GROUP 1A ELEMENTS [ALKALI METALS]

Contents

• Introduction
• Group 1 elements
• Occurrence of Alkali metals
• Physical properties of alkali metals
• Chemical properties of alkali metals
• Lithium
• Sodium
• Potassium
• Rubidium
• Caesium
• Francium
• Summary
• Recommended Videos
• Test Questions
• Discuss And Explain

Learning Objectives

By the end of this section, you should be able to:

• Describe the alkali metals.
• Describe the physical and chemical properties of alkali metals.

Introduction

I believe that this is not your first time to hear the word 'Periodic Table'. The Periodic Table contains all known elements; and these elements are arranged in columns and rows - technically known as groups and periods.

Elements in the same group share a common characteristic - presence of the same number of electrons in their outermost or valence shell. On the other hand, elements in the same period have only the same number of electron shells. Right now we're going to be discussing the elements in Group 1 of the periodic table.

Periodic Table of Elements

Group 1 Elements

Group 1 Elements are elements that have only one electron in their valence shells. They're very ready to lose their valence electron to achieve a stable state, hence rapidly forming compounds. This makes us to say that group 1 Elements are highly reactive, and infact, it is difficult (or nearly impossible) to find them freely in nature.

Elements in Group 1 of the periodic table include the non-metal: hydrogen, and six metals that are called alkali metals.

Hydrogen has the smallest and lightest atoms of all elements. Pure hydrogen is a colorless, odorless, tasteless gas that is non-toxic but highly flammable (can catch fire). Hydrogen gas exists mainly as diatomic (“two-atom”) molecules (H₂). Hydrogen is the most abundant element in the universe and the third most abundant element on Earth, occurring mainly in compounds such as water etc

The Alkali metals include: Lithium, Sodium, Potassium, Rubidium, Caesium and Francium.

The word 'Alkali' comes from an Arabic word "al-qalyah" which is translated as "plant ashes" which contain many basic compounds. Hence, alkali metals are so - called because of a character; they form strongly alkaline or basic compounds when reacted with water.

Occurrence of Alkali metals

Like I have said earlier, alkali metals are very reactive and they do not occur in free State in nature. They happen to react with non-metals that occur in nature, hence they exist in combined form.

This means that the pure form of alkali metals are extracted from their minerals or ores (we discussed extraction of metals in the previous topic).

Pure form of alkali metals are stored under inert conditions like in kerosene or oil or vacuum; to avoid reaction with moisture or gases in air.

Sodium and potassium ores are found in large amounts in nature while Lithium, Rubidium and Caesium are discovered in relatively low amounts.

Francium on the other hand is radioactive, and very hard to discover in nature. Once you see the Francium element, enjoy its presence, because in the next 21 minutes, or maximum of 22 minutes, it would disintegrate.

Physical Properties of Alkali Metals

All Alkali metals share some common physical properties:

• Alkali metals are solids at room temperature.

• Alkali metals have low density (or light, low mass per volume) and some of them even float on water (as they react with it).

• Alkali metals are relatively soft. Some are soft enough to be cut with a knife. When freshly cut they are bright and shiny, but they soon take on a dull look because of their reaction with air - do not try this for Caesium or Rubidium because they will explode in your face.

Alkali metals are very soft and can be easily cut with knife

• All Alkali metals have low melting and boiling temperature. This is due to the weak bonds in the atoms of alkali metals; and the farther electrons get away from the nucleus (of an atom), the faster the atom (or element) melts and boils. Hence, melting and boiling points of alkali metals decrease down the group.

• When alkali metals are heated (in hot Bunsen flame) they produce characteristic flame colours, which is usually used in flame tests and spectroscopy (or qualitative analysis) to identify alkali metals. Some people usually love to know why metals give off coloured flames, a simple explanation is because of the movement of electron(s) that has/have been excited to a higher energy level, back into a lower energy level. In other words, due to low Ionization Energies (=energy required to excite or remove an electron) of group 1 metals, electrons are easily excited to higher energy levels. When they drop back to the lower energy states, they emit visible light of characteristic colours. E.g Yellow colour of sodium flame is produced by the transition of an electron that has been excited into 3p orbital, back into 3s orbital.

• Alkali metals are good conductors of heat and electricity.

