilicate with metal, marking the first isolation of this non-metal. The discovery of in 1894 by English physicist Lord Rayleigh and Scottish chemist profoundly impacted atomic theory, as the —isolated from atmospheric air through and chemical removal of other components—did not fit into the existing periodic due to its chemical unreactivity and monatomic nature. This anomaly challenged prevailing views of atomic weights and valency, prompting Ramsay's subsequent isolation of other and leading to the recognition of group 18 as a new family of inert elements by the early . Rayleigh and Ramsay's work earned them Nobel Prizes in 1904, underscoring argon's role in expanding the periodic table. Dmitri Mendeleev's 1869 periodic table highlighted the significance of period 3 elements by predicting undiscovered ones based on gaps, such as eka-aluminum (later , discovered in 1875), which he forecasted to have properties closely analogous to aluminum, including a low around 6 g/cm³ and a near 210°C—predictions remarkably validated by gallium's actual traits. The sequential filling of 3s and 3p orbitals in period 3 elements, from sodium through , provided key evidence supporting the theory in the early , particularly through Niels Bohr's 1913 atomic model and subsequent quantum developments, which explained the periodicity observed in their chemical behaviors and confirmed the third shell's capacity for eight electrons.

Atomic and electronic properties

Electron configurations

The electron configurations of period 3 elements, which span atomic numbers 11 to 18 (sodium through ), are determined by adding electrons to the core ([Ne], or 1s² 2s² 2p⁶) in the third principal . These configurations follow the standard notation and are as follows:
ElementAtomic NumberElectron Configuration
Sodium (Na)11[Ne] 3s¹
Magnesium (Mg)12[Ne] 3s²
Aluminum (Al)13[Ne] 3s² 3p¹
(Si)14[Ne] 3s² 3p²
(P)15[Ne] 3s² 3p³
(S)16[Ne] 3s² 3p⁴
(Cl)17[Ne] 3s² 3p⁵
(Ar)18[Ne] 3s² 3p⁶
The governs this sequential filling, directing electrons to occupy orbitals of lowest energy first, starting with the 3s subshell (capacity: 2 electrons) followed by the 3p subshell (capacity: 6 electrons), resulting in a total of 8 valence electrons by . In period 3, there are no deviations from this order, unlike in the d-block transition metals where electron promotion can occur for stability. The further constrains this process, stipulating that no two electrons in an atom can share identical sets of four quantum numbers (n, l, m_l, m_s), thus limiting each orbital to a maximum of two electrons with opposite spins. The number of valence electrons in these configurations increases progressively from 1 in sodium to 8 in , which underpins the transition from metallic to non-metallic character across the period. In textual representation of orbital diagrams, the 3s subshell for magnesium features two electrons with paired spins (↑↓), while the 3p subshell fills sequentially: for aluminum, one in one p orbital (↑ --- ---); for , three across three p orbitals per Hund's rule (↑ ↑ ↑); and for , four electrons with one pair (↑↓ ↑ ↑). This pairing in p orbitals becomes complete in , with all six electrons occupying three filled orbitals (↑↓ ↑↓ ↑↓). Across period 3, the decreases from sodium (186 pm) to (94 pm). This trend arises because the nuclear charge increases from left to right while the provided by the inner 2p electrons remains constant, leading to a higher (Zeff) that pulls the valence electrons closer to the nucleus./08%3A_Periodic_Properties_of_the_Elements/8.06%3A_Periodic_Trends_in_the_Size_of_Atoms_and_Effective_Nuclear_Charge) The first generally increases across period 3, from 496 kJ/mol for sodium to 1521 kJ/mol for , reflecting the stronger attraction of valence electrons to the nucleus due to rising Zeff and decreasing atomic size. However, there are notable dips: aluminum (578 kJ/mol) has a lower value than magnesium (738 kJ/mol) because its occupies a higher-energy 3p orbital that experiences less penetration toward the nucleus and thus weaker attraction; similarly, (1000 kJ/mol) has a slightly lower value than (1012 kJ/mol) due to increased electron-electron repulsion in the paired 3p orbital of , making electron removal easier. Electronegativity, measured on the Pauling scale, increases from 0.9 for sodium to 3.0 for across period 3, paralleling the gain in non-metallic character as atoms become better able to attract electrons in bonds due to higher Zeff. Argon has no defined value, as it does not form covalent bonds. This trend underscores how the constant shielding allows the increasing nuclear charge to enhance electron-attracting power without added screening from new shells.

