INORGANIC CHEMISTRY

1. INORGANIC CHEMISTRY

2. Modern Periodic Table

3. Features of the Periodic Table • The Periodic Table is an arrangement of elements in order of increasing atomic number. • Each element in a horizontal row (Period) differs from the preceding element by addition of an electron to the electron shell and a proton to the nucleus. • Elements are arranged such that elements in a particular vertical column (Group) have the same number of electrons in its valence shell

4. ns2np6 ns1 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 d10 d1 d5 4f 5f Ground State Electron Configurations of the Elements 8.2

5. (1) The chemical and physical properties of the elements are periodic functions of the atomic number (number of protons in the nucleus = number of electrons in the neutral atom). (2) The elements can be arranged in groups (columns) of elements that possess related chemical and physical properties. (3) The elements can be arranged in periods (rows) of elements that possess progressively different physical and chemical properties

6. Orbitals Being Filled 1 8 Groups 2 1s 1 3 4 5 6 7 1s 2s 2 2p 3s 3p 3 4p 3d Periods 4s 4 4d 5p 5s 5 La 5d 6p 6 6s Ac 6d 7 7s 4f Lanthanide series 5f Actinide series Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 345

7. Elements show gradual changes in certain physical properties as one moves across a period or down a group in the periodic table. These properties repeat after certain intervals. In other words they are PERIODIC Periodic properties include: --Atomic Radius -- Ionization Energy --Electronegativity -- Electron Affinity -- Ionic Radius

8. Atomic Radii Radius Atomic Radius = half the distance between 2 nuclei of a diatomic molecule/adjacent atoms.

9. Trends in Atomic Radii • Influenced by three factors: 1. Energy Level • Higher energy level is further away. 2. Charge on nucleus • Higher charge pulls electrons in closer. 3. Shielding effect - electron repulsion

10. Group trends H • As we go down a group... • each atom has another energy level, • so the atoms get bigger. Li Na K Rb

11. Periodic Trends • As you go across a period, the radius gets smaller. • Electrons are in same energy level. • More nuclear charge. • Outermost electrons are closer. Na Mg Al Si P S Cl Ar

12. The radius decreases across a period owing to increase in the positive charge from the protons. • Each added electron feels a greater and greater + charge because the protons are pulling in the same direction, whereas the electrons are scattered. Small Large All values are in nanometers .

13. Atomic Radius

14. I1 + X (g) X+(g) + e- I2 + X+(g) X2+(g) + e- I3 + X2+(g) X3+(g) + e- Ionization Energy Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state. I1 first ionization energy I2 second ionization energy I3 third ionization energy I1 < I2 < I3

15. Ionization Energy Cont’d Ionization energy is the energy required to remove an electron from an isolated gaseous atom • Ionization energy is always endothermic, that is energy is added to the atom to remove the electron. • The larger the atom is, the easier it is to remove its electrons. • The energy required to remove an electron from an atom reduces as the size of the atom increases • Ionization energy and atomic radius are inversely proportional. .15

16. Factors Affecting Ionization Energy Nuclear Charge The larger the nuclear charge, the greater the ionization energy. Shielding effect The greater the shielding effect, the less the ionization energy. Radius The greater the distance between the nucleus and the outer electrons of an atom, the less the ionization energy. Sublevel An electron from a full or half-full sublevel requires additional energy to be removed.

17. Group trends in IE • As you go down a group, first IE decreases because... • Atomic radius of the atoms increases • More shielding.

18. First Ionization Energies(in kilojoules per mole) H 1312.1 He 2372.5 Li 520.3 Be 899.5 B 800.7 C 1086.5 N 1402.4 O 1314.0 F 1681.1 Ne 2080.8 Na 495.9 Mg 737.8 Al 577.6 Si 786.5 P 1011.8 S 999.7 Cl 1251.2 Ar 1520.6 K 418.9 Ca 589.9 Ga 578.6 Ge 761.2 As 946.5 Se 940.7 Br 1142.7 Kr 1350.8 Rb 402.9 Sr 549.2 In 558.2 Sn 708.4 Sb 833.8 Te 869.0 I 1008.7 Xe 1170.3 Smoot, Price, Smith, Chemistry A Modern Course 1987, page 188

19. Periodic Trends in IE • Atoms in the same period have valence electrons in the same energy level. • Same shielding. • But, nuclear charge increases across the period • So IE generally increases from left to right. • Exceptions at full and 1/2 full orbitals.

