Quantum Numbers, Atomic Orbitals, and Electron Configurations PDF

Quantum Numbers, Atomic Orbitals, and Electron Configurations Each electron in an atom is described by four different quantum numbers. The first three (n, l, ml) specify the particular orbital of interest, and the fourth (ms) specifies how many electrons can occupy that orbital. 1. Principal Quant...

Quantum Numbers, Atomic Orbitals, and Electron Configurations Each electron in an atom is described by four different quantum numbers. The first three (n, l, ml) specify the particular orbital of interest, and the fourth (ms) specifies how many electrons can occupy that orbital. 1. Principal Quantum Number (n): n = 1, 2,…,8. This number describes the average distance of an electron from the nucleus, like the innermost electron shell, which has a principal quantum of 1.This specifies the energy of an electron and the size of the orbital. All orbitals that have the same value of n are said to be in the same shell (level). For a hydrogen atom with n=1, the electron is in its ground state; if the electron is in the n=2 orbital, it is in an excited state. The total number of orbitals for a given n value is n2. Orbitals for which n=2 are larger those for which n=1, for example. Because they have opposite electrical charges, electrons are attracted to the nucleus of the atom. Energy must therefore be absorbed to excite an electron from an orbital in which the electron is close to the nucleus (n=1) into an orbital in which it is further from the nucleus (n=2). The principal quantum number therefore indirectly describes the energy of an orbital. As n increases, the energies of the orbitals also increase. The principal quantum number n may have any positive integral value between n=1 and n= ∞. The n corresponds to the orbital energy level or ‘‘shell”. The shell with n=1 is called the 1st shell, the shell with n=2 is called the second shell and so forth. For a particular energy level, there may be subshell, the number of which is defined by the quantum number l. 2. Angular or orbital Quantum Number (l): l = 0,..., n-1. Specifies the shape of an orbital with a particular principal quantum number. The angular quantum number divides the shells into smaller groups of orbitals called subshells (sublevels). Usually, a letter code is used to identify l to avoid confusion with n: l 0 1 2 3 4 5... Letter s p d f g h... The subshell with n=2 and l=1 is the 2p subshell; if n=3 and l=0, it is the 3s subshell, and so on. The value of l also has a slight effect on the energy of the subshell; the energy of the subshell increases with l (s < p < d < f) 3. Magnetic Quantum Number (ml): ml = -l,..., 0,..., +l. Specifies the orientation in space of an orbital of a given energy (n) and shape (l). This number divides the subshell into individual orbitals which hold the electrons; there are 2l+1 orbitals in each subshell. Thus the s subshell has only one orbital, the p subshell has three orbitals, and so on 4. Spin Quantum Number (ms): ms = +½ or -½. Specifies the orientation of the spin axis of an electron. An electron can spin in only one of two directions (sometimes called up and down). The Pauli exclusion principle (Wolfgang Pauli, Nobel Prize 1945) states that no two electrons in the same atom can have identical values for all four of their quantum numbers. What this means is that no more than two electrons can occupy the same orbital, and that two electrons in the same orbital must have opposite spins. Because an electron spins, it creates a magnetic field, which can be oriented in one of two directions. For two electrons in the same orbital, the spins must be opposite to each other; the spins are said to be paired. These substances are not attracted to magnets and are said to be diamagnetic. Atoms with more electrons that spin in one direction than another contain unpaired electrons. These substances are weakly attracted to magnets and are said to be paramagnetic  Derive possible sets of quantum numbers for n=2 and explain the significance of these sets of numbers Let n=2 n-denotes an energy level l lies in the range 0 to (n-1) Therefore for n=2 l=0----(n-1) l=0-----(2-1) l=0----1 l=0 and 1. This means that the n=2 level gives rise to 2 sub-levels, one with l=0 and one with l=1 Now determine the possible values of quantum number ml Values of ml lie in the range -l ----0---+l When l=0 ml=-0----+0= 0 When l=1 ml=-1--0--+1=-1, 0, +1 Write down two possible sets of quantum numbers that describe an electron in a 2s atomic orbital. What is the physical significance of these unique sets? The 2s atomic orbital is defined by the set of quantum numbers n=2, l = 0, ml = 0. An electron in a 2s atomic orbital may have one of two sets of four quantum numbers: n=2, l = 0, ml = 0; ms = + 1/2 or n=2, l = 0, ml = 0; ms = - 1/2 If the orbital were fully occupied with two electrons, one electron would have ms = 1 + /2, and the other electron would have ms = - 1/ , i.e. the two electrons 2 would be spin paired. The Ground-state electronic configuration is the most probable or the most energetically favoured configuration. The Aufbau process Consider the following hypothetical process- the building up of more complex atom starting with the simplest atom, hydrogen. This hypothetical process is called Aufbau process (meaning building up in German) In this process we proceed from an atom of one element to the next by adding a proton and the requisite number of neutrons to the nucleus and one electron to the appropriate orbital. We pay particular attention to this added electron, called the differentiating electron. Hydrogen, Z =1. The lowest energy state for the electron in a hydrogen atom is the 1s orbital. The electronic configuration is 1s 1 Helium, Z =2. In the helium atom a second electron goes into the 1s orbital. The two electrons must have opposing spins, 1s2. Lithium Z=3. The differentiating electron cannot be accommodated in the 1s orbital (Pauli exclusion principle). It must be placed in the next lowest energy orbital 2s. The electron configuration is 1s22s1 Beryllium Z=4  SUMMARY Electrons in an atom are located in defined regions called electron shells, which surround the nucleus. This arrangement of electrons is referred to as the electron configuration.  There are ‘rules’ which determine how electron shells are filled, and how many electrons they can contain: Inner shells begin filling first; they are smaller and can hold less electrons.  A maximum of 2 electrons can occupy the first shell. A maximum of 8 electrons can occupy the second shell.  A maximum of 18 electrons can occupy the third shell, but the fourth shell will begin to fill once the third shell contains 8 electrons.  A maximum of 8 electrons can occupy the valence shell (outermost shell) of any atom, unless the valence shell is the only shell, in which case there can be a maximum of 2 electrons. The electron configuration of an atom can be written as the numbers of electrons in each shell, separated by a comma.   Periodicity Periodicity refers to trends or recurring variations in element properties with increasing atomic number. Periodicity is caused by regular and predictable variations in element atomic structure. Mendeleev organized elements according to recurring properties to make a periodic table of elements. Elements within a group (column) display similar characteristics. The rows in the periodic table (the periods) reflect the filling of electrons shells around the nucleus, so when a new row begins, the elements stack on top of each other with similar properties. For example, helium and neon are both fairly unreactive gases that glow when an electric current is passed through them. Lithium and sodium both have a +1 oxidation state and are reactive, shiny metals.  Elements were first arranged in order of increasing atomic weight by Dimitri Mendeleev. He observed that elements with similar properties appeared periodically at regular intervals. The earliest version of the periodic table was constructed so that elements with similar properties fell into vertical columns. In the modern form of the periodic table, elements are arranged in order of increasing atomic numbers (i.e. the number of protons in the atom). The periodic table illustrates the periodic law. When elements are arranged in ascending atomic number, similar chemical and physical properties recur periodically.  The reason why the properties of elements are related periodically to their atomic numbers is that, the atomic number of an element determines its electronic configuration, which in turn determines its atomic properties (like atomic radius, ionization enthalpy, electronegativity e.t.c.). Furthermore, the atomic properties directly affect the type of bonding and structure of the element. The periodic table is divided into 7 horizontal rows and 8 vertical columns. The horizontal rows are called periods which are numbered from the top downwards (period 1, period 2,……… period 7). Elements in the same period have the same number of occupied electron shell(s). The vertical columns are called groups and they are numbered from left to right (Group 1A, IIA,…….Group 0). Elements in a particular group have the same number and arrangement of the outermost shell electron(s).  The periodic table can also be divided into four blocks namely, s-, p-, d-, and f- blocks. S-block Group IA and IIA elements form the s-block as their outermost shell electrons are located in the s-subshell. P-block Elements of groups IIIA to O are known as p-block elements. It is because of their outermost shell electrons are located in the p- subshell. Elements in the s- and p- blocks are known collectively as the representative elements. However, this is not applicable to elements in group O. They are called noble gas elements. D-block D-block elements have their highest energy electrons in the inner d-subshell. They are also called transition elements. F-block F-block elements have electrons filling the inner f- subshell. There are two series of f-block elements, the lanthanide series {Z=58-71} and actinide series {Z= 90- 103}. They are also called the inner-transition elements. Uses of Periodicity Periodicity was helpful to Mendeleev because it showed him gaps in his periodic table where elements should be. This helped scientists find new elements because they could be expected to display certain characteristics based on the location they would take in the periodic table. Properties That Display Periodicity Periodicity can include many different properties, but the key recurring trends are: Ionization Energy - This is the energy needed to completely remove an electron from an atom or ion. Ionization energy increases moving left to right across the table and decreases moving down a group. Electronegativity- A measure of how readily an atom forms a chemical bond. Electronegativity increases moving left to right across a period and decrease moving down a group.  