Physical properties

Densities and phase states

The period 3 elements display distinct phase states at (STP, defined as 0 °C and 1 atm). The metals sodium, magnesium, and , along with the metalloids and the non-metals phosphorus and , exist as solids under these conditions. In contrast, chlorine and argon are gases, reflecting their weak intermolecular forces and low boiling points. Allotropes play a significant role in the physical properties of and , particularly their . White , the most common allotrope at , has a of 1.82 g/cm³, while the more stable red allotrope is denser at 2.16–2.34 g/cm³ due to its polymeric structure. For , the rhombic form (stable below 95.5 °C) exhibits a of 2.07 g/cm³, whereas the monoclinic form (stable between 95.5 °C and 119 °C) has a slightly lower of 1.96 g/cm³, arising from differences in packing. Densities across the period show an irregular progression, influenced by increasing offset by decreasing atomic radii and shifts in from metallic to covalent network to molecular structures./Descriptive_Chemistry/Elements_Organized_by_Period/Period_3_Elements/Physical_Properties_of_Period_3_Elements) Early metallic elements are relatively low-density due to loosely packed body-centered cubic structures, while non-metals transition to lighter forms as atomic size contracts but packing efficiency varies. The table below summarizes representative densities (using standard allotropes where applicable) and phase states at STP.
ElementSymbolDensity (g/cm³)Phase at STP
SodiumNa0.97Solid
MagnesiumMg1.74Solid
Al2.70Solid
Si2.33Solid
P1.82 (white)Solid
S2.07 (rhombic)Solid
Cl0.0032Gas
Ar0.0018Gas
Data sourced from standard reference tables; densities for gases are at STP.

Melting and boiling points

The melting and boiling points of period 3 elements exhibit distinct trends influenced by their atomic structures and bonding mechanisms. Sodium and magnesium, as alkali and alkaline earth metals, display relatively low values due to weak metallic bonding involving delocalized electrons from few valence shells. Aluminium shows higher points, reflecting stronger metallic bonds from additional electrons. Silicon reaches a peak with its giant covalent network structure. In contrast, phosphorus, sulfur, chlorine, and argon have low points owing to simple molecular or atomic forms held by weak intermolecular forces. The following table summarizes the melting and boiling points for these elements, using standard conditions and noting allotropes where relevant:
ElementMelting Point (°C)Boiling Point (°C)Notes
Sodium97.8883Metallic solid
Magnesium6501090Metallic solid
6602520Metallic solid
14143265Covalent network solid
44 (white)281Molecular (P₄)
115 (rhombic)445Molecular (S₈)
-102-34Diatomic gas (Cl₂)
-189-186
(Data sourced from Royal Society of Chemistry periodic table entries: Na, Mg, Al, Si, P, S, Cl, Ar. Values rounded for clarity; precise measurements vary slightly by source.) Across the metals from sodium to , melting and s increase progressively because the strengthens with more delocalized electrons available per atom—sodium contributes one, magnesium two, and aluminium three—leading to greater electrostatic attraction in the lattice. This trend peaks at , where the giant covalent structure involves strong directional bonds throughout an extended diamond-like lattice, requiring substantial energy to disrupt. 's notably high , in particular, arises from its robust with highly delocalized electrons, allowing it to vaporize only at elevated temperatures despite a similar to magnesium. From phosphorus onward, the points drop sharply as the elements form discrete molecules: white as P₄ tetrahedra, sulfur as S₈ rings, chlorine as Cl₂ diatomic units, and argon as isolated atoms. These are linked solely by weak van der Waals forces, which provide minimal resistance to thermal disruption, resulting in low melting and boiling points; for instance, 's anomalously low stems directly from its molecular P₄ form rather than a polymeric structure. This shift from extended bonding to molecular isolation underscores the transition from metals to non-metals in the period.

