S-Block Elements

• The outermost orbital of s block elements consists of either one or two electrons and next to the outer most shell (penultimate shell) has either 2 or 8 electrons.
• s-block elements show a fixed valency which depends on the number of electrons present in the outermost shell. The valency of alkali metals is 1 and that of alkaline earth metals 2.
• Except hydrogen , all s-block elements have low values of ionisation potential decreases in the case of alkali metals and alkaline earth metals as the atomic number increases . Ionisation potential increases on moving horizontally from IA to IIA. On account of low values of ionisation potentials , these elements are highly electropositve , i.e., easily lose valency electrons and form cations.
• The ions M⁺ ( alkali metals ) and M²⁺ ( alkaline earth metals ) have stable configuration of inert gases . These ions do not lose further electrons , hence s-block elements do not show variable oxidation states.
• Alkali metals have the largest atomic radii in their corresponding periods followed by alkaline earth metals . Values of atomic and ionic radii increase from top to bottom in groups. Atomic volume follows the same trend .
• Except hydrogen , s-block elements have low values of electronegativity which decreases from top to bottom in both the groups .
• The standard reduction potentials of s- block elements are highly negative. These are on the top the electrochemical series. The alkali metals and alkaline earth metals cannot be prepared by doing electrolysis of aqueous solutions of their salts.
• The binding energies of s-block elements are low, since one (Group IA) or two ( Group IIA ) electrons are available for bond formation in the crystal lattice of the metal . The binding energy decreases from IA to IIA .
• On account of low values of binding energies , the melting and boiling points are low in the case of s-block elements . These decrease in a group from top to bottom while increase on moving horizontally from IA to IIA .
• s-block elements are soft and malleable . IIA elements are harder than alkali - metals.
• s-block elements are good conductors of heat and electricity .
• These are strong reducing agents as ionisation potential values are low. Reducing property increases in a group from top to bottom while decreases moving horizontally from IA to IIA .
• The reducing property of these elements in solution depend on standard reduction potentials . Lithium having maximum negative E° value , is the strongest reducing agent amongst all alkali metals in solution. Reducing nature increases from Na to Cs in alkali group and from Be to Ba in IIA group .
• The ions of s-block elements are highly hydrated in aqueous solution . Smaller the ion , greater shall be its hydration. The degree of hydration decreases from Li⁺ to Cs⁺ in group IIA . During hydration , energy is released . This is known as heats of hydration. This decreases from top to bottom in groups . The radii of the hydrated ions are in the order of :
  
  IA Li⁺ (aq) > Na⁺ (aq) > K⁺ (aq) > Rb⁺ (aq) > Cs⁺ (aq)
  II A Be²⁺ (aq) > Mg²⁺ (aq) > Ca²⁺ (aq) > Sr²⁺ (aq) > Ba²⁺ ( aq)
• The mobility of these ions under the influence of electric current depends on the size i.e., conductivity of the aqueous solution of cesium salts is the highest while that of lithium salts is minimum in the case of alkali metals .
• When s-block elements or their salts are heated in flame , they characteristic colour to the flame except Be and Mg.
  
  Li Na K Rb Cs
  crimson yellow violet violet violet
  Be Mg Ca Sr Ba
  - - brick red crimson apple green
• When these are heated, electrons are excited to higher energy states by absorption of energy . These electrons when return to the original ground state, the energy is given out in the form of light.
• Reactivity of these elements increases from top to bottom in a group , Li , Be and Mg show somewhat abnormal properties . s-block elements show the following chemical properties .
  (i)They decompose water readily and evolve hydrogen .
  (ii) They displace hydrogen from for oxygen and form corresponding salts.
  (iii) They have great affinity for oxygen and non-metals .
  (iv) Their hydroxides are strong alkalies and oxides are highly basic.
• The alkali metals are extremely reactive metals. They all readily lose one electron, the outermost s-electron, to form a single positively charged ion that has the same electronic configuration as one of the noble gases. The relatively simple chemistry of alkali metals is probably due to the ease with which their single outermost electron can be lost in order to achieve stable rare gas electronic configuration. As a group, the alkali metals have the lowest first ionization energies among all the elements in the periodic table. Because of their strong chemical activity, the alkali metals occur in nature only in the form of compounds. Alkali metals never occur in free state in nature. Compounds of sodium and potassium are the most abundant, while those of remaining alkali metals occur in nature rarely. Francium is radioactive element and short lived. These metals are readily oxidised in air. The less electropositive such as lithium is found as silicate and the most electropositive elements of IA occur as chlorides. To prevent oxidation they are stored in closed vessels of low density (0.53 g/cc). Lithium is the lightest or least dense of all metals, with a density roughly half that or water. Lithium is kept wrapped in paraffin wax.
• Because of low ionisation energies and strong electropositive character, alkali metals such as K, Rb and Cs show photoelectric effect .
• The alkali metals must be protected from exposure to air because they react with O₂, H₂O and CO₂ . When burnt in air, lithium forms monoxide (Li₂ O), sodium forms monoxide in limited O₂ and peroxide in excess O₂. Other alkali metals form largely the super oxides. In the presence of atmospheric oxygen we will get the mixture of oxides.
  