Atomic Radius- This is half the distance between the middle of two atoms just touching each other. Atomic radius decreases moving left to right across a period and increases moving down a group.  Moving Left → Right Ionization Energy Increases Electronegativity Increases Atomic radius Decreases  Moving Top → Bottom Ionization Energy Decreases Electronegativity Decreases Atomic Radius Increases Occurrence and extraction of metals Metals and their alloys are extensively used in our day- to-day life. They are used for making machines, railways, motor vehicles, bridges, buildings, agricultural tools, aircrafts, ships etc. Therefore, production of a variety of metals in large quantities is necessary for the economic growth of a country. Only a few metals such as gold, silver, mercury etc. occur in free state in nature. Most of the other metals, however, occur in the earth's crust in the combined form, i.e., as compounds with different anions such as oxides, sulphides, halides etc. In view of this, the study of recovery of metals from their ores is very important. Some of the processes of extraction of metals from their ores are called metallurgical processes. Occurrence of Metals Metals occur in nature in free as well as combined form. Metals having low reactivity show little affinity for air, moisture, carbon dioxide or other non-metals present in nature. Such metals may remain in elemental or native (free) state in nature. Such metals are called "noble metals" as they show the least chemical reactivity. For example gold, silver, mercury and platinum occur in free state. On the other hand, most of the metals are active and combine with air, moisture, carbon dioxide and non-metals like oxygen, sulphur, halogens, etc. to form their compounds, like oxides, sulphides, carbonates, halides and silicates. i.e., they occur in nature in a combined state. A naturally occurring material in which a metal or its compound occurs is called a mineral. An ore is a material that contains a sufficiently high concentration of a mineral to constitute an economically feasible source from which the metal can be recovered. The main active substances present in nature, especially in the atmosphere are oxygen and carbon dioxide. In the earth's crust, sulphur and silicon are found in large quantities. Sea-water contains large quantities of chloride ions (obtained from dissolved sodium chloride). Most active metals are highly electropositive and therefore exist as ions. It is for this reason that most of the important ores of these metals occur as (i) oxides (ii) sulphides (iii) carbonates (iv) halides and (v) silicates. Some sulphide ores undergo oxidation by air to form sulphates. This explains the occurrence of sulphate ores. Ores are invariably found in nature in contact with rocky materials. These rocky or earthy impurities accompanying the ores are termed as gangue or matrix Type of Ore Metals (Common Ores) Native Metals Gold (Au), silver (Ag) Oxide ores Iron (Haematite, Fe2O3 ); Aluminium (Bauxite, Al2O3. 2H2O); Tin (Cassiterite, SnO2); Copper (Cuprite, Cu2O); Zinc (Zincite, ZnO); Titanium (Ilmenite, FeTiO3 , Rutile, TiO2) Type of Ore Metals (Common Ores) Sulphide ores Zinc (Zinc blende, ZnS); Lead (Galena, PbS); Copper (Copper glance, Cu2S); Silver (Silver glance or Argentite, Ag2S); Iron (Iron pyrites, FeS2) Carbonate ores Iron (Siferite, FeCO3); Zinc (Calamine, ZnCO3) , Lead (Cerrusite, PbCO3) Type of Ore Metals (Common Ores) Sulphate ores Lead (Anglesite, PbSO4) Halide ores Silver (Horn silver, AgCl); Sodium (Common salt or Rock salt, NaCl); Aluminium (Cryolite, Na3AlF6) Silicate ores Zinc(Hemimorphite, 2ZnO.SiO2.H2O) The process of extracting the metals from their ores and refining them is called metallurgy. The choice of the process depends upon the nature of the ore and the type of the metal. The metal content in the ore can vary depending upon the impurities present and chemical composition of the ore. Some common steps involved in the extraction of metals from their ores are : (i) Crushing and pulverization (ii) Concentration or dressing of the ore (iii) Calcination or roasting of the ore (iv) Reduction of metal oxides to free metal (v) Purification and refining of metal. Crushing and Pulverization The ore is generally obtained as big rock pieces. These big lumps of the ore are crushed to smaller pieces by using jaw- crushers and grinders. It is easier to work with crushed ore. The big lumps of the ore are brought in between the the plates of a crusher forming a jaw. One of the plates of the crusher is stationary while the other moves to and fro and the crushed pieces are collected below. The crushed pieces of the ore are then pulverized (powdered) in a stamp mill. The heavy stamp rises and falls on a hard die to powder the ore. The powdered ore is then taken out through a screen by a stream of water. Concentration or Dressing of the Ore Generally, the Ores are found mixed with earthy impurities like sand, clay, lime stone etc. These unwanted impurities in the Ore are called gangue or matrix. The process of removal of gangue from powdered Ore is called concentration or dressing. There are several methods for concentrating the Ores. The choice of method depends on the nature of the Ore. Some important methods are : (i) Gravity separation (Hydraulic washing) : In this method, the light (low specific gravity) impurities are removed from the heavier metallic Ore particles by washing with water. It is therefore, used for the concentration of heavier oxide Ores, like haematite (Fe2O3), tinstone (SnO2 ) and gold (Au). In this method, as shown in the diagram below, the powdered Ore is agitated with water or washed with a strong current of water. The heavier Ore settles down rapidly in the grooves and the lighter sandy and earthy materials (gangue particles) are washed away (ii) Magnetic separation method : By this method, those ores can be concentrated which either contain impurities which are magnetic or are themselves magnetic in nature. For example, the tin Ore, tin stone (SnO2) itself is non- magnetic but contains magnetic impurities such as iron tungstate (FeWO4) and manganese tungstate (MnWO4) The finely powdered Ore is passed over a conveyer belt moving over two rollers, one of which is fitted with an electromagnet. The magnetic material is attracted by the magnet and falls in a separate heap. In this way magnetic impurities are separated from non-magnetic material. (iii) Froth floatation method : This method is especially applied to sulphide ores, such as galena (PbS), zinc blende (ZnS), or copper pyrites (CuFeS2). It is based on the different wetting properties of the surface of the Ore and gangue particles. The sulphide ore particles are wetted preferentially by oil and gangue particles by water. In this process, finely powdered ore is mixed with either pine oil or eucalyptus oil. It is then mixed with water. Air is blown through the mixture with a great force. Froth is produced in this process which carries the wetted ore upwards with it. Impurities (gangue particles) are left in water and sink to the bottom from which these are drawn off. (iv) Chemical method : In this method the Ore is treated with a suitable chemical reagent which dissolves the Ore leaving behind insoluble impurities. The ore is then recovered from the solution by a suitable chemical method. This is applied for extraction of aluminium from bauxite (Al2O3.2H2O). Bauxite is contaminated with iron (III) oxide (Fe2O3), titanium (IV) oxide (TiO2) and silica (SiO2). These impurities are removed by digesting the powdered Ore with aqueous solution of sodium hydroxide at 420 K under pressure. Aluminium oxide dissolves in sodium hydroxide, whereas, iron (III) oxide, silica and titanium (IV) oxide remain insoluble and are removed by filtration. Al2O3 + 6NaOH→ 2Na3AlO3 + 3H2O Calcination and Roasting of the Ore The concentrated Ore is converted into metal oxide by calcination or roasting. (A) Calcination : Calcination involves heating of the concentrated ore in a limited supply of air so that it loses moisture, water of hydration and gaseous volatile substances. The Ore is heated to a temperature so that it does not melt. Two examples of calcination are given below: (i) Removal of water of hydration Al2O3.2H2O →Al2O3+2H2O (ii) Expulsion of CO2 from carbonate ZnCO3 →ZnO + CO2 (B) Roasting : Roasting is a process in which the concentrated Ore is heated in a free supply of air at a temperature insufficient to melt it. The following changes take place during roasting (i) Drying of the Ore. (ii) Removal of the volatile impurities like arsenic, sulphur, phosphorus and organic matter. 4As + 3O2 → 2As2O3(g) S + O2 → SO2(g) 4P + 5O2 → P4O10(g) (iii) Conversion of the sulphide ores into oxides 2PbS + 3O2 →2PbO + 2SO2 2ZnS + 3O2 →2ZnO + 2SO2 Calcination and roasting are generally carried out in a reverberatory furnace or in a multiple hearth furnace Reduction of the Metal Oxides to Free Metal This process is carried out after calcination or roasting of ores. In this process called smelting, the oxide ores are converted into the metallic state by reduction  (A) Smelting : Smelting is a process in which the oxide ore in molten state is reduced by carbon or other reducing agents to free metal (i) by using carbon as a reducing agent : This method is used for the isolation of iron, tin and zinc metals from their respective oxides. The oxide ores are strongly heated with charcoal or coke. Reduction occurs by the action of carbon and/or carbon monoxide which is produced by the partial combustion of coke or charcoal. Fe2O3 + 3C→ 2Fe + 3CO Fe2O3 + CO→ 2FeO + CO2 FeO + CO→ Fe + CO2 SnO2 + 2C →Sn + 2CO ZnO + C → Zn + CO Although the ore has been concentrated in an earlier step, it is still contaminated with some gangue material which is finally removed in the reduction process by the addition of flux during smelting. Flux is a chemical substance which combines with impurities at higher temperatures to form slag. which is not soluble in the molten metal. Flux are of two types Basic flux and acidic flux Basic Flux : On heating, lime stone is converted into calcium oxide used as basic flux which combines with acidic impurities like silica in metallurgy of iron and forms fusible calcium silicate CaSiO3 CaCO3 → CaO + CO2(g) (Limestone) CaO + SiO2 → CaSiO3 (Basic flux) (Acid gangue) Slag Acidic flux : SiO2 is used as acidic flux to remove basic impurity of FeO in metallurgy of Cu.  