Chemical properties and reactivity

Oxidation states and bonding

The oxidation states of period 3 elements reflect their group positions and configurations, with s-block elements showing limited positive states corresponding to ns electron loss, while p-block elements display a wider range due to np involvement and potential octet expansion. Sodium (Na) exhibits only the +1 oxidation state, achieved by loss of its single 3s . Magnesium (Mg) is restricted to +2, losing both 3s electrons. Aluminium (Al) commonly adopts +3 by emptying its 3s and 3p orbitals. Silicon (Si) primarily shows +4 in oxides and halides, with a rare -4 state in silanes like silane (SiH₄). Phosphorus (P) has oxidation states of -3 (e.g., , PH₃), +3 (e.g., phosphite), and +5 (e.g., ). Sulfur (S) features -2 (sulfides), +4 (sulfites), and +6 (sulfates). Chlorine (Cl) displays -1 (chlorides), +1 (), +5 (), and +7 (). (Ar), as a , maintains 0 in its elemental form.
ElementCommon Oxidation States
Na+1
Mg+2
Al+3
Si+4 (rare -4)
P-3, +3, +5
S-2, +4, +6
Cl-1, +1, +5, +7
Ar0
Bonding behaviors in period 3 elements transition from metallic on the left to covalent and intermolecular forces on the right, influenced by increasing electronegativity and decreasing metallic character across the period. Sodium, magnesium, and aluminium form metallic lattices in their elemental states, characterized by delocalized valence electrons providing conductivity and malleability. Silicon adopts a covalent network structure in its diamond-like solid form, with each atom tetrahedrally bonded to four others via sp³ hybrid orbitals, resulting in high hardness and melting point. Phosphorus exists as covalent molecular P₄ tetrahedra, sulfur as S₈ crowns, and chlorine as Cl₂ diatomic molecules, all linked by weak van der Waals forces in the solid phase. Argon atoms interact solely through van der Waals forces, yielding a face-centered cubic lattice with minimal interatomic attraction. The valence electrons dictate bonding preferences: s-block elements (Na, Mg) have low ionization energies and electronegativities, favoring in compounds by complete electron transfer to highly electronegative elements like or oxygen. In contrast, p-block elements exhibit higher electronegativities, promoting covalent bonding through electron sharing; later members (P, S, Cl) access multiple oxidation states via d-orbital participation, enabling expanded octets beyond eight electrons, as seen in hypervalent molecules. For instance, (NaCl) exemplifies with a +1 Na and -1 Cl, forming a rock-salt lattice stabilized by electrostatic forces. (SiO₂) demonstrates covalent network bonding, with Si in +4 state linked to O in a tetrahedral arrangement of SiO₄ units. (PCl₅) illustrates covalent bonding and the +5 state, featuring trigonal bipyramidal geometry around P with five Cl atoms. The reactivity of period 3 elements exhibits a clear trend from left to right, transitioning from highly reactive metals to moderately reactive metalloids, increasingly reactive nonmetals, and finally an inert . On the metallic side, sodium demonstrates extreme reactivity, vigorously reacting with to produce and gas via the equation 2Na+2H2O→2NaOH+H22\mathrm{Na} + 2\mathrm{H_2O} \rightarrow 2\mathrm{NaOH} + \mathrm{H_2}, often igniting due to the exothermic nature of the process. Magnesium, while still reactive, shows reduced vigor compared to sodium, burning in air to form and reacting with more slowly, particularly when hot. Aluminum's reactivity further diminishes; it rapidly forms a thin, protective layer (Al2O3\mathrm{Al_2O_3}) upon exposure to air, which passivates the surface and prevents further oxidation under normal conditions, though the oxide is amphoteric, dissolving in both acids and bases. Silicon, as a , displays mild reactivity overall, primarily forming a stable, impervious (SiO2\mathrm{SiO_2}) layer when exposed to oxygen, which protects the underlying material and contributes to its use in semiconductors. Among the nonmetals, reactivity escalates from to to . Phosphorus and sulfur ignite spontaneously in air and form acidic oxides, while chlorine is a potent that readily reacts with metals like sodium to form : 2Na+Cl2→2NaCl2\mathrm{Na} + \mathrm{Cl_2} \rightarrow 2\mathrm{NaCl}, releasing significant energy. This progression culminates in , the that remains chemically inert due to its stable , showing no tendency to form compounds under standard conditions. Overall, this progression reflects a shift in reducing power from the left (metallic elements readily lose electrons) to oxidizing power on the right (nonmetals gain electrons), influenced by increasing and decreasing atomic size across the period. Diagonal relationships arise due to similarities in ionic radii and charge densities, such as between and magnesium, and between and , leading to comparable chemical behaviors like the formation of similar compounds. Regarding oxide and hydride formation, metallic elements produce basic s (e.g., Na2O\mathrm{Na_2O}, MgO\mathrm{MgO}) and hydrides that react with to yield hydroxides, while nonmetallic s (e.g., SO2\mathrm{SO_2}, P4O10\mathrm{P_4O_{10}}) are acidic, dissolving in to form oxoacids; aluminum serves as an amphoteric bridge.

