Inorginic Chemistry for 2nd class Main Group Chemistry INORGANIC CHEMISTRY for 2nd class Chemistry of Main Group Elements by Dr. Khalil K. Abid Department of Chemistry , College of Science University of Mustansiriyah, Baghdad – IRAQ Email: [email protected]; [email protected] 1 Inorginic Chemistry for 2nd class Main Group Chemistry THE s-BLOCK ELEMENTS 1 – Characteristic properties of the s-block elements: （a）Metallic character and low electronegativity 2 Inorginic Chemistry for 2nd class Main Group Chemistry The members of Group I and II are all metals. They are silvery coloured and tarnish in air. They show relatively weak metallic bonding because they have only one or two valence electrons. Alkali metals are all soft, silvery metals, with its lower member (like Cs) softer than its upper member and tarnishing more rapidly in air (therefore they are often kept immersed in paraffin). They all crystallize with body-centred cubic lattice, so that all of them have low densities as there are more space in the metallic lattice. The low core charge(+1) of the atoms results in very poor control of the outer shells electrons, which are easily delocalized or lost from the atoms. As a result, metallic bond among atoms is weak and the metals can be easily cut with a knife, besides having low melting and boiling points. Alkaline earth metals are grey metals and are harder than alkali metals, although they can still be cut with a knife. The lower member barium is very soft, and it tarnishes rapidly. Metallic bond among atoms is stronger than their Group I counterparts; and the melting points and boiling points of the elements are all higher, as a result of the effect of increased core charge (+2) and decreased atomic radius. Beryllium and magnesium have hexagonal close packed structures, whilst calcium and strontium have face-centred cubic close packed structures -- the more efficient packing of the atoms accounts for the higher density of alkaline earth metals. All s-block elements have low electronegativity values, i.e., they are electropositive, with a tendency to lose their outer electrons relatively easily. As each group is descended, the elements become more and more electropositive -- the elements tend to lose electrons more readily. This is because the outer shell is further from the nucleus as each group is descended. Group II elements are more electro- negative (less electropositive) than Group I elements in the same period, as the nuclear charge has increased by one. （b）Formation of basic oxides and hydroxides 3 Inorginic Chemistry for 2nd class Main Group Chemistry Group I and II metal oxides have ionic lattices and are highly basic. Some oxides react exothermically with water to give the hydroxides. For example : CaO(s) + H O(l) → Ca(OH) (aq) 2 2 Magnesium oxide is slightly soluble in water and dissolves in acids to give salts : MgO(s) + 2H+(aq) → Mg2+(aq) + H O(l) 2 All Group I metal hydroxides are soluble in water. They are among the strongest bases known. The properties of Group II metal hydroxides are summarized as in table : (c）Predominantly ionic bonding with fixed oxidation state in their compounds The s-block elements form compounds that are predominantly ionic in nature, showing fixed oxidation states of +1 in Group I and +2 in Group II. Alkali metals have one electron in their outermost s-orbital, with completely filled inner orbitals. Hence they show only +1 oxidation state because their first ionization enthalpies are low. Consider the enthalpy change for the formation of NaCl (s) from Na metal and 2 chlorine gas in the standard state. Since the second electron of sodium must be removed from a noble gas electronic structure [Ne], its second ionization enthalpy is very high, about 8000 kJ mol-1. This amount of energy cannot be repaid by the exothermic processes, and therefore sodium is restricted to the chemistry of a monovalent ion with a fixed oxidation state of +1. Alkali earth metals have two electrons in their outermost s-orbital, with completely filled inner orbitals. Hence they show +2 oxidation state only as the sum of first and second ionization enthalpies is also relatively low. In all their compounds the elements of Group II have the oxidation state of +2. The considerable endothermic contribution from the double ionization of a Group II metal M to M+2 is compensated by the very large lattice enthalpy. However, the third ionization of M, corresponding to the removal of an electron from an inner shell, is so large that the energy required cannot be recovered. Hence these elements do not occur with oxidation state +3. 4 Inorginic Chemistry for 2nd class Main Group Chemistry （d）Characteristic flame color of salts S-block elements give characteristic colors in the flame test (except Be and Mg). In the flame test, their electrons are excited to a higher energy level. When the electron drops back to the original lower level, the extra energy is given out. In the case of s-block elements, the amount of energy given out is small; hence it lies in the visible light region. These elements, therefore, have colored flames : Group I Flame color Group II Flame color Li Red Be None Na Golden yellow Mg None K Lilac Ca Brick red Rb Red Sr Crimson Cs Blue Ba Apple green （e）Weak tendency to form complexes A complex is a polyatomic ion or molecule formed when molecular or ionic groups ( called ligands ) form dative covalent bonds with a central metal atom or cation. Complex formation is a common feature with d-block metal ions. The formation of such complexes is due to the presence of low energy empty d-orbitals in the d-block metal ions that can accept the lone pairs from the surrounding ligands. S- block metal ions rarely form complexes because they do not have low energy vacant orbitals available for the bonding with the lone pairs of surrounding ligands. The alkali metal ions have least tendency towards complex formation. The relatively large size of their ions, completely filled inner orbitals, and low charge (+1) are responsible for such behavior. The alkaline earth metal ions have a relatively greater tendency to form complexes than alkali metal ions. It can be explained in terms of the higher charge (+2) and smaller size of the Group II cations. Beryllium forms a particular wide range of complexes for this reason. 