Alkali and Alkaline Earth Metals PDF

Alkali Metals Alkaline earth Metals 1 of 32 © Boardworks Ltd 2005  The s-Block Elements Elements of Groups IA* (the alkali metals) and IIA* (the alkaline earth metals)  constitute the s-block elemen...

Alkali Metals Alkaline earth Metals 1 of 32 © Boardworks Ltd 2005  The s-Block Elements Elements of Groups IA* (the alkali metals) and IIA* (the alkaline earth metals)  constitute the s-block elements  their outermost shell electrons are in the s orbital *Note: In the following, Groups IA and IIA are abbreviated as Groups I and II respectively. 2 of 32 © Boardworks Ltd 2005  The s-block elements 3 of 32 © Boardworks Ltd 2005  Contents Alkali Metals and Alkaline earth Metals Electron structure and reactivity Physical properties Reactions Uses Summary activities 4 of 32 © Boardworks Ltd 2005  Electron structures Group 1 Group 2 (Alkali Metals) Alkaline Earth Metals) Elements Symb. Elec. Struc. Elements Symb. Elec. Struc. Lithium Li [He], 2s1 Beryllium Be [He], 2s2 Sodium Na [Ne], 3s1 Magnesium Mg [Ne], 3s2 Potassium K [Ar], 4s1 Calcium Ca [Ar], 4s2 Rubidium Rb [Kr], 5s1 Strontium Sr [Kr], 5s2 Cesiumm Cs [Xe], 6s1 Barium Ba [Xe], 6s2 Francium Fr [Rn], 7s1 Radium Ra [Rn], 7s2 5 of 32 © Boardworks Ltd 2005  General Properties Group I (Alkali Metals) Group II (Alkaline Earth Metals) 1- Largest size in their 1- Smaller than gp. I. periods. 2- Density = m/v 2- Density = m/v Denser than gp. I Low density. 3- I.P. is higher than gp. I. 3- Low ionization potential. 4- Electronegativity is very 4- Electronegativity is very low but higher than gp. I. low. 6 of 32 © Boardworks Ltd 2005  General properties Alkali metals are different to typical (transition) metals, such as iron and copper. Unlike typical metals, alkali metals:  are soft and can be cut by a knife – softness increases down the group;  have a low density – lithium, sodium and potassium float on water;  have low melting and boiling points. However, alkali metals do share a few properties with typical metals, because:  they are good conductors of heat and electricity;  they are shiny – this is only seen when they are freshly cut. 7 of 32 © Boardworks Ltd 2005  Electron structure and reactivity The reactivity of alkali metals increases down the group. What is the reason for this?  The size of each element’s atoms, and the number of full electron shells, increases down increase in reactivity Li the groups. Na  This means that, down the groups, the electron in the outer shell gets further away from the K nucleus and is shielded by more electron shells. Rb  The further an electron is from the positive attraction of the nucleus, the easier it can be Cs lost in reactions.  This means that reactivity increases as the size of the atom increases. 8 of 32 © Boardworks Ltd 2005  1- Effect of light: as a result of low I.P. of gp. I Dec. I.P. Group I (Alkali Metals) Inc. Size Emit electrons when irradiated with light. The electron emitted by this way called photoelectron. The emission becomes more easier from up to down, Why?. Cs and K are used in photoelectric cells. Also they give characteristic flame colorations when the excited electrons go back to its levels in the flame test. 9 of 32 © Boardworks Ltd 2005  Flame colour When alkali metals are heated and added to a jar of oxygen, they burn fiercely with a coloured flame.  lithium  sodium  potassium burns with a burns with an burns with a red flame golden yellow lilac (violet) flame flame 10 of 32 © Boardworks Ltd 2005  2- Solubility, Hydration and conductivity: All the simple salts of gr I are soluble in water (polar solvent) The decrease in hydration from LiCs All lithium salts are hydrated Many sodium salts Few potassium salts No rubidium salts or cesium salts are hydrated The solvation number, n, is 4 for Li+ and Be2+ 6 for elements in periods 3 and 4 of the periodic table The strength of the bonds between the metal ion and water molecules in the primary solvation