SiO2 + FeO → FeSiO3 (Acidic flux) (Basic gangue) Slag The slag, such as calcium silicate formed during smelting floats over the molten metal and is thus easily removed. Another advantage is that the slag provides a covering to the molten metal thus preventing it from getting oxidized by air. (ii) Other reducing agents : Oxide ores which cannot be reduced by carbon or metals which show affinity to carbon by forming metal carbides, are reduced by reducing agents like aluminium, sodium, magnesium or hydrogen. Oxide like chromium oxide (Cr2O3) or manganese oxide (Mn3O4) are reduced by aluminium powder is a highly exothermic reaction. This process is known as Goldschmidt's Alumino- thermite reduction method. Cr2O3 + 2Al→2Cr + Al2O3 + Heat 3Mn3O4 + 8Al→ 9Mn +4Al2O3 + Heat Heat is generated in the process due to the formation of Al2O3 which is a highly exothermic reaction. Titanium is obtained by the reduction of TiCl4 (produced by the action of carbon and chlorine on TiO2) by Mg in an inert atmosphere of argon (Kroll process). Heat TiCl4 +2Mg → Ti + 2MgCl2 1103K Titanium can also be obtained by the reduction of TiO2 by sodium TiO2 + 4Na→ Ti + 2Na2O Tungsten and molybdenum can be obtained by the reduction of their oxides by Hydrogen MoO3 + 3H2 → Mo + 3H2O (iii) Self-reduction : This is applied to the sulphide ores of copper, mercury and lead. The ores are heated in air, a part of these sulphide ores is changed into the oxide or sulphate which then reacts with the remaining part of the sulphide ore to give the metal and sulphur dioxide.  The reactions showing their extraction are given below : (i)2Cu2S + 3O2 → 2Cu2S + Copper glance 2SO2 2Cu2O + Cu2S → 6Cu + SO2 Copper produced at this stage is called blister copper. The evolution of sulphur dioxide produces blisters on the surface of solidified copper metal. (2). 2HgS + 3O2 → 2HgO + Cinnabar 2SO2 2HgO + HgS → 3Hg + SO2 (3). 2PbS + 3O2 → 2PbO + Galena 2SO2 PbS + 2O2 → PbSO4 PbS + 2PbO → 3Pb + SO2 PbS + PbSO4 → 2Pb + 2SO2 (B). Reduction of concentrated ores by other methods: Some metals cannot be obtained from their ores by using common reducing agents such as C, CO, H2 etc. Other methods of reduction are used for such cases. (i) Reduction by precipitation : Noble metals like silver and gold are extracted from their concentrated ores by dissolving metal ions in the form of their soluble complexes. The metal ions are then regenerated by adding a suitable reagent. For example, concentrated argentite ore (Ag2S) is treated with a dilute solution of sodium cyanide (NaCN) to form a soluble complex Ag2S+4NaCN→ 2Na[(AgCN)2] + Na2S This solution is decanted off and treated with Zn to precipitate silver 2Na[(AgCN)2] + Zn → Na2[(ZnCN)4] + 2Ag↓ (ii) Electrolytic Reduction : Active metals like sodium, potassium and aluminium etc., are extracted by the electrolysis of their fused (molten) salts. For example, sodium is obtained by the electrolysis of fused sodium chloride (Down's process). The reactions taking place in the electrolytic cell are NaCl + Na + Cl− Na+ ions move towards the cathode and Cl− ions move towards the anode. Following reactions take place at the electrodes : At the Cathode Na+ + e−→Na(Reduction) (negative electrode) (metal) At the Anode Cl−→ Cl + e− (Oxidation) (positive electrode) Cl + Cl→ Cl2 Aluminium is extracted from molten alumina (Al2O3) by electrolysis. The melting point of alumina is quite high (2323K) which is inconvenient for electrolysis. It dissolves in molten cryolite (Na3AlF6) at around 1273 k. The reactions  which take place in the cell are:. At the Cathode Al3+ +3e−→ Al(metal) At the Anode C + 2O2− →CO2 + 4e− Refining of Metals Except in the electrolytic reduction method, metals produced by any other method are generally impure. The impurities may be in the form of (i) other metals (ii) unreduced oxide of the metal (iii) non-metals like carbon, silicon, phosphorus, sulphur etc. and (iv) flux or slag. Crude metal may be refined by using one or more of the following methods : (i) Liquation : Easily fusible metals like tin, lead etc. are refined by this process. In this method, the impure metal is poured on the sloping hearth of a reverberatory furnace and heated slowly to a temperature little above the melting point of the metal. The pure metal drains out leaving behind infusible impurities. Liquidation (ii) Poling : Poling involves stirring the impure molten metal with green logs or bamboo. The hydrocarbons contained in the pole reduce any metal oxide present as impurity. Copper and tin are refined by this method. Polling (iii) Distillation : Volatile metals like zinc and mercury are purified by distillation. The pure metal distils over, leaving behind non-volatile impurities. (iv) Electrolytic Refining : A large number of metals like copper, silver, zinc, tin etc. are refined by electrolysis. A block of impure metal is made the anode and a thin sheet of pure metal forms the cathode of the electrolytic cell containing suitable metal salt solution which acts as an electrolyte. On passing current, pure metal deposits at the cathode sheet while more electropositive impurities are left in solution. Less electropositive metals do not dissolve and fall away from the anode to settle below it as anode mud. Electrolytic Refining For example, in the electrolytic refining of crude copper (blister copper), a large piece of impure copper is made anode and a thin piece of pure copper is made the cathode. An acidified solution of copper sulphate is used as an electrolyte. On passing an electric current of low voltage through the solution copper (II) ions obtained from copper sulphate solution go to the cathode where they are reduced to the free copper metal and get deposited Cu2+ + 2e → - Cu (at cathode) An equivalent amount of the metal from the anode dissolves into the electrolyte as Cu2+ ions Cu → Cu2+ + 2e- (at anode) As the process goes on, anode becomes thinner while the cathode becomes thicker. The impurities like silver, gold settle down at the bottom of the cell as 'anode mud'.  Group one-The Alkali metals Element Symbol Electronic Structure Lithium Li 1s22s1 Sodium Na 1s22s22p63s1 Potassium K 1s22s22p63s23p64s1 Rubidium Rb 1s22s22p63s23p64s24P65s1 Caesium Cs 1s22s22p63s23p64s24P65s15P66s1 Francium Fr [Rn]7s1  Occurrence and Abundance Despite their close chemical similarity, the elements do not occur together, mainly because their ions are of different size. Sodium is the most abundant, followed by potassium, rubidium, lithium, and caesium. Francium is intensely radioactive and very rare.  Lithium is the thirty-fifth most abundant element by weight and is mainly obtained as the silicate minerals, spodumene LiAl(SiO3)2 and lepidolite Li2Al2(SiO3)3(FOH)2.  Sodium and potassium are the seventh and eighth most abundant elements by weight in the earth’s crust. The largest source of sodium is rock salt (NaCl). Various salts including NaCl, Na2B4O7.10H2O (borax), NaNO3 (saltpetre) and Na2SO4 (mirabilite) are obtained from deposits formed by the evaporation of ancient seas.  Potassium occurs mainly as deposits of KCl (sylvite), a mixture of KCl and NaCl (Sylvinite), and the double salt KCl.MgCl2.6H2O (carnalite).  There is no convenient source of rubidium and only one of caesium and these elements are obtained as by-products from lithium processing.  All of the elements heavier than bismuth (atomic number 83) 83Bi are radioactive. Thus francium (atomic number 89) is radioactive and has a short half-life period of 21 minutes it does not occur appreciably in nature.  Extraction of metals The metals may all be isolated by electrolysis of a fused salt, usually the fused halide, often with impurity added to lower the melting point. Sodium is made by the electrolysis of a molten mixture of about 40% NaCl and 60%CaCl2 in a Downs cell.This mixture melts at about 600C compared with 803C for pure NaCl. The small amount of calcium formed during the electrolysis is insoluble in the liquid sodium, and dissolves in the eutectic mixture.  There are three advantages to electrolyzing It lowers the melting point and so reduces the fuel bill. The lower temperature results in a lower vapour pressure for sodium, which is important as sodium vapour ignites in air.  At lower temperature the liberated sodium metal does not dissolve in the melt, and this is important  because if it dissolved it would short-circuit the electrodes and thus prevent further electrolysis. A Downs cell comprises a cylindrical steel vessel lined with firebrick, measuring about 2.5m in height and 1.5m in diameter.  The anode is a graphite rod in the middle, and is surrounded by a cast steel cathode. A metal gauze screen separates the two electrodes, and prevents the Na formed at the cathode from recombining with Cl2 produced at the anode. The molten sodium rises, as it is less dense than the electrolyte, and it is collected in an inverted trough  and removed, and packed into steel drums. A similar cell can be used to obtain potassium by electrolyzing fused KCl. However, the cell must be operated at a higher temperature because the melting point of KCl is higher, and this results in the vapourization of the liberated potassium. Since sodium is readily available, the modern method is to reduce molten KCl with sodium vapour at 850 C in a large fractionating tower. This gives K of 99.5% purity. Na +KCl →NaCl +K Rb and Cs are produced in a similar way by reducing the chlorides with Ca at 750C under reduced pressure. All react with water to give hydrogen gas and the metal hydroxide. They also react with the oxygen in the air to give either an oxide, peroxide or superoxide depending on the metal. These metals almost always form ions with a positive (+1) charge.  Most of the alkali metals glow with a characteristic color when placed in a flame; lithium is crimson, sodium gives off an intense yellow, potassium is Lilac, rubidium is a red-violet, and caesium gives off blue light.  Electronic structure Group 1 elements all have one valence electron in their outer orbital- an s electron, which occupies a spherical orbital. The single valence electron is a long distance from the nucleus, is only weakly held and is readily removed. In contrast the remaining electrons are closer to the nucleus, more tightly held, and are removed only with great difficulty. Because of similarities in the electronic structures of these elements, many similarities in chemical behaviour would be expected.  Size of atoms and ions Group 1 atoms are the largest in their periods in the periodic table. When the outer electron is removed to give a positive ion, the size decreases considerably. There are two reasons for this. 1). The outermost shell of electrons has been completely removed. 2). Having removed an electron, the positive charge on the nucleus is now greater than the charge on the remaining electrons, so that each of the remaining electrons is attracted more strongly towards the nucleus. This reduces the size further.  Positive ions are always smaller than the parent atom. Even so, the ions are very large, and they increase in size from Li+ to Fr+ as extra shells of electrons are added. The Li+ is much smaller than the other ions. For this reason, Li only mixes with Na above 380°C and it is immiscible with the metals K, Rb and Cs, even when molten. In contrast the other metals Na, K, Rb and Cs are miscible with each other in all proportions. Density The atoms are large, so group 1 elements have remarkably low densities.  