CHM 101 Inorganic Chemistry Past Paper

**CHM 101 Dr. ONIPEDE's Part** **Recommended texts: 1. Concise Inorganic Chemistry by J.D. Lee, 2. Inorganic Chemistry by Catherine E. Housecroft and Alan G. Sharpe, Modern Inorganic Chemistry by C. Chamber and A.K. Holliday** **Comparative chemistry of group IA, group IIA and group 14.** **Group...

**CHM 101 Dr. ONIPEDE's Part** **Recommended texts: 1. Concise Inorganic Chemistry by J.D. Lee, 2. Inorganic Chemistry by Catherine E. Housecroft and Alan G. Sharpe, Modern Inorganic Chemistry by C. Chamber and A.K. Holliday** **Comparative chemistry of group IA, group IIA and group 14.** **Group IA (Hydrogen \[H\] (1s^1^) Lithium \[Li\] (1s^2^2s^1^) Sodium \[Na\] (1s^2^2s^2^2p^6^3s^1^), Potassium \[K\] (1s^2^2s^2^2p^6^3s^2^3p^6^4s^1^), Rubidium \[Rb\] (1s^2^2s^2^2p^6^3s^2^3p^6^3d^10^4s^2^4p^6^5s^1^) Cesium \[Cs\] (1s^2^2s^2^2p^6^3s^2^3p^6^3d^10^4s^2^4p^6^4d^10^5s^2^5p^6^6s^1^), Francium \[Fr\] (\[Rn\] 7s^1^),** **Rn: 1s^2^2s^2^2p^6^3s^2^3p^6^3d^10^4s^2^4p^6^4d^10^5s^2^5p^6^4f^14^5d^10^6s^2^6p^6^** Group IA are called the alkali metals; they are all silvery-white metals but Cs is golden yellow metal and excellent conductors of electricity, they are soft and highly reactive metals, (the softness and low melting point is as a result of the elements contributing only one electron to the molecular orbital). they have one electron in their outer valence electron and thus form univalent ionic and colourless compounds the oxides and hydroxides are strong bases and the oxosalts are very stable. Because of the one valence electron in their outmost shell, they have the largest atomic size in their period. Nonetheless on losing the single valence electron, the outermost shell is completely lost in ionization; hence they have the smallest ionic size in their period. They also have the lowest ionization energy in their respective period. They are very reactive and form simple ionic compound which are soluble in water, they are reactive with air and water, and thus they are stored under inert solvent (hydrocarbon solvent). They are found combined in nature as a result of their reactivity. They readily form alloy among themselves (e.g. Na/K) and with other metals (e.g. Na/Hg), they tarnish rapidly in air giving a layer of oxide, peroxide and dioxide. At ambient temperature all group IA, adopt a body centre- cubic structure but at low temperature lithium forms hexagonal close-packed structure. Lithium is the first member of the group and show marked difference from other members of the group. But the head element usually has the diagonal relationship with the diagonally next element to it. Sodium and potassium make about 4% of the earth crust. Despite their close chemical similarities the elements do not occur together primarily because their ions are of different sizes. The ores of group IA include tourmaline (Na~3~Li~3~Al~6~(BO~3~)~3~(SiO~3~)~6~F~4~) spodumene LiAl(SiO~3~)~2~ petalite (LiAl(Si~2~O~5~)~2~, borax (Na~2~B~4~O~7~.10H~2~O), mirabilite (Na~2~SO~4~), sylvite (KCl) carnallite (KCl.MgCl~2~.6H~2~O), lepidolite K(Li,Al)~3~(Al,Si,Rb)~4~O~10~(F,OH)~2~ avogadrite (K,Cs)BF~4~), londonite ((Cs,K)Al~4~B~4~(B,Be)~12~O~28~ while francium is found in trace amount in ores of uranium and thorium. The metals are all reactive and hence cannot be extracted by the furnace, but could only be extracted by electrolysis. In flame lithium gives crimson, sodium gives yellow colour, potassium gives lilac, rubidium gives red-violet and caesium gives blue colour. Some selected properties of group IA metal are as tabulated below **Lithium** **Sodium** **Potassium** **Rubidium** **Cesium** **Francium** ------------------------------------------------------------------------------------------------------------------------------------ ------------- ------------ --------------- -------------- ------------ -------------- atomic symbol Li Na K Rb Cs Fr atomic number 3 11 19 37 55 87 atomic mass 6.94 22.99 39.10 85.47 132.91 223 valence electron