shell increases with ↑ Electrical charge, z, on the metal ion ↓its radius, r. 11 of 32 © Boardworks Ltd 2005  The primary shell of water molecules which hydrate a metal ion is forming a complex. Li+ is tetrahedrally surrounded by 4 H2O molecules but Na+ is octahedrally surrounded by 6 H2O molecules. A secondary layer of water molecules further hydrated the ions, are only held by weak ion- dipole attractive forces. Strength of such forces is ↑ when the size of metal ion ↓ 12 of 32 © Boardworks Ltd 2005  In general → → so than larger ions Smaller the cation, higher the charge density, bigger the size of the solvated ion, and lesser the mobility and so conductivity takes order: Fr+>Cs+>Rb+>K+>Na+>Li+ 13 of 32 © Boardworks Ltd 2005  2- Reducing Properties Group I (Alkali Metals) Group II (Alkaline Earth Metals) These elements all These elements all have one valence s have two valence s electron, and like to electrons, and like to form +1 ions. This form +2 ions. This makes them excellent makes them excellent reducing agents. reducing agents. 14 of 32 © Boardworks Ltd 2005  3- Cohesive Energy: It is the force holding the atoms or ions together in the solid state and determines the hardness. It depends on: 1- Number of electrons participate in the bonds. 2- Strength of the bonds formed. Group I (Alkali Metals) Group II (Alkaline Earth Metals) One electron participate in Two electron participate in bond. bonds. Low cohesive energy. High cohesive energy. Soft and the softness Hard and the hardness increase down the gp. decrease down the gp. …………..? …………..? 15 of 32 © Boardworks Ltd 2005  4- Melting and Boiling Points Group I (Alkali Metals) Group II (Alkaline Earth Metals) Low melting and boiling High melting and boiling point and decrease down point, but there is the gp. irregularity down the gp. Because the metals adopt different crystal structure. 181 oC 1287 oC 649 oC 63 oC 839 oC 788 oC 28.5 oC 16 of 32 © Boardworks Ltd 2005  5- Color and magnetism of compounds Group I (Alkali Metals) Group II (Alkaline Earth Metals) Compounds of group I and II are diamagnetic and colorless……………….?!? except those where the acid radical is colored, e.g. chromates and permanganates. KMnO4 purple K2Cr2O7 orange Na2CrO4 yellow Why…..?!? Ligand to metal charge transfers. In each case, oxygen transfers electrons to the empty d orbitals on the metal atom. 17 of 32 © Boardworks Ltd 2005  Contents Alkali Metals and Alkaline earth Metals Electron structure and reactivity Physical properties Reactions Uses Summary activities 18 of 32 © Boardworks Ltd 2005  1- Reactions with water How do alkali and alkaline earth metals react with water? 19 of 32 © Boardworks Ltd 2005  1- Reactions with water All alkali metals react readily with water. The reaction becomes more vigorous down the group, and creates a lot of heat. 2M + 2H2O  2MOH + H2 - Li + H H O O H Li Li + + - + H H H H Li + O H O This reaction creates alkaline hydroxide ions. This is why the group 1 elements are called the alkali metals. The reaction also produces a gas that can be ignited by a lighted splint. What is this gas? 20 of 32 © Boardworks Ltd 2005  1- Reaction with water Lithium is the least reactive of the alkali metals. When added to water, it fizzes and moves around slowly across the surface of the water. lithium + water  lithium + hydrogen hydroxide 2Li (s) + 2H2O (l)  2LiOH (aq) + H2 (g) Sodium fizzes more than lithium, and moves quickly 2Na (s) + 2H2O (l)  2NaOH (aq) + H2 (g) 21 of 32 © Boardworks Ltd 2005  1- Reaction with water Group II (Alkaline Earth Metals)  Less reactive than gp. I.  Be: doubt whether reacts with steam water to form the oxide BeO, or fails to react at all.  Mg: reacts with hot water.  