• All Alkali metals have a body centered cubic crystal structure.

• All Alkali metals have an ionic radius that is smaller than their atomic radius. Let me explain this well, I will be using sodium as an example; take a look at the image below:

Atomic structure of sodium, Na

Ionic structure of sodium, Na+

When sodium was still in its ground state, the outer valence electron makes sodium's atomic radius to be wider; but once sodium loses its valence electron to become an ion, the outer shell obliterates, the protons (in the nucleus) are now more than the electrons and they pull the available electrons closer to the nucleus, hence, sodium becomes a shell less and has a smaller (ionic) radius.

I believe you got that!

Chemical properties of Alkali metals

• Alkali metals are the most electropositive (=tendency to give out or lose electron).

• Alkali metals react with air at room temperature. Lithium reacts the least with air; it slowly tarnishes in damp air. Sodium and potassium tarnish rapidly. Rubidium and Caesium catch fire on exposure to air. Therefore Lithium, Sodium and potassium are usually stored in oil or kerosene while the rest are kept in a vacuum.

• Alkali metals react with all non-metals except the noble gases and nitrogen (which can only react with lithium).

• Alkali metals react with acids to yield metallic salt and hydrogen gas.

• All alkali metals dissolve in pure liquid ammonia to give solutions that are blue if dilute and bronze if concentrated.$$2{\text{M}}_{(s)}+2{\text{NH}}_{3(l)}\to 2{\text{MNH}}_{2(aq)}+{\text{H}}_{2(g)}$$ Where M=Li, Na, K, Rb, Cs

Lithium

Discovery

Lithium was discovered in 1817 by Johan Arfwedson August who identified it as a constituent of a naturally occuring mineral; petalite. At that time, Lithium could not be studied properly, until some time later when it was finally isolated by Sir Humphry Davy and W.T. Brande.

Occurrence

Lithium occurs in nature as:

• Petalite LiAlSi₄O₁₀
• Lepidolite K(Li Al)₃(Al, Si, Rb)₄ O₁₀(F,OH)₂
• Spodumene LiAlSi₂O₆
• And other subsurface brines.

Lithium ores account for 0.002% of the earth's crust.

Extraction

Lithium is extracted from its ores (majorly spodumene or petalite) by electrolysis, which is pre-initiated by treatment of the ore. Some modern extraction processes involve the use of hydrometallurgy + electrolysis.

Physical Properties

• Lithium is a silvery white solid at room temperature.

Lithium

• Lithium is the lightest known metal. Its density is so small that it would float on oil that is floating on water. Density of lithium is 0.534 g/cm³.

• Lithium is a soft and flammable metal.

• Although Lithium is soft, still it is the hardest alkali metal. Its melting and boiling point is the highest of all Alkali metals. Melting point temperature of Lithium is 180.5°C and its boiling point temperature is 1342°C.

• Lithium gives off a crimson coloured flame when heated in a hot Bunsen burner.

Crimson flame produced by heated lithium in Bunsen burner

Chemical Properties

• Lithium has an atomic number of 3, hence it has 3 electrons and 3 protons in ground state.

• The atomic mass of Lithium is 6.941. A Lithium atom will therefore have; 6.941 – 3 ≈ 4 neutrons.

• Lithium is the least reactive alkali metal, it is however the strongest reducing agent; its reducing power is drastically reduced when solvation (or dissolution) can't take place.

• Lithium can lose electron to become Li⁺ and it can also gain electron to become Li^- , but tendency of losing electron is higher.

• Electronic Configuration of Lithium is 1s² 2s¹ or [He]2s¹

Reaction Of Lithium

• On exposure to air, Lithium reacts with the atmospheric water vapour and oxygen which after sometime turns into a black tarnish - Lithium nitride.

Black tarnish of Lithium Nitride

• Lithium burns in air to form Lithium monoxide (or Lithium oxide) Li₂O. Other alkali metals form peroxides or superoxides you will find them out later in this post. This means that Lithium is the only alkali metal that forms a normal oxide when reacted with oxygen.

Lithium Oxide Li₂O

• Li₂O reacts less violently with water to give Lithium hydroxide.

$${\text{Li}}_2\text{O}+{\text{H}}_2\text{O}\to 2\text{LiOH}$$

PRO TIP: Lithium hydroxide is a weak base while hydroxides of other alkali metals behave as strong bases.