configuration 2*s*^1^ 3*s*^1^ 4*s*^1^ 5*s*^1^ 6*s*^1^ 7*s*^1^ melting point/boiling point (°C) 180.5/1342 97.8/883 63.5/759 39.3/688 28.5/671 27/--- density (g/cm^3^) at 25°C 0.534 0.97 0.89 1.53 1.93 --- atomic radius (pm) 167 190 243 265 298 --- first ionization energy (kJ/mol) 520 496 419 403 376 393 most common oxidation state +1 +1 +1 +1 +1 +1 ionic radius (pm)\* 76 102 138 152 167 --- electron affinity (kJ/mol) −60 −53 −48 −47 −46 --- electronegativity 1.0 0.9 0.8 0.8 0.8 0.7 standard electrode potential (*E*°, V) −3.04 −2.71 −2.93 −2.98 −3.03 --- product of reaction with O~2~ Li~2~O Na~2~O~2~ KO~2~ RbO~2~ CsO~2~ --- type of oxide basic basic basic basic basic --- product of reaction with N~2~ Li~3~N none none none none --- product of reaction with X~2~ LiX NaX KX RbX CsX --- product of reaction with H~2~ LiH NaH KH RbH CsH --- **\*The values cited are for four-coordinate ions except for Rb^+^ and Cs^+^, whose values are given for the six-coordinate ion.** **Comparative relationship between lithium and magnesium** 1. 2. 3. 4. 5. 6. 7. 8. 9. **Chemical properties of group IA metals** 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. **Differences between lithium and other group IA metals** 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. **Group IIA (Beryllium \[Be\] (1s^2^2s^2^), Magnesium \[Mg\] (1s^2^2s^2^2p^6^3s^2^), Calcium \[Ca\] (1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^), Strontium \[Sr\] (1s^2^2s^2^2p^6^3s^2^3p^6^3d^10^4s^2^4p^6^5s^2^) Barium \[Ba\] (1s^2^2s^2^2p^6^3s^2^3p^6^3d^10^4s^2^4p^6^4d^10^5s^2^5p^6^6s^2^), Radium \[Ra\] (\[Rn\]7s^2^)** **Rn: 1s^2^2s^2^2p^6^3s^2^3p^6^3d^10^4s^2^4p^6^4d^10^5s^2^5p^6^4f^14^5d^10^6s^2^6p^6^** They are called alkaline earth metals; they are all silvery in colour but Be and Mg are greyish metal they are highly reactive metals; but are less reactive than group IA counterparts. They are divalent and form colourless ionic compound, the oxides and hydroxides are less basic than their group IA counterpart. Thus their oxosalts such as carbonates, sulphates, and nitrates are less stable to heat. Their atomic and ionic sizes are smaller than their group IA counterpart but they have higher densities than group IA metals. The group IIA metals have two valence electrons which participate in bonding, hence are harder, have higher cohesive energy as well as higher melting and boiling point than their group IA counterpart. Beryllium is the head element in the group and hence shows marked difference from other members of the group. They all dissolve in liquid ammonia like their group IA counterpart to give a blue solution. The solutions decompose very slowly forming amide and hydrogen. Evaporation of the ammonia from group IA metals yields the metals, but with group IIA metals evaporation of ammonia gives hexammoniates of the metals; they slowly decompose to give amides. M(NH~3~)~6~ M(NH~2~)~2~ + 4NH~3~ + H~2~ They are highly reactive metals and hence cannot be extracted by oxidation or reduction in the furnace, but can only be purified by electrolysis. The ores of alkaline earth metals include beryl; Be~3~Al~2~Si~6~O~18~, phenacite ; Be~2~SiO~4~, kieserite: MgSO~4~.H~2~O, carnallite; KCl.MgCl~2~.6H~2~O, Limestone; CaCO~3~, Gypsum; CaSO~4~.2H~2~O, celestite; SrSO~4~ and strontianite SrCO~3~, Baryte BaSO~4~, radium is radioactive and found in uranium ore pitch blende. The flame test of beryllium is white, magnesium is brilliant white, calcium; brick red, strontium crimson, barium apple-green, radium; crimson red. Some selected properties of alkaline earth metals are as tabulated below **Beryllium** **Magnesium** **Calcium** **Strontium** **Barium** **Radium** ---------------------------------------- --------------- --------------- ------------- --------------- ------------ ------------ atomic symbol Be Mg Ca Sr Ba Ra atomic number 4 12 20 38 56 88 atomic mass 9.10 24.31 40.08 87.62 137.34 226 valence electron configuration 2*s*^2^ 3*s*^2^ 4*s*^2^ 5*s*^2^ 6*s*^2^ 7*s*^2^ melting point/boiling