Ca, Sr, and Ba react rapidly with cold water. M+2H2O  M(OH)2+H2 M = Ca, Sr or Ba 22 of 32 © Boardworks Ltd 2005  2- Reactions with air All alkali metals react with air to form metal oxides. This produces a layer of dull oxide on the surface of the metal, called tarnish. The speed with which alkali metals react with air increases down the group:  lithium – tarnishes slowly;  sodium – tarnishes quickly;  potassium – tarnishes very quickly. Why are alkali metals stored in oil? The oil prevents them from reacting with air and tarnishing 23 of 32 © Boardworks Ltd 2005  Equations for reaction with air The reaction between an alkali metal and air is an example of an oxidation reaction: lithium + oxygen  lithium oxide 4Li (s) + O2 (g)  2Li2O (s) What are the word and chemical equations for the reaction of sodium and air? sodium + oxygen  sodium oxide 4Na (s) + O2 (g)  2Na2O (s) 24 of 32 © Boardworks Ltd 2005  2- Reaction with Air ( O2, N2) Group I (Alkali Metals) Group II (Alkaline Earth Metals) Oxide Oxide, Nitride 25 of 32 © Boardworks Ltd 2005  2- Reaction with Air ( O2, N2) Group I (Alkali Metals) Group II (Alkaline Earth Metals) All gp. I are very reactive All gp. II burn in air to and tarnish rapidly in air form mixture of simple to form the oxide oxide and nitrides. M + Air different types of M + Air MO + M3N2 oxides  Be is relatively Li is the only metal that unreactive. It does not react below 600 oC. But forms a mixture of oxide and nitride. powder is much more Li + Air Li O + Li N reactive and burns. 2 3 Li3N+ 3H2O3LiOH + NH3 26 of 32 © Boardworks Ltd 2005  Reaction with Oxygen Group I (Alkali Metals)  Types of oxides: Li forms Li2O O2- Na forms Na2O2 [-O-O-]2- Others form MO2 O2- O2- is monoxides, diamagnetism and colorless. [-O-O-]2- is peroxides, colored due to defects in crystal lattice, and diamagnetic. O2- is superoxides, colored due to defects in crystal lattice and the presence of unpaired electron, and paramagnetic. 27 of 32 © Boardworks Ltd 2005 c) Why do small metals form normal oxides but large metals form peroxides on heating in oxygen? Small 2+ ion Peroxide ion 2 Oxygen atoms each carrying 2+ a negative charge, and joined by a covalent bond The small 2+ ion close to the peroxide ion attracts electrons in the peroxide ion strongly towards the positive ion. This is on the way to form a simple oxide ion if the right-hand oxygen atom breaks off. This means the small positive ion polarizes the negative ion. Ions of the metals at the top of the Group have such a polarizing power that any peroxide ion near them falls to pieces to give an oxide and oxygen. As you go down the Group and the positive ions get bigger, they don't have so much effect on the peroxide ion. 28 of 32 © Boardworks Ltd 2005  Reaction with Nitrogen Group I (Alkali Metals) Li is the only metal that reacts with N2, to form nitride. 6Li + N2 2Li3N  + 3H2O 6Li+ N2 3LiOH+ NH3 29 of 32 © Boardworks Ltd 2005  Reaction with Nitrogen Group II (Alkaline Earth Metals) Be burns in N2 and form volatile, covalent nitride. 3Be + N2 Be3N2 Other gp. II burns in N2 and form nonvolatile, ionic nitride. 3Ca + N2 Ca3N2  + 6 H2 O 3Ca+ N2 3Ca(OH)2+ 2NH3 30 of 32 © Boardworks Ltd 2005  Why do Li and group II metals form nitrides on heating in air?  Formation of the nitride needs energy to break the strong triple bonds in N2 to form N3- and to form the metal ions.  This energy is provided by the lattice energy (Energy evolved when the ions come together to produce the crystal lattice).  The lattice energy increases as the attraction between the ions increases, i.e: by decreasing the size and increasing the charge.  In Group II, the attractions between the 2+ metal ions and the 3- nitride ions are big enough to produce very high lattice energies enough to form the nitride. In case of Li, its small size is enough to produce high lattice energy to form Li3N. 31 of 32 © Boardworks Ltd 2005  3- Hydroxides. Group I (Alkali Metals) Group II (Alkaline Earth Metals) Hydroxides of group I  Be(OH)2 is amphoteric. are the strongest bases, 2NaOH 2HCl called (Caustic Alkali). Na2BeO2 + 2H2O BeCl 2H O Li2O+ H2O  2LiOH Sodium beryllate 2+ 2 Na2O2+2H2O2NaOH + H2O2  MgO reacts with H2O KO2+2H2OKOH+H2O2+1/2 O2 forming Mg(OH)2 (weak base). Decrease Basisty  CaO reacts very readily with water forming Ca(OH)2 (lime Increase water, moderately strong base)  Ba hydroxides are even stronger bases. 