  2Li + 1/2 O₂ Li₂ O
  2Na + O₂ Na₂ O₂
  2Na + 1/2 O₂ Na₂ O
  Rb + O₂ RbO₂ (Superoxide , containing O₂^- ion)
  
• The alkali metals replace hydrogen from those acids which are not oxidizing agents (eg. HCI, dilute H₂ SO₄). While reacting with oxidizing acids such as conc. H₂ SO₄, conc. HNO₃ and dilute HNO₃ , these metals do not replace hydrogen but H₂ S , N₂ O and NH₃ are respectively formed.
  
  5 H₂SO₄ + 8M 4 M₂SO + H₂S + 4H₂O
  10 HNO₃ (conc.) + 8M 8 MNO₃ + N₂O + 5H₂ O
  9 HNO₃ (dil.) + 8M 8MNO₃ + NH₃ + 3H₂
• The fact that lithium reacts with water less vigorously may probably be due to its higher melting point. Other alkali metals have low melting points (below 373 K) and melt due to the heat liberated in their initial reaction with water, thus exposing more metal surface to attack. The reaction, therefore, proceeds much more rapidly than with lithium which remains in solid state.
• Despite its high ionisation energy, lithium is the only alkali metal, which reacts directly with nitrogen and carbon to form the nitride and carbide respectively. Other alkali metals do not react directly with either nitrogen or carbon. This is probably due to the fact that small size of the nitride and carbide ions coupled with the small size of Li⁺ ions results in high lattice energies for these compounds. Thus heats of formation of these compounds are also high resulting in more stability.
• All alkali metals form amalgams with mercury and the reaction is exothermic. Sodium amalgam is less reactive than the sodium metal. All alkali metals combine directly With P, As, P and S.
• When heated with NH₃ gas at 300-400°C in presence of Fe, alkali metals form amides of the type MNH₂. The expected order of reactivity of the metal and the corresponding stability of the salt formed is Cs > Rb > K > Na > Li.
• Except lithium nitrate, the nitrates of alkali metals are thermally stable. Alkali metal sulphates form double salts called alums. Lithium does not form these alums because of its small size.
• The only monoxide obtained by direct combination of the metal with oxygen is that of lithium, Li₂ O. Monoxides are ionic and colourless compounds and contain the oxide ion O^2-. They are strongly basic. With the exception of lithium monoxide, all others can be produced by reducing the nitrite or nitrate with free metal.
  
  2NaNO₂ + 10Na 6Na₂O + N₂
• Superoxides are the direct oxidation products of K, Rb and Cs, in excess of oxygen. Sodium super oxide can also be prepared by heating sodium peroxide with oxygen under pressure in a sealed vessel.
  