Metallic Ionic Density radius radius (gcm-3) (Å) M+(Å) Li 1.52 0.76 0.54 Na 1.86 1.02 0.97 K 2.27 1.38 0.86 Rb 2.48 1.52 1.53 Cs 2.65 1.67 1.90 Ionization Energy The first ionization energies for the atoms in this group are appreciably lower than those for any other group in the periodic table. The atoms are very large so the outer electrons are only held weakly by the nucleus.  Hence the amount of energy needed to remove the outer electron is not very large.  On descending the group from Li to Na to K to Rb to Cs, the size of the atoms increases; the outermost electrons become less strongly held, so the ionization energy decreases. Electronegativity The electronegativity values for the elements in this group are very small- in fact the smallest values of any element. Thus when these elements react with other elements to form compounds, a large electronegativity difference between the two atoms is probable, and ionic bonds are formed. Li- 1.0, Na-0.9, K-0.8, Rb-0.8, Cs-0.7(Pauling’s electronegativity). Chemical properties Some reactions of Group 1 metals 2M +2H2O →2MOH +H2 The hydroxides are the strongest bases known. With excess dioxygen 4Li +O2 →2Li2O Monoxide is formed by Li and to a small extent by Na.  2Na+O2→Na2O2 Peroxide is formed by Na and to a small extent by Li K +O2 →KO2 Superoxide formed by K, Rb, Cs.  2M +H2 →2MH ionic ‘salt-like’ hydrides. 6Li +N2 →2Li3N Nitride formed only by Li. 3M +P →M3P All the metals form phosphides  3M +As →M3As All the metals form arsenides 3M +Sb →M3Sb All the metals form stibnides  2M +X →M2X (X=S,Se,Te) All the metals form sulphides, selenides, and tellurides.  2M +X2 →2MX (X=F, Cl, Br, I) All the metals form fluorides, chlorides, bromides, and iodides 2M + 2NH3 →2MNH2 + H2 All the metals form amides  Uses of Lithium: Lithium is used to make electrochemical cell (both primary and secondary batteries). Lithium is used in lubricants, in glass industries, and in alloys of lead, aluminum, and magnesium to make them less dense and stronger. Uses of sodium: Liquid sodium metal is used as a coolant in fast breeder nuclear reactor. Sodium has many biological uses like nerve signal transmission. Sodium nitrite is a principal ingredient in gunpowder. The pulp and paper industry uses large amounts of sodium hydroxide, sodium carbonate, and sodium sulphate.  Sodium carbonate is used by power companies to absorb sulfur dioxide, a serious pollutant, from smokestack gases (locomotive chimneys or ship chimneys). Sodium carbonate is also used in the glass and detergent industries. Sodium chloride is used in foods and to soften the water. Sodium bicarbonate (baking soda) is used in the food industry as well.  Uses of potassium: Potassium is an essential element for life. Roughly 95% of Potassium compounds are used as fertilizers for plants.  Potassium hydroxide is used in detergent. Potassium chlorate is used in explosive.  Potassium carbonate is used in ceramics, colour TV tubes and fluorescent light tubes. Potassium bromide is used in photography industries. Uses of Rubidium and caesium: Rubidium is used almost exclusively for research, but caesium is used in special glasses and radiation detection equipment GROUP TWO-THE ALKALINE EARTH METALS Element Symbol Electronic Structure Beryllium Be 1s22s2 Magnesium Mg 1s22s22p63s2 Calcium Ca 1s22s22p63s23p64s2 Strontium Sr 1s22s22p63s23p63d104s24p65s2 Barium Ba 1s22s22p63s23p63d104s24p64d105s25p66s2 Radium Ra [Rn]7s2  Alkaline earth metals make up the second group of the periodic table. This family includes the elements beryllium, magnesium, calcium, strontium, barium, and radium (Be, Mg, Ca, Sr, Ba, and Ra, respectively). These metals are silver and soft, much like the alkali metals of Group 1.  Each alkaline earth metal has two valence electrons. They will easily give these electrons up to form cations. These metals become increasingly more reactive as you go down the periodic table. This is concurrent with general periodic trends.  The group two elements show the same trends in properties as were observed with Group 1. However, beryllium stands apart from the rest of the group and differs much more from them than lithium does from the rest of Group 1.  The main reason for this is that the beryllium atom and Be2+ are both extremely small, and the relative increase in size from Be2+ to Mg2+ is four times greater than the increase between Li+ and Na+.  Beryllium and barium compounds are all very toxic. The elements form a well-graded series of highly reactive metals, but are less reactive than Group 1. They are typically divalent and generally form colourless ionic compounds.  The oxides and hydroxides are less basic than those of Group 1: hence their oxosalts (carbonates, sulphates, nitrates) are less stable to heat. Occurrence and Extraction These elements are all found in the Earth’s crust, but not in the elemental form as they are so reactive. Instead, they are widely distributed in rock structures.  Beryllium, like its neighbours Li and B is relatively not very abundant in the earth’s crust. It occurs to the extent of about 2ppm and is thus similar to Sn (2.1 ppm), Eu (2.1 ppm) and As (1.8 ppm). Beryllium is found in small quantities as the silicate minerals beryl Be3Al2Si6O18 and phenacite Be2SiO4.  Magnesium is the sixth most abundant element in the earth’s crust (27640 ppm ). The main minerals in which magnesium is found are carnellite (KCl.MgCl2.6H2O), magnesite (MgCO3) and dolomite (MgCO3.CaCO3).  Calcium is the fifth most abundant element in the earth’s crust. Hence the third most abundant metal after Al and Fe. Calcium is found in gypsum (CaSO4.2H2O), anhydrite (CaSO4), fluorite (CaF2), apatite (Ca5(PO4)3F) and limestone (CaCO3) There are two crystalline forms of CaCO3, calcite and aragonite.  Strontium (384ppm) and barium (390ppm) are much less abundant, but are well known because they occur as concentrated ores, which are easy to extract. They are respectively the fifteenth and fourteen element in abundance.  Strontium is mined as celestite SrSO4 and strontianite (SrCO3). Ba is mined as Barytes BaSO4. Radium is extremely scarce and is radioactive. It was first isolated by Pierre and Marie Curie by processing many tons of the uranium ore pitchblende.  Of the elements in this Group only magnesium is produced on a large scale. It is extracted from sea-water by the addition of calcium hydroxide, which precipitates out the less soluble magnesium hydroxide. This hydroxide is then converted to the chloride, which is electrolysed in a Downs cell to extract magnesium metal.  The metals of this group are not easy to produce by chemical reduction because they are themselves strong reducing agents, and they react with carbon to form carbides. They are strongly electropositive and react with water, and so aqueous solutions cannot be used for displacing them with another metal, or for electrolytic production.  All the metals can be obtained by electrolysis of the fused chloride, with sodium chloride added to lower the melting point, although strontium and barium tend to form a colloidal suspension. Properties of the elements  The alkaline earth metals are silvery white, lustrous and relatively soft. Their physical properties when compared with those of group 1A metals, show that they have a substantially higher melting point., boiling point, enthalpies of fusion and vapourization.  They have two valency electrons which may participate in metallic bonding, compared with one electron for Group 1 Metals. Consequently Group 2 metals are harder, have higher cohesive energy and much higher melting points and boiling points than Group 1 elements.  The melting points do not vary regularly, mainly because the metals adopt different crystal structures. Melting points of Group 1 and 2 Melting Pt(°C) Be 1287 Li 181 Mg 649 Na 98 Ca 839 K 6 Sr 768 Rb 39 Ba 727 Cs 28.5  This can be understood in terms of the size factor and the fact that two valency electrons per atom are now available for bonding. Again, Be is notable in melting more than 1100°C above Li and being nearly 3.5 times as dense; its enthalpy of fusion is more than 5times that of Li.  Size of atoms and ions Group 2 atoms are large but are smaller than the corresponding group 1 elements as the extra charge on the nucleus draws the orbital electrons in. Similarly the ions are large, but smaller than those of Group 1, especially because of the removal of two orbital electrons increases the effective nuclear charge further.  Thus, these elements have higher densities than group 1 metals.  Size and Density metallic Ionic Density radius(Å) Radius(Å) (gcm-3) Be 1.12 0.31 1.85 Mg 1.60 0.72 1.74 Ca 1.97 1.00 1.55 Sr 2.15 1.18 2.63 Ba 2.22 1.35 3.62 Ra 1.48 5.5  Comparison with group 1A shows the substantial increase in the ionization energies; this is related to their smaller size and higher nuclear charge and is particularly notable for Be. Ionization energy 1st 2nd 3rd Be 899 1757 14847 Mg 737 1450 7731 Ca 590 1145 4910 Sr 549 1064 Ba 503 765 Ra 509 979  The third ionization energy is so high that M3+ ions are never formed. The ionization energy for Be2+ is high and its compounds are typically covalent, Mg also forms some covalent compounds. However, the compounds formed by Mg, Ca, Sr and Ba are predominantly divalent and ionic.  Since the atoms are smaller than those in Group 1, the electrons are more tightly held so that the energy needed to remove the first electron (1st ionization energy) is greater than those in Group 1.  Once one electron has been removed, the ratio of charges on the nucleus to orbital electrons is increased, so that the remaining electrons are more tightly held. Hence the energy needed to remove a second electron is nearly double that required for the first. ELECTRONEGATIVITY The electronegativity values of Group 2 elements are low, but are higher than the values for Group 1. The electronegativity difference between Group 2 metals (Mg, Ca, Sr, and Ba) and the halogens or oxygen is large and the compounds are ionic. The value for Beryllium is higher than for others. HYDRATION ENERGIES The hydration energies of the Group 2 ions are four or five times greater than for group 1 ions. This is largely due to their smaller size and increased charge. The crystalline compounds of Group 2 contain more water of crystallization than the corresponding Group 1 compounds.  Thus NaCl and KCl are anhydrous but MgCl2.6H2O, CaCl2.6H2O and BaCl2.2H2O all have water of crystallization. Note that the number of molecules of water of crystallization decreases as the ions become larger. Differences between Beryllium and the other Group 2 elements Beryllium is anomalous in many of its properties and shows diagonal relationship to aluminum in Group 3. It is extremely small and has a high charge density and so by Fajans rules it has a strong tendency to covalency.  Beryllium hydride is electron deficient and polymeric with multicentre bonding like aluminium hydride. The halides of beryllium are electron deficient, and polymeric with halogen bridges. BeCl2 usually forms chains but also exists as the dimer. AlCl3 is dimeric. Be forms many complexes – not typical of Groups 1 and 2.  