• Lithium reacts with hydrogen to yield Lithium hydride.

$$2\text{Li}+{\text{H}}_2\to 2\text{LiH}$$

PRO TIP: LiH is the most stable hydride formed by the alkali metals.

• Lithium hydride on reaction with AlCl₃ in ether solution forms lithium aluminium hydride which is a useful reducing agent in organic chemistry.

$$4\text{LiH}+{\text{AlCl}}_3\to {\text{LiAlH}}_4+3\text{LiCl}$$

• Pure Lithium reacts less violently with water to give Lithium hydroxide and hydrogen gas.

Lithium reacts less violently with water

• Lithium reacts with carbon to form ionic Lithium carbide, Li₂C₂, whereas similar carbides of other alkali metals are not formed by direct reaction with carbon.

• Lithium is the only alkali metal which reacts with nitrogen gas to form Lithium nitride. The other alkali metals do not combine with nitrogen at all.

$$6{\text{Li}}_{(s)}+{\text{N}}_{2(g)}\to 2{\text{Li}}_3{\text{N}}_{(s)}$$

PRO TIP: Reaction between Lithium and nitrogen is slow at room temperature, but faster at higher temperatures.

• Lithium nitride can react directly with water to form ammonia.

$${\text{Li}}_3\text{N}(s)+3{\text{H}}_2\text{O}(l)\to 3\text{LiOH}(aq)+{\text{NH}}_3(g)$$

• Lithium also reacts with sulphur to form Lithum sulphide.

$$2\text{Li}+\text{S}\to {\text{Li}}_2\text{S}$$

• Lithium reacts with Halogens to yield Lithium Halide.

$$\text{Li}+\text{F}\to \text{LiF}$$

• Although Lithium shows a strong resemblance to other alkali metals, still it has many similarities to Magnesium. This is due to its ionic radius [Li⁺(0.68Å)] which is closer to Mg²⁺(0.65Å) than to Na⁺(0.97Å). So Lithium is more similar to Magnesium than to Sodium in many different ways. This phenomenon is called diagonal relationship.

PRO TIP: In general, the elements of the 2nd row or 2nd period of the periodic table show differences from the other elements in their respective groups, the first three members of the 2nd period exhibit many similarities to those elements located diagonally below them in the periodic table.

Diagonal relationship in the periodic table

• Lithium compounds exhibit higher covalent character than ionic, which is responsible for their solubility in organic solvents (like petrol, methyl alcohol, etc).

• Lithium carbonate, lithium fluoride and lithium phosphate are sparingly soluble in water whereas the corresponding salts of other alkali metals are soluble in water.

• Lithium hydrogen carbonate, LiHCO₃, is not obtained in a solid form, whereas other alkali metals form solid hydrogen carbonates.

• Lithium chloride (LiCl) is deliquescent (absorbs moisture from the atmosphere and dissolves in it). Other alkali metal chlorides do not form hydrates.

• Lithium nitrate, LiNO₃, when heated, gives lithium oxide, Li₂O. The other alkali metal nitrates decompose to give their corresponding nitrites and oxygen.

$$4{\text{LiNO}}_3\to 2{\text{Li}}_2\text{O}+4{\text{NO}}_2+{\text{O}}_2$$

$$2{\text{NaNO}}_3\to 2{\text{NaNO}}_2+{\text{O}}_2$$

Uses of Lithium

• Lithium is combined with other metals to form alloys that are used for different purposes. Li-Pb alloy is used for making hard bearings, Li-Al alloy is used for aircraft construction, Li-Mg alloy is used for armour plate and aerospace components.

• Lithium is used for producing thermonuclear energy which is required to propel rockets and controlled missiles.

• LiBr is used in medicine as a sedative while LiCl is used in air conditioners to regulate humidity.

• Lithium is used in the manufacturing of dry cell batteries like mobile phone batteries, Electric vehicle batteries etc 

• Compounds of lithium are used to make certain kinds of glass and porcelain products.

Sodium

Discovery

Sodium, often called soda, has been noticed in compounds many years ago until the year 1807 when Sir Humphry Davy was able to successfully isolate Sodium by the electrolysis of sodium hydroxide.