  Na₂ O₂ + O₂ 2NaO₂
• Super oxides contain O₂^- ion which has one electron more than the oxygen molecule.
• The stability of peroxides and Superoxides increases with increase with increase in size of cation and shows stabilisation of large anions by large cation through lattice energy effects. The inability of lithium to form a peroxide or superoxide may probably be due to strong polarisation power of its ion. The Li⁺ ion is strong enough to restrict the oxidation of an oxide ion,O^2- to the peroxide ion ,O₂^2-. Similarly Na⁺ prevents the oxidation of O₂^2- to O₂^-.
• Nickel and cast iron are resistant to fused caustic alkalis. Elements having electronegativity values less than 1.5 are not attacked by alkali hydroxides. Those elements having electronegativity values between 1.5 and 2 .0 dissolve in hydroxides and H₂ is evolved with the formation of alkali metal salt of the oxide . Elements having electronegativities greater than 2 dissolve without evolution of hydrogen forming their hydride or the alkali metal derivative of that hydride as one of the products.
• Hydroxides of all alkali metals are thermally very stable except LiOH , which decomposes to the oxide when calcined.
  2LiOH Li₂ O + H₂ O
• Lithium hydride is by fact the most stable hydride which dissociates above 1273 K. Sodium and potassium hydrides dissociate below 773K. The stability decreases from LiH to CsH.
• The Li-H bond has been fond to be only 25% ionic. As the ionic character of M-H bond increases with increase in size of the metal, it would appear that M⁺ H^- is less stable than the covalent M-H bond . The alkali metal hydrides are strong reducing agents and reducing property increases with decrease in stability. Lithium also forms some complex of the Li AlH₄ and LiBH₄ , which like simple hydrides , are also good reducing agents.
• For a substance to dissolve , the hydration energy must be greater than lattice energy . Due to small size of Li⁺ ion , the hydration energy of LiF is considerably high , but it has low solubility in water because its lattice energy is even higher . The hydration and lattice energies of LiF are -1034 and - 1039 kJ/ mole respectively.
• The gain in lattice energy resulting from the substitution of a smaller oxide ion from larger CO₃^2- ion enables the decomposition. Bicarbonates of all alkali metals , except that of Li, can be obtained in the solid form. All carbonates and bicarbonates are soluble in water . Lithium carbonate is sparingly soluble.
• Lithium is extracted by the electrolysis of a fused mixture of LiCl and KCl ( 1:1) . KCl reduces the m.p of of LiCl from 883K to 673 K . Sodium is extracted by the electrolysis of a mixture of NaCl and CaCl₂ (Down's process) . CaCl₂ decreases the m.p of NaCl from 1076K to 778K. Potassium cannot be prepared by the electrolysis due to its low melting point and ready vapourisation. It can be prepared by passing sodium vapour over molten KCl in a counter current fractionating tower. Rubedium and caesium are made by similar methods.
• The alkali metals are best purified by distillation .
• The alkali metals are soluble in liquid ammonia giving a solution which is paramagnetic , highly conducting, highly reducing and deep blue in colour . The solution is blue when dilute and acquires and intense blue colour as more metal is added . Colour of these solutions is independent of the metal involved. Dilute solutions are paramagnetic and blue solutions have a broad absorption band at 1450nm. The blue solution has remarkably high electrical conductivity, which varies with concentration in an anomalous manner. Properties of these solutions strongly suggest that alkali metal atoms ionise in liquid ammonia forming solvated cations and solvated electrons. The solutions are strong reducing agents due to the presence of free electrons. It reduces (a) Metal halides to metals (b) N₂O to N₂ (c) O₂ to O₂^- and then to O₂^2- . The solution of a metal in liquid ammonia causes a large increase in volume (decrease in density). This solution of a metal in liquid ammonia causes a radius 300-400 pm. As a result , the solutions occupy far greater volume than expected from the sum of volumes of metal and solvent. This causes a decrease in density .
  
• Most abundant alkali metal in earth's crust is sodium . It was first isolated by Davy in 1807 . It is one of the most reactive metals and so it does not occur in free combines vigorously with water and many other elements and compounds, where it acts as a strong reducing agent . It occurs in combined form as rock salt (NaCl) , chile saltpetre or caliche (NaNO₃ ), borax (Na₂B₄O₇, 10H₂O), cryolite (Na₃AlF₆) and felspar (NaAiSi₃ O₈). Sodium is extracted by Castener's process and Down's process. Sodium is soft silvery white metal, lighter than water .Its density is 0.972 g./cc. It is malleable as well as ductile.
• There are number of important sodium compounds such as NaOH (caustic soda), sodium carbonate (Na₂ CO₃. 10 H₂ O ) , Na₂ CO₃ , (soda ash) and sodium bicarbonate , NaHCO₃ (baking soda). Sodium hydroxide is hygroscopic. When it is dissolved in water, a large amount of heat is evolved (exothermic). NaOH is prepared by the electrolysis of NaCl solution (brine). Industrially it is prepared by mixing lime with sodium carbonate. Sodium carbonate is obtained from mines or by Solvay process. Sodium chloride can be obtained from sea or underground deposits. Sodium carbonate is widely used in the manufacture of soaps, detergents and glass etc. Sodium bicarbonate or baking soda is used to make breads and cakes light and fluffy because when it is heated, CO₂ is evolved. 2NaHCO₃ NaCO₃+ CO₂ + H₂O , Sodium bicarbonate is also used in fire extinguishers . Zinc uranyl acetate is one of the few fairly insoluble compounds of sodium. Its formula is NaZn (UO₂)₃ (CH₃ COO)₉ . It is pale yellow in colour . Sodium metal is used in Wurtz reaction. Sodium is also used as a coolant in atomic reactors.
  