point (°C) 1287/2500 649/1105 839/1494 768/1381 727/1850 700/1700 density (g/cm^3^) at 25°C 1.85 1.74 1.55 2.62 3.62 5.5 atomic radius (pm) 112 160 197 215 222 --- first ionization energy (kJ/mol) 899 737 590 549 503 509 most common oxidation state +2 +2 +2 +2 +2 +2 ionic radius (pm)\* 31 72 100 118 135 148 electron affinity (kJ/mol) ≥ 0 ≥ 0 -- 2 -- 5 -- 14 -- electronegativity 1.5 1.2 1.0 1.0 0.9 standard electrode potential (*E*°, V) -- 1.85 -- 2.38 -- 2.87 -- 2.89 -- 2.90 -- 2.92 **Chemical properties of group IIA metals** 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. **Comparative relationship between beryllium and aluminum** 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. **Group 14 (Carbon \[C\] (1s^2^2s^2^2p^2^), Silicon \[Si\] (1s^2^2s^2^2p^6^3s^2^3p^2^), Germanium \[Ge\] (1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^10^4p^2^), Tin \[Sn\] (1s^2^2s^2^2p^6^3s^2^3p^6^3d^10^4s^2^4p^6^5s^2^4d^10^5p^2^) Lead \[Pb\] (1s^2^2s^2^2p^6^3s^2^3p^6^3d^10^4s^2^4p^6^4d^10^5s^2^5p^6^4f^14^5d^10^6s^2^6p^2^) Flerovium \[Fl\] (1s^2^2s^2^2p^6^3s^2^3p^6^3d^10^4s^2^4p^6^4d^10^5s^2^4f^14^5d^10^6s^2^6p^6^5f^14^6d^10^7s^2^7p^2^)** **Rn: 1s^2^2s^2^2p^6^3s^2^3p^6^3d^10^4s^2^4p^6^4d^10^5s^2^5p^6^4f^14^5d^10^6s^2^6p^6^** The group members varies from non-metal (C), to metalloid (Si, Ge) to metal (Sn, Pb) , they are the most abundant element found on earth as they constitute the earth structure as well as biological system (such as proteins, carbohydrates, and fats), all except lead exhibit allotropy. Silicon is the second most abundant element by weight in the earth crust and it constitute 24% of the earth crust. The ores of group 14 elements include diamond, graphite, fullerene, charcoal, peat, Quartz; SiO~2~, zeolite;, Na~2~(Al~2~Si~3~O~10~)2H~2~O, Cassiterite; SnO~2~, teallite; PbSnS~2~, Galena; PbS, Anglesite; PbSO~4~, [boulangerite](https://en.wikipedia.org/wiki/Boulangerite); Pb~5~Sb~4~S~11~, Winklers' ore (75% Ag, 15% S 7% Ge) Some selected properties of group 14 elements +---------+---------+---------+---------+---------+---------+---------+ | | **Carbo | **Silic | **Germa | **Tin** | **Lead* | **Flero | | | n** | on** | nium** | | * | vium** | +=========+=========+=========+=========+=========+=========+=========+ | atomic | C | Si | Ge | Sn | Pb | Fl | | symbol | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | atomic | 6 | 14 | 32 | 50 | 82 | 114 | | number | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | aomic | 12.01 | 28.09 | 72.64 | 118.71 | 207.2 | 289.19 | | mass | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | valence | 2*s*^2^ | 3*s*^2^ | 4*s*^2^ | 5*s*^2^ | 6*s*^2^ | 7s^2^7p | | electro | 2p^2^ | 3p^2^ | 4p^2^ | 5p^2^ | 6p^2^ | ^2^ | | n | | | | | | | | configu | | | | | | | | ration | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | melting | 4100/51 | 1420/32 | 945/285 | 232/262 | 327/175 | 61/150 | | point/b | 00 | 80 | 0 | 3 | 1 | | | oiling | | | | | | | | point | | | | | | | | (°C) | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | density | 2.2 | 2.33 | 5.32 | 7.27(Wh | 11.30 | 11.7 | | (g/cm^3 | (Graphi | | | ite) | | | | ^) | te) | | | | | | | at 25°C | | | | | | | | | 3.5 | | | | | | | | (diamon | | | | | | | | d) | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | atomic | 77 | 117 | 122 | 162 | 175 | 180 | | radius | | | | | | | | (pm) | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | first | 1086 | 786 | 760 | 707 | 715 | 832.2 | | ionizat | | | | | | | | ion | | | | | | | | energy | | | | | | | | (kJ/mol | | | | | | | | ) | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | most | +4 | +4 | +4 | +4 | +4 | +4 | | common | | | | | | | | oxidati | | | | | | | | on | | | | | | | | state | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | ionic | 29 | 40 | 53 | 69 | 77.5 | | | radius | | | | | | | | (pm)\* | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | electro | -- 122 | -- 134 | -- 119 | -- 107 | -- 35 | | | n | | | | | | | | affinit | | | | | | | | y | | | | | | | | (kJ/mol | | | | | | | | ) | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | electro | 2.6 | 1.9 | 2.0 | 2.0 | 1.8 | | | negativ | | | | | | | | ity | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ | standar | 0.21 | -- 0.86 | -- 0.18 | -- 0.12 | 0.79 | | | d | | | | | | | | electro | | | | | | | | de | | | | | | | | potenti | | | | | | | | al | | | | | | | | (*E*°, | | | | | | | | V) | | | | | | | +---------+---------+---------+---------+---------+---------+---------+ Differences between carbon silicon and other elements 1. 