GROUP 3A ELEMENTS [BORON FAMILY]

Contents

• Introduction
• Group 13 / 3A Elements
• Occurrence of Group 3A Elements
• Physical Properties of Group 3A Elements
• Chemical Properties of Group 3A Elements
• Boron
• Aluminium
• Gallium
• Indium
• Thallium
• Nihonium
• Summary
• Recommended Videos
• Test Questions
• Discuss And Explain

Learning Objectives

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

• List and explain the characteristics of each Group 3A elements.
• Explain inert Pair effect in heavier group 3A elements.
• Describe the physical and chemical properties of group 3A elements.

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. In our last discussion, we talked on elements in Group 2A, right now we're to discuss the elements in Group 13 or 3A of the periodic table.

Periodic table of elements

Group 13 / 3A Elements

Group 13 / 3A Elements are elements that have three electrons in their valence shell. Two of these valence electrons can be found in the s orbital of the valence shell, while the third electron is in a p orbital. Hence, group 13 / 3A elements usually have an electronic configuration that ends with ns² np¹ where n = period number.

As you can see, I used Group 13 / 3A; this is because this group is rather far from the group 2 that we discussed in the previous post. After group 2, we have the transition elements, that span from group 3 to group 12. Group 13 follows, and it is more commonly referred to as Group 3A, this is because of the 3 valence electrons (possessed by the elements in this group).

The elements lose their 3 valence electrons to form compounds with +3 oxidation state, and moving down the group, the elements start to rather form compounds with +1 oxidation state.

The elements in Group 3A of the Periodic Table are: Boron, Aluminium, Gallium, Indium, Thallium, and synthetic Nihonium (also called Ununtrium).

The Group 3A of the periodic table is sometimes referred to as The Boron Family or Triels or First Post-Transition elements.

Occurrence of Group 3A Elements

The Group 3A elements are found in combined states in nature. Boron minerals are very rare in the earth's crust, accounting for about 0.001%. 

Aluminium is commonly found and is even the third most abundant element in the earth's crust (after oxygen and silicon), accounting for about 8%. Gallium and indium are found in moderate quantities in other minerals in the earth's crust while thallium is a bit rare. Nihonium has never been found in nature, it is made in the lab and it is therefore a synthetic element.

Physical Properties of Group 3A Elements

• Except for Boron (which is classified as a metalloid), Group 3A elements are silvery solids at room temperature.

• Group 3A elements are soft metals except Boron which is relatively hard. The characteristic softness of the elements increases down the group.

• The atomic and ionic radii of group 3A elements are usually less than that of the group 2A elements in each of their respective periods.

• The melting points of group 3A elements are high; Boron has a higher melting point than Beryllium, Aluminium has almost the same melting point as magnesium (due to diagonal relationship). The trend is however, not constantly maintained as we move down the group. Also, the densities of group 3A elements are usually higher than the group 2 metals.

• The group 3A elements are toxic metals, and they are good conductors of heat and electricity. Boron requires heat before it can be able to conduct electricity.

• The group 3A elements give colour to flame when heated.

Chemical Properties of Group 3A Elements

• Group 3A elements are relatively reactive elements except Boron which would only react at high temperature.

• Metallic property of group 3A elements increases down the group.

• Electropositivity increases from Boron to Aluminium, but decreases from Aluminium to Thallium.

Why? 

The increase in electropositive character from Boron to Aluminium is associated with the addition of an extra electron shell, which is farther from the central nucleus, hence, lesser force of attraction (between the valence electrons and the protons in the nucleus) leads to a more electropositive character. However, Ga, In and Tl have electrons arranged in d orbitals, and electrons present in d orbitals do not shield the nucleus very effectively; therefore, the protons in the nucleus are able to attract the valence electrons more strongly, making them harder to lose; leading to a lesser electropositive character. This is evidenced by the increase of Ionization Energy (which estimates the amount of energy needed to lose an electron) from Al to Ga.