• Electrolysis of aqueous NaCl gives H₂ gas at cathode . Metallic sodium is released at cathode, if aqueous NaCl is electolysed-using mercury as cathode.
• Soda lye is solution of NaOH in water. Common name for KOH is caustic potash.
• NaOH can be prepared by (a) Causticizing process or lime soda process (b) Nelson's method (c) Castner Kellner methid . NaOH is deliquescent and so a standard solution ofNaOH cannot be prepared simply by weighing.
  
• Al , Zn, Sn , Pb, Cr and Sb are the metals which dissolve in NaOH solution and liberate H₂ .
• The amphoteric substances such as Al₂O₃ , ZnO, Al (OH)₂ dissolve in excess of NaOH solution forming aluminates and zincates.
  Al (OH)₃ + 3NaOH NaAlO₃ + 3H₂O
  Zn(OH)₂ + 2 NaOH NaZnO₂ + H₂O
• Aqueous solution of NaOH and KOH can CO₂ from the atmosphere.
• Alcoholic KOH is used as dehydrohalogenating agent.
• Na₂ CO₃ is called soda ash and Na₂ CO₃. 10H₂O is called washing soda.
• Na₂ CO₃ can be prepared by Le Blanc process and Solvay process or ammonia soda process.
• Solvay process is based on low solubility of NaHCO₃ . The end product of Solvay process in NaHCO₃ . In soda process, the raw materials usedare Ca CO₃ , NH₃ and NaCl .
• K₂ CO₃ , known as pearl ash , can not be prepared by Solvay process due to high solubility of KHCO₃ .
• NaHCO₃, known as baking soda, is less soluble in water than Na₂ CO₃ because of polymeric anions formed by hydrogen bonding .
• All baking powders contain NaHCO₃ and acid salt such as potassium hydrogen tartrate or calcium dihydrogen phosphate .
• An equimolar mixture of Na₂ CO₃ and K₂ CO₃ is called fusion mixture .
• Sodium bicarbonate is used to make breads and cakes light and fluffy because when it is heated CO₂ is evolved. Sodium bicarbonate is also used in cold drinks, as antacid and in baking powders.
• Na₂CO₃ is widely used in the manufacture of soaps, detergents and glass. Both NaOH and Na₂ CO₃ are among the top 20 chemicals used in industries .
• A deliquescent compound such as NaOH is that which absorbs water readily from the atmosphere and thereby form concentrated solutions containing the alkali metal cation (Na⁺) and OH^- ions.
• Na₂O is used as a dehydrating agents as well as polymerising agent in organic chemistry.
• Anhydrous sodium sulphate is called salt cake and hydrated sodium sulphate is called Glauber's salt (Na₂SO₄. 10 H₂ O). Glauber's salt shows efflorescence in dry air and it is used as purgative in medicine. Na₂ SO₄. 10H₂O becomes anhydrous (Na₂SO₄), above 32.5°C (efflorescence).
• Na₂ O₂ is a yellow powder (Na₂ O₂. 8H₂O). It is capable of absorbing gases like CO, CO₂, SO₂, NO₂, etc. Hence it is used to purify air in submarines.
• Under trade name oxone, Na₂ O₂ is used as a bleaching agent.
• K₂O₂ is a white hygroscopic solid . KO₂ is orange paramagnetic powder . K₂O₂ is a yellow hygroscopic solid and highly poisonous substance. It is more poisonous than NaCN.
• Potassium was first isolated by Davy in 1808 . It is soft, low malting (mp = 63°C) and more reactive than sodium . Carnallite , KCl . MgCl₂. 6H₂O is the important mineral source of potassium . In India potassium is found largely as saltpetre (KNO₃ ). The chemical properties of potassium closely resemble those of sodium, but its reactions are more vigorous. Potassium is one of the few elements that form a super oxide (O₂^-) on reaction with air.
• An alloy of potassium with sodium (Na = 78% , K = 22%) is a liquid resembling mercury, It is low melting (mp = 13°C ) and high boiling (b.p 1450° C ) . It is used as a high temperature heat exchange fluid for use in power plants. At 100°C , it is less corrosive than steam and has a much higher thermal conductivity . Since potassium attack the glass electrodes, the same can not prepare potassium those of sodium.
• The potassium ion (K⁺ ) is chemically inert . The compounds of potassium are, in general more soluble than those of sodium. The main uses of potassium compounds are as fertilizers, where KCl and KNO₃ are commonly used.
• Potassium salts of fatty acids are used in making soft soaps, because they are more soluble than those of sodium . Potassium is mainly obtained from deposits of NaCl and form brines in various lakes.
  