5 Inorginic Chemistry for 2nd class Main Group Chemistry 2 – Variation in properties of the s-block elements and their compounds （a）Variations in atomic radii, melting points hydration and ionization enthalpies Group I atoms and ions are larger than their Group II neighbors in the same period. As each Group II atom has one more electron in the outer shell and one more proton in the nucleus than has the Group I atom in the same period. Since the screening of the outer s- electrons from the nucleus is the same for both atoms, the nucleus of the Group II atoms exerts a stronger attraction on the outer electrons because it has one more proton. This greater ‘effective nuclear charge’ of Group II atoms reduces their atomic radii. The similarity in size of the lithium and magnesium ions leads to a number of chemical similarities in the compounds of the two elements. The hydration enthalpy is the enthalpy change when polar water molecules cluster around a metal ion : M+(g) + aq. → M+(aq) or M+2(g) + aq. → M+2(aq) As each group is descended, the enthalpy of hydration of an ion becomes less negative. The larger the ion, the weaker the electric field around it and water molecules are attracted less strongly. This means that fewer water molecules are attracted to the ion and the enthalpy of hydration, resulting from this electrostatic attraction is smaller. Group II ions have hydration enthalpies higher than their Group I counter- parts, as they are doubly charged positive ions of smaller ionic radius. This higher charge density results in a greater electrostatic attraction between water molecules and Group II ions and more negative enthalpy of hydration. Ionization enthalpies: For each element in Group I, the first ionization enthalpy is so much smaller than the other two. The outer s-electron is relatively easy to remove because it is well shielded from the nucleus by the inner electrons. Removal of a second electron involves the removal of an electron with stable noble gas electronic configuration from an inner shell, which is closer and not so well shielded from the nucleus; it is therefore more strongly held by the nucleus. The two outer s-electrons are relatively easy to remove because they are well shielded from the nucleus by the inner electrons. Removal of a third electron is relatively difficult because the inner electrons, which have stable noble gas electronic configuration, are closer and not so well shielded from the nucleus. Ionization enthalpy decreases as each group is descended. This is because the outer electrons are further out and progressively better shielded from the nucleus as each group is descended. 6 Inorginic Chemistry for 2nd class Main Group Chemistry Melting points: Most of the differences in melting points can be interpreted in terms of the strength of metallic bonds. The strength of these bonds is dependent on the following factors: 1. atomic size, 2. number of electrons in the outer shell of each atom, 3. metallic crystal structure. Melting points generally decrease down each group. A metal is regarded as a fairly closely-packed assembly of positive ions held together by a mobile electron ‘cloud’. As each group is descended, the number of electrons in the cloud remains the same but the ionic size increases. Therefore, the electron cloud becomes more diffuse resulting in weaker attractive forces between the cloud and the ions. The melting points in Group II are higher than for the corresponding Group I elements because the ions of Group II are doubly charged and smaller than their Group I neighbors. Thus, the alkaline-earth elements contribute two electrons per atom to the mobile electron ‘cloud’ whereas the alkali metals only contribute one electron per atom. The greater number of electrons and the smaller ionic size of the Group II elements help to increase the density of the electron cloud and thus increase the attractive forces between the mobile electrons and the cationic lattice. In Group I all the metals have the same crystal structure whereas those of Group II are of more than one type. Since crystal structure affects physical properties, the irregularity in melting points can be explained in terms of change of crystal structures. （b）Reactions of the elements with oxygen and water Alkali metals: The elements of Group I are called alkali metals because they react with water to form alkaline solutions. They are very reactive and have to be stored under paraffin oil to protect them from air and moisture. The alkali metals have low first ionization enthalpies and form ions easily by losing one electron thereby forming a univalent ion with a noble gas like structure. They react readily with water, dilute acids and most non-metals. The order of the reactivity of the elements is : Cs > Rb > K > Na > Li 7 Inorginic Chemistry for 2nd class Main Group Chemistry Alkaline earth metals: The Group II elements are called alkaline earth metals because many of their compounds are found as minerals in rocks. This is possible because the compounds are insoluble in water, so they are not washed away by rain. The alkaline earth metals are less reactive than alkali metals due to greater ionization enthalpies. The order of the reactivity of the elements is: Ba > Sr > Ca > Mg > Be Reaction with oxygen: With the exception of lithium and beryllium, all s-block elements