• Elements (that are above) in the group, have tendency of forming compounds with +3 oxidation states, but as we move down the group, tendency to form compounds with +1 oxidation state increases. The univalent thallium compounds are the most stable. This monovalency occurs because electrons that are in the valence s orbital remain paired and therefore would not participate in bond formation. This usually occurs among heavy elements in the p-block and is called Inert Pair Effect. Inert Pair Effect is the reluctance of s-electrons to become unpaired, or take part in covalent bonding; only the electrons available in p orbital are available for bonding. In our case here, only one electron is available in p orbital, thus, the monovalency of heavier elements in group 3A.

• Due to the small sizes of their ions, their high charge and large values for the sum of the first ionization energies make us to believe that group 3A elements rather form covalent compounds. However, with the increase in ionic sizes (as we move down the group), more ionic compounds are formed. Boron compounds are always covalent while aluminium stands in between covalent and ionic.

Boron

Discovery

Boron was first discovered by Joseph Gay-Lussac and Louis Thénard in 1808, confirming Boron in its compounds; few days later, Sir Humphry Davy was able to isolate Boron. Jöns Jacob Berzelius identified Boron as an element in 1824, whilst several scientists argue that pure Boron was not gotten until around 1900s.

Occurrence

Boron occurs in nature as:

• Borax, Na₂B₄O_7·10H₂O
• Kernite, Na₂B₄O_7·2H₂O
• Ulexite, NaCaB₅O₉·8H₂O
• Colemanite, Ca₂B₆O_11·5H₂O
• Orthoboric Acid, H₃BO₃

Boron ores account for about 0.001% of the earth's crust.

Extraction

Boron preparation is a tough job to do.

Boron is prepared from Borax by thermal reduction using Na or Mg or Al as reducing agent. However, this preparation method does not usually yield pure Boron.

Pure Boron is produced by reduction of Boron halides (e.g, BBr₃ vapour), with Hydrogen gas, H₂.

$$2{\text{BBr}}_{3(g)}+3{\text{H}}_{2(g)}\to 2{\text{B}}_{(g)}+6{\text{HBr}}_{(g)}$$

Ultrapure boron is produced by the decomposition of diborane (B₂H₆) at high temperatures, the product is then purified by the zone melting or Czochralski processes.

Physical Properties

• Boron is allotropic (=exists in various physical forms). Crystalline Boron is a brittle, dark and lustrous (=shiny) metalloid while powder (or amorphous) Boron appears to be dark brown.

Pure Boron | photo credit: Wikimedia Commons

• Boron has a density of about 2.3 g/cm³.

• The melting point temperature of Boron is 2300°C and its boiling point is 3930°C.

• Boron burns with a bright green flame when heated in a hot Bunsen burner.

Chemical Properties

• Boron is the fifth element on the periodic table. In ground state, a Boron atom has 5 electrons, 5 protons.

• Boron has an atomic mass number of 10.8; hence Boron has approximately 6 neutrons in one atom.

• The electronic configuration of Boron is 1s² 2s² 2p¹ or [He] 2s² 2p¹.

• Boron forms majorly covalent compounds.

In covalent compounds (or bonding), there's electron sharing (or pairing). Prior to bond formation, Boron usually promote one electron from the (paired electrons in the) s orbital, into a vacant p orbital, and it afterwards, pairs up its (now) three unpaired electrons (with unpaired electrons from atoms of other elements or radicals) to form a total of stable 6 electrons in the outer shell of the newly formed molecule, this is disobedient to the octet rule. Due to this electron-deficient nature (because valence electrons are not complete 8), Boron compounds are known to behave as Lewis acids or electron acceptors.

• Boron hardly undergoes any reaction unless the reaction is subjected to higher temperatures.

• Boron is said to possess similar characteristics (or diagonal relationship) to Si, than it is to Aluminium.

Reaction of Boron

• On exposure to air, Boron will not react with air at room temperature, but when subjected to high temperatures, Boron burns to form Boron trioxide.

$$4\text{B}+3{\text{O}}_2\to 2{\text{B}}_2{\text{O}}_3$$

These are usually referred to as sesquioxides. A sesquioxide is an oxide that has three atoms of oxygen and two atoms of another element or its radical; M₂O₃

Boron trioxide | photo credit: Materialscientist via Wikimedia Commons

• Boron will also react with Nitrogen to form Boron nitride.

$$2\text{B}+{\text{N}}_2\to 2\text{BN}$$

Boron nitride is very hard and it exists in different forms and shapes.