Individual elements

Sodium

Sodium is a soft, silvery-white with an of approximately 23.0 u, making it one of the lightest metals. Due to its high reactivity with oxygen and moisture in the air, metallic sodium must be stored under oil or in an inert atmosphere to prevent spontaneous ignition or tarnishing. This reactivity aligns with broader trends in alkali metals, where sodium exhibits vigorous reactions with , producing hydrogen gas and . Among sodium's most important compounds are (NaCl), commonly known as table salt, which serves as a fundamental dietary and industrial chemical; (NaOH), or caustic soda, widely used in production and pH regulation; and (Na₂CO₃), referred to as soda ash, essential for glass manufacturing and . Metallic sodium is produced industrially through the of molten in a Down's cell, where the molten salt is heated to about 600°C, and an decomposes it into sodium metal at the and gas at the . Sodium carbonate, in turn, is primarily obtained from natural deposits or via the , contributing to its global production of over 60 million tons annually. Sodium finds diverse applications, including in high-pressure sodium-vapor lamps, which emit a warm light efficient for street and industrial lighting due to their high . In , liquid sodium serves as a in fast breeder reactors because of its excellent thermal conductivity and low neutron absorption, enabling efficient at high temperatures without pressurization. Biologically, sodium functions as a major extracellular , crucial for maintaining osmotic balance and facilitating impulses through the sodium-potassium mechanism. Notably, sodium exhibits a with magnesium, sharing similarities in compound solubility and reactivity patterns, such as the formation of sparingly soluble carbonates and hydroxides, despite their adjacent positions in period 3. Sodium is particularly abundant in , where it constitutes about 10.8 grams per kilogram, primarily as NaCl, accounting for roughly 30% of the total dissolved ions.

Magnesium

Magnesium is a with 12 and standard of 24.305 u. It appears as a lightweight, silvery-white metal with low density of 1.738 g/cm³, making it the lightest structural metal. In Period 3, magnesium exhibits trends in physical properties such as increasing density and compared to sodium, reflecting its position among the alkaline earth metals. When ignited, magnesium burns with an intense white flame due to its high reactivity with oxygen, producing . Key compounds of magnesium include magnesium oxide (MgO), known as magnesia, which serves as a refractory material in high-temperature applications like crucibles and furnace linings owing to its high melting point of 2852°C and thermal stability. Magnesium sulfate (MgSO₄·7H₂O), commonly called Epsom salt, is used in medicine as a laxative and for muscle relaxation baths. In organic synthesis, Grignard reagents (RMgX, where R is an alkyl or aryl group and X is a halogen) are organomagnesium compounds formed by reacting magnesium with alkyl halides in ether solvents; these act as strong nucleophiles for carbon-carbon bond formation in reactions like the addition to carbonyls./Aldehydes_and_Ketones/Synthesis_of_Aldehydes_and_Ketones/Grignard_Reagents) Magnesium finds extensive applications in alloys, particularly in and automotive industries, where its low weight enhances ; for example, the AZ31 (composed of magnesium with 3% aluminum and 1% ) is used in fuselages and car components like gearbox housings. Its pyrophoric nature makes it ideal for flares and , providing bright white illumination through rapid . Additionally, magnesium is essential in dietary supplements to support nerve function, , and , with recommended daily intakes of 310–420 mg for adults. A unique property of magnesium is its ability to burn in carbon dioxide atmospheres, as demonstrated by the reaction where ignited magnesium reduces CO₂ to carbon and forms magnesium oxide: 2Mg+CO2→2MgO+C2\text{Mg} + \text{CO}_2 \rightarrow 2\text{MgO} + \text{C} This exothermic process highlights magnesium's strong affinity for oxygen. Biologically, magnesium is central to chlorophyll, the green pigment in plants, where the Mg²⁺ ion at the core of the porphyrin ring enables light absorption for photosynthesis; deficiency leads to chlorosis in plants.

Aluminium

Aluminium (Al), atomic number 13, is a in period 3 of the periodic table with a standard of 26.982 u. It appears as a silvery-white, lightweight solid that is highly ductile and malleable, allowing it to be readily formed into sheets, wires, and complex shapes. A key feature is its natural formation of a thin, adherent oxide layer (Al₂O₃) on exposure to air, which passivates the surface and imparts excellent corrosion resistance in most environments. Aluminium is the most abundant metal in Earth's crust, comprising approximately 8.2% by weight, primarily occurring in the mineral bauxite, a hydrated oxide ore. Extraction involves refining bauxite to alumina (Al₂O₃) via the Bayer process, followed by electrolysis in the Hall-Héroult method, where purified alumina is dissolved in molten cryolite (Na₃AlF₆) and reduced at carbon anodes to produce molten aluminium. Alumina itself appears as corundum in its crystalline form and exhibits amphoteric behavior, dissolving in strong bases to form aluminates, as in the reaction: Al2O3+2NaOH→2NaAlO2+H2O\text{Al}_2\text{O}_3 + 2\text{NaOH} \rightarrow 2\text{NaAlO}_2 + \text{H}_2\text{O} This property distinguishes it from purely basic oxides of earlier period 3 metals. Common aluminium compounds include aluminium chloride (AlCl₃), a versatile Lewis acid that accepts electron pairs due to its electron-deficient aluminium center, widely employed as a catalyst in organic reactions like Friedel-Crafts acylations. Another is potassium aluminium sulfate (KAl(SO₄)₂·12H₂O), or alum, utilized in water treatment for coagulation and in textile dyeing as a mordant. The metal's combination of low density (about one-third that of ), high thermal and electrical conductivity, and non-toxicity enables diverse applications, including lightweight foils and beverage cans for , structural components in , and alloys in automotive and industries.