Be like Al is rendered passive by nitric acid. Be is amphoteric, liberating H2 with NaOH and forming beryllates. Al forms Aluminates. Be(OH)2 like Al(OH)3 is amphoteric. Be salts are extensively hydrolysed. Be salts are among the most soluble known.  Beryllium forms an unusual carbide Be2C which like Al4C3 reacts with water to give methane whereas magnesium carbide and calcium carbide give propyne and ethyne(formerly called acetylene) respectively. Be2C+4H2O → 2Be(OH)2 + CH4 Mg2C3 + 4H2O → 2Mg(OH)2 + C3H4 CaC2 + 2H2O → Ca(OH)2 + C2H2  Beryllium metal is relatively unreactive at room temperature, particularly in its massive form. It does not react with water or steam even at red heat and does not oxidize in air below 600°, though powdered Be burns brilliantly on ignition to give BeO and Be3N2.  The halogens (X2) react above 600°C to give BeX2 but the chalcogens (S, Se, Te) require higher temperatures to form BeS, e.t.c. Ammonia reacts above 1200°C to give Be3N2 and carbon forms Be2C at 1700°C. In contrast with the other group IIA metals, Be does not react directly with Hydrogen, and BeH2 must be prepared indirectly.  Cold concentrated HNO3 passivates Be but the metal dissolves readily in dilute aqueous acids (HCl, H2SO4, HNO3) with the evolution of hydrogen. Beryllium is sharply distinguished from the other alkaline earth metals in reacting with aqueous alkalis(NaOH, KOH) with evolution of hydrogen.  Magnesium is more electropositive than the amphoteric Be and reacts more readily with most of the non metals. It ignites with the halogens, particularly when they are moist, to give MgX2 and burns with dazzling brilliance in air to give MgO, Mg3N2.  It also reacts directly with the other elements in Group V and VI (and Group IV); when heated and even forms MgH2 with hydrogen at 570 and 200 atm. Steam produces MgO, or Mg(OH)2 plus Hydrogen and ammonia reacts at elevated temperature to give Mg3N2.  The heavier alkaline earth metals, Ca, Sr, Ba (and Ra) react even more readily with non metals and again the direct formation of nitrides M3N2 is notable. Other products are similar though the hydrides are more stable and the carbides less stable than for Be and Mg.  There is also a tendency, previously noted for the alkali metals to form peroxides MO2 of increasing stability in addition to the normal oxides MO.  Some reactions of Group 2 metals M +2H2O →M(OH)2 +H2 Mg with hot water, and Ca, Sr and Ba react rapidly with cold water. 2M +O2 →2MO Normal oxide formed by all group members. With excess dioxygen Ba +O2 →BaO2 Ba also forms the peroxide. M +H2 →MH2 Ionic ‘salt–like’ hydrides formed at high temperatures by Ca, Sr and Ba. 3M +N2 →M3N2 All form nitrides at high temperatures.  3M +2P →M3P2 All the metals form phosphides at high temperatures. M +X →MX (X=S,Se,Te) All the metals form sulphides, selenides, and tellurides. M +X2 →MX2 (X=F, Cl,Br, I) All the metals form fluorides, chlorides, bromides and iodides.  M +2NH3 →M(NH2)2 + H2 All the metals form amides at high temperatures. Hydroxides Be(OH)2 is amphoteric, but the hydroxides of Mg, Ca, Sr and Ba are basic. The basic strength increases from Mg to Ba and Group 2 shows the usual trend that basic properties increase on descending a group. SULPHATES The solubility of the sulphates in water decreases down the group, Be  Mg  Ca  Sr  Ba. Thus BeSO4 and MgSO4 are soluble but CaSO4 is sparingly soluble and the sulphates of Sr, Ba and Ra are virtually insoluble.  The significantly higher solubilities of BeSO4 and MgSO4 are due to the high enthalpy of solvation of the smaller Be2+ and Mg2+ ions. The sulphates all decompose on heating, giving the oxides: heat MgSO4 → MgO+SO3  MgSO4 and CaSO4 cause permanent hardness in water while the presence of Mg(HCO3)2 and Ca(HCO3)2 causes temporary hardness in water.  HYDRIDES The elements Mg, Ca, Sr and Ba all react with hydrogen to form hydrides MH2. Beryllium hydride is difficult to prepare, and less stable than the others. Hydrides are all reducing agents and are hydrolysed by water and dilute acids with the evolution of hydrogen.  CaH2 +2H2O→Ca(OH)2+2H2 NITRIDES The alkaline earth elements all burn in dinitrogen and form ionic nitrides M3N2. This is in contrast to Group 1 where Li3N is the only nitride formed. 3Ca +N2 →Ca3N2 All the nitrides are all crystalline solids, which decompose on heating and react with water, liberating ammonia and forming either the metal oxide or hydroxide e.g. Ca3N2 +6H2O →3Ca(OH)2 +2NH3  Compounds The predominant divalence of the Group IIA metals can be interpreted in terms of their electronic configuration, ionization energies and size. Further ionization to MX3 is impossible[14847kJmol-1 for Be, 7731kJmol-1 for Mg and 4910kJmol-1 for calcium].  USES In its elemental form, magnesium is used for structural purposes in car engines, pencil sharpeners, and many electronic devices such as laptops and cell phones. In a biological sense, magnesium is vital to the body's health: the Mg2+ ion is a component of every cell type.  Calcium metal is used to make alloys with Aluminium for bearings. It is used in the iron and steel industry to control carbon in cast iron and as a scavenger for P, O and S. Other uses are as a reducing agent in the production of other metals-Zr, Cr, Th and U- and for removing traces of N2 from argon. Chemically CaH2 is sometimes used as a source of H2. Radium was used for radiotherapy treatment of cancer at one time: other forms of radiation are now used.  GROUP 4A ELEMENTS Element Symbol Electronic Structure Carbon C [He] 2s22p2 Silicon Si [Ne] 3s23p2 Germanium Ge [Ar] 3d104s24p2 Tin Sn [Kr] 4d105s25p2 Lead Pb [Xe] 4f145d106s26p2  OCCURRENCE OF THE ELEMENTS The elements are all well known, apart from germanium. Carbon is the seventeenth and silicon the second most abundant element by weight in the earth’s crust. Germanium minerals are very rare. Germanium occurs as traces in the ores of  other metals and in coal, but it is not well known. The abundances of tin and lead are comparatively low, they occur as concentrated ores which are easy to extract.  Carbon occurs in large quantities combined with other elements and compounds mainly as coal, crude oil, and carbonates in rocks such as calcite CaCO3, magnesite MgCO3 and dolomite [MgCO3. CaCO3].  Silicon occurs very widely, as silica SiO2 (sand and quartz) and in a wide variety of silicate minerals and clays. Germanium is only found as traces in some silver and zinc ores. Tin is mined as cassiterite SnO2. lead is found as the ore galena PbS.  Extraction Carbon black (soot) is produced in large amounts. It is made by the incomplete combustion of hydrocarbons from natural gas or oil. Natural graphite is usually found as a mixture with mica, quartz and silicates, which contains 10-60% C.  Silicon is made by reducing SiO2 and scrap iron with coke. SiO2 + Fe + 2C→ FeSi + 2CO There must be an excess of SiO2 to prevent the formation of the carbide SiC. High purity silicon is made by converting Si to SiCl4, purifying this by distillation, and reducing the chloride with Mg or Zn.  SiO2 + 2C→ Si + 2CO Si + 2Cl2→ SiCl4 SiCl4 + 2Mg→ Si + 2MgCl2 The only important ore of tin is SnO2(Cassiterite). SnO2 is reduced to the metal using Carbon at 1200-1300C in an electric furnace. The product often contains traces of iron, which make the metal hard.  Iron is removed by blowing air through the molten mixture to oxidize the iron to FeO, which then floats to the surface. The main oxide of lead is galena PbS. There are two methods of extracting the element: 1). Roast in air to give PbO and then reduce with coke or CO in a blast furnace. 2PbS + 3O2→ 2PbO + 2SO2 +C 2Pb(liquid) + CO2(gas)  2). PbS is partially oxidized by heating and blowing air through it. The air is turned off after some time, heating continues and the mixture undergoes self-reduction. heat in heat in 3PbS → PbS + 2PbO → 3Pb(liquid) + air air SO2(gas)  The Group 4A elements are found in the p-block. Each of these elements has only two electrons in its outermost p orbital with electron configuration ns2np2. The Group 4A elements tend to adopt oxidation states of +4 and, for the heavier elements, +2 due to the inert pair effect.  The inert pair effect is the phenomenon of electrons remaining paired in valence shell. It can be defined as the reluctance of the outermost shell s-electrons to participate in bonding.  Members of this group conform well to general periodic trends. The atomic radii increase down the group, and ionization energies decrease.  Metallic properties increase down the group. Carbon and silicon are non- metals, germanium has some metallic properties, tin and lead are metals.  The elements in this group are relatively unreactive, but reactivity increases down the group. MII oxidation state becomes increasingly stable on descending the group. Pb often appears more noble(unreactive) than expected from its standard electrode potential of -0.13volts.  The unreactiveness is partly due to a surface coating of oxide. Carbon, silicon and germanium are unaffected by water, tin reacts with steam to give SnO2 and H2. Pb is unaffected by water, probably because of a protective oxide film.  Carbon is the fourth most abundant element in the known universe but not nearly as common on the earth, despite the fact that living organisms contain significant amounts of the element.  Common carbon compounds in the environment include the gases carbon dioxide (CO2) and methane (CH4). Allotropes Carbon exists in several forms called allotropes. Diamond is one form with a very strong crystal lattice, known as a precious gem from the most ancient records.  Graphite is another allotrope in which the carbon atoms are arranged in planes which are loosely attracted to one another (hence its use as a lubricant). The recently discovered fullerenes are yet another form of carbon.  Inorganic carbon may come in the form of diamond as transparent, isotropic crystal. It is the hardest natural occurring material on this earth. Diamond has four valence electrons, and when each electron bonds with another carbon it creates a sp3-hybridized atom. The boiling point of diamond is 4827°C.  Unlike diamond, graphite is opaque, soft, dull and hexagonal. Graphite can be used as a conductor (electrodes) or even as pencils.  Germanium, categorized as a metalloid in group 4A, the Carbon family, has five naturally occurring isotopes. Germanium, abundant in the Earth's crust has been said to improve the immune system of cancer patients. It is also used in transistors, but its most important use is in fiber-optic systems and infrared optics.  The name for silicon is taken from the Latin silex which means "flint". The element is second only to oxygen in abundance in the earth's crust and was discovered by Berzelius in 1824. The most common compound of silicon, SiO2, is the most abundant chemical compound in the earth's crust, which we know it better as common beach sand.  