20. He • He has a greater IE than H. • same shielding • greater nuclear charge H First Ionization energy Atomic number

21. He • Li has lower IE than H • Outer electron further away • outweighs greater nuclear charge H First Ionization energy Li Atomic number

22. He • Be has higher IE than Li • same shielding • greater nuclear charge H First Ionization energy Be Li Atomic number

23. He • Breaks the pattern, because the outer electron is paired in a p orbital and experiences inter-electron repulsion. N O C H First Ionization energy Be B Li Atomic number

24. He F N O C H First Ionization energy Be B Li Atomic number

25. Ne He F N • Ne has a lower IE than He • Both are full, • Ne has more shielding • Greater distance O C H First Ionization energy Be B Li Atomic number

26. PROBLEM: Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1: Question Arrangement of Elements by First Ionization Energy (a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs PLAN: IE increases as you proceed up in a group; IE increases as you go across a period. SOLUTION: (a) He > Ar > Kr Group 8A(18) - IE decreases down a group. (b) Te > Sb > Sn Period 5 elements - IE increases across a period. (c) Ca > K > Rb Ca is to the right of K; Rb is below K. (d) Xe > I > Cs I is to the left of Xe; Cs is further to the left and down one period.

27. Electronegativity • Electronegativity is a measure of the ability of an atom in a molecule to attract electrons to itself. • This concept was first proposed by Linus Pauling (1901-1994) who later won a Nobel Prize for his efforts.

28. X (g) + e- X-(g) Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond Electronegativity - relative,

29. Group Trends in Electronegativity • So as you go down a group in the periodic table, the electrons are further away from the nucleus and better shielded from the nuclear charge and thus not as attracted to the nucleus. • For that reason the electronegativity decreases as you go down the periodic table. .29

30. Periodic Trends in Electronegativity: • The atoms have same energy levels but size decreases across the period. • Hence, Electronegativity increases from left to right across a period • F is highest - or most electronegative element

31. Summary of Periodic Trends

32. Periodicity of Period 3 Elements

33. The Period 3 elements 1 2 3 4 5 6 7

34. Major properties which change across the period are: • Structure and bonding • Elements change from metal through metalloid to non metals. • Acid-base properties • Redox properties • Solubility and complexing properties The changes are related to Change in size of the atom Change in nuclear charge Increasing number of valence electrons

35. The elements show graduation in properties with exception of argon.

36. Reactions with oxygen 4Na(s) + O2(g) → 2Na2O(s) The reactions of the period 3 elements with oxygen are redox reactions. In each reaction, the oxidation state of the elements increases and the oxidation state of the oxygen decreases. For example, when sodium is burned in oxygen the oxidation state of the sodium increases from 0 to +1 (oxidation), while the oxidation state of the oxygen decreases from 0 to -2 (reduction). 0 0 +1 -2

37. Reactions with oxygen: summary Element Description Equation Na Mg Al Si P S burns vigorously with a yellow flame 4Na(s) + O2(g) → 2Na2O(s) burns vigorously with a bright white flame 2Mg(s) + O2(g) → 2MgO(s) burns vigorously with a bright white flame 4Al(s) + 3O2(g) → 2Al2O3(s) burns with a bright white flame and white smoke Si(s) + O2(g) → SiO2(s) burns spontaneously with a bright white flame and smoke 4P(s) + 5O2(g) → P4O10(s) burns with a blue flame S(s) + O2(g) → SO2(g)

38. 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) Mg(s) + H2O(g) → MgO(s) + H2(g) Cl2(aq) + H2O(l) HClO(aq) + HCl(aq) The reactions with water are all redox reactions. 0 +1 -2 +1 -2 +1 0 0 +1 -2 +2 -2 0 0 +1 -2 +1 +1 -2 +1 -1 The metals are oxidised and their oxidation state increases. The hydrogen is reduced and its oxidation state decreases. The chlorine is both oxidized and reduced.