Silicon

Silicon is a element in period 3 with 14 and of 28.1 u. In its pure form, it appears as a hard, brittle, gray crystalline solid with a metallic luster, exhibiting a crystal structure where each silicon atom is tetrahedrally coordinated to four others via covalent bonds./Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14:The_Carbon_Family/Z014_Chemistry_of_Silicon(Z14)) This structure contributes to its properties, enabling controlled electrical conductivity through doping with impurities to create n-type or p-type materials essential for electronic devices. is the second most abundant element in , comprising approximately 28% by mass, primarily occurring in rather than as the free element. The most common compounds of silicon include (SiO₂), known as silica, which forms the basis of , , and . Silica is a network solid with a three-dimensional tetrahedral arrangement of SiO₄ units, providing structural integrity to many rocks and soils. Silicates, derived from silica by incorporating metal cations, constitute the majority of the and include chain and ring structures found in clays and various minerals like feldspars and micas. Unlike carbon, which readily forms diverse catenated chains and multiple bonds for organic versatility, silicon prefers oxygen-bridged networks in silicates, limiting its to less stable forms./Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14:The_Carbon_Family/Z014_Chemistry_of_Silicon(Z14)/Silicon_and_Group_14_Elements) Silicones represent another class of silicon compounds, consisting of synthetic polymers with a repeating -Si-O-Si- backbone and organic substituents, such as methyl groups, granting them flexibility, thermal stability, and water repellency for use in lubricants, sealants, and medical implants. Silicon's semiconductor characteristics make it indispensable in modern technology, particularly through doping processes that introduce or to alter its electrical properties for fabricating integrated circuits and microchips. These doped silicon wafers form the foundation of transistors and diodes in computers and . Additionally, high-purity is crucial for photovoltaic solar cells, converting sunlight to with efficiencies up to 25% in commercial panels. (SiC), a compound formed by reacting with carbon, serves as a hard material for grinding and metals, ceramics, and stones due to its high thermal conductivity and wear resistance.

Phosphorus

Phosphorus is a non-metal element in period 3 of the periodic table with 15 and of 31.0. It exists in several allotropes, each with distinct physical and chemical properties, reflecting its versatility as a reactive non-metal essential to both industrial applications and biological systems. Discovered in by German alchemist during experiments aimed at isolating the from , phosphorus was named for its glowing appearance, evoking the "morning star" (from Greek phosphoros, meaning light-bearer). The most common allotropes are white, red, and black . White appears as a waxy, translucent solid that is highly toxic and spontaneously ignites in air, producing a characteristic garlic-like odor; it exhibits , glowing faintly in the dark due to slow oxidation. To prevent ignition, white is stored underwater. Red is an amorphous, non-toxic powder that is more stable in air and less reactive than its white counterpart. Black , resembling in structure, is the most thermodynamically stable allotrope, dense and semiconducting with a layered morphology. Key phosphorus compounds include (\ceH3PO4\ce{H3PO4}), a tribasic widely used in and rust removal, and various phosphates derived from it, such as fertilizers that enhance for growth. (\cePH3\ce{PH3}) is a toxic, flammable gas employed as a fumigant and in . These compounds underscore phosphorus's role in , where phosphates constitute about 80-90% of global phosphorus consumption to support . Applications of phosphorus leverage its allotropes and compounds across industries. Red phosphorus serves as a safe, non-toxic ingredient in safety matches, providing the striking surface that ignites via friction. Phosphates appear in detergents as water softeners and in flame retardants to inhibit combustion in textiles and plastics. Biologically, phosphorus is an essential macronutrient, integral to nucleic acids like DNA and energy carriers like ATP, supporting cellular functions in all living organisms; however, excess phosphate runoff from fertilizers can trigger eutrophication in water bodies, leading to algal blooms that deplete oxygen and harm aquatic ecosystems.