• Potassium salts are occasionally called potash, independent of the associated anion.
• Potassium ion gives a characteristic purple colour in a Bunsen flame.
• Yellow potassium sodium hexa nitro cobaltate (III) , K₂ Na [ Co(NO₂)₆ ] and yellow potassium hexachloroplatinate , K₂ Pt Cl₆ are the less soluble potassium salts.
• Black ash is impure Na₂ CO₃ produced in Leblanc process when salt cake is reduced with coke
• Hypo is Na₂ S₂O₃ . 5H₂O.
• Fire extingushers contain H₂ SO₄ , Na₂ CO₃ and NaHCO₃.
• Soda lime is NaOH + CaO .
• When compared with alkali metals
  (a) They are less reactive than alkali metals
  (b) They are less electropositive than alkali metals. Hence they are less metallic than alkali metals
  (c) Their reducing power is much less than those of alkali metals
  (d) They are less basic than alkali metals.
  
• Alkaline earth metals become more reactive as the outer electrons are further remote from the nucleus. Thus barium is the most reactive member of the family, excluding radium, which is radioactive. Be is less electropositive and less and less likely to lose its two outer electrons than the other elements because of its small size. Many of the compounds of Be are not ionic, but polar covalent.
• All the elements react with oxygen when heated in the gas forming the monoxide of the type MO. Be and Mg react less readily because of the formation of a protective oxide film on their surface. Barium also forms some peroxide alongwith monoxide. If the reaction is carried out under pressure, peroxides are formed both by Sr and Ba. On heating in air the elements form a mixture of both oxides and nitrides.
• On heating with nitrogen all the elements of the group form nitrides of the type M₃ N₂. The nitrides contain N^3- ion and are ionic nature.
• With the exception of Be, all alkaline earth metals react directly with hydrogen to form ionic hydrides of the type MH_2.
• Since the alkaline earth metal ions have rare gas configuration , they are colorless and diamagnetic.
• All alkaline earth metals form halides of the type MX₂. All halides of Be are covalent and electron deficient. Anhydrous beryllium halides are polymeric and contain three central bonds. MgBr₂ and MgI₂ are soluble in acetone because of their covalent nature.
• The hydroxides may be obtained by slaking the oxides with water. MgO slakes very slowly, but the oxides of Ca, Sr and Ba slake readily.
• BeO as well as Be (OH) ₂ are amphoteric, Mg(OH) ₂ is mild base and its aqueous suspension, known as milk of magnesia is used as an antacid. The Ca(OH) ₂ and Sr(OH)₂ are moderately strong bases, while Be(OH)₂ is almost as strong as the alkali hydroxides. All the oxides, BeO, MgO, CaO, SrO and BaO have
  NaCl - type structure (4: 4 Coordination).
• Beryllium halide are covalent, hygroscopic and fumes in air because of hydrolysis. BeF₂ is one of the few metal fluorides, which does not ionize completely in solution. BeF₂ is soluble in water, MgF₂ is sparingly soluble, while CaF₂ , SrF₂ and BaF₂ are insoluble . All other halides are soluble. The solubility decreases from BeF₂ to BaF₂ due to difference in their structures. BeF₂ has a linear polymeric chain structure, while others have crystal lattice structures. In the vapor state BeCl₂ molecule is linear and has no dipole moment. This corresponds to Be in sp-valenece state. Other halides are hygroscopic and form hydrates of the type MgCl₂. 6H₂O, BaCl₂. 2H₂O, CaCl₂. 6H₂O etc.
  