32 of 32 © Boardworks Ltd 2005  4- Hydrides. Group I (Alkali Metals) All metals react with H2 and form ionic hydrides M +H -  The reaction decrease in the order LiCs.  Reducing agents that react with water liberating hydrogen (source of hydrogen). LiH + H2O  LiOH + H2 33 of 32 © Boardworks Ltd 2005  4- Hydrides. Group II (Alkaline Earth Metals)  All except Be and Mg form ionic hydrides MH2  Beryllium and magnesium hydrides are covalent and polymeric (Intermediate Hydrides).  Structure of (BeH2)n Three-center bonds are formed, in which a banana- shaped molecular orbital covers three atoms Be-H- Be, and contains two electrons. H H H Be Be Be Be H H H 34 of 32 © Boardworks Ltd 2005  H H H Be Be Be Be H H H Three-center two-electron bonds 35 of 32 © Boardworks Ltd 2005  5- Halides. Group I (Alkali Metals)  Li forms trihydrates LiX.3H2O, but the other alkali metal halides form anhydrous crystals.  All the halides adopt a NaCl structure with a coordination number of 6.  CsBr and CsI have a CsCl structure with a coordination number of 8. 36 of 32 © Boardworks Ltd 2005  5- Halides. Group II (Alkaline Earth Metals)  Beryllium halides are covalent, hygroscopic and fume in air due to hydrolysis.  In the gas phase, beryllium halides are monomeric and dimeric. Cl Cl Be Cl Cl Be Be Cl Cl 37 of 32 © Boardworks Ltd 2005   In the solid states, beryllium halides polymerize in order to achieve 4 coordination. 2BeO + CCl4 2BeCl2 + CO2 800 OC Cl Cl Cl Be Be Be Be Cl Cl Cl (Coordinate-Covalent Bond)  The other halides are ionic.  The halides are hygroscopic and form hydrates. CaCl2 is a well known drying agent. 38 of 32 © Boardworks Ltd 2005   Acidity of Be Salts:  Say Why…BeCl2, BeSO4 and Be(NO3)2 are ionic and have acidic character when dissolved in water...? Beryllium halides are covalent. While, hydrated Be salts are ionic with acidic character.  1- Ionic Character: Be atom 1s 2s 2p Be2+ ion 1s 2s 2p 1s 2s 2p Be2+ ion having gained 4 lone pairs from 4 oxygen atoms of 4 H2O 39 of 32 © Boardworks Ltd 2005   Forming the hydrated complex increases the effective size of 2+ the Be ion, thus spreading the charge over a large area, according to Fajan’s rule stable ionic salts are formed. 2- Acidic Character: H 2+ + + H2O + [ (H2O)3Be O ] H [Be (H2O)3(OH)] + 3O H  The bond between Be and O is strong bond, the strength of this bond weakens the O-H bonds, hence there is a tendency to lose protons. 40 of 32 © Boardworks Ltd 2005  6- Reaction with Carbon Group I (Alkali Metals) Li is the only metal that reacts directly with C.  Li + 2C Li2C2 The other metals react with acetylene. Na+ C2H2  NaHC2 Na2C2 Sodium hydrido Sodium carbide carbide Na2C2 + 2H2O 2NaOH + C2H2 [C  C]2- [C  C  H] - Carbide Ion Hydrido Carbide Ion 41 of 32 © Boardworks Ltd 2005  6- Reaction with Carbon Group II (Alkaline Earth Metals) Be form carbide of formula BeC2 1900- 2000 OC BeO + C BeC2 + O2 Be form carbide of formula BeC2  Be+ C2H2 BeC2 + H2 Mg, Ca, Sr, Ba, form carbides of formula MC2, by direct reaction with carbon. Ca + 2C CaC2 1100OC 42 of 32 © Boardworks Ltd 2005  7- Reaction with organic Compounds Group I (Alkali Metals) The alkali metals replace hydrogen in organic acids forming salts such as sodium acetate CH3COONa and potassium benzolate C6H5COOK. CH3COOH + NaOH → CH3COONa + H2O Group II (Alkaline Earth Metals)  Magnesium forms Grignard reagents, which are probably the most versatile reagents in organic chemistry. Mg + RBr ether RMgBr Grignard Reagent 43 of 32 © Boardworks Ltd 2005  8- Formation of Complexes Complexes are favored by positive ions that: 1- Small size. 2- Great charge. 