• Boron carbide is an extremely hard material (almost as diamond) which is obtained by reacting B₂O₃ with carbon in an electric furnace.

$$2{\text{B}}_2{\text{O}}_3+9\text{C}\to {\text{B}}_4{\text{C}}_3+6\text{CO}$$

Boron carbide | photo credit: preslav

• Boron is prepared with hydrogen (in the lab) to form Boranes. Boranes don't occur naturally and they usually have chemical formula of B_xH_y. 

• They are electron-deficient, colourless and diamagnetic compounds. The simplest type of borane is diborane, B₂H₆. Diborane reacts spontaneously with air, usually violently, and with a green flash. We have lots of higher and complex Boranes; The CHM students will learn more on this in higher classes.

• Boron reacts with halogens to form Boron trihalides.

$$2\text{B}+3{\text{X}}_2\to 2{\text{BX}}_3$$

 Preparation of Boron halides in the lab follow different method(s).

The halides of Boron are strongly covalent compounds. They are volatile, highly reactive, exist as gases, and they act as Lewis Acids (due to 6 electrons in outer shell instead of 8, causing their electron-deficient nature; a vacant p orbital, thus, they act as electron acceptors;).

Boron trifluoride reacts with water to form fluoroborates; while other halides of Boron react with water to form Boric acids.

$${\text{BCl}}_3+3{\text{H}}_{2}O\to {\text{H}}_3{\text{BO}}_3+3\text{HCl}$$

The Lewis acidity strength increases as we move down the halogen group.

All Boron halides have trigonal planar molecular shape and are also monomeric (as opposed the polymeric structure of Boron hydrides or Boranes).

• Boron acts as a non-metal and reacts with other metals, forming metal boride.

Uses of Boron

• The highly exothermic reaction of Boranes with oxygen make them considerable as rocket fuels by space program.

• Borax is a water softener in washing powders.

• Duran, Pyrex, are types of Borosilicate glass containing Boron- (I can never forget that this unexpectedly appeared as one of my exam questions long time then, and I happily failed it🙄😂).

• Boron is used in the production of fibre glasses and ceramic, together with super-hard materials.

• Boron is used in agricultural fertilisers since plants require a trace amount of Boron as nutrients.

• Boron compounds are used in bleaches, mild antiseptics and detergents.

Aluminium

Discovery

Aluminium was discovered by Hans Christian Ørsted in 1825.

Occurrence

Aluminium occurs in many minerals in nature, some of these minerals are:

• Bauxite, Al₂O₃·2H₂O
• Alumina, Al₂O₃
• Cryolite, Na₃AlF₆
• Kaolin, Al₂O₃6SiO₂·2H₂O
• Feldspar KAlSi₃O₅
• Mica (Muscovite)
• Gibbsite, Al(OH)₃
• Boehmite, AlO(OH)
• Diaspore

Aluminium minerals account for about 8.3% of the earth's crust.

Extraction

Aluminium is obtained through the electrolytic 'Bayer–Hall Héroult processes' in which Bauxite ore (which is first calcinated in caustic soda, NaOH, to remove impurities) is dissolved in molten cryolite, Na₃AlF₆, at a temperature of about 1000°C, then electricity is made to pass through this mixture which reduces the Al³⁺ ions to neutral Al atoms. The molten Aluminium is then drained off and cast into large blocks.

Physical Properties

• Aluminium is a shiny, silvery metal at room temperature.

Aluminium

• The density of Aluminium is 2.73 g/cm³.

• The melting point temperature of Aluminium is about 660°C and its boiling point is 2470°C.

• Aluminium has a face centered cubic crystalline structure and it is diamagnetic.

• Aluminium burns with a colourless flame when heated in a hot Bunsen burner.

Chemical Properties

• Aluminium is the 13th element on the periodic table, thus, one atom has 13 protons and 13 electrons.

• Aluminium has an atomic mass number of 26.9, hence, one Aluminium atom has about 14 neutrons.