produce more than one oxide. The three types of oxides are all ionic, and the ions are related as follows: O-2 → O -2 → O - 2 2 Lithium will melt when heated and eventually burns to form the monoxide and some lithium nitride: 2Li(s) + O (g) → Li O(s) 2 2 3Li(s) + N (g) → Li N(s) 2 3 Sodium burns in air with a yellow flame. The white product remaining is not pure sodium oxide. In fact, it is mainly sodium peroxide, Potassium, rubidium and Cesium ignite spontaneously to form peroxides and superoxides. All alkali earth metals burn brilliantly when heated in air to form a mixture of oxide and nitride. The rate of reaction and the proportion of nitride formed increases down the group: Mg(s) + O (g) → MgO(s) 2 3Mg(s) + N (g) → Mg N (s) 2 3 2 There are no Group II superoxides as the cations have too great polarizing power. When barium peroxide is heated, it decomposes to produce a colorless gas. Reaction with water: The reactivity of alkali metals increases markedly down the group. Lithium reacts quietly or very slowly with water. Sodium reacts vigorously with water. Potassium reacts so vigorously with water that the hydrogen formed catches fire because sufficient heat is generated in the reaction between potassium and water to ignite the hydrogen. Both rubidium and caesium react explosively. Beryllium does not react with water even when hot. Magnesium reacts very slowly with cold water but reacts vigorously with steam at red heat. Calcium reacts vigorously with cold water, producing hydrogen and whitish suspension of calcium hydroxide, some of which dissolves. Barium react very vigorously to produce hydrogen and barium hydroxide. 8 Inorginic Chemistry for 2nd class Main Group Chemistry （c）Reaction of the oxides with water, acids and alkalis All alkali metal oxides are basic oxides and thus react with water to form hydroxides, and with acids to form salt and water. O -(s) + H O (l) → 2OH-(aq) 2 2 O - (s) + 2H+ (aq) → H O (l) 2 2 Alkali metal oxides are decomposed by cold water. They react with acid readily and have no reaction with alkalis. Na O(s) + H O(l) → 2 Na+ (aq) + 2OH- (aq) 2 2 Li O(s) + 2HCl(aq) → 2 Li+ (aq) + 2Cl- (aq) + H O(l) 2 2 Group II metal oxides are much less soluble than those of Group I due to their higher lattice enthalpies. Except beryllium oxide, which is amphoteric, the other alkaline earth metal oxides are basic. The basic strength increases from magnesium oxide to barium oxide. This is because the ionic size increases from Mg+2 to Ba+2. Solubility of the oxides increases from magnesium to barium due to the decreasing magnitude of lattice enthalpies. Magnesium oxide is slightly soluble in water producing an insoluble hydroxide : MgO(s) + H O(l) → Mg+2 (aq) + 2 OH- (aq) 2 All the other oxides slake rapidly in water to give alkaline solutions of their hydroxides with considerable evolution of heat : CaO(s) + H O(l) → Ca+2 (aq) + 2 OH- (aq) 2 BaO(s) + H O(l) → Ba+2 (aq) + 2 OH- (aq) 2 All Group II oxides react with acids to form salts and water: MgO(s) + H SO (aq) → MgSO (aq) + H O(l) 2 4 4 2 Beryllium oxide is virtually insoluble. It is amphoteric and reacts with both acids and alkalis （d）Relative thermal stabilities of the carbonates and hydroxides The thermal stability of an ionic compound depends on the charge and the size of the ions. If the cation has a greater polarizing power (smaller size and higher charge), it will distort the electron cloud of the neighbouring large anion to a greater extent, and hence the compound becomes less stable. Most carbonates and hydroxides readily undergo 9 Inorginic Chemistry for 2nd class Main Group Chemistry decomposition on heating to give oxides. Down each group, as the size of the cation increases, the polarizing power of the cation decreases. The compound with large cation become more stable. Hence, the thermal stability of carbonates and hydroxides of both group I and Group II metals increases down the group. Group II metals are much smaller than Group I metal ions and have a higher charge. Consequently, their polarizing power is greater. Carbonates and hydroxides of Group II metals are less stable to thermal decomposition than those of Group I metals. Carbonates: Except for lithium carbonate, the alkali metal carbonates are stable to heat. The relative instability of lithium carbonate can be explained in terms of the gain in lattice enthalpy when the very small cation, Li+ is combined with the smaller oxide ion instead of the much larger carbonate ion. All Group II carbonates are decomposed on heating. MCO (S) → MO(s) + CO (g) 3 2 The thermal stability of the carbonates increases down the group because as the group is descended, the polarizing power of the cation decreases. Thus, cations of the upper elements distort the anion electron clouds far more than those of the lower elements do. Hydroxides: Lithium hydroxide is the only Group I metal hydroxides which decomposes at Bunsen burner temperatures because the small Li+ ion has considerable polarizing power which polarizes the electron cloud of hydroxide ion to such extent that decomposition occurs on heating. The thermal stability of Group II hydroxides increases down the group because polarizing power decreases with increasing size of the cation. Δ M(OH) (s) → MO(s) + H O(g) 2 2 Questions 1. Explain the relatively large difference between the second and the third ionization enthalpies for all Group II elements. 2. Describe and explain the trend in ionization enthalpies as each group is descended. 3. How do you expect the reactivity of the elements to change on going down each group 4. How would you expect the reactivity of the elements to change going across from Group I to Group II in the same period ? 5 – Group I & II elements seldom form complex. 6 – Ionic radius of any Group I or II element is smaller than the atomic radius? 10