3- Empty orbital. Group I (Alkali Metals) Group II (Alkaline Earth Metals)  Too large to form  Be forms many complexes readily. complexes [BeF4]2- and  Polydentate ligands Ba very few. such as crown ethers  Chlorophyll (green form complexes with plant pigment, complex alkali metal ions. of Mg+2 and porphyrin). O  Ca, Ba form O O M+ complexes with O O strong ligands such O as EDTA 18-crown-6 alkali metal complex 44 of 32 © Boardworks Ltd 2005  Structure of [BeF4]2- complex Electronic structure of Be atom (ground state) 2s2 2p Electronic structure of Be in excited state 2s1 2p1 Electronic structure of Be in gaseous BeF2 sp hybridization Electronic structure of Be in [BeF4 ]-2. 2 F- ions donate 2 electron pairs and a tetrahedral complex is formed. 2 sp3 hybridization F Be F F F 45 of 32 © Boardworks Ltd 2005  Chlorophyll (green plant pigment, complex of Mg+2 and porphyrin). Chlorophyll 6CO2 + 6H2O C6H12O6 + 6O2 sunlight 46 of 32 © Boardworks Ltd 2005  HO O C-CH2 CH2C O OH N-CH2-CH2-N HO O C-CH2 CH2C O OH EDTA 47 of 32 © Boardworks Ltd 2005  Anomalous or abnormal behavior of Lithium A. Differences between lithium and the other Group I Elements Reasons of differences: 1. Its small size 2. High electronegativity 3. Absence of d orbitals Examples of differences: 1- The melting and boiling points of Li are high. 2- Li is much harder. 3- Li reacts the least readily with oxygen forming the normal oxide and the higher oxides being unstable. 4- Li is much less electropositive; therefore many of its compounds are less stable, 5- Li compounds are more heavily hydrated than those of the rest. 6- Li has a greater tendency to form complexes than the heavier elements, and ammoniated salts such as [Li(NH3)]4I exist as solids.. 48 of 32 © Boardworks Ltd 2005  Anomalous or abnormal behavior of Beryllium A. Differences between Be and the other Group 2 Elements Reasons of differences: 1. Its small size 2. High electronegativity 3. Absence of d orbitals Examples of differences: 1- Be forms complexes not typical of Group 1 and 2. 2- Be is amphoteric. 3- Be salts are extensively hydrolyzed to give the tetrahedral complex ion [Be(H2O)4]+2 and stable ionic. 4- Be salts are acidic in pure water (why?). 5- Be salts are among the most soluble known. 49 of 32 © Boardworks Ltd 2005  Contents Alkali Metals Electron structure and reactivity Physical properties Reactions Uses Summary activities 50 of 32 © Boardworks Ltd 2005  Uses of lithium Lithium and its compounds are used in:  batteries – elemental lithium is used in non-rechargeable batteries. Lithium compounds are used in lithium-ion batteries, which are rechargeable.  alloys – with other metals, such as aluminium, copper and manganese, for use in aircraft parts.  medical treatment – lithium carbonate is sometimes used to treat mental illnesses such as depression. 51 of 32 © Boardworks Ltd 2005  Uses of sodium Elemental sodium is used in:  street lights – sodium vapour gives them their yellow glow.  nuclear reactors – used as a coolant due to its good conductivity and low melting point. Sodium compounds are in many household products:  sodium chloride – table salt  sodium hydrogencarbonate – bicarbonate of soda  sodium hydroxide – oven cleaner 52 of 32 © Boardworks Ltd 2005  Uses of potassium Potassium compounds are used in:  fertilizers – potassium is an essential element for plants. It is usually added as a chloride, sulfate, nitrate or carbonate.  fireworks and explosives – as potassium nitrate and potassium chlorate.  food preservation – as potassium nitrate. 53 of 32 © Boardworks Ltd 2005  Problems a. Why are beryllium salts acidic when dissolved in water? b. Why do small metals form normal oxides but large metals form peroxides on heating in oxygen? c. Discuss in brief the anomalous behavior of Li. d. Discuss briefly the abnormal behavior of Be. e. Compare between group 1 and 2 in the following points. 1. Reducing properties. 2. Softness. 3. Reactions with water, oxygen and hydrogen. f. Comment on the structure of beryllium hydride and beryllium halides. g. Uses of alkali metals and its applications 54 of 32 © Boardworks Ltd 2005