Sulfur

Sulfur is a nonmetallic element in period 3 of the periodic table, appearing as a bright solid at with an atomic mass of 32.065 u. It exists in multiple allotropes, the most stable being rhombic sulfur (α-sulfur), which consists of S8 crown-shaped rings in an orthorhombic , and monoclinic sulfur (β-sulfur), stable between 95.5°C and 119°C with a similar but differently arranged ring structure. Another notable form is plastic sulfur, an amorphous, rubbery produced by rapidly cooling molten sulfur, demonstrating sulfur's capacity for —forming long chains of sulfur atoms similar to carbon but with weaker S-S bonds. Sulfur forms a variety of compounds, including oxides and sulfides central to its chemistry and environmental role. (SO2), a colorless gas produced when sulfur burns in air, acts as a key pollutant contributing to and respiratory issues by reacting with atmospheric to form . (H2SO4), often dubbed the "king of chemicals" for its vast industrial production exceeding 200 million tons annually, is a strong diprotic acid used in fertilizers, batteries, and refining. Common sulfides include iron (FeS2), a cubic that serves as a major natural sulfur source and exemplifies sulfur's role in metal ores. Industrially, sulfur is essential for applications like the of rubber, where it cross-links chains to enhance durability, a process invented in 1839 and still foundational to tire production. It also features in black gunpowder as an oxidizer alongside charcoal and potassium nitrate, enabling combustion for historical explosives and fireworks. Most is manufactured via the , involving catalytic oxidation of SO2 to SO3 using vanadium pentoxide, followed by absorption in concentrated H2SO4 to produce the acid. Unique geological occurrences include vibrant deposits around volcanic fumaroles, such as crystalline formations at Kīlauea Volcano in , where sulfur gases condense into yellow "flowers" or encrustations resembling crowns.

Chlorine

Chlorine is a halogen element in period 3 of the periodic table, with atomic number 17 and standard atomic weight of 35.45. It exists as a diatomic molecule, Cl₂, under standard conditions, appearing as a pale yellow-green gas that is denser than air and possesses a pungent, irritating odor. This gas is highly toxic, causing severe respiratory irritation and damage upon inhalation even at low concentrations, and it liquefies under moderate pressure or cooling. Chlorine exhibits the highest standard among period 3 elements at +1.36 V for the Cl₂/Cl⁻ couple, making it the strongest oxidant in the period and enabling it to readily accept electrons in reactions. Common compounds include (HCl), a strong acid also known as muriatic acid, used in industrial cleaning and ; (NaClO), the active ingredient in for disinfection; and organochlorine compounds such as (PVC), a durable , and various pesticides derived from chlorinated hydrocarbons. Industrially, chlorine is primarily produced via the chlor-alkali process, an electrolytic method that decomposes ( solution) to yield chlorine gas, , and hydrogen. Key applications of chlorine leverage its oxidizing and disinfecting properties, including where it inactivates pathogens in and ; production of PVC for pipes, flooring, and packaging; and synthesis of pharmaceuticals, such as antibiotics and antiseptics. Historically, gas was deployed as a during , notably in the 1915 German attack at , causing thousands of casualties through and choking. Additionally, chlorine-containing chlorofluorocarbons (CFCs), once widely used as refrigerants and propellants, contribute to stratospheric by releasing chlorine atoms that catalytically destroy ozone molecules.

Argon

Argon is a colorless, odorless that constitutes approximately 0.93% of Earth's atmosphere by volume, making it the third most abundant atmospheric gas after and oxygen. With an of 39.948 u, argon exhibits remarkable chemical inertness due to its stable electronic configuration, featuring a full outer shell of eight electrons that resists forming bonds under standard conditions. This inertness positions argon as the concluding element in period 3 of the periodic table, exemplifying the trend toward increasing stability in atomic structure across the period. Discovered in 1894 by British scientists Lord Rayleigh and Sir William Ramsay through meticulous analysis of atmospheric gases, argon was isolated as a non-reactive residue after removing oxygen, , and other known components from air. This marked argon as the first identified, challenging existing periodic table frameworks and paving the way for the recognition of group 18 elements. The element's name derives from the Greek word argos, meaning "inactive" or "lazy," a fitting descriptor for its reluctance to participate in chemical reactions. Despite its inert nature, argon can form rare and unstable compounds under extreme conditions, such as the (HArF) molecule, which remains stable only below -246°C and serves primarily in fundamental research on chemistry. More practically, argon participates in transient states, like the ArF* species in argon fluoride lasers, which emit light at 193 nm for applications in and medical treatments. Argon's inert properties underpin diverse applications, including its role as a shielding gas in to prevent oxidation of molten metals. In lighting, it fills incandescent and fluorescent bulbs, extending filament life by inhibiting reactions with oxygen. Medically, liquid argon enables , where extreme cold from expanding argon gas destroys abnormal tissues, such as in cancer treatments. Additionally, argon's isotopic stability facilitates potassium-argon (K-Ar) dating, a geochronological method that measures the accumulation of ⁴⁰Ar from the of ⁴⁰K in volcanic rocks to determine ages up to billions of years.