Occurrence

Sodium occurs in nature as:

• Sodium carbonate (soda ash), Na₂CO₃
• Sodium nitrate (Chile salt Peter), NaNO₃
• Sodium chloride (rock salt / halite), NaCl

Sodium ores account for about 2.36% of the earth's crust.

Extraction

Sodium is extracted from its ores (majorly Sodium chloride) by electrolysis, using a special type of electrolytic cell called Down's Cell.

Physical Properties

• Sodium is a silvery-white solid at room temperature.

Sodium

• Sodium is a light metal, its density is only slightly less than that of water, therefore it floats on water. Density of sodium is 0.97g/cm³.

• Sodium is a soft metal and can be cut with a sharp knife.

• The melting and boiling point of sodium is less than that of Lithium. The Melting point temperature of Sodium is 97.7°C and its boiling point temperature is approximately 883°C.

• Sodium gives off a golden-yellow coloured flame when heated in a hot Bunsen burner.

Golden-yellow flame produced by sodium heated in Bunsen burner

Chemical Properties

• Sodium has an atomic number of 11, hence it has 11 electrons and 11 protons in ground state.

• The atomic mass of Sodium is 22.989. Sodium will therefore have; 22.989 - 11 ≈ 12 neutrons.

• Sodium is highly reactive, more than Lithium but less than other alkali metals below it in the group.

• Sodium readily loses its valence electron to become Na⁺ in compounds.

• Electronic Configuration of sodium is 1s² 2s² 2p⁶ 3s¹ or [Ne] 3s¹

Reaction Of Sodium

• On exposure to air, sodium reacts with atmospheric oxygen to form sodium oxide, ${\text{Na}}_2\text{O}$ which quickly reacts with water vapour to give grayish/whitish sodium hydroxide film.

• Sodium burns in oxygen (at higher temperatures) to form sodium peroxide Na₂O₂.

Sodium peroxide

• Sodium peroxide, Na₂O₂ reacts violently with water to form sodium hydroxide.

$${\text{Na}}_2{\text{O}}_2(s)+2{\text{H}}_2\text{O}(l)\to 2\text{NaOH}(aq)+2{\text{H}}_2{\text{O}}_2(g)$$

• Pure sodium also reacts violently with water to form sodium hydroxide and hydrogen gas.

$$2\text{Na}(s)+2{\text{H}}_2\text{O}(l)\to 2\text{NaOH}(aq)+{\text{H}}_2(g)$$

A pound of Sodium explodes in water | courtesy: Gizmodo Australia

PRO TIP: The basic strength of Sodium hydroxide (NaOH) is more than Lithium hydroxide (LiOH) but less than potassium hydroxide (KOH). In general, basic strength of alkali metal hydroxides increases down the group.

• Sodium reacts with hydrogen at high temperatures to yield sodium hydride.

$$2\text{Na}(l)+{\text{H}}_2(g)\to 2{\text{Na}}^{+}{\text{H}}^{-}(l)$$

• Sodium hydride will react with water to evolve hydrogen gas.

$$\text{NaH}(l)+{\text{H}}_2\text{O}(l)\to \text{NaOH}(aq)+{\text{H}}_2(g)$$

• Sodium does not react directly with carbon, sodium carbide can only be formed by heating sodium together with ethyne.

• Sodium does not react with nitrogen at all.

• Sodium reacts with sulphur to form sodium sulphide. Sodium polysulphides can also be formed.

• Sodium reacts with halogens to yield sodium halides. 

• Sodium is reacted together with mercury to form sodium amalgam.

Sodium hydroxide

Uses of Sodium

• Bright yellow lights on streets are sodium vapour lamps.

• Sodium compounds are used for wide range of industrial purposes like production of glass, paper, soap, and textiles.

• Sodium is an important element of biological life.

• Sodium chloride is used as a preservative and for several other uses, sodium hydrogen trioxocarbonate(IV) (baking soda) is used in baking as a raising agent.

• Sodium is also used for production of tetraethyl lead and as a reductant during extraction of titanium metal.

• Sodium is used in the manufacture of sodium cyanide which is used in extraction of gold.

• Sodium is used in nuclear reactions to absorb some of the heat produced during the reactions.