• The electronic configuration of aluminium is 1s² 2s² 2p⁶ 3s² 3p¹ or [Ne] 3s² 3p¹

• Aluminium compounds act as partly ionic and partly covalent.

• Aluminium exhibits diagonal relationship with Beryllium, hence, it has many properties of Beryllium (that we discussed in the previous topic). Some of them are:

- Aluminium forms a protective oxide coating on its surface when exposed to air- this passivation layer prevents further oxidation of the Aluminium metal. Due to the ubiquity of Aluminium, this characteristic makes it a preferable coating material to prevent rusting of metals like iron.

- Aluminium oxide and Aluminium hydroxide behave as amphoteric substances; just as Beryllium.

- Be and Al do not impart any colour to flame.

- Beryllium carbide and Aluminium carbide react with water to form methane.

$${\text{Be}}_2\text{C}+2{\text{H}}_2\text{O}\to 2\text{BeO}+{\text{CH}}_4$$

$${\text{Al}}_4{\text{C}}_3+2{\text{H}}_2\text{O}\to 2{\text{Al}}_2{\text{O}}_3+{\text{CH}}_4$$

- Aluminium chloride, AlCl₃, is a volatile solid, existing as a polymer (in solid phase), a dimer (in molten phase and gas phase) and a monomer (in hot vapour phase).

Aluminium trichloride 3D structures

- Both Aluminium and Beryllium are resistant to Nitric acid, HNO₃.

Reaction of Aluminium

• When a piece of aluminium sheet is exposed to moist air, it acquires a thin, continuous coating of aluminium oxide, which prevents further attack on the metal by atmospheric oxygen and water under normal conditions.

However, if aluminium powder is heated to 800°C and above, the metal will react with air to form aluminium oxide, Al₂O₃, and aluminium nitride, AIN.

$$4\text{Al}+3{\text{O}}_2\to 2{\text{Al}}_2{\text{O}}_3$$

$$2\text{Al}+{\text{N}}_2\to 2\text{AlN}$$

The reaction is usually accompanied by the evolution of heat and intense white light. This property of aluminium is exploited in flash light photography.

Aluminium oxide | photo credit: Aariuser

• Heated aluminium combines with the halogens, sulphur, nitrogen, phosphorus and carbon, accompanied by the evolution of heat.

• Aluminium on heating with hydrogen forms aluminium hydride.

$$2\text{Al}+3{\text{H}}_2\to 2{\text{AlH}}_3$$

• Aluminium is amphoteric. It dissolves in both acids and bases. Aluminium reacts slowly with dilute hydrochloric acid and more rapidly with concentrated hydrochloric acid to displace hydrogen.

$$2\text{Al}+6\text{HCl}\to 2{\text{AlCl}}_3+3{\text{H}}_2$$

Aluminium does not react with dilute sulphuric acid. However, it reacts with hot concentrated sulphuric acid to liberate sulphur(IV) oxide gas.

$$2\text{Al}+6{\text{H}}_2{\text{SO}}_4\to {\text{Al}}_2({\text{SO}}_4{)}_3+3{\text{SO}}_2+6{\text{H}}_2\text{O}$$

Aluminium metal does not react with nitric acid at any concentration, because of the formation of protective layer of aluminium oxide. The acid is said to render the aluminium passive. Nitric acid is, therefore, frequently transported in aluminium containers. 

Aluminium dissolves in both sodium and potassium hydroxides to form a soluble aluminate, with the evolution of hydrogen.

$$2\text{Al}+2\text{NaOH}+6{\text{H}}_2\text{O}\to 2{\text{NaAl(OH)}}_4+3{\text{H}}_2$$

• When Aluminium oxide is heated with carbon, Al₄C₃, is formed. This on hydrolysis evolves methane.

$${\text{Al}}_4{\text{C}}_3+2{\text{H}}_2\text{O}\to 2{\text{Al}}_2{\text{O}}_3+{\text{CH}}_4$$

Uses of Aluminium

• Most of the aluminium that are produced are destined to be alloyed; Aluminium alloys have a wide range of uses; aluminium foils, beverage cans, automobiles, aircraft, trucks, railway cars, marine vessels, bicycles, spacecrafts, Building and construction (such as windows, doors, siding, building wire, sheathing, roofing, etc.), a wide range of household items, from cooking utensils to furniture, machinery and equipment (processing equipment, pipes, tools), casing of many metallic phones and laptops.