Occurrence and biological significance

Natural abundance and production

In the cosmos, period 3 elements exhibit varying abundances shaped by processes, with magnesium and ranking among the most prevalent heavy elements due to their production in massive stars via silicon burning and explosive . Relative to in the solar photosphere, their number abundances (log ε, where ε = N(X)/N(H) × 10¹²) are approximately 7.60 for Mg, 7.51 for Si, 7.12 for S, 6.45 for Al, 6.40 for Ar, 6.24 for Na, 5.50 for Cl, and 5.41 for . These values reflect the solar system's primordial composition, with Mg and Si particularly enriched compared to lighter elements like Na and heavier ones like . On , the distribution of period 3 elements is dominated by geochemical processes, concentrating them in the crust through igneous, sedimentary, and metamorphic activities. and aluminum are by far the most abundant, forming the backbone of , while sodium and magnesium are moderately common, and , , , and occur at trace levels. , as a , is primarily atmospheric rather than crustal. The table below summarizes their mass abundances in the continental crust:
ElementSymbolAbundance (wt%)
SiliconSi27.7
AluminumAl8.1
SodiumNa2.8
MagnesiumMg2.1
PhosphorusP0.07
SulfurS0.03
ChlorineCl0.01
ArgonAr<0.001
Natural sources for these elements are tied to geological formations. Sodium primarily occurs in evaporite deposits such as halite (NaCl) and in feldspar minerals within igneous rocks. Magnesium is abundant in seawater (about 1.3 g/L) and minerals like dolomite (CaMg(CO₃)₂) and magnesite (MgCO₃). Aluminum is concentrated in bauxite ores, which are weathered residuals rich in aluminum hydroxides. Silicon is ubiquitous in quartz (SiO₂) and sands derived from it. Phosphorus is mainly in sedimentary phosphate rocks, particularly apatite [Ca₅(PO₄)₃(F,Cl,OH)], often of marine origin. Sulfur appears in native elemental form, sulfides (e.g., pyrite), gypsum (CaSO₄·2H₂O), and volcanic emissions. Chlorine is extracted from seawater (1.9% by weight as NaCl) and evaporite salts. Argon comprises 0.934% of the atmosphere by volume, originating from the radioactive decay of potassium-40 in the crust. Industrial production methods leverage these sources through energy-intensive processes, often electrochemical or thermal. Sodium metal is produced commercially via electrolysis of molten sodium chloride in the Downs cell, where NaCl is mixed with CaCl₂ to lower the melting point, yielding sodium at the cathode and chlorine at the anode. Magnesium is obtained either by electrolyzing magnesium chloride derived from seawater or brine (Dow process) or by thermal reduction of calcined dolomite with ferrosilicon in the Pidgeon process. Aluminum is extracted from bauxite via the Bayer process to produce alumina (Al₂O₃), followed by electrolysis in the Hall-Héroult process using cryolite flux. Metallurgical-grade silicon is manufactured by carbothermic reduction of quartz in an electric arc furnace: SiO₂ + 2C → Si + 2CO. White phosphorus (P₄) is produced by heating phosphate rock with coke and silica:
4 \mathrm{Ca_5(PO_4)_3F + 18 SiO_2 + 30 C \rightarrow 3 P_4 + 30 \mathrm{CO} + 18 \mathrm{CaSiO_3} + \mathrm{CaF_2}
Sulfur is recovered using the Frasch process, which involves injecting superheated water into underground deposits to melt and pump it to the surface. Chlorine gas is generated by electrolyzing aqueous NaCl brine in the chlor-alkali process (e.g., membrane cell), co-producing sodium hydroxide. Argon is isolated as a byproduct of air separation through fractional distillation of liquefied air, exploiting its boiling point between oxygen and nitrogen. These extraction activities carry environmental consequences, particularly habitat disruption. For example, mining activities in the Brazilian , including mining for aluminum, have caused significant , with one study showing mining-related forest loss extending up to 70 km beyond boundaries and totaling 11,670 km² of cleared land between 2005 and 2015.