Potassium

Discovery

Potassium was first isolated in 1807 by Sir Humphry Davy during electrolysis of molten KOH. Infact, potassium was the first metal to be isolated by electrolysis.

Occurrence

Potassium occurs in nature as:

• Potash
• Orthoclase
• Nitre (salt peter) KNO₃
• Carnalite KCl·MgCl₂·6H₂O
• Langbeinite
• Sylvite
• Polyhalite

Potassium ores account for about 2.09% of the earth's crust.

Extraction

Potassium is extracted from its hydroxides by electrolysis. It is extracted from its chloride by thermal reduction with sodium. It can also be extracted from its fluoride by griesheimer process which involves reacting potassium fluoride with calcium carbide to produce potassium.

Thermal reduction method:

$$\text{Na}+\text{KCl}\to \text{NaCl}+\text{K}$$

Griesheimer process:

$$2\text{KF}+{\text{CaC}}_2\to 2\text{K}+{\text{CaF}}_2+2\text{C}$$

Physical Properties

• Potassium is a silvery solid at room temperature.

Potassium

• Potassium density is less than sodium, potassium is the second least dense metal to exist, after Lithium. Its density is approximately 0.89 g/cm³.

• Potassium is a soft solid and can be easily cut with a knife.

• The melting and boiling point of potassium is less than that of sodium. The Melting point temperature of potassium is 63.5°C and its boiling point temperature is 759°C.

• Potassium gives off a lilac coloured flame when heated in a hot Bunsen burner.

Potassium produces a lilac flame when heated in a hot Bunsen burner

Chemical Properties

• Potassium has an atomic number of 19, hence it has 19 electrons and 19 protons in its ground state.

• The atomic mass of Potassium is 39. Potassium will therefore have; 39 - 19 = 20 neutrons.

• Potassium is highly reactive than sodium.

Potassium loses its valence electron to become K⁺ ion in compounds.

• Electronic configuration of Potassium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ or [Ar] 4s¹

Reaction Of Potassium

• On exposure to air, potassium quickly reacts with atmospheric oxygen and water vapour to give a dull finish which is flammable. You pressurise or scratch it, it catches fires.

• Potassium burns in oxygen to form potassium superoxide KO₂.

• Potassium superoxide, KO₂ reacts very violently with water to form Potassium hydroxide.

$$2{\text{KO}}_2+2{\text{H}}_2\text{O}\to 2\text{KOH}+{\text{H}}_2{\text{O}}_2+{\text{O}}_2$$

• Pure potassium explodes in water to form potassium hydroxide and hydrogen gas.

$$2\text{K}+2{\text{H}}_2\text{O}\to 2\text{KOH}+{\text{H}}_2$$

Bulk Potassium explodes in water producing a lilac flame

• Potassium reacts with hydrogen to yield ionic potassium hydride which will form Potassium hydroxide and flammable hydrogen when put in water.

• Potassium reacts with halogens to yield potassium halides, higher halogens can form polyhalides e.g KI₃

• Potassium is inert to nitrogen.

Uses Of Potassium

• Potassium is an important component in plants nutrition and its compounds is used in the production of fertilisers for agriculture.

• Potassium compound is used in medical facet to treat a type of kidney stone disease.

• Potassium is used in textiles, baking and leather industry.

• Potassium is also used to saponify fat and oils to produce potash soap.

• Potassium compound is used as an oxidant in gun powder.

• A type of potassium compound is used as an artist's pigment.

• Potassium superoxide acts as a portable source of oxygen and absorber of carbon(IV)oxide. It is widely used in the respiration systems in mines, submarines and spacecrafts.

$$4{\text{KO}}_2+2{\text{CO}}_2\to 2{\text{K}}_2{\text{CO}}_3+3{\text{O}}_2$$

Rubidium

Discovery

Rubidium was discovered in 1861 by Robert Bunsen and Gustav Kirchhoff through flame spectroscopy of lepidolite. Because of the bright red lines in its emission spectrum, they chose a name derived from the Latin word rubidus, meaning "deep red".

Occurrence

Rubidium occurs in nature as:

• Lepidolite
• Pollucite
• Leucite
• Zinnwaldite
• Carnalite

Rubidium ores account for about 0.009% of the earth's crust.

