Alkali metals: alkali metals periodic table facts

What are the alkali metals?

The alkali metals are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). Together they fill the left-hand column of the periodic table, labelled group 1 in the modern IUPAC system and group IA in the older American CAS system. Anyone studying the Pa-231 decay chain and why the sample sizes studied are tiny should consult the protactinium atomic mass page, which also covers its rarity and uses.

Because lithium leads the column, chemists sometimes call the family the lithium family. The name "alkali" comes from the Arabic al-qalī, plant ashes, because the hydroxides these metals form dissolve into strongly alkaline solutions.

Caesium is the IUPAC spelling; in the United States the same element is written cesium, so US safety data sheets list cesium hydroxide where British texts say caesium hydroxide. Hydrogen also sits at the top of group 1, but it is a gas and a nonmetal under ordinary conditions, so it is not counted among the alkali metals.

On the other side of each alkali metal in the periodic table you find the alkaline earth metals of group 2 (beryllium, magnesium, calcium and so on), and one row up and across sits a noble gas. Each alkali metal atom has exactly one electron more than the noble gas before it, which is why it gives that electron up so easily.

The halogens of group 17, on the far side of the periodic table, are the mirror image: they need one electron to reach a noble-gas shell, which is why an alkali metal and a halogen combine so violently into a salt. Among all the chemical elements, the alkali metals show the clearest group trends, so the alkali metals are usually the first group of elements a chemistry course studies in detail.

Electron configuration of the group 1 elements

Every alkali metal has the same outer electron configuration: a full noble-gas core plus one electron in an s-orbital, written ns¹. That single valence electron sits far from the nucleus and is shielded by all the inner electrons, so the atoms lose it readily to form a cation with a +1 charge.

The +1 oxidation state is the only one these metals show in ordinary compounds. That rule holds for all six elements, which is why alkali metal chemistry is so easy to predict.

Because the group 1 elements all lose the same single electron, the resulting cations all have a noble-gas arrangement of electrons. That shared shape is why a sodium salt and a caesium salt often look and behave alike, and why their cations are almost colourless in solution.

Physical properties of alkali metals

As metals go, the alkali metals are unusual: they are far softer and lighter than the alkaline earth metals next door. All of them are silvery (caesium has a faint golden tint), and all are soft enough to cut with a knife; sodium cuts like cold butter. For readers asking about the rarest stable rare earth metal, the thulium element covers its price and reserves.

A fresh cut face is bright and shiny for only a moment before it tarnishes to a dull grey or white film as oxygen and water vapour attack it. The tarnish on lithium is often darker, because it also reacts with nitrogen in the air.

In the solid state each alkali metal packs its atoms in a body centered cubic lattice. Like other metals they are good conductors of heat and electricity, and they are malleable and ductile. The physical properties of these elements change steadily down the group, which makes them a classic example of how position in the periodic table predicts behaviour.

Melting point and boiling point down the group

Every one of the alkali metals has a low melting point for a metal, and the value falls with each step down the group. Caesium melts at 28.5 °C, so it turns liquid on a warm day or in your hand (if you were ever foolish enough to hold it). Francium is thought to melt even lower, but so little of it has ever existed at once that its bulk values are estimates.

The boiling point follows the same downward slope, from 1,342 °C for lithium to 671 °C for caesium. Sodium's short liquid range near 100 °C and its high boiling point are why liquid sodium works as a coolant in some fast nuclear reactors: it stays liquid over a wide temperature span without needing high pressure.

Density and metallic bonding

Lithium, sodium and potassium all have a density below 1 g/cm³, so they float on water while they react. Lithium is the least dense of all the solid elements, light enough to float on mineral oil. The general trend is upward down the group, because mass grows faster than atomic size, but potassium breaks the pattern and is slightly less dense than sodium.

The soft feel, low melting points and low densities all come from weak metallic bonding. Each atom gives just one electron to the shared "sea" of electrons that holds a metal together, and the effective nuclear charge holding that electron is small. As the atoms grow larger down the group, the electron sea is spread over more space and the bonds weaken further, so each of the heavier alkali metals is softer and melts lower than the one above it.

Periodic trends in the lithium family

Few groups of elements in the periodic table show periodic trends as cleanly as the alkali metals. Going from lithium to caesium, each new element adds a complete electron shell, so the outer electron sits further from the nucleus every time.

Almost every property of the alkali metal group follows from that one change. The heavier alkali metals are larger, give up their electron more easily and react more violently; the lighter alkali metals are smaller, harder and more restrained.