Properties Silicon is a crystalline semi-metal or metalloid. One of its forms is shiny, grey and very brittle (it will shatter when struck with a hammer). It is a group 4A element in the same periodic group as carbon, but chemically behaves distinctly from all of its group counterparts.  Silicon shares the bonding versatility of carbon, with its four valence electrons, but is otherwise a relatively inert element. However, under special conditions, silicon can be made to be a good deal more reactive.  Silicon exhibits metalloid properties, is able to expand its valence shell, and is able to be transformed into a semiconductor; distinguishing it from its periodic group members.  27.6% of the Earth's crust is made up of silicon. Although it is so abundant, it is not usually found in its pure state, but rather its dioxide and hydrates. SiO2 is silicon's only stable oxide, and is found in many crystalline varieties.  Its purest form being quartz, but also as jasper and opal. Silicon can also be found in feldspar, micas, olivines, pyroxenes and even in water. Silicon is most commonly found in silicate compounds.  Applications Carbon has a very high melting and boiling point and rapidly combines with oxygen at elevated temperatures. In small amounts it is an excellent hardener for iron, yielding the various steel alloys upon which so much of modern construction depends.  Activated carbon is used extensively in sugar industry as decolourizing agent. It is also used in purification of chemicals and gases. It is used as catalyst.  An important (but rare) radioactive isotope of carbon, C-14, is used to date ancient objects of organic origin. It has a half-life of 5730 years but there is only 1 atom of C-14 for every 1012 atoms of C-12 (the usual isotope of carbon).  Silicon is a semiconductor with a clear shiny bluish grey metallic lustre. It is used in the production of transistors while an isotope 29Si, is used in NMR spectroscopic studies.  Si/steel alloys are used for the construction of electric motors. Silicon dioxide and Silicon (in the form of clay or sand) are important components of bricks, concrete and Portland cement. Silicon  Silicon parts are used in computers, transistors, solar cells, Liquids Crystal Display (LCD) screens and other semiconductors devices. Silicates are used to make pottery and enamel. Sand which contains silicon is an important component of glass. Silicones are used in high temperature greases, waxes, breast implants, contact lenses, explosives and pyrotechnics.  Germanium is transparent to infra-red light and is therefore used for making prisms and lenses and windows in infra-red spectrophotometer.  Germanium is used in transistor technology and in optics. Magnesium germanate is used in the special alloy, strain gauges and in superconductors. Germanium is used in electronic application when doped with arsenic, gallium or other elements.  Germanium oxide has a high refractive index of refraction and dispersion, making it suitable for use in wide-angle camera lenses and objectives lenses for microscopes. It is also used as an alloying agent (adding 1% germanium to silver stops if from tarnishing) in fluorescent lamps and so a catalyst.  Both Ge and germanium oxide are transparent to infrared radiation and so are used in the manufacture of infrared spectroscopes.  Tin is used for electroplating steel to make tin- plate and alloys. Lead is used to make lead/acid storage batteries. Lead is used as protective shielding against X- ray and radiation from nuclear reactors  GROUP IV ELEMENTS C Si Ge Sn Pb  CARBON 12C isotope (predominantly) 13C (smaller amt) 14C (radioactive) – archeological dating – radioactive tracers 1st 4 Ionization Energies (much higher)  CARBON Diamond (extremely unreactive @ rt) Graphite (layer structure) reacts more readily – oxidized to mellitic acid C6(CO2H)6 {hot conc. HNO3} – active reducing agent – reacts readily with many oxides liberate the element form carbide.  CARBIDES 3 types which are similar Most polar (ionic) {electropositive metals} Most covalent (molecular) {electronegative non – metals} Somewhat complex (interstitial) {elements in the center of the d-block}  Salt – like carbides Be2C and Al4C3 - sometimes called “methanides” – yield predominantly CH4 on hydrolysis C2 units are well known e.g. acetylides (ethynides) of alkali metals MI2C2, alkaline earth metals MIIC2 and the lanthanides LnC2 and Ln2C3 i.e. Ln4(C2)3 Gp. IB (Cu, Ag, Au) are explosive Gp. IIB (Zn, Cd, Hg) are poorly characterized  Salt – like carbides MI2C2 - colourless crystalline compounds – react violently with water – oxidized to carbonate on being heated in air CaC2 - most important compound – production of ethyne (chemical industry & oxy - acetylene welding) – fixing N2 from air to give calcium cyanamide {CaCN2} widely used as a fertilizer  Interstitial Carbides Lanthanides form – metal rich carbides M3C {normal carbides (LnC2)} Actinides form – monocarbides MC (several early T. elements) Which are infusible, extremely hard, refractory materials that retain many characteristic properties of metals  Interstitial Carbides C atoms occupy octahedral interstices in a closed-packed lattice of metal atoms Carbides of Cr, Mn, Fe, Co and Ni – profuse in number – complicated in structure – great industrial importance e.g. cementite Fe3C (important constituent of steel) – Much more reactive than the interstitial carbides of the earlier T.S.  HYDRIDES AND HALIDES Catenation - best illustrated in the hydrides (organic chemistry) Replacement of H by F as in CF4 – greatly increases both thermal stability & chemical inertness great strength of the C-F bond Thus, fluorocarbons – resistant to attack by acids, alkalis, o. agents, r. agents and most chemicals up to 600 oC – Immiscible with both water & hydrocarbon solvents – when combined with other groups confer water-repellance and stain resistance – paper, textiles & fabrics  OXIDES Forms 2 extremely stable oxides - CO & CO2 3 oxides of considerably lower stability (C3O2, C5O2 and C12O9) number of unstable or poorly characterized oxides like C2O, C2O3  OXIDES Tricarbon dioxide C3O2 – foul-smelling gas – from dehydration of malonic acid @ reduced pressure over P4O10 at 140oC – polymerizes at rt to a yellow solid O=C=C=C=O – readily rehydrates to malonic acid, CH2(CO2H)2 – reacts with NH3&HCl to give amide & acid chloride CH2(CONH2)2 and CH2(COCl)2  CYANIDES Cyanogen (CN)2 – colourless poisonous gas (like HCN) – possessing considerable thermal stability (800 oC) when pure – presence of trace impurities normally results in polymerization at 300-500 oC to paracyanogen with a condensed polycyclic structure  SILICON 28Si (92%) with ~5% 29Si & 4% 30Si more volatile than C – substantially lower energy of vaporization (smaller Si - Si bond energy) Semiconductor – distinct shiny, blue – grey metallic lustre  SILICON Massive crystalline form – relatively unreactive except at high t. – Formation of very thin continuous, protective surface layer of SiO2 makes rxn with O2, water and steam, less effective – Oxidation in air - not measurable below 900 oC – between 950-1160 oC rate of formation of vitreous SiO2 rapidly increases – at 1400 oC, N2 in air gives SiN & Si3N4 – Sulphur vapour reacts at 600 oC – P vapour reacts at 1000 oC – unreactive towards aq. Acids – conc. HNO3 / HF oxidizes & fluorinates the element  SILICON Molten Si (in contrast) – extremely reactive forms alloys or silicides with most metals rapidly reducing most metal oxides – very large heat of formation of SiO2 (900 kJ mol-1) Silicon does not form binary compounds with Ge, Sn, & Pb compound with C - SiC – outstanding academic & practical interest  SILICON CARBIDE Greater thermal stability than any other binary compound of Si – decomposition by loss of Si appreciable @ 2700 oC resists attack by most aq. acids (HF inclusive but not H3PO4) oxidized in air only above 1000 oC {protective SiO2 layer} – removed by molten hydroxides or carbonates thereby accelerating oxidation. SiC + 2NaOH + 2O2 = Na2SiO3 + H2O + CO2  SILICIDES Formulae cannot be rationalized by application of simple valence rules Bonding varies – essentially metallic to ionic & covalent Observed stoichiometries include – M6Si, M5Si, M4Si, M15Si4 – M3Si, M5Si, M2Si, M5Si3 – M3Si2, MSi, M2Si3 – MSi2, MSi3, MSi6  SILICON HYDRIDES (SILANES) Silanes, SinH2n+2 – known as unbranched & branched chains (up to n = 8) – no cyclic compounds or unsaturated analogues of alkenes and alkynes – colourless gases or volatile liquids – extremely reactive spontaneously igniting or exploding in air – Thermal stability decreases with increasing chain length only SiH4 is stable indefinitely at rt.  SILICON HYDRIDES (SILANES) Much more reactive than the corresponding C compounds – (a) larger radius of Si which would facilitate nucleophilic attack – (b) great polarity of Si-X bonds – (c) presence of low-lying d – obitals which permit formation of 1:1 & 1:2 adducts thereby lowering the activation energy of the rxn  SILICON HYDRIDES (SILANES) Pure silanes do not react with pure water or dilute acids in silica vessels but even traces of alkali dissolved out of glass apparatus catalyze the hydrolysis – which is rapid and complete (SiO2.nH2O + 4H2) Solvolysis with MeOH can be controlled – several products SiH4-n (OMe)n (n=2,3,4) Explodes in the presence of Cl2 or Br2 rxn with Br2 can be moderated at –80o – good yields of SiH3Br & SiH2Br2  SILICON HALIDES Si & SiC react readily with all halogens – colourless volatile reactive products SiX4 Two different tetrahalides heated together – they equilibrate to form a random distribution of silicon halides nSiX4 + (4-n)SiY4 = 4SiXnY4-n Higher homologues SinX2n+2 – volatile liquids or solids Catenation in Si compounds reaches its maximum in the halides rather than the hydrides (contrary to C)  SILICON HALIDES Fluoropolysilanes up to Si16F34 up to at least Si6Cl14 and Si4Br10 are known  SILICA Chemically resistant to all acids except HF Dissolves slowly in hot conc. alkali & more rapidly in fused MOH or M2CO3 – M2SiO3 Among halogens, only F2 attacks SiO2 readily – SiF4 + O2 Reaction of SiO2 with the oxides of metals & semimetals – great interest & importance in glass technology & ceramics  SILICATES Alkali metal carbonates fused with silica (~1300o) – CO2 + complex mixture of alkali silicates Basic unit of structure is SiO4 tetrahedron – Occur singly or by sharing oxygen atoms, in small groups, in small cyclic groups, in infinite chains or in infinite sheets few simple discrete SiO44- orthosilicate anions are known e.g. Be2SiO4  SILICATES Simplest condensed silicate anion - pyrosilicate ion Si2O76- formed by combining two SiO4 tetrahedra with sharing of O atoms e.g. Sc2Si2O7 Cyclic silicates are known Si3O96- and Si6O1812- e.g. BaTiSi3O9 and Be3Al2 Si6O18 Infinite chain anions such as (SiO32-)n and (Si4O116-)n are known.  