• Aluminium has affinity for oxygen, it is therefore used as a reducing agent in chemical and steel industries.

• Aluminium is used for making petrol and milk storage tanks because it reflects heat and prevents them of being over heated in the sun.

• Aluminium is used in electrical wiring; this is because it is much cheaper than copper which is a better conductor.

• Potassium Aluminium sulphate, commonly known as alum, is used in the dyeing industry as a mordant (or binder), what it does is to fix dye to fabrics and render (the dye) as insoluble in water. It is also used as a chemical flocculant in the purification of water, and in medicine as an astringent (which causes vasoconstriction= contraction of blood vessels) to stop bleeding.

Gallium

Discovery

Gallium was discovered in 1875 by Paul-Emile Lecoq de Boisbaudran.

Occurrence

Gallium is the 34th most abundant metal on earth's crust. It is only found majorly as an impurity in the ores of other metals, like bauxite, coal, diaspore, germanite, zinc blende, etc

Extraction

Gallium is produced by electrolytic reduction of its aqueous solutions. For semiconductor use, it is further purified by zone melting or Czochralski process.

Physical Properties

• Gallium is many times called a weird metal because it displays some weird metallic properties.

Gallium | photo credit: Wikimedia Commons

• It is a shiny silvery solid at room temperature, but soon melts into a silvery-white liquid when the room's temperature is higher.

• Metallic Gallium has a very low melting point of about 29°C - 30°C, but the liquid requires a very much higher temperature to boil, around 2400°C. Gallium is then said to have an extremely wide liquid temperature range.

• Metallic Gallium has a density of about 5.9 g/cm³. Surprisingly, Gallium metal usually floats on its liquid, which means that liquid Gallium has a higher density than metallic Gallium- this is weird for a metal.

• When Liquid Gallium solidifies to metallic Gallium, expansion is usually noticed- which is again weird for a metal.

• Liquid Gallium wets glass and form an adhesive metallic layer.

• Scientists say that Gallium does not have a simple crystalline structure like other metals, its structure is said to be ortho-rhombic.

• Gallium burns with a violet flame when heated in a hot Bunsen burner.

Chemical Properties

• Gallium has an atomic number of 31 and an atomic mass number of 69.7.

• The electronic configuration of Gallium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p¹ or [Ar] 4s² 3d¹⁰ 4p^1.

• Gallium is the first metal to come after the first set of the d block elements. The d block elements have electrons in their d orbitals, and electrons in d orbitals do not shield the nucleus very effectively (like s and p orbitals). Since the protons and electrons are of the same number, and the valence electrons are not properly shielded, the protons in the nucleus are able to attract the valence electrons more strongly, which leads to the dropped reactivity of Gallium, because it will hardly lose its valence electrons.

• Gallium is usually found in its compounds as having a +3 oxidation state. In some other compounds, Gallium has +1 oxidation state, but it is somekind of rare because such compounds tend to be unstable. It is not impossible for Gallium to gain electron(s) and have a negative oxidation state.

• Gallium, unlike other elements above it, forms more ionic compounds than covalent.

• Gallium forms alloys with most metals.

• The oxide and hydroxide of Gallium are said to possess amphoteric nature.

Reaction of Gallium

• At room temperature, gallium metal does not react with air and water because it forms a passive, protective oxide layer. At higher temperatures, however, it reacts with atmospheric oxygen to form gallium(III) oxide, Ga₂O₃.

$$4\text{Ga}+3{\text{O}}_2\to 2{\text{Ga}}_2{\text{O}}_3$$

• Gallium reacts with ammonia at 1050°C to form gallium nitride, GaN.

• Gallium also forms binary compounds with phosphorus, arsenic, and antimony: gallium phosphide (GaP), gallium arsenide (GaAs), and gallium antimonide (GaSb).

• Gallium(III) oxide reacts with fluorine or fluorine compounds to form gallium(III) fluoride, GaF₃. It is an ionic compound that is strongly insoluble in water.