Atomic radius and ion size

The atomic radius grows sharply down the group, from 152 pm for lithium to 265 pm for caesium (metallic radii). The alkali metals have the largest atomic radii in each period of the periodic table. Their ions are much smaller than the atoms, because losing the outer electron removes the whole outer shell: a Na⁺ ion is roughly half the radius of a sodium atom.

Ion size matters in practice. The tiny Li⁺ ion pulls hard on neighbouring water molecules and anions, which is why lithium compounds are often the odd ones out in solubility, while large alkali metal ions such as Cs⁺ slip easily into big crystal lattices.

Ionisation energy and electronegativity

The first ionisation energy is the energy needed to strip the outer electron from a gaseous atom. For the alkali metals it is the lowest of any group of elements, and it falls from 520 kJ/mol for lithium to 376 kJ/mol for caesium.

(American sources spell it ionization energy; the quantity is the same.) A low ionisation energy is the direct reason these metals form +1 ions so readily.

Electronegativity, the pull an atom exerts on shared electrons, is very low across the group: 0.98 for lithium down to 0.79 for caesium on the Pauling scale. With electronegativity this low, an alkali metal bonded to a nonmetal almost always forms an ionic bond rather than sharing electrons, especially with elements such as chlorine and oxygen.

Francium is the exception worth knowing. Its measured ionisation energy (393 kJ/mol) is slightly higher than caesium's, not lower. Its electrons move so fast that relativistic effects pull the 7s orbital inward, holding the outer electron a little tighter than the simple trend predicts.

Why reactivity increases down the group

Each of the alkali metals is highly reactive, and reactivity climbs down the group because the outer electron is easier to remove each time. All of them are strong reducing agents, which means they push their electron onto whatever they react with.

Lithium is the gentlest; caesium (cesium in American texts) is the most reactive of all the stable metals, and rubidium and caesium ignite on their own in air at room temperature. As powerful reductants, the alkali metals never occur free in nature.

One subtlety: lithium has the most negative standard electrode potential of the group, so on paper it is the strongest reducing agent in water. It still reacts more slowly than sodium, because a small, high-melting atom releases its energy less quickly. Thermodynamics says how much energy a reaction gives out; kinetics says how fast.

How alkali metals react with water and air

All the alkali metals react with water, and the reaction is strongly exothermic. The heavier the metal, the faster and hotter the reaction. Heat from the reaction can ignite the hydrogen it releases, so with the heavier metals the result is often an explosion rather than a fizz.

In air, the same chemistry plays out more slowly. Oxygen and moisture attack the surface, and the oxidation forms oxides and hydroxides that leave the metal coated in a white layer. In both cases the solution or residue left behind is strongly alkaline, which is how the group got its name.

Hydrogen gas and a metal hydroxide

Every reaction of an alkali metal M with water follows the same balanced equation:

With any of the alkali metals, two atoms of metal give one molecule of hydrogen gas and two units of metal hydroxide. Here is a quick worked example with your own numbers.

Suppose you drop 0.85 g of lithium into excess water. Lithium's molar mass is 6.94 g/mol, so

At room temperature and pressure one mole of gas fills about 24.0 L, so that small piece releases about 1.47 L of hydrogen and leaves 2.93 g of lithium hydroxide in solution. Repeat the calculation with the same 0.85 g of sodium and you get only 0.444 L; with potassium, 0.261 L. The lightest metal reacts most calmly, yet gram for gram it produces the most hydrogen, because each gram contains far more atoms.

Alkali metal compounds and their oxidation state

Because every alkali metal holds the +1 oxidation state in its compounds, the formulas are simple and predictable: MCl, MOH, M₂CO₃, M₂SO₄. Nearly all alkali metal compounds are ionic, white or colourless, and soluble in water. Those chemical properties, and that simple +1 chemistry, make sodium and potassium salts some of the most widely used chemicals in industry, from soap making to glass and fertiliser.

Oxides, peroxides and superoxides

Burning the alkali metals in air shows the size trend in a neat way. Lithium forms a normal oxide, Li₂O. Sodium mainly forms the peroxide, Na₂O₂.

Potassium, rubidium and caesium form the superoxide, MO₂. The larger the cation, the better it can stabilise the larger oxygen anion next to it. All these oxides react with water to give hydroxide solutions, and potassium superoxide is used in breathing equipment because it releases oxygen while absorbing carbon dioxide.