GERMANIUM, TIN AND LEAD Notable Pair wise similarities in I.Es – Si & Ge (filling of 3d10 shell) – Sn & Pb (4f14 shell) Ge – brittle, grey –white lustrous crystals with diamond structure – metalloid with similar electrical resistivity to Si at rt – mp, bp & associated enthalpy changes are also lower than for Si Trend continues for Sn & Pb – very soft, and low-melting metals  CHEMICAL REACTIVITY AND TRENDS Ge – more reactive and more electropositive than Si – dissolves slowly in hot conc. H2SO4 & HNO3 – does not react with water or with dil. acids or alkalis unless an oxid. agent like H2O2 or NaOCl is present – fused alkalis react to yield germanates – oxidized in air at red heat GeO2 – H2S & S(g) yield GeS2 – Cl2 & Br2 yield GeX4 on moderate heating – HCl gives both GeCl4 & GeHCl3 – Alkyl halides react with heated Ge (as with Si) to give organogermanium halides  CHEMICAL REACTIVITY AND TRENDS Sn – notably more reactive & electropositive than Ge though still markedly amphoteric in its aq chemistry – stable towards both water & air at ordinary temps – reacts with steam to give SnO2 + H2 – heating in air or O2 gives SnO2 – Dil HCl and H2SO4 show little, if any, rxn – dil HNO3 produces Sn(NO3)2 & NH4NO3  CHEMICAL REACTIVITY AND TRENDS Sn – Hot conc HCl yields SnCl2 + H2 – hot conc H2SO4 forms SnSO4 + SO2. Occurrence of SnII compounds in these rxns is notable By contrast, hot aq alkali yield hydroxostannate (iv) compounds e.g Sn + 2KOH + 4H2O →K2[Sn(OH)6] + 2H2 Reacts readily with Cl2 and Br2 in the cold, F2 & I2 on warming to give SnX4 reacts vigorously with heated S & Se, to form SnII & SnIV chalogenides – depending on the proportions used with Te gives SnTe.  CHEMICAL REACTIVITY AND TRENDS Finely divided Pb powder is pyrophoric reactivity is greatly diminished by formation of a thin, coherent protective layer of insoluble prd like oxide, oxocarbonate, sulphate or chloride F2 reacts at rt to give PbF2 Cl2 gives PbCl2 on heating Molten Pb reacts with the chalcogens to give PbS, PbSe & PbTe  CHEMICAL REACTIVITY AND TRENDS steady trend towards increasing stability of MII rather than MIV compounds (Ge, Sn, Pb) - “inert pair effect” {heavier B subgroup metals} Notable exception is the organometallic chemistry of Sn & Pb which is almost entirely confined to the MIV state Catenation is also impt for Ge, Sn, Pb chemistry thought less so than for C & Si Ability of both Sn & Pb to form polyatomic cluster anions of very low formal oxidation state (e.g M52-, M94- etc) – reflects the tendency of heavier B subgroup elements to form chain, ring or cluster homo-polyatomic ions  HYDRIDES AND HYDROHALIDES Germanes, GenH2n+2 – colorless gases or volatile liquids for n=1-5 – prepn , physical propts and chemical rxns very similar to those of silanes – though they are less volatile and noticeably less reactive – GeH4 does not ignite in contact with air (cf. SiH4 & SnH4) & unaffected by aq acid or 30% aq NaOH – acts as an acid in liq NH3 -> NH4+ + GeH3- ions – alkali metals in liq NH3 -> MGeH3  HYDRIDES AND HYDROHALIDES Germanuim hydrohalides GeHxX4-x (X = Cl, Br, I; x = 1,2,3) colourless, volatile, reactive liquids – valuable synthetic intermediates (cf SiH3I) e.g. hydrolysis of GeH3Cl yields O(GeH3)2 SnH4 decomposes slowly to Sn + H2 at rt unattacked by dil aq acids or alkalis but decomposed by more conc solns Potent reducing agent  HYDRIDES AND HYDROHALIDES Sn2H6 – even less stable higher homologues not known By contrast, organotin halides are more stable catenation up to H(SnPh2)6H has been achieved PbH4 least well characterized Gp IVB hydride unlikely that it has ever been prepd  HALIDES Ge, Sn & Pb forms two series: MX2 & MX4 PbX2 more stable than PbX4 whereas reverse is true for Ge steady increase in stability thus: CX2 OS COMPLEXES – Coordination Compounds OH2 2+ H2O OH2 Fe Examples OH2 H2O OH2 [Fe(H2O)6]2+ OH2 [Co(NH3)6]3+ [Cr(OH)6]3- NH3 3+ H3N NH3 [CuCl4]2- Co NH3 H3N NH3 NH3  COMPLEXES – Coordination Compounds Most transition metals have vacant d-orbitals o accept electron pairs o Forming coordinate covalent bonds in coordination compounds Cationic [Cr(H2O)6]3+ , [Co(NH3)6]3+ , [Ag(NH3)2]+ Neutral [Fe(CO)5] , [PtCl2(NH3)2] Anionic [Ni(CN)4]2- , [Fe(CN)6]3-  COMPLEXES – Coordination Compounds Many are very stable – low dissociation consts Include important biological substances – Haemoglobin & Chlorophyll O2 carrying protein, contains Fe2+ bound to large porphyrin rings – Transport involves coordination & subsequent release of O2 Necessary for photosynthesis, contains Mg2+ bound to large porphyrin rings Vit. B-12 is a large complex of Cobalt  COMPLEXES – Coordination Compounds Many practical applications – Water treatment – Soil and plant treatment – Protection of metal surfaces – Analysis of trace amounts of metals – Electroplating – Textile dyeing  ATOMIC RADII Sc Ti V Cr Mn Fe Co Ni Cu Zn 1.44 1.32 1.22 1.17 1.17 1.17 1.16 1.15 1.17 1.25 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd 1.62 1.45 1.34 1.29 --- 1.24 1.25 1.28 1.34 1.41 La Hf Ta W Re Os Ir Pt Au Hg 1.69 1.44 1.34 1.30 1.28 1.26 1.26 1.29 1.34 1.44  ATOMIC RADII Decreases from left to right – increase in nuclear charge Increases near end of series – increase in effective nuclear charge outweighed by greater repulsions (d e- in nearly filled orbitals) Increases down the group but 3rd row TMs have nearly same radii as 2nd row TS due to LANTHANIDE CONTRACTION A similar trend is observed with the ionic radii  ATOMIC RADII LANTHANUM (57La) f electrons less shielding than d < p < s Higher effective nuclear charge is felt by outermost electron –Smaller radii than expected  IONIC RADII Sc3+ Ti4+ V3+ 0.745 0.605 0.64 Y3+ Zr4+ Nb3+ 0.90 0.72 0.72 La3+ Hf4+ Ta3+ 1.032 0.71 0.72 Lanthanide elements  CATALYTIC PROPERTIES TiCl3 – Ziegler-Natta catalyst in production of polythene V2O5 – Contact process for H2SO4, converts SO2 to SO3 MnO2 – Catalyst to decompose KClO3 to give O2 Fe – Haber-Bosch process for making NH3 FeCl3 – Production of CCl4 from CS2 and Cl2 PdCl2 – Wacker process, in the reaction C2H4 + H2O + PdCl2 → CH3CHO + 2HCl + Pd  CATALYTIC PROPERTIES Pt – Used in 3-stage convertors for cleaning car exhaust fumes Pd – Hydrogenation e.g. Phenol to cyclohexanone CuCl2 – Deacon process for HCl to Cl2 Ni – Raney nickel, reduction processes  CHE 126 INTRODUCTION TO TRANSITION METAL CHEMISTRY CONTINUES…..  Classification into Sub-groups Subdivided into 8 groups – I → VIII Elements form many compounds of similar stoichiometry as the MGMs o though dissimilar chemical properties I II III IV V VI VII VIII NaCl MgBr2 Al(NO3)3 CCl4 POCl3 SO42– Cl2O7 KNO3 CaCl2 Ga(OH)3 PbO2 PO43– H2S2O7 HClO4 CuCl ZnBr2 Sc(NO3)3 TiCl4 VOCl3 CrO42– Mn2O7 AgNO3 CdCl2 Y(OH)3 ZrO2 VO43– H2Cr2O7 HMnO4  Classification into Sub-groups Gp VIII consists of 3 gps of metals with 3 horizontal triads o Fe → Fe Co & Ni o Pd → Ru Rh & Pd o Pt → Os Ir & Pt I II III IV V VI VII VIII NaCl MgBr2 Al(NO3)3 CCl4 POCl3 SO42– Cl2O7 Fe Co Ni KNO3 CaCl2 Ga(OH)3 PbO2 PO43– H2S2O7 HClO4 Ru Rh Pd CuCl ZnBr2 Sc(NO3)3 TiCl4 VOCl3 CrO42– Mn2O7 Os Ir Pt AgNO3 CdCl2 Y(OH)3 ZrO2 VO43– H2Cr2O7 HMnO4  Common Trends in Sub-groups Corresponding compds in same OS, covalent character decreases & ionic character increases down the group Increasing electrical conductivity in aq solns & increasing mps/bps for the heavier compounds V2O 5 = 690o Nb2O 5 = 1460o Ta2O 5 = 1800o melting point Ionic character Compounds in difft proportions, lower OS usually more ionic o TiCl2 & TiCl3 are ionic solids whereas TiCl4 is molecular liquid For a given TM, O2– & OH– of lower OS are basic, intermediate OS are amphoteric, high OS are acidic  Manganese oxides & hydroxides Acidity increases as OS increases o Variation in chemical properties as Oxidation State increases is typical of many TMs +2, +3, +4, +6 & +7 in its simple compounds +2 state is most stable KMnO4 (+7), stable, very strong oxidizing agent o common laboratory reagent  Manganese oxides & hydroxides +1 HClO (hypochlorous) ------- +2 ------- Mn(OH)2 +3 HClO2 (chlorous) Mn(OH)3 +4 ------- H2MnO3 (manganous) +5 HClO3 (chloric acid) ------- +6 ------- H2MnO4 (manganic) +7 HClO4 (perchloric acid) HMnO4 (permanganic acid) Permanganic & perchloric acids – very strong acids & oxid. ag. Former exists only in soln but neutralized by bases – KMnO4 (very strong OA) MnO4- + 8H+ + 5e- → Mn2+ + 4H2O E° = +1.51V  Manganese oxides & hydroxides Mn = metal ; Cl = nonmetal ✓ Properties in high OS are very much alike ✓ but trends in acidity & basicity as a function of OS are similar o Mn2O7 (dark brown explosive liquid) o Cl2O7 (colourless explosive liquid) o Both anhydrides are formed by dehydrating HMnO4 and HClO4 respectively Chemical properties are usually not as similar for corresponding A & B group elements  Manganese oxides & hydroxides Other group VII elements form compounds with similar formulae o Heptoxides react with water → strongly acidic solns Mn2O7 + 3H2O → 2 (H3O+ + MnO4–) permanganic Tc2O7 + 3H2O → 2 (H3O+ + TcO4–) pertechnetic Re2O7 + 3H2O → 2 (H3O+ + ReO4–) perrhenic Mn2O7 dangerously explosive o Tc2O7 & Re2O7 are stable enough to be sublimed Evidence of increasing stability of higher O States as transition metal group is descended INTRODUCTION TO NUCLEAR CHEMISTRY INTRODUCTION TO NUCLEAR CHEMISTRY Electrons physical and chemical properties of atoms, ion, molecules etc, in every day life Nuclear chemistry phenomena in the nuclei of atoms Atoms are not conserved but nucleons are, Nucleons are rearranged to form different nuclei Main diff between a chemical reaction and a nuclear reaction. 