39. Reactions with water: summary Element Description Equation Na Mg Al Si P S Cl Cl2(aq) + H2O(l) HClO(aq) + HCl(aq) Ar vigorous 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) Mg(s) + 2H2O(l) → Mg(OH)2(aq) + H2(g) Mg(s) + H2O(g) → MgO(s) + H2(g) slow with cold water;vigorous with steam no reaction – no reaction – no reaction – no reaction – dissolves to formchlorine water no reaction –

40. OXIDES OF PERIOD 3 ELEMENTS Oxides are binary compounds of an element with oxygen. E + O EO Elements can form three types of oxides depending on oxidation state of oxygen: Normal oxide - oxidation state of –II O2- Peroxide- oxidation state –I ion is O22- Super oxide = oxidation state -½ O2-

41. NORMAL OXIDES The normal oxides of Periods 3 elements can be grouped into 3 types according to the nature of their bonding: 1.Ionic oxides; Na and Mg 2. Ionic oxides with high covalent character; Al 3. Covalent oxides Si-Cl

42. Periodicity in nature of bonding in the oxides of Periods 3 elements

43. Group 1 and 2 Oxides Group 1 and 2 Oxides are BASIC Na and Mg are metals (form cations) ; they bond with O2- to form ionic oxides. The oxide ion can bond with H+ ions and they act as bases dissolving in water to give alkaline solutions. Na2O(s) + H2O(l) 2Na+(aq) + 2OH-(aq) They will also neutralize acids to produce salt and water. MgO(s) + 2HCl(aq)  Mg2+(aq) + 2Cl-(aq)

44. Aluminum Oxide Aluminum oxide does not dissolve in water easily . It is AMPHOTERIC which means it will react with (and dissolve in) acids and bases. Acting like a base: Al2O3(s) + 6H+(aq)  2Al3+(aq) + 3H2O(l) Al2O3(s) + 3H2SO4(aq)  Al2(SO4)3(aq) + 3H2O(l) Acting like an acid: Al2O3 (s) + 3H2O(l) + 2OH-(aq)  2Al(OH)4-(aq) Al2O3(s) + 2OH-(aq)  3H2O(l) + 2Al(OH)4-(aq)

45. Acidic Oxides The remaining oxides of period 3 (Si – Cl) form acidic solutions. Silicon dioxide has little acid-base activity but it shows weakly acidic properties by slowly dissolving in hot concentrated alkalis to form silicates. SiO2(s) + 2OH-(aq) → SiO32-(aq) + H2O(l)

46. Acidic Oxides Phosphorus (V) oxide reacts to form a solution of phosphoric (V) acid, a weak acid P4O10(s) + 6H2O(l) → 4H+(aq) + 4H2PO4-(aq) Phosporus (III) oxide reacts with water to produce phosphoric (III) acid: P4O6(s) + H2O(l)  4H3PO3(aq)

47. Acidic Oxides Sulphur (VI)oxide reacts with water to make sulphuric acid: SO3(l) + H2O(l)  H2SO4(aq) Suphur (IV) oxide reacts with water to produce sulphurous acid: SO2(g) + H2O(l)  H2SO3(aq) Cl2O7 reacts with water to produce perchloric acid: Cl2O7(l) + H2O(l)  2HClO4(aq) Cl2O reacts with water to produce chlorous acid: Cl2O(l) + H2O(l)  2HClO(aq)

48. Period 3 oxides and water: summary Oxide Bonding Ions after H2O reaction Type of solution pH Na2O MgO Al2O3 SiO2 P4O10 SO2 SO3 ionic Na+(aq),OH-(aq) strongly alkaline 13–14 Mg2+(aq),OH-(aq) moderately alkaline 10 ionic ionic/covalent – (insoluble) – 7 – 7 covalent – (insoluble) covalent H+(aq),H2PO4-(aq) strongly acidic 0–1 covalent H+(aq),HSO3-(aq) weakly acidic 2–3 H+(aq),HSO4-(aq) strongly acidic 0–1 covalent

49. Properties of the Third Period Oxides

50. CHLORIDESGroup 1 and 2 NaCl and MgCl2, are ionic crystalline solids with high melting points. NaCl dissolves in water to form a neutral solution NaCl(s) Na+(aq) + Cl-(aq) MgCl2 dissolves to form a slightly acidic solution: MgCl2(s)  Mg2+(aq) + 2Cl-(aq) The resulting solutions can conduct electricity due to the free moving ions.