What Are Alkali Metals On The Periodic Table?

What Are Alkali Metals?

The alkali metals consist of six elements positioned in Group 1, the far-left column of the periodic table. These include lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). Although hydrogen (H) also possesses a single electron in its outermost shell, it stands apart from this group; being a gas at room temperature rather than a metal, it is not considered an alkali metal.

This group earns its name from the distinctive way these metals react with water: the reaction produces substances known as alkalies. Alkalies are simply hydroxide compounds derived from alkali metals, such as sodium hydroxide and potassium hydroxide.

These compounds are notably strong bases and quite caustic in nature. A familiar example is lye, which is sodium hydroxide. When alkalies encounter acids, the resulting chemical reaction forms salts.

In their pure state, alkali metals appear as soft, lustrous metals with relatively low melting points. What stands out about them is their remarkable reactivity, especially with air and moisture. For this reason, handling and storage require particular care.

Chemically, alkali metals are quick to lose their outermost electron, forming cations with a single positive charge (+1). Their softness is such that one can cut them with a knife, revealing a gleaming surface that quickly tarnishes when exposed to air, primarily due to oxidation from moisture and oxygen, and in the case of lithium, even nitrogen.

Given how easily these metals react, they are typically stored under oil to keep them away from air. They do not occur freely in nature, but instead are always found as part of salts. Among the group, cesium is recognized as the most reactive metal.

It’s worth noting that all alkali metals will react with water, but the intensity increases as you move down the group, with heavier members reacting even more vigorously than their lighter counterparts.

Alkali Metals on Periodic Table

The alkali metals—namely lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr)—form a distinctive group within the periodic table’s s-block. What sets these elements apart is that each one has a single electron occupying its outermost s-orbital.

If you’re looking for a clear illustration of how elemental properties can shift within a group, the alkali metals are a prime example. They consistently display similar chemical and physical characteristics, and as you trace the group from top to bottom, some striking patterns start to emerge:

• Increasing atomic radius,
• Decreasing electronegativity
• Increasing reactivity
• Decreasing melting and boiling points

In general, their densities increase when moving down the table, with the exception of potassium, which is less dense than sodium.

Properties Of Alkali Metals

Alkali metals are a group of chemical elements in the periodic table with the following physical and chemical properties:

• Shiny
• Soft
• Silvery
• Highly reactive at standard temperature and pressure
• Readily lose their outermost electron to form cations with a charge of +1

The alkali metals stand out for their notably low melting points. For instance, lithium melts at 180.5 °C (356.9 °F), while cesium transitions to a liquid state at just 28.4 °C (83.1 °F). In addition to this, these metals are recognized for their excellent ability to conduct both heat and electricity.

However, their high reactivity means they are rarely found in pure form in nature; instead, they tend to form compounds with other elements—common examples include sodium chloride (NaCl), better known as table salt, and potassium chloride (KCl).

A distinctive trait of alkali metals is their softness; in fact, you can cut them with something as simple as a plastic knife. Yet, once exposed to air, their shiny surfaces don’t last long—they oxidize quickly, leading to rapid tarnishing.

This pronounced reactivity also means special precautions are needed: to avoid unwanted reactions with air, alkali metals are typically kept under oil.

According to the modern IUPAC classification, the alkali metals make up group 1 on the periodic table, with the exception of hydrogen. One striking behavior they share is their reaction to water. Notably, as you move down the group, the metals become increasingly reactive with water, with the heavier members reacting more vigorously than their lighter counterparts.

Properties of the alkali metals:

Properties	Lithium	Sodium	Potassium	Rubidium	Cesium	Francium
Atomic number	3	11	19	37	55	87
Atomic weight (or stablest isotope)	6.941	22.99	38.098	86.468	132.905	223
Color of element	Silver	Silver	Silver	Silver	Silver	—
Melting point (°c)	180.5	97.72	63.38	39.31	28.44	27
Boiling point (°c)	1,342	883	759	688	671	677
Density at 20 °c (grams per cubic centimeter)	0.534	0.971	0.862	1.532	1.873	—
Volume increase on melting (percent)	1.51	2.63	2.81	2.54	2.66	—
Valence	1	1	1	1	1	1
Mass number of most common isotopes (terrestrial abundance, percent)	6 (7.59), 7 (92.41)	23 (100)	39 (93.2581), 40 (0.0117), 41 (6.7302)	85 (72.17), 87 (27.83)	133 (100)	—
Color imparted to the flame	Red	Yellow	Violet	Yellow Violet	Blue	—
Main spectral emission lines (wavelength, angstroms)	6,708; 6,104	5,890; 5,896	7,699; 7,665	4,216; 4,202	4,593; 4,555	—
The heat of fusion (calories per mole/kilojoules per mole)	720 (3)	621 (2.6)	557 (2.33)	523 (2.19)	500 (2.09)	500 (2)
Specific heat (joules per gram kelvin)	3.582	1.228	0.757	0.363	0.242	—
Electrical resistivity at 293–298 k (microhm centimeters)	9.5	4.9	7.5	13.3	21	—
Magnetic susceptibility (cgs units)	14.2 (10−6)	16 (10−6)	20.8 (10−6)	17 (10−6)	29 (10−6)	—
Crystal structure	Body-Centered Cubic	Body-Centered Cubic	Body-Centered Cubic	Body-Centered Cubic	Body-Centered Cubic	—
Radius: atomic (angstroms)	1.67	1.9	2.43	2.65	2.98	—
Radius: ionic (+1 ion, angstroms)	0.9	1.16	1.52	1.66	1.81	1.94
Radius: metallic (angstroms, 12-coordinate)	1.57	1.91	2.35	2.5	2.72	2.8
First ionization energy (kilojoules per mole)	520.2	495.8	418.8	403	375.7	380
Oxidation potential for oxidation from the 0 to +1 oxidation state at 25 °c (volts)	3.04	2.71	2.93	2.92	2.92	2.92
Electronegativity (Pauling)	0.98	0.93	0.82	0.82	0.79	0.7

List of Alkali Metals

The alkali metals are:

• Lithium (Li)
• Sodium (Na)
• Potassium (K)
• Rubidium (Rb)
• Cesium (Cs)
• Francium (Fr)

According to the International Union of Pure and Applied Chemistry (IUPAC), hydrogen (H) is not classified as an alkali metal. The main reason is that, under normal conditions of temperature and pressure, hydrogen exists as a gas rather than a solid metal.

That said, hydrogen does share a number of characteristics with the elements found in the alkali metal group. Interestingly, if subjected to extremely high pressures, hydrogen actually behaves much like an alkali metal.

Examples Of Alkali Metals

The alkali metals are six chemical elements in Group 1, the leftmost column in the periodic table. They are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr).

Lithium

Lithium stands out as the lightest metal identified so far and is unique among metals for its ability to react with nitrogen. Its oxide form, lithium oxide, displays amphoteric behavior—meaning it can act as both an acid and a base.

Interestingly, while the other alkali metals tend to form predominantly ionic compounds, lithium sets itself apart in this regard as well. Its relatively high charge density, compared to its alkali counterparts, leads to a much greater degree of hydration.

Tracing lithium’s history takes us back to 1817, when Johan Arfvedson first discovered it while analyzing the mineral petalite (LiAlSi₄O₁₀). The name “lithium” itself has roots in the Greek word ‘lithos,’ which translates to “stone.”

Today, lithium can be found in a range of sources—pegmatitic minerals, certain clays, brine deposits, and even in the vast waters of our oceans. Remarkably, it is also present in all living organisms.

Uses of Lithium

• Lithium is used in heat-resistive ceramics and glasses.
• An alloy of Lithium is used in aircraft building.
• Lithium Deuteride is used as a fusion fuel in thermonuclear weapons.
• Lithium batteries are packed with a lot of energy as compared to other metals. Revolutionized devices like cell phones, and computers use lithium batteries.
• Lithium salts are used as a mood-stabilizing drug.
• Lithium-6 is the main source of tritium production.
• Lithium is used in deoxidizing copper and copper alloys.
• Lithium compounds are used as pyrotechnic colorants in fireworks that produce red luminance.
• Lubricating greases are produced from Lithium.

Sodium

Sodium is a chemical element most of us encounter daily, usually as sodium chloride, or table salt, in our meals. Industrially, sodium is obtained by electrolyzing sodium chloride, a process that highlights just how reactive this element is.

Historically, the isolation of sodium dates back to 1806, when the chemist Sir Humphry Davy managed to extract it by passing an electric current through molten sodium hydroxide. Interestingly, sodium’s origins stretch far beyond our laboratories.

In the universe, this element is formed inside stars. It comes about either when two carbon atoms fuse together through nuclear fusion or, in some cases, when a neon atom picks up a proton.

Uses of Sodium

• Sodium is used as luster in metals.
• Liquid Sodium is used as a coolant in nuclear reactors.
• The sodium salt of fatty acids is used in soap.
• NaK, an alloy of sodium and potassium, is an important heat transfer agent.
• Sodium compounds are used in paper, textile, petroleum, and chemical industries.
• Sodium Iodide is used to treat extensive ringworm.
• Sodium is used in street lights and sodium vapor lamps as it can give a yellow glow with bright luminance.
• Sodium hydroxide is used as an oven cleaner.

Potassium

Potassium, which holds the nineteenth position on the periodic table, is fundamental to human health. As an essential mineral, it must be carefully regulated within the body, since both excess and deficiency, known as hyperkalemia and hypokalemia, respectively, can have serious consequences.

When exposed to air, potassium quickly takes on a grayish appearance due to oxidation. For this reason, it is typically stored under petroleum, which protects it from moisture and further chemical reactions.

Uses of Potassium

• Potassium chloride is essential for the growth of plants. It is used in fertilizers.
• Potash improves water retention, yield, nutrient value, taste, color, texture, and disease resistance of food crops.
• Potassium chlorate and potassium nitrate are used in explosives and fireworks.
• Potassium nitrate is used as a food preservative.
• Potassium maintains blood pressure and acidity levels in our bodies.
• Potassium chromate is used in the tanning of leather and in the manufacture of inks, gun powder, dyes, safety matches, etc.,
• Potassium is essential for normal cell respiration and electrolyte function as 95% of our cells are made of potassium.
• Potassium hydroxide is used to make detergents.
• Potassium helps to pump fluids inside the heart and the nerves.

Rubidium

Rubidium is a radioactive element. It is derived from the Latin word rubius meaning deepest red.

Uses of Rubidium

• Rubidium 82 is used in myocardial perfusion.
• Rubidium is used in the manufacture of atomic clocks, electronic tubes, and photocells.
• Rubidium is used as a working fluid in vapor turbines.
• It is used as a component in the engines of space vehicles.
• Rubidium vapor is used in laser cooling.
• Rubidium chloride is used to induce cells to take up DNA.
• It is used in thermoelectric generators.
• Rubidium Carbonate is used in making optical glasses.
• Due to the hyperfine structure of rubidium’s energy levels, it is used in atomic clocks.
• A compound made up of rubidium, silver and iodine has certain electrical characteristics and is used in making thin-film batteries.

Cesium

Cesium stands out as a notably active metal, recognized for being the most electropositive element among all known metals. This remarkable characteristic means cesium readily forms compounds by bonding with various anions.

However, cesium is highly toxic and requires careful handling. Among its compounds, cesium hydroxide deserves mention as it represents the strongest base identified so far.

When considering isotopes, cesium offers quite a range, but cesium-133 is especially significant. This stable isotope serves as the foundation for precise timekeeping in devices such as atomic clocks, which have become essential standards in measuring time. Interestingly, cesium also remains in a liquid state at or just above room temperature, setting it apart from many other metals.

Uses of Cesium

• Cesium-134 is used in the nuclear power industry.
• Used in photoelectric cells due to their quick electron emission.
• Cesium is used as a catalyst for the hydrogenation of certain organic compounds.
• It is used in propulsion systems.
• It removes air traces from vacuum tubes.
• Cesium is used in photovoltaic cells, television image devices, and night-vision equipment.
• Cesium vapor is used in the magnetometer.
• Cesium-137 is used in brachytherapy to treat cancers. (Brachytherapy is a cancer treatment method using radioactive elements)
• Cesium chloride solution is used in molecular biology for density gradient ultracentrifugation, primarily for the isolation of viral particles, subcellular organelles and fractions, and nucleic acids from biological samples.
• Cesium is used as a standard in spectrophotometry
• It is used in military aircraft.

Francium

Francium stands out as the element with the lowest electronegativity among all those currently known. Not only is it highly radioactive, but it also holds the distinction of being the heaviest metal in its group. In terms of production, francium is typically obtained by bombarding thorium with protons or, alternatively, by exposing radium to neutrons.

This element is extremely scarce, and its practical applications are quite limited. For the most part, francium finds its use within laboratory settings, mainly for scientific investigation. Its utility is further restricted by its rapid decay, as francium has a notably short half-life.

When it comes to alkali metals in general, their significance in various fields is well established. However, it’s important to emphasize that these substances should only be handled under professional supervision. Their reactions can become dangerously vigorous, and their toxic nature means that proper precautions are essential.

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