2. 3. 4. Comparative relationship between B and Si 1. 2. 3. Reactions of group 14 elements 1. 2. 3. 4. 5. Atoms, molecules and structures To predict structures of atoms in a molecule a law must be obeyed which depicts the orientation of the atoms in the molecule. One of such rules is the Lewis octet rule. The Lewis octet rule state that elements tend to bond in such a way that each atom has eight electrons in its valence shell giving it the same configuration as a noble gas. The law could be obeyed in the following steps 1. 2. 3. 4. Another law that determines the shape of molecules is the Valence Shell Electron-Pair Repulsion (VSEPR) theory the principle is that each group of valence electrons around a central atom is located as far away as possible from the others in order to minimize repulsions. Rule 1; Regions of high electron concentration (bonds and no lone pairs on the central atom) repel one another and to minimize their repulsions, these regions move as far apart as possible while maintaining the same distance from the central atom. e.g. A molecule with two atoms attached to the central atom BeCl~2~, there is no lone pair on the central atom. To be as far as possible, the two bonding pair lie on the opposite sides of the Be atom and so the electron arrangement is linear and hence VSEPR model predicts a linear shape with angle 180^o^ Boron trifluoride molecule BF~3~; there are three bonding pairs attached to the central atom and no lone pairs, to be as far as possible, the three bonding pair must lie at the corners of an equilateral triangle. Because a F atom is attached to each bonding pair, the BF~3~ molecule predicted by VSEPR to be trigonal planar with angle 120^o^  Methane CH~4~; there are four bonding pairs on the central atom, to be far apart as possible, the four pairs the molecule must take up a tetrahedral arrangement around the C atom. Because the electron arrangement is tetrahedral and an H atom is attached to each bonding pair, the molecule is predicted by VSEPR to be tetrahedral with angle 109.5^o^ Phosphorus pentachloride PCl~5~: there are five bonding pairs and no lone pairs on the central atom. VSEPR predicts they be far apart as possible hence three atoms lie in the corners of equilateral triangle and two lie above and below the plane of the triangle; bond on the equatorial plane are 120^o^ the angle between axial and equatorial atoms is 90^o^ , the axial ClPCl bond angle is 180^o^ and the shape is trigonal bipyramidal.  Sulfur hexafluoride SF~6~; there are six atoms attached to the central S atom and no lone pairs on the central atom; VSEPR predicts that four pair of the atom be at the corner of a square on the equator and the two pairs above and below the plane of the square, All its bond angles are either 90^o^ or 180^o^ and all the F atoms are equivalent. The shape is predicted to be octahedral. Rule 2: there is no distinction between single and multiple bonds; a multiple bond it treated as a single region of high electron concentration. Two electron pairs in a double bond stay together and they repel other bond or lone pair as a unit. The three electron pairs in a triple bond also stay together and act like a single region of high electron concentration. e.g. CO~2~ has similar structure to BeCl~2~ even though CO~2~ has double bonds, also C[*O*~3~^2−^]{.math.inline} , the two pairs of electron in the double bond are treated as a unit and hence the shape is as given by VSEPR is given as trigonal planar. When there is more than one central atom, we consider the bonding about each atom independently, e.g. ethylene CH~2~=CH~2~ each carbon has three atoms attached but no lone pairs, the angle around each carbon is trigonal planar with angle 120^o^ for HCH and HCC Nitrate ion has nitrogen as the central atom and all oxygen attached to the central atom and it has a trigonal planar structure.  