Hydroxides, halides, hydrides and nitrides

• Alkali metal hydroxides such as sodium hydroxide (caustic soda) and potassium hydroxide are strong bases used to make soap, paper and biodiesel.
• Halides are the classic ionic salts; sodium chloride, common table salt, is the best known, and potassium chloride is the main source of potash fertiliser.
• Hydrides such as lithium hydride and sodium hydride contain the H⁻ ion and serve as strong bases and hydrogen stores.
• Nitrides: lithium is the only alkali metal that reacts directly with nitrogen at room temperature, forming dark red lithium nitride, Li₃N; no other alkali metal forms a stable nitride this way.
• Alkali metals also form intermetallic compounds with each other and with metals of groups 13 and 14, such as the liquid sodium-potassium alloy NaK.

Ammonia solutions and organolithium reagents

Dissolve sodium in liquid ammonia at −33 °C and the solution turns a deep blue. The metal gives up its electron, and that electron sits free in a cage of ammonia molecules.

At high concentration the liquid turns bronze and conducts like a metal. Chemists use these solutions as powerful reducing agents in the Birch reduction, a staple of synthetic chemistry.

Lithium also forms organolithium compounds, where lithium bonds directly to carbon. Butyllithium is one of the most used reagents in organic chemistry for making carbon-carbon bonds, and alkali metal cations can be held in crown ethers and cryptands, the coordination compounds that won the 1987 Nobel Prize in Chemistry.

Discovery of alkali metals by Davy, Bunsen and Kirchhoff

Sodium and potassium compounds have been used since ancient times, as salt, soda and potash (wood-ash extract), but for centuries nobody could separate the alkali metals from them. In 1807 Humphry Davy passed a current from a large voltaic pile through melted caustic potash and saw globules of a new metal burst into flame: potassium.

Days later he isolated sodium from caustic soda the same way. It was the first time electrolysis had isolated new elements.

Lithium was the next of these elements to be found, in 1817, when Johan August Arfwedson found it in the mineral petalite; the name comes from the Greek lithos, stone. Rubidium and caesium came next, found through light rather than chemistry, and francium last of all: Marguerite Perey identified it in 1939 at the Curie Institute as a decay product of actinium.

Flame test colours and emission spectra

In 1860 Robert Bunsen and Gustav Kirchhoff used their new spectroscope to examine the light from evaporated mineral water and saw two blue lines no known element produced. They named the element caesium, from the Latin caesius, sky blue.

A year later deep red lines in the mineral lepidolite gave them rubidium, from rubidus, deep red. Both elements take their names straight from the emission spectra, and the same physics is behind the flame test you can still run in a school lab:

Lithium

Crimson red

Sodium

Bright yellow-orange (the 589 nm line)

Potassium

Lilac, best seen through blue cobalt glass

Rubidium

Red-violet

Caesium

Blue-violet

Sodium's yellow is so intense that a trace of it masks the others, which is why a dirty wire gives a misleading flame test for potassium.

Where free alkali metals occur and how they are made

You will never find free alkali metals lying in a rock. They are too reactive, so in nature they always occur as compounds in minerals and dissolved salts. Sodium and potassium are each over 2% of Earth's crust by mass, ranking sixth and seventh in abundance among all elements.

Rubidium is surprisingly common at about 90 parts per million, more than lithium at about 20 ppm. Caesium is rarer at about 3 ppm, and francium is so radioactive that well under a kilogram exists on Earth at any moment.

Seawater holds vast amounts of dissolved salt, and evaporated ancient seas left thick beds of rock salt (the mineral halite) and potassium minerals such as sylvite. Lithium comes from hard-rock spodumene and from brines in the salt flats of Chile, Argentina and Bolivia.

Electrolysis in the Downs cell

Sodium is made industrially by the Downs process: molten sodium chloride, mixed with calcium chloride to lower its melting point to about 600 °C, is electrolysed in a Downs cell. Sodium metal collects at the iron cathode and floats to the top, while chlorine gas bubbles off at the graphite anode.

A steel mesh keeps the two products apart, because they would otherwise recombine explosively. Lithium comes from its chloride the same way, while potassium is usually produced by reducing hot potassium chloride with sodium vapour.

Uses of alkali metals in batteries, clocks and lamps

The applications of the alkali metals and their compounds reach from your kitchen to the definition of time itself:

• Caesium atomic clocks define the second as 9,192,631,770 oscillations of a caesium-133 transition; rubidium clocks are compact backups in telecom networks and GPS satellites.
• The sodium-vapour lamp gives the orange glow of older street lights, turning electricity into visible light very efficiently.
• Salt (NaCl) is a seasoning, a food preservative and a road de-icer; potassium salts feed most fertilisers.
• Lithium carbonate is a long-established medicine for bipolar disorder, and lithium is also added to aluminium alloys for aircraft and to lubricating greases.
• Caesium formate brines are dense drilling fluids for deep oil and gas wells.

Lithium batteries and dendrite growth on anodes

Batteries show the group's traits at work. Lithium is the lightest member of the family, it gives up its single outer electron easily, and it has the most negative electrode potential of the six, so it packs the most energy per gram into batteries. Lithium-ion cells move Li⁺ ions between a graphite anode and a metal-oxide cathode, and they power almost every phone, laptop and electric car.

Primary lithium batteries use lithium metal directly. The long-standing problem with metal anodes is dendrite growth: as the cell charges, lithium plates out in needle-like branches that can pierce the separator and short the cell. Research into solid electrolytes aims to fix that, while sodium-ion chemistry swaps in the next metal down the group: sodium carries the same +1 charge, so it works the same way, with a little less energy per gram but far cheaper and more plentiful supply.

Alkali metals in the body: biological role

Of the alkali metals, only sodium and potassium are essential elements for every animal. As electrolytes, Na⁺ and K⁺ carry the electrical signals in your nerves and muscles.

The sodium-potassium pump in each cell membrane uses energy to push three sodium ions out for every two potassium ions it brings in, and that gradient powers nerve impulses and controls water balance. Too much sodium raises blood pressure; too little potassium causes muscle weakness and an irregular heartbeat.

Lithium has no known biological role, but in carefully controlled doses it stabilises mood. Rubidium behaves like potassium and is taken up by the body without known harm at trace levels.

Caesium is not needed by the body either, and the radioactive isotope caesium-137, released by nuclear accidents, is a health hazard because the body absorbs it as if it were potassium. Francium is too short-lived to matter outside a laboratory.

Storage and safety for pure alkali metals

The same reactivity that makes the alkali metals interesting makes them dangerous. The main hazards are fire, burns from the caustic hydroxides and oxides they form, and the flammable hydrogen they release on contact with water. Any laboratory handling pure alkali metals should have a written procedure, trained people and no one working alone.

• Work in a glove box filled with inert gas where possible, or otherwise in a fume hood kept free of water and ignition sources.
• Wear safety glasses, chemical-resistant gloves and a flame-resistant lab coat.
• Cut and weigh small pieces quickly, under oil, and return the rest to its container at once.
• Quench scraps slowly with isopropanol or tert-butanol, which react far more gently than water, and treat residues as hazardous waste.

Mineral oil, argon and inert atmosphere storage

Sodium and potassium are kept submerged in mineral oil or dry kerosene to keep out air and moisture. Lithium floats on oil, so it is stored coated in oil, petroleum jelly or paraffin wax, or under argon; nitrogen is not inert toward lithium. Rubidium and caesium are sealed in glass ampoules under argon or vacuum, because they ignite in air.

Good storage also means watching old stock: potassium kept for years can grow a yellow crust of potassium superoxide, which can explode when cut or pressed against oil. Containers should be airtight and stored away from every other chemical group.

Putting out metal fires with a class D extinguisher

Never use water, foam or a carbon dioxide extinguisher on alkali metal fires. Water feeds the reaction with more hydrogen, and these metals keep burning in carbon dioxide. Use a Class D fire extinguisher, which smothers the fire with a dry powder such as sodium chloride or graphite, or cover a small fire with dry sand.

Hydrogen, francium and element 119 at the edges of group 1

Group 1 has three awkward elements at its edges. Hydrogen sits at the top because it also has one s electron, but it is a gas, forms covalent bonds, and can even gain an electron like a halogen. Only under the crushing pressures inside Jupiter and Saturn is hydrogen thought to become metallic.

Francium, at the bottom, is real but almost impossible to study in bulk. Its longest-lived isotope, francium-223, has a half-life of about 22 minutes, and its intense radioactivity would vaporise any visible sample with its own heat. Everything known about its chemistry comes from experiments on a few thousand atoms, which behave like a heavier version of caesium.

The next member would be element 119, provisionally named ununennium (Uue). Laboratories in Japan, Russia and the United States have tried to make it; none has succeeded so far.

Calculations suggest that relativistic effects would make ununennium less reactive than caesium or potassium, with a higher ionisation energy, so it may break the group trend. Whether ununennium counts as a true alkali metal is one of the open questions in superheavy-element chemistry, and any ununennium atoms made will survive for only microseconds.

Pseudo-alkali metals: ammonium, thallium and others

Some ions are not in group 1 but behave like alkali metal ions. Ammonium, NH₄⁺, has a similar size to K⁺ and forms salts much like potassium's.

Thallium in its +1 state forms compounds that resemble those of rubidium, and the cobaltocenium ion behaves like a large alkali metal cation. The coinage elements copper, silver and gold also have one s electron and a +1 state, but their filled d shell makes them far less reactive, which is why they are coinage metals rather than alkali metals.