Rubidium element (Rb): rubidium periodic table facts

Rubidium (Rb) basics: atomic number, atomic weight and periodic table place

The element rubidium occupies group 1, period 5 and the s-block of the periodic table, directly after the noble gas krypton and just before Sr. Its single outer electron is what drives almost all of rubidium's chemistry. The core identifiers you will see in any reference work are collected here. In radiology and healthcare imaging, contrast agents for magnetic resonance take centre stage in the gadolinium element for curious readers of medicine.

Those values make it the heaviest alkali element most chemists will ever handle in quantity, since the next one down, francium, is too short-lived to collect.

Physical properties of Rb

Freshly cut, rubidium metal is silvery-white and so ductile that it can be pressed like cold butter. It dulls within seconds once exposed, so most photographs show it sealed in an ampoule. When reading about Sharpless dihydroxylation, which earned a Nobel Prize and used the tetroxide, see element 76, osmium for more notes.

Melting point, boiling point and density

Rubidium melts at 39.3 °C, only a couple of degrees above body temperature, which is why it liquefies on a warm summer day or in the palm of your hand. It stays liquid across a wide range before boiling at 688 °C. With a value of 1.53 g/cm³ rubidium is the first alkali element denser than water, so a lump sinks rather than skating across the surface the way sodium does.

Hardness, crystal structure and thermal data

The solid packs into a body-centered cubic lattice and rates about 0.3 on the Mohs hardness scale, softer than talc. Its specific heat capacity is roughly 363 J/(kg·K) and its bulk modulus only 2.5 GPa, so it compresses easily. The vapor pressure climbs quickly once it melts, a trait that later sections show is useful in physics labs.

Chemical properties of the alkali metal

Rubidium is the second most electropositive of the stable alkali metals, beaten only by caesium. Its lone 5s electron sits far from the nucleus in the outer shell, so it gives that electron up with very little push, and that is the root of its fierce reactivity.

How Rb reacts with water, air and halogens

Drop a pea-sized piece of rubidium into a beaker and it reacts violently, forming rubidium hydroxide and hydrogen gas; the heat released usually ignites that gas. Left in the open, rubidium also ignites spontaneously in air. With oxygen it forms a family of oxides, from the monoxide to an orange superoxide when oxygen is plentiful, and it combines readily with the halogens to give white crystalline salts.

Oxidation states and electronegativity

In nearly every compound rubidium carries a +1 charge, existing as a monovalent ion. A −1 state is known only in unusual rubidide salts, so tables list its oxidation states as +1 and −1. Its electronegativity is 0.82 on the Pauling scale, which makes it a powerful reducing agent that hands electrons to almost anything it touches.

Rb electron configuration and ionization energies

Written in full, rubidium's arrangement is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s¹. The gap between its ionization energies tells the whole story of its chemistry:

• First ionization energy: about 403 kJ/mol, among the lowest of any element.
• Second ionization energy: about 2,633 kJ/mol, more than six times higher, because it means breaking into the stable krypton-like core.
• Electron affinity: about 46.9 kJ/mol.
• Atomic radius (empirical): 248 pm, with a covalent radius near 220 pm and a van der Waals radius of 303 pm.

That jump of more than 2,200 kJ/mol is why you never meet a doubly charged form of this element in ordinary chemistry.

Common Rb compounds and salts

Because the free element is so reactive, almost all practical work uses rubidium compounds instead. The most important are:

Rubidium chloride (RbCl)

The most widely used salt; molecular biologists use it to make bacterial cells take up DNA, and it serves as a tracer because living tissue normally holds very little of it.

Rubidium hydroxide (RbOH)

A strongly corrosive base and the usual starting point for making other salts.

Rubidium carbonate

Added to certain optical glasses to lower conductivity and improve stability.

Rubidium silver iodide

An ionic crystal with exceptionally high room-temperature conductivity, explored for thin-film batteries.

Halide salts

Rubidium fluoride, rubidium bromide and rubidium iodide, all soluble and colourless.

Rb isotopes and half-life

About 30 synthetic rubidium isotopes are known, but nearly all disappear within weeks. Only two exist in nature, and one of them makes every natural sample slightly radioactive, enough radioactivity to fog photographic film over a few months.

The two natural Rb isotopes

Because rubidium-87 lives far longer than the age of the universe, it is a primordial nuclide and the sample in your lab is only mildly radioactive. To put that slow clock in perspective: running the exponential decay law over Earth's 4.54 billion years shows that only about 6.2% of the heavier isotope present when the planet formed has turned into ⁸⁷Sr through beta decay since then.