Molecule with lone pairs on the central atom VSEPR use AX~n~E~m~ to identify different combinations of atoms and lone pairs attached to the central atom. We let A represent a central atom, X an attached atom and E a lone pair. The sulfite ion S[*O*~3~^2−^]{.math.inline} is an example of AX~3~E. If there are no lone pairs on the central atom (an AX~n~ molecule), each region of high electron concentration corresponds to an atom and so the molecular shape is the same as electron arrangement. However if lone pairs are present, the molecular shape differs from the electron arrangement because only positions of the atoms are considered when naming the shape. Example is the four regions of high electron concentration in S[*O*~3~^2−^]{.math.inline} are farthest apart if they adopt a tetrahedral arrangement. However the shape of the ion is described by the location of the atoms, not the lone pair. Because only three of the tetrahedral location are occupied by atoms, the shape of an S[*O*~3~^2−^]{.math.inline} ion is trigonal pyramidal. This gives the rule; Rule 3: All regions of high electron density, lone pairs and bonds are included in a description of the electronic arrangement, but only the position of atoms are considered when reporting the shape of a molecule. A single unpaired electron on the central atom also is a region of high electron density and is treated like a lone pair when determining molecular shape. e.g. NO~2~ has a single nonbonding electron, a "lone half-pair." Thus NO~2~ has a trigonal planar electron arrangement (including the unpaired electron on N), but its shape is angular (V shaped).  (V shape) Because bond angles in molecule with lone pairs are typically smaller than expected, lone pairs are treated in the VSEPR model as having a strong repelling effect than do electrons in bonds. i.e. the lone pairs push the atoms bonded to the central atom closer together. We rationalize that the electron cloud of a lone pair can spread over a larger volume than a bonding pair, because a bonding pair (or several bonding pairs in a multiple bond) is pinned down by two atoms not one. This leads to the rule; Rule 4: The strengths of repulsions are in the order lone pair-lone \> lone pair-atom \> atom-atom. The lowest energy is achieved when lone pairs are as far from each other as possible. The energy is also lowest if the atoms bonded to the central atom are from lone pairs, even though that might bring the atoms closer to other atoms. We can predict electron arrangement in an AX~4~E molecule or ion, such as I[*F*~4~^+^]{.math.inline} is trigonal bipyramidal, but there are two different possible locations for the lone pair. An axial lone pair lies on the axis of the molecule, where it repels three electron pairs strongly, an equatorial lone pair lies on the molecule's equator on the plane perpendicular to the molecular axis, where it repels only two electron pair strongly. Therefore the lowest energy is achieved when a lone pair is equatorial, producing a seesaw-shaped molecule An AX~3~E~2~ molecule such as ClF~3~, also has a trigonal bipyramidal arrangement of electron pairs, but two of the electron pairs are lone pair. These two pairs are farthest apart if they occupy two of the three equitorial positions but move away from each other slightly. This gives a T-shaped molecule.  An AX~4~E~2~ molecule such as I[Cl~4~^−^,]{.math.inline} which has an octahedral arrangement of electron pairs, two of which are lone pairs, the two lone pairs are farthest apart when they lie opposite each other and so the molecule is square planar.