• Gallium reacts with other halogens to form Gallium trihalides which normally exist as dimeric molecules.

Uses of Gallium

Gallium is majorly used in the semiconductor industry, for manufacturing of electronic components such as LEDs, transistors, laser diodes etc.

Indium

Discovery

Indium was discovered in 1863 by Ferdinand Reich and Hieronymus T. Richter by Spectroscopy.

Occurrence

Indium is the 68th most abundant element in the earth's crust. Its minerals are very rare, but it is found in more common ores like Zinc blende and chalcopyrite.

Extraction

Indium is usually gotten as a by-product during smelting of zinc blende, or ores containing Gallium as an impurity. Further purification can be done by electrolysis.

Physical Properties

• Indium is a shiny silvery metal at room temperature.

Indium

• Indium is soft and can be cut with a knife.

• Density of indium is about 7.3 g/cm³.

• Indium has a body centered tetragonal crystalline structure and is diamagnetic.

• The melting and boiling points of Indium are 157°C and 2080°C respectively.

• Indium burns with a dark blue or indigo flame when heated in a hot Bunsen burner.

Chemical Properties

• Indium has an atomic number of 49 on the periodic table. It has an atomic mass number of about 114.8.

• Electronic configuration of indium is [Kr] 5s² 4d¹⁰ 5p¹.

• Indium exists majorly as +3 oxidation state in many of its compounds. In some other cases, the 5s electrons are not donated (due to inert Pair effect) and Indium forms compounds with +1 oxidation state, but the compounds are usually less stable.

• Indium is chemically similar to Gallium, though usually intermediate between Gallium and thallium.

• Indium oxides and hydroxides are amphoteric; though amphoteric character is less than those above it in the group.

• Indium has 39 known isotopes.

Reaction of Indium

• Like others above it, indium does not react with air at room temperature, but at higher temperatures, powdered indium will react with oxygen in air to form In₂O₃.

• Indium oxide is amphoteric, but most times it is used as a base.

• Indium reacts with water to reproduce soluble indium(III) hydroxide, which is also amphoteric.

$$2\text{In}+6{\text{H}}_2\text{O}\to 2{\text{In(OH)}}_3+3{\text{H}}_2$$

• Indium reacts with halogens to form trihalides. 

• Indium fluoride is polymeric.

• Check reaction of Gallium for other reactions of Indium

Uses of Indium

Indium and its compounds are most commonly used in the semiconductor industry. It also found use in dentistry, usually used in form of alloys.

Thallium

Discovery

Thallium was discovered in 1861 by Sir William Crookes by flame spectroscopy. Its discovery was made by two contemporary scientists, in the same year. The second scientist was Claude-Auguste Lamy who discovered thallium from residues of the production of sulphuric acid.

Occurrence

Thallium is not found in nature, it is usually found in potassium-based ores or heavy metal sulphide ores.

Extraction

Thallium is produced by electrolytic purification of the by-products of smelting of heavy metal sulphide ores.

Physical Properties

• Thallium is a very soft silvery solid at room temperature. It is soft enough to be cut with a knife.

Pieces of thallium in ampoule under argon, via W. Oelen

• Thallium has a density of 11.85 g/cm³.

• Thallium has melting and boiling points of 304°C and about 1460°C respectively.

• Thallium burns with a bright green flame when heated in a hot Bunsen burner.

Chemical Properties

• Thallium has an atomic number of 81 and an atomic mass number of 204.4.

• Thallium has an electronic configuration of [Xe] 6s² 4f¹⁴ 5d¹⁰ 6p¹

• Thallium does not readily lose its 6s electrons due to inert Pair effect, it only loses the electron that is available in 6p shell, thus, having +1 oxidation state in its compounds. In some other cases, strongly electronegative non-metals forcefully collect the 6s electrons and make thallium have a +3 oxidation state, but these compounds are less stable because thallium can spontaneously reduce itself back to +1 oxidation state under standard conditions; compounds having thallium in +1 oxidation state are more stable than the ones with +3 oxidation state, unlike other elements above it in the group.

• Thallium forms strongly ionic compounds.

Soluble thallium salts are regarded to as highly toxic.