• The alkaline earth metals generally dissolve in dilute mineral acids, but Be is rendered passive by nitric acid. With acid solutions all of them react to yield the respective salts. Most salts of alkaline earth metals (Ca, Sr, and Ba) such as fluorides, sulphates and phosphates are sparingly soluble in water.
• Beryllium dissolves in cold conc. aqueous alkali evolving hydrogen and forming alkaline earth metal beryllate.
• Be + OH^- + H₂O HBeO₃^- + H₂ ­
• Other elements of the group do not dissolve in alkali.
• Sulphates, carbonates and hydroxides of these metals decompose on heating to yield oxides.
• Nitrates of these metals are water soluble and decompose to oxide, NO₂ and O₂ gases.
• When Ca, Sr and Ba are dissolved in liquid ammonia, a paramagnetic, highly reducing and highly conducting deep blue colored solution is obtained . These properties are due to the presence of solvated electrons in solution.
• Marie and Pierre Curie discovered radium in 1898. They produced pure radium in 1911. It was the first radioactive element to be discovered. Radium is extracted from pitchblende, a uranium ore. About 7 tonnes of pitchblende are required to produce 1g of pure radium. All isotopes of radium decay spontaneously to other elements, emitting dangerous radiations. The final product of the nuclear disintegration is lead.
• Barium and strontium are found frequently as sulphates, BaSO₄ and SrSO₄ or the carbonates, BaCO₃ and SrCO₃. Magnesium occurs in the earth crust as magnesite (MgCO₃), dolomite (CaCO₃), carnallite (KCl, MgCl₂. 6H₂O), Epsom salt (MgSO₄. 7H₂O) and asbestos CaMg₃ (SiO₃) ₄. Magnesium ions are present in seawater to the extent of 2%. Magnesium is also present in chlorophyll.
• Beryllium forms the carbide Be₂C while other metals give ionic carbides of the MC₂ type. When reacted with water Be₂C liberates methane, while other carbides liberate acetylene.Mg₂C₃ liberates propyne.
• The nitride of Be, (Be₃N₂) is volatile, while nitrides of other metals are non-volatile.
• The tendency of common salt (NaCl) to get sticky in humid summer is partly due to the presence of small amounts of MgCl₂ in it.
• Heating gypsum at 120° C 5°C gives Plaster of Paris (CaSO₄. 1/2H₂O) which combines vigorously with water and sets to a hard mass. Setting of plaster of Paris is accompanied by evolution of heat (exothermic reaction), catalyzed by common salt (NaCl) and retarded by borax and alum. Setting of plaster of Paris
  is due to hydration as well as transition. Hardening of plaster of Paris is because of transition of orthorhombic gypsum to monoclinic gypsum. Strength of plaster of Paris on hardening is because of interlacing of needles of monoclinic gypsum On setting, plaster of Paris undergoes expansion .
• Gypsum or alabaster or selenite is calcium sulphate dihydrate, CaSO₄. 2H₂O. Heating gypsum at 200°C gives dead burnt (CaSO₄). Heating gypsum at 400°C gives CaO, O₂ and SO₂.
• When Ca CO₃ is heated to high temperature, it decomposes to form CO₂ and CaO, commonly known as lime or quick lime. Quick lime reacts vigorously with water to form a strong base, Ca (OH) _2, which is much less soluble in water than Ba (OH) ₂.The formation of Ca (OH)₂ is an exothermic reaction . Ca (OH) ₂ is known as slaked lime because it is formed when CaO has slaked its thirst for water. When mixed with sand, Ca (OH) ₂ hardens as mortar (1 part lime + 3parts sand) and cement, by absorbing CO₂ from air.
• Barium is not found free in nature. In combined state it occurs as Barytes or heavy spar, BaSO₄ and BaCO₃. Barium burns in air producing mostly BaSO₄. It also burns in CO₂ to form BaO and carbon or CO. Barium is strong reducing agent (SRP Ba²⁺ Ba = -2.90 V ) . It is silvery white metal and m.p. is 850° C . Barium is good conductor of heat and electricity. Barium forms alloys with other metals and these alloys are used in vacuum tubes and sparking plugs, where the low ionization energy of Ba facilitates of good sparks. Barium forms two oxides, BaO and BaO₂. BaO₂ is peroxide, containing O₂^2- ion.
• Magnesium is used as a de-oxidizer in metallurgy and as a fuse in aluminothermite process. It is also used in flash bulbs, pyrotechnics and in fireworks. Magnalium is an alloy of Al and Mg. Elecktron is an alloy of Mg and Zn.