Biological roles

Period 3 elements play diverse roles in biological systems, with several serving essential functions in cellular processes, structural components, and metabolic pathways across organisms, while others exhibit or minimal involvement. Sodium, magnesium, , , and are macronutrients critical for life, whereas has trace benefits in certain , aluminum is neurotoxic, and is biologically inert. Sodium is vital for maintaining electrolyte balance and nerve function in animals, primarily through the Na⁺/K⁺ pump that generates membrane potentials essential for signal transmission. In plants, it supports osmotic regulation, sometimes substituting for potassium. Magnesium acts as a cofactor in over 300 enzymes, facilitating ATP hydrolysis and stabilizing ribosomes, and forms the core of chlorophyll in photosynthesis. Phosphorus is indispensable for all life, forming the backbone of DNA and RNA, and serving as the key element in ATP for energy transfer; it also contributes to bone and teeth structure as calcium phosphate. Sulfur is incorporated into amino acids like cysteine and methionine, enabling protein folding and redox reactions via cofactors such as coenzyme A. Chlorine, as chloride ions, regulates osmotic pressure and fluid balance in cells, and is a component of gastric hydrochloric acid (HCl) for digestion. Silicon plays a trace role in forming silica shells (frustules) in diatoms, enhancing structural integrity in aquatic algae, and aids synthesis in connective tissues of animals. Aluminum lacks any essential biological function and is toxic, accumulating in the and bones to cause , including links to dialysis encephalopathy and potential contributions to pathology. Argon, a , has no known biological role due to its chemical inertness. These elements participate in key biogeochemical cycles influencing ecosystems. The involves uptake by organisms from soils and , with human activities like application leading to runoff that causes ; is often limited in oceans, constraining primary productivity. In the , reduce (SO₄²⁻) to (H₂S), supporting and recycling in sediments. Sodium and , as NaCl, regulate in marine environments, where elevated levels can stress freshwater organisms by disrupting balance. Deficiencies and excesses highlight their physiological impacts. Magnesium deficiency leads to muscle cramps and impaired nerve function due to disrupted enzyme activity. Excess chlorine intake via salt (NaCl) contributes to hypertension by increasing blood volume and vascular resistance. Phosphorus remains essential but its scarcity in ocean surface waters limits phytoplankton growth, affecting global carbon cycling.

Comparative data

Property summary table

ElementSymbolAtomic NumberElectron ConfigurationAtomic Radius (pm)1st Ionization Energy (kJ/mol)Electronegativity (Pauling)Density (g/cm³)Melting Point (°C) / Boiling Point (°C)Common Oxidation States
SodiumNa11[Ne] 3s¹186 (metallic)4960.930.9798 / 883+1
MagnesiumMg12[Ne] 3s²160 (metallic)7381.311.74650 / 1090+2
AluminiumAl13[Ne] 3s² 3p¹143 (metallic)5781.612.70660 / 2519+3
SiliconSi14[Ne] 3s² 3p²111 (covalent)7861.902.331410 / 3260+4, -4
PhosphorusP15[Ne] 3s² 3p³107 (covalent, white)10122.191.82 (white)44 / 277+5, +3, -3
SulfurS16[Ne] 3s² 3p⁴105 (covalent)10002.582.07 (rhombic)115 / 445+6, +4, -2
ChlorineCl17[Ne] 3s² 3p⁵99 (covalent)12513.160.003 (gas)-101 / -34-1, +1, +5, +7
ArgonAr18[Ne] 3s² 3p⁶N/A (van der Waals 188)1520N/A0.0018 (gas)-189 / -1860
Data compiled from authoritative sources including NIST Atomic Properties for configurations and ionization energies (converted from eV using 1 eV ≈ 96.5 kJ/mol), and Royal Society of Chemistry periodic table pages for physical properties such as density, melting and boiling points , and Webelements for atomic radii and electronegativities . Oxidation states from standard chemical references. Values are approximate standard conditions; for non-metals, radii are covalent where applicable, densities for common forms. The elements in period 3 display a pronounced trend from metallic to non-metallic character moving from left to right, reflecting increasing effective nuclear charge and decreasing atomic radius. Sodium, magnesium, and aluminum behave as metals with delocalized electrons enabling electrical conductivity, whereas phosphorus, sulfur, and chlorine act as non-metals with localized electrons that insulate. Silicon occupies an intermediate position as a metalloid and semiconductor, facilitating controlled electron flow essential for modern electronics. Key physical properties underscore this periodicity: atomic radii decrease across the period as protons pull valence electrons closer, while first ionization energies generally rise due to the same electrostatic attraction, though subshell effects cause minor dips. Melting points increase to a maximum at , driven by the shift from weaker metallic bonds to robust giant covalent networks, before declining with simple molecular structures on the right. Chemically, reactivity evolves from strong reducing behavior in metals, which readily lose electrons, to oxidizing tendencies in non-metals that gain electrons to achieve stability. These patterns affirm Mendeleev's periodic law, wherein elemental properties recur with atomic number, providing a framework for predicting behavior. Anomalies, such as the diagonal similarity between aluminum and silicon stemming from comparable charge densities that enhance polarizing power, deviate slightly from strict horizontal trends but reinforce overall coherence. In contemporary contexts, such trends inform compound design; for instance, sodium chloride forms an ionic lattice due to low charge density in Na⁺, while aluminum chloride exhibits covalent bonding from Al³⁺'s high charge density, affecting solubility and reactivity in industrial processes.

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