Reaction of Thallium

• On exposure to air, thallium quickly oxidizes and forms a bluish-grey oxide which is not a passivation layer, hence, Thallium must be stored in oil or under inert conditions.

• Thallium readily dissolves in water to form Thallium hydroxide which is basic in nature. Thallium oxide is also basic, unlike the amphoteric character of other elements above it in the group.

• Thallium readily dissolves in sulphuric and nitric Acids to form Thallium sulphate and Thallium nitrate salts. Reaction with hydrochloric acid is usually slow due to the formation of insoluble thallium(I) chloride layer.

• Thallium reacts with halogens to form Thallium halides, and sometimes, less stable Thallium trihalide. Thallium chloride and Thallium bromide have caesium chloride structure, while thallium fluoride and thallium iodide have distorted sodium chloride structures.

• Thallium reacts with sulphur, Phosphorus and other non-metals.

• Thallium also forms many double salts.

Uses of Thallium

• The toxicity of soluble thallium compounds, especially thallium sulphate which is odourless and tasteless, were used in rat poisons, but later banned due to safety reasons.

• Thallium(I) bromide and Thallium(I) iodide crystals are used in infrared optic materials.

• Thallium(I) sulphide and thallium selenide are used in the semiconductor industry.

• Radioactive thallium-201 is injected into patients taking stress test, used in the diagnosis of heart disease.

Nihonium

Nihonium is a synthetic element created by some team of Japanese scientists by bombarding two different nuclei (according to their publication, they used Bismuth + Zinc) together to form a heavier element. Nihonium has not been discovered in nature, and the most stable isotope created in the lab has a half life of max 10 seconds. Thus, Nihonium has not been studied and all its properties are only based on predictions and/or assumptions. IUPAC recommended the name Ununtrium (Uut) as a placeholder name until the element is actually discovered; this is however, usually ignored by scientists and in most text books.

Key Points

📌 Group 3A elements have 3 electrons in their outermost shell.

📌 Group 3A elements lose their valence electrons to form compounds with +3 oxidation states. However, when moving down the group, inert Pair effect causes heavier elements to lose only one electron, making them have oxidation state of +1.

📌 Except for Boron, Group 3A elements are relatively soft metals.

📌 Group 3A elements are relatively reactive, reactivity increases from Boron to Aluminium, but then reduces from Aluminium to Gallium, increases again from Gallium to indium, but finally reduces from Indium to Thallium.

📌 Group 3A elements all have high boiling points, and their melting points are relatively high, except for Gallium which is characterised as a weird metal.

📌 Boron exhibits diagonal relationship with Silicon of Group 4A, while Aluminium exhibits diagonal relationship with Beryllium of Group 2A.

📌 The oxides and hydroxides of Group 3A elements behave as amphoteric but basic character increases down the group.

📌 Group 3A elements are: Boron, Aluminium, Gallium, Indium, Thallium, and synthetic Nihonium.

Recommended Videos

Group 3A Elements

Don't mind the accent, follow the explanation in this video, it contains what you need to know.

Boron

Aluminium

Gallium

Indium

Thallium

Nihonium

Test Questions

Discuss And Explain

1. Aluminium liberates _ gas when it reacts with conc. H₂SO₄. Explain why it doesn't liberate hydrogen gas instead.

2. Boron forms _ bond in its compounds. Explain the reason for this.

3. Why is Aluminium not found as a free element? Explain its extractive processes.

4. Give the names, electronic configurations and Occurrence of Group 3A elements of the periodic table.

5. Explain why Thallium(I) compounds are rather stable than the Thallium(III) compounds.

6. Explain inert Pair effect and the reason behind it.

7. Explain why the reactivity of Gallium is less that of Aluminium since reactivity was meant to increase down the group.

8. Explain how Boron carbide is produced. Why is it used as a material in ceramics?

9. Explain why Nitric Acids are usually transported in Aluminium containers, without the acid eating up the aluminium component.

10. List out the different flame colours emitted by group 3A elements.

11. Which of the group 3A elements has an orthorhombic crystalline structure? Describe its other characteristics.

12. Which Group 3A element would its metal float on its liquid? Explain its reactivity when compared with Aluminium and Thallium.