10 INTRODUCTION TO NUCLEAR CHEMISTRY Protons Neutrons Electrons Charge +1 0 –1 Mass (amu) 1 1 0 Location Nucleus Nucleus Orbital 11  The Mass Defect of a nuclide Difference between the actual mass of an atom’s nucleus and the sum of the masses of its constituent nucleons Mass defect = calcd mass – actual mass It is mass of the energy binding the nucleus Evidence – Sum of the masses of individual neutrons and protons always greater than the mass of the atom 12  Nuclear Binding Energy Energy required to split a nucleus of an atom into its component nucleons Measure of the stability of the nucleus Strength of the nuclear bond depends on the number of n + p 13  Nuclear Binding Energy Einstein equation describes the equivalence and interconvertibilty of mass and energy Law of conservation of mass-energy “the total of the mass and energy of the universe is a constant” Every energy change must be accompanied by change in mass ΔE = Δmc2 Nuclear binding energy, calcd by converting mass defect to energy Missing mass = NUCLEAR BINDING ENERGY o energies in nuclear processes 14  Nuclear Binding Energy ΔE = Δmc2 Units of E = Joules, Joules/nucleon, Joules/mole To convert to Joules/mole, Joules x Avogadro’s number. To convert to Joules per nucleon, Joules/nucleons To convert to MeV, Joules x 6.2415064799632 x 1012 MeV 1 Mev = 1.6021892 x 10-13 J Unit of mass = kg, mass (kg) = m (u) x 1.66053886 x 10–27 kg c = is the speed of light in vacuum = 3 x 108 m/s 15  Calculation Examples (i) Calculate the average binding energy per mole of a U- 235 isotope. Show your answer in kJ/mole and in J/nucleon. (U-235 has 92 protons, 143 neutrons, and has an observed mass of 235.04393 u). Solution Mass defect = 235.04393 u – (1.007825 x 92 + 1.008665 x 143) = 1.915065 u Convert to kg: 1.915065 u x 1.66053886 x 10–27 kg = 3.18004 x 10–27 kg Calc the energy: E = Δmc2 E = 3.18004 x 10-27 kg x (3 x 108)2 = 2.862 x 10-10 J = 2.862 x 10-13 kJ = 2.862 x 10-13 x 6.02 x 1023 kJ/mole Note: 1 u = 931.5 MeV/c2 16  Calculation Examples Atomic mass of 3919K is 38.96371u. (i) Calculate the binding energy for this nuclide (n = 1.008665 u, p = 1.007825 u, c = 3 x 108 m s-1, 1 u = 1.6605655 x 10-27 kg). (ii) Calculate the total binding energy of one mole of 3919K atoms. Solution 39 K has 19 protons, 20 neutrons 19 Calcd mass = (19 x 1.007825) + (20 x 1.008665) = 19.148675 u +20.1733 u = 39.32197 u Mass defect = calculated mass – actual mass = 39.32197 – 38.96371 = 0.35826 u 17  Calculation Examples This loss of mass is equivalent to (0.35826 x 1.6605655 x 10-27) = 5.9493 x 10-28 kg Energy equivalent of this mass = mc2 = (5.9493 x 10-28)( 2.9979 x 108)2 = 5.3469 x 10-11 J For 1 mole, total binding energy is this energy x Avogadro number i.e. (-5.3469 x 10-11)(6.022 x 1023) = -3.2199 x 1013 J/mol In MeV/nucleon; E = (0.35826 u x 931.5 MeV/c2)c2 /39 = 8.56 MeV/nucleon 18  Binding energy per nucleon Average energy required to remove an individual nucleon from a nucleus Binding energy per nucleon = nuclear binding energy (nucleus)/no of nucleons in nucleus – Useful for comparing one nucleus with another – Analogous to the ionization energy of an electron in an atom. – If relatively large, the nucleus is relatively stable. – The values are estimated from nuclear scattering experiments – increases rapidly for lighter elements – As no of nucleons increases, the nucleons are held together more strongly – Curve occurs @ Fe (nucleons are most strongly bound together) 19  Binding energy per nucleon Average energy required to remove an individual nucleon from a nucleus Binding energy per nucleon = nuclear binding energy/no of nucleons in nucleus – Values are estimated from nuclear scattering experiments – Measure of the stability of nuclei, large energy = stability – Useful for comparing one nucleus with another – Analogous to IE of an electron in an atom – Values ≈ 6–10 MeV to remove a nucleon from a nucleus – cf. 13.6 eV needed to ionize an electron in the ground state of hydrogen Nuclear force is strong 20  Binding energy per nucleon Increases rapidly for lighter elements Maximum occurs around atomic mass 56 Fe (nucleons are most strongly bound together) Suggests nuclei with mass numbers in the region of 56, i.e. nickel, iron are the most stable This is why nuclear fusion in the cores of stars ends with Fe 22  Binding energy per nucleon The shape of the graph has to do with competing forces in the nucleus As no of nucleons increases, the nucleons are held together more strongly At low atomic masses, attractive nuclear forces between nucleons dominate over repulsive electrostatic forces between protons 23  Binding energy per nucleon At high atomic masses, repulsive electrostatic forces between forces begin to dominate, and these forces tend to break apart the nucleus rather than hold it together Nuclei divided (fission) or combined (fusion) release an enormous amount of energy Energy used to generate electrical and solar power 24  Nuclear Reactions The four main types of nuclear reactions: Fission – of unstable heavy nuclei Fusion - of light nuclei o occurs naturally only in the sun and other stars Decay – spontaneous decay of radioactive nuclides Transmutation - Bombardment reactions 1  Nuclear Fission The nucleus of a large atom is split into smaller and lighter nuclei o when bombarded with neutrons o Release of a large amount of energy o plus neutrons and photons (in form of γ rays) Mass of original nucleus is slightly higher than the masses of its individual nuclei o This difference is the mass defect The final products are more stable The case of decay process is called spontaneous fission o very rare process. 2  Nuclear Fission Most important is uranium-235 fission: 235 U + 1 n → 89 Kr + 144 Ba + 3 1 n 92 0 36 56 0 For every neutron consumed, three neutrons are produced Hence, used in both nuclear power plants o to provide electricity to an entire city Used in nuclear bombs - to destroy an entire city E.g., 1 kg of uranium can release sufficient energy for combusting 4 billion kg of coal. Used to produce steam to generate electricity 3  Nuclear Fission Most of energy released appears as kinetic energy of the fission fragments Fission fragments interact strongly with the surrounding atoms/molecules traveling at high speed, causing them to ionize 4  Nuclear Fusion Two or more lighter nuclei collide at a very high energy and fuse together into a new heavy nucleus. Fusion reactions release even greater amounts of energy, They only occur at unfathomably high temperatures They occur in stars in outer space The sun is basically one giant fusion reactor Hydrogen nuclei fuse together into helium nuclei, releasing the light and heat that warms our planet 5 Nuclear Fusion 6  Nuclear Fusion Two pairs of protons fuse, forming two deuterons. 4 1H → 2 2D Each deuteron fuses with an additional proton to form helium-3; 2D + 1H → 3He 3He + 3He → 6Be → 2 1H + 4He + 2 v + 2 e+ + γ Since the helium-4 atom has less energy or resting mass than the 4 protons which initially came together, energy is radiated outside the core and across the solar system. Sun uses up about 600 million tons of hydrogen nuclei per second Convert into helium releasing 384.6 trillion trillion Joules of energy per second equivalent to the energy released in the explosion of 91.92 billion megatons of TNT per second. 7  RADIOACTIVITY Decomposition of one nuclide to form a different nuclide Emission by unstable nuclei o particles or electromagnetic radiation or both Natural or artificial Approx. 33% of the elements has natural radionuclides radio isotopes or radioactive nuclides All known isotopes of elements heavier than Bismuth (Z > 83) are radioactive 8  RADIOACTIVITY 1st natural transmutation of an element (Rutherford and Soddy 1902) 226 Ra 88 → 4 He 2 + 222 Rn 86 1st artificial (Rutherford 1919) 14 N 7 + 4 He 2 → 17 O 8 + 1 p 1 9  RADIOACTIVITY Discovery of the neutron (Chadwick 1932) 9 Be 4 + 4 He 2 → 12 C 6 + 1 n 0 conservation of mass number & nuclear change in both reactants and products 10  RADIOACTIVITY Artificial radionuclides (more than 350) – nuclear bombs testing (1955-62) –operation of nuclear power plants –research laboratories and facilities like hospitals 11  RADIOACTIVE EMISSIONS α-particle now known as helium nucleus 42He Emission of α- particle by radionuclide AX Z → A-4 X Z-2 + 4 He + γ 2 234 U → 230 Th + 4 He + γ 92 90 2 Parent nuclide Daughter nuclide α-particle emission leaves nuclide in an excited state, then γ-radiation is emitted to reach stable (or ground) state. Total energy of emitted α-particle + γ-radiation is equal to the reaction energy. 12  RADIOACTIVE EMISSIONS α-particle Heavy nuclides with low neutron/proton ratio and those with Z > 83 undergo α-decay 185 Au 79 → 181 Ir + 77 4 He 2 18579 Au with 106n & 79p (106/79 = 1.34) lies below stability band and o the daughter nuclide with 104n & 77p is closer to the band and hence more stable (104/77 = 1.35). 13  RADIOACTIVE EMISSIONS β-particle - Electron emission (10n → 1 p 1 + 0 e) –1 No change in mass number but increases atomic no. by 1 Net effect - conversion of neutron to proton (n → p) AX → A Y + 0 e z Z+1 –1 n/p ratio decreases most common mode of decay for radioactive nuclides since they lie above stability band (too many neutrons) 14 C → 14 N + 0 e 6 7 –1 15  RADIOACTIVE EMISSIONS Positron emission β+ 1 p 1 → 1 n 0 + 0 e +1 (p → n) Positron = electron with equal mass but opposite charge (+1 instead of –1) causes decrease of 1 in atomic no but no change in mass number AX z→ AZ-1Y + 0+1e increases n/p ratio, isotopes with too many protons (below stability band) decay by this process 16  RADIOACTIVE EMISSIONS Positron emission (p → n) Only artificial radionuclides have been observed to undergo position emission When positron and electron interact, they annihilate each other all their masses are converted to energy Two 0.51 MeV traveling in opposite directions 17  RADIOACTIVE EMISSIONS Electron capture (11p + 0–1e → 10n) Positive charge of an unstable nucleus can be decreased by electron capture of one of its own orbital electrons Proton captures an electron to produce neutron (p → n) 18  RADIOACTIVE EMISSIONS Electron capture As electrons rearrange themselves to compensate for the electron pulled into the nucleus, x-rays are emitted. A zX+ 0-1e → AZ-1Y 202 Tl → 202 Hg (followed by x-rays) 81 80 50 V + 0 e → 50 Ti + x-rays. 23 -1 22 19  RADIOACTIVE EMISSIONS γ-decay Radioactive decay sometimes leaves nucleus at unstable excited nuclear energy level, emission of γ-rays → ground state nuclear energy level (excess energy electron magnetic radiations) AX * Z → AX Z + 0 γ0 This takes place within a nanosecond following particle emission. 20 BAND OF STABILITY (NEUTRON/PROTON RATIO) Why are some nuclides radioactive and others are not? A plot of the number of neutrons versus the number of protons in stable nuclei will provide an answer The plot shows nuclei with an equal number of n and p fall along the ‘broken’ line, n/p = 1:1 Isotopes of lighter elements up to 4020Ca fall on or quite close to the n=p line while in heavier elements n increases faster than p and the n/p ratio eventually reaches about 5:3 21  BAND OF STABILITY (NEUTRON/PROTON RATIO) The additional neutrons provide the additional nuclear force necessary to hold larger no of protons close together within the nucleus Once Atomic No. reaches 83, even extra neutrons cannot maintain stability Hence, all known nuclides of Z > 83 are unstable and radioactive For each nuclear charge, only isotopes with a neutron/proton ratio within a specific range (band of stability) are stable and not radioactive 23  BAND OF STABILITY (NEUTRON/PROTON RATIO) Essentially, radioactivity is the spontaneous transformation of unstable nuclei to nuclei with more favorable n/p ratios Nuclides with too many protons fall below the stability band and decay so that there is a decrease in the number of protons relative to the number of neutrons (greater n/p ratio) Nuclides with too many neutrons are above the stability band and decay so that there is a decrease in the number of neutrons relative to the number of protons (smaller n/p ratio) 25  BAND OF STABILITY (NEUTRON/PROTON RATIO) Thus, the elements are β-emitters, since beta decay has the overall effect of losing a neutron and gaining a proton: 1 n → 1 p + 0 e 0 1 -1 E.g., 209F → 2010Ne + 0-1e β decay of F-20 reduces the n/p ratio from 11/9 to 10/10 thereby moving the surviving nucleus (Neon-20) closer to the center of the stability band.