Extraction

Initial extraction method for Rubidium was electrolysis. Rubidium is not so widely used, so Rubidium is extracted from Rubidium-Caesium alum by fractional crystallization - remember electrolysis is expensive. Two other methods are reported, the chlorostannate process and the ferrocyanide process - hydrometallurgy.

Physical Properties

• Rubidium is a silvery-white solid at room temperature.

Rubidium

• Rubidium is the first alkali metal that has a density that is more than water. This makes Rubidium sinks in water, unlike the other alkali metals above it in the group. The density of Rubidium is 1.53 g/cm³.

• Rubidium is a very soft and ductile solid.

• The melting and boiling point of Rubidium is less than that of potassium. The Melting point temperature of rubidium is 39.3°C and its boiling point temperature is 688°C.

• Rubidium gives off a red coloured flame when heated in a hot Bunsen burner.

Rubidium gives off a red coloured flame when heated in a hot Bunsen burner | photo credit: Wikimedia commons

Chemical Properties

• Rubidium has an atomic number of 37, hence it has 37 electrons and 37 protons in its ground state.

• The atomic mass of Rubidium is 85.5 . Rubidium will therefore have; 85.5 - 37 ≈ 49 neutrons.

• Rubidium is highly reactive than Potassium.

• Rubidium loses its valence electron to become Rb⁺ ion in compounds.

• Electronic configuration of Rubidium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s¹ or [Kr] 5s¹

Reaction of Rubidium

• On exposure to air, Rubidium will ignite spontaneously and catch fire.

• Rubidium burns in oxygen to form Rubidium Superoxide RbO₂.

• Rubidium superoxide, Rubidium peroxide and Rubidium oxide react very violently with water to form Rubidium hydroxide.

$$2{\text{RbO}}_2+2{\text{H}}_2\text{O}\to 2\text{RbOH}+{\text{H}}_2{\text{O}}_2+{\text{O}}_2$$

• Pure Rubidium also explodes massively in water to form Rubidium hydroxide and hydrogen gas.

$$2\text{Rb}+2{\text{H}}_2\text{O}\to 2\text{RbOH}+{\text{H}}_2$$

Rubidium explosion in water | photo credit: gifsforum.com

Rubidium does not react directly with Nitrogen nor Carbon.

Uses Of Rubidium

• Rubidium is sometimes used in colourful fireworks to give a purple colour.

• Rubidium is used in thermoelectric generator to generate electricity.

• Rubidium is used in the resonant element present in atomic clocks.

• Rubidium is used as a getter in vacuum tubes- to remove traces of gases.

• Rubidium is used as a component of photoelectric cells.

• Rubidium is used during production of some special types of glass.

• Radioactive Rubidium is used in the treatment of cancers.

Caesium

Discovery

Caesium was discovered in 1860 by German scientists; Robert Bunsen and Gustav Kirchhoff; by flame spectroscopy of mineral water from a particular part of Germany- It was the first element to be discovered by a spectroscope, and its name was given because of blue lines in its emission spectrum- caesius means sky-blue. It was not until 1882 before it was successfully isolated by Carl Setterberg.

Occurrence

Caesium occurs in nature as:

• Pollucite Cs(AlSi₂O₆)
• Avogadrite ((K,Cs)BF₄)
• Pezzottaite (Cs(Be₂Li)Al₂Si₆O₁₈)
• Londonite ((Cs,K)Al₄Be₄(B,Be)₁₂O₂₈)

Caesium ores account for about 0.0001% of the earth's crust.

Extraction

Pure Caesium can be obtained by electrolysis / chemical dissolution + reduction / thermal decomposition.

Physical Properties

• Caesium is a silvery-golden solid at room temperature, however it quickly melts into liquid form in a warm room because of its very low melting point.

Caesium

• Caesium density is 1.93 g/cm³ when solid and approximately 1.84 g/cm³ when liquid.

• Caesium is known as the softest solid element at room temperature.

• The Melting point temperature of caesium is 28.5°C and its boiling point temperature is 671°C - the second known metal with the lowest melting point, after mercury.

• Caesium gives off a blue coloured flame when heated in a hot Bunsen burner.

Caesium gives off a blue coloured flame when heated in a hot Bunsen burner

Chemical Properties

• Caesium has an atomic number of 55, hence it has 55 electrons and 55 protons in its ground state.

