CHAPTER 5 PERIODIC TABLE for fundamental

Introduction to fundamental Chemistry

1. INTRODUCTION TO PERIODIC TABLE CHAPTER 5

2. Structure of the period table Elements are arranged in order of atomic number Column =  group Row =  period All elements in the same period have the same number of electron shells i.e. the same quantum number (e.g. n=3 for period 3) Elements in each group or period show similar characteristics in either chemical or physical behaviour

3. CLASSIFICATION OF THE ELEMENTS Alkali Metal Alkali Earth Metal Period Group Halogen Noble Gas

4. GROUP NO 1A ELEMENTS (ns 1 , n  2)

5. GROUP NO 2A ELEMENTS (ns 2 , n  2)

6. GROUP NO 3A ELEMENTS (ns 2 np 1 , n  2)

7. GROUP NO 4A ELEMENTS (ns 2 np 2 , n  2)

8. GROUP NO 5A ELEMENTS (ns 2 np 3 , n  2)

9. GROUP NO 6A ELEMENTS (ns 2 np 4 , n  2)

10. GROUP NO 7A ELEMENTS (ns 2 np 5 , n  2)

11. ns 1 ns 2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 5 ns 2 np 6 d 1 d 5 d 10 4f 5f GROUND STATE ELECTRON CONFIGURATIONS OF THE ELEMENTS

12. Valence electrons Valence electron is the outermost electron of an element. The chemical reactivity of the element s is largely determined by outermost electrons. Core electrons All non-valence electrons in an atom. Electrons are in the inner orbital of an element.

13. METALS Shiny and silvery Solids are malleable and ductile Malleable (can be shaped with hammer) Ductile (can be drawn into wires) Good conductors of heat and electricity Most metal oxides are basic Tend to form cations NON - METALS Not luster, many colours Often brittle; some are hard, some soft Brittle (hard but easily broken or cracked) Poor conductors (graphite is an exception) Most non-metallic oxides are acidic Tend to form anions or oxyanions in solutions PROPERTIES OF METALS vs NON METALS

14. METALS All are solids at 25ºC (except Hg) Low ionization energies Form positive ions ( cation ) Oxides are basic CaO (s) + H 2 O(l)  Ca(OH) 2 ( aq ) Metal oxide + acid  salt + water MgO (s) + 2HCl( aq )  MgCl 2 ( aq ) + H 2 O(l)

15. NON-METALS Vary greatly in appearance. Group 7A exist as diatomic atoms Example: H 2 ( colourless gas), Cl 2 (green gas), Br 2 (red liquid), I 2 (purple volatile solid). Tend to gain electrons to form anions Oxides are acidic CO 2 + H 2 O  H 2 CO 3 ( aq ) Non-metal oxide + acid  salt + water SO 3 + 2KOH  K 2 SO 4 ( aq ) + H 2 O(l )

16. METALLOIDS