Rubidium–strontium dating of rocks

That same slow clock makes rubidium–strontium dating one of geology's standard tools for dating rocks. As magma cools, fractional crystallization leaves the element enriched in the remaining melt, so different minerals start with different parent-to-daughter ratios. Measuring today's ⁸⁷Sr/⁸⁶Sr ratio alongside the rubidium content of several minerals from one rock gives its age, provided the rock has not been altered since it formed.

How Rb was discovered by flame spectroscopy

The history of rubidium begins in 1861, when the German chemists Robert Bunsen and Gustav Kirchhoff found it in Heidelberg using flame spectroscopy, the technique they had just invented.

The deep red emission spectrum behind the name

Working with the new spectroscope, Bunsen and Kirchhoff heated the mica lepidolite and saw two bright red lines in its emission spectrum that matched no known substance. They named the element from the Latin rubidus, meaning deepest red.

In an ordinary flame test the colour looks red-violet and is hard to tell apart from potassium by eye, which is exactly why the spectroscope mattered. Rubidium was only the second element, after caesium a year earlier, to be identified this way.

Isolating the free element

Getting metallic rubidium out of its salts proved harder. Electrolysis of molten chloride gave a blue mixture rather than shiny metal, so Bunsen turned to heating charred rubidium tartrate.

The product caught fire almost as soon as it formed, yet he still measured its properties remarkably close to modern values. Truly pure samples were not prepared until 1928.

Rb occurrence in lepidolite, pollucite and seawater

Rubidium is more common than you might guess: at roughly 90 parts per million of Earth's crust it outranks copper and zinc. It never forms a mineral of its own, though. Its large ion slips into potassium sites, so it is scattered through lepidolite, carnallite, leucite and zinnwaldite, rarely above a few percent.

Seawater and some mineral springs carry it in trace amounts. The richest concentrations sit in pegmatite ore bodies, where the element gathered in the last liquid to crystallize; pollucite deposits at Bernic Lake in Manitoba, Canada, are the best-known source.

Rb production and refining

World output is tiny, only a few tonnes of compounds a year, and there is no dedicated mine. The usual path to pure rubidium is:

1. Recover it as a by-product of caesium refining or of lithium refining.
2. Separate it from potassium and caesium by repeated crystallization or by the chlorostannate and ferrocyanide processes.
3. Finish with reduction of rubidium chloride by metallic calcium at about 750 °C under low pressure, then distil off elemental rubidium.

Because supply follows lithium and cesium demand, the industry has little control over how much rubidium reaches the market in a given year.

Rb uses and applications

Rubidium's applications lean heavily toward research and precision electronics. Smaller niches include photocells (a photocell turns light into current), special types of glass, fireworks with a red-purple tint, and atomic magnetometers.

Atomic clocks and GPS timing

The hyperfine transition of rubidium atoms ticks at a steady 6.8 GHz, which makes it a compact, low-cost frequency standard. Rubidium atomic clocks are less accurate than cesium fountains on their own, but when steered by GPS they deliver excellent timing at a fraction of the price, which is why telecommunications networks and cell-site transmitters rely on them.

Laser cooling and bose-Einstein condensate research

Rubidium vapor absorbs light at wavelengths that cheap diode lasers produce, so it became a workhorse for laser cooling. In 1995 a cloud of the heavier isotope's atoms chilled to billionths of a degree formed the first Bose–Einstein condensate, earning a share of the 2001 Nobel Prize in physics.

Rb-82 in medical imaging

Rubidium-82, made on site from a ⁸²Sr generator, is used in positron emission tomography. Heart muscle takes it up like potassium, so myocardial perfusion imaging can show which areas get poor blood flow, and it also helps locate brain tumors in nuclear medicine. It vanishes fast: our own calculation shows only about 6.5% remains five minutes after elution and roughly 0.4% after ten, so the dose must be produced beside the scanner.

Rb health effects and safe handling

Rubidium salts are only mildly toxic; your body treats the dissolved cation much like potassium ions and stores it in the intracellular fluid, with a biological half-life of roughly one to one and a half months. Researchers have even tested it in bipolar disorder and depression.

The bare element is a different matter: contact causes caustic skin and eye damage, and a burning sample causes thermal burns. Store it under dry mineral oil or sealed in glass under an inert atmosphere, keep it away from moisture, and use a class D extinguisher rather than water on a fire.

How Rb compares with other alkali metals

Lining element 37 up against its group 1 neighbours shows the trends you learn in chemistry class: going down the group, atoms get bigger, the outer electron is held more loosely, melting temperatures fall and reactivity climbs.

Element 37 is where the reaction with moisture shifts from lively to dangerous, and it is the first metal in the sequence that is denser than the liquid it reacts with.