• The atomic mass of Caesium is 132.9 . Caesium will therefore have; 132.9 - 55 ≈ 78 neutrons.

• Caesium is dangerously reactive. It is referred to as an hazardous element and must not be handled just anyhow because of its high reactivity.

• Caesium is the least electronegative element, this means it is the most electropositive element, and this makes it the most reactive metal.

Caesium, many of the time, will lose its valence electron to become Cs⁺ ion in compounds.

• Electronic configuration of Caesium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s¹ or [Xe] 6s¹

Reaction of Caesium

• On exposure to air, caesium will instantly ignite and catch fire.

• Caesium burns in oxygen to form Caesium Superoxide; CsO₂.

• Caesium superoxide, caesium peroxide, caesium oxide explode in water to form Caesium hydroxide.

$$2{\text{CsO}}_2+2{\text{H}}_2\text{O}\to 2\text{CsOH}+{\text{H}}_2{\text{O}}_2+{\text{O}}_2$$

• Pure Caesium explodes violently and instantly on contact with water to form Caesium hydroxide and hydrogen gas. It can be seen that the end fire is usually less than when you react a considerable amount of sodium with water; this is because caesium explodes instantly in water when a good amount of hydrogen gas has not yet been built up, unlike sodium where a considerable amount of hydrogen gas has built up and the heat from the reaction will then ignite the flammable hydrogen, leading to explosion.

$$2\text{Cs}+2{\text{H}}_2\text{O}\to 2\text{CsOH}+{\text{H}}_2$$

Caesium explosion in water | photo credit: gifsforum.com

• Caesium reacts with halogens, non-metals except noble gases, other metals to form alloys; but caesium does not react with Nitrogen- remember, Lithium is the only alkali metal that reacts with Nitrogen because nitrogen behaves like an inert gas.

• The chemistry of caesium is same as other alkali metals that we've discussed above.

Uses of Caesium

• Caesium is used in the most accurate atomic clocks.

• An isotope of Caesium is used in nuclear reactors.

• Caesium, mixed with formic acid is used as a drilling fluid for oil extractive industries.

• Caesium is used in thermionic generators (=they convert heat energy to electrical energy).

• Caesium is also used in photoelectric cells (=they convert light energy to electrical energy).

• Caesium halide crystals are employed during research to detect gamma and x-ray radiation.

• Caesium compounds are also useful in molecular biology for density gradient ultracentrifugation (=used to separate deeply mixed constituents of a fluid).

• Caesium vapour is used in many common magnetometers (=device that measures magnetic field and magnetic dipole moments; magnetic strength and orientation of a magnet or any object that produces a magnetic field).

• Like other alkali metals, Caesium can also be used as a getter in vacuum tubes.

• Caesium and rubidium carbonates are added to glass because they reduce electrical conductivity and improve stability and durability of fibre optics and night vision devices.

• Caesium salts are also sometimes used in medicine, but potassium or rubidium salts are usually better choices.

Francium

Discovery

Francium had a very funny discovery if you read the encyclopaedia very well. After the discovery of Caesium, about 10 years later, scientists believed that there should be another alkali metal after caesium. Research teams then attempted to locate and isolate the missing element. There were four false claims of the element's discovery until the real and actual discovery by Marguerite Perey in 1939, France. The element was given several names but in the end, it was named after France, Francium. She discovered it by the alpha decay of actinium 227 and she initially thought it was a known element, then tests eliminated the possibilities of it being the already known elements, further tests that were carried out showed that this element had properties of an alkali metal, until it was finally said to be the missing, alkali metal.

Occurrence

Francium is a radioactive element and it is formed by radioactive alpha decay of actinium, which is found only in very minute amounts in uranium minerals.

Francium from uranium minerals are very rare and are regarded to as trace.

Extraction

Francium is not extracted; like said earlier, it is gotten by radioactive decay. However, modern researchers believe Francium can be synthesized, or in other words, Francium can be synthesized by bombarding gold isotope - ¹⁹⁷Au with a beam of oxygen isotope - ¹⁸O. However, only very little amount of Francium is gotten.

Physical Properties

• Bulk Francium has never been seen or collected, therefore we currently cannot say if Francium would be a solid or liquid at room temperature. Researchers however tell that Francium appears to be a solid at STP.

• It is estimated that Francium has a density of about 2.48 g/cm³.

• The melting point temperature of Francium is estimated to be 27°C and its boiling point is 677°C.

Chemical Properties

• Francium has an atomic number of 87.

• Francium has a mass number of 223.

• Francium is radioactive, and its most stable isotope has an half life of about 21 minutes and 46 seconds.

• It is presumed that if enough Francium were to be collected, it would be as highly reactive as the other alkali metals in its group. Still, Francium is believed to be the second most electropositive element, after caesium.

•Electronic configuration of Francium is [Rn] 7s¹

•All isotopes of Francium decay into astatine, radium or radon.

Reaction Of Francium

• Due to Francium's instability, not so many compounds of it are known.

• It is expected that Francium would react with halogens to yield Francium halides.

$$\text{Fr}+\text{Cl}\to \text{FrCl}$$

• Francium halides are all soluble in water and are expected to be white solids.

• Francium perchlorate is produced by the reaction of francium chloride and sodium perchlorate.

• It is expected that Francium would form Francium Superoxide FrO₂ when burnt in oxygen.

• Francium nitrate, sulfate, hydroxide, carbonate, acetate, and oxalate, are all soluble in water, while the iodate, picrate, tartrate, chloroplatinate, and silicotungstate are insoluble.

Uses of Francium

Due to Francium's instability and rarity, there's no currently known use of Francium. It is only used for research purposes in the fields of Chemistry.

Key Points

📌 Group 1 elements are elements that have one electron in their valence shells.

📌 Group 1 elements include Hydrogen, Lithium, Sodium, Potassium, Rubidium, Caesium and Francium.

📌 Hydrogen is a non-metal while the rest are known as alkali metals.

📌 Alkali metals are very reactive, and they exist as soft solids at room temperature.

📌 Alkali metals readily lose their valence electrons and they react with non-metals to form ionic solids.

📌 Alkali metals are pyrophoric= meaning they react explosively with water, and can ignite in air.

📌 Alkali metals are soft, low dense solids; and they have low melting and boiling points which systematically reduces down the group.

📌 They all have body centered cubic crystal lattice structure and they give out coloured flames when heated in a hot Bunsen burner.

📌 Lithium gives off a Crimson coloured flame, Sodium gives off a yellow coloured flame, Potassium gives off a lilac coloured flame, Rubidium gives off a red coloured flame and Caesium gives off a blue coloured flame; Francium is unstable.

📌 Lithium exhibits some anomalous characters when compared with other alkali metals, this is because of its diagonal relationship with magnesium.

📌 Alkali metals are produced by electrolysis while others beneath in the group are produced using several other methods involving hydrometallurgy.

📌 All Alkali metals react on exposure to air.

📌 All Alkali metals react with oxygen to give oxides. Lithium forms a normal oxide, sodium forms a peroxide while others form superoxides.

📌 Lithium is the only alkali metal that reacts with Nitrogen and directly with carbon.

📌 All Alkali metals react with hydrogen to yield metal hydride, which can be dropped in water to form metal hydroxide and hydrogen gas.

📌 Alkali metals react explosively with water to form metal hydroxide.

📌 Alkali metals form strong alkali bases, and the basic strength increases down the group.

📌 All alkali metals dissolve in pure liquid ammonia to give blue electric conducting solutions if dilute and bronze if concentrated.

Recommended Videos

Lithium

Sodium

Potassium

Rubidium

Caesium

Francium

Test Questions

Discuss and Explain

1. What are alkali metals?

2. Why is hydrogen, a nonmetal, placed in the same group as the alkali metals?

3. Explain why group 1 elements often form compounds with elements in group 17.

4. Compare and contrast hydrogen and francium.

5. Why must the alkali metals be stored under an inert liquid like kerosene or an inert gas like argon?

6. How are the alkali metals produced commercially?

7. Why do you think sodium metal is the least expensive metal per unit volume?

8. What is the only element that reacts directly with nitrogen gas at room temperature?

9. Explain why the alkali metals cannot be stored in water.

10. What are the chemical formulas for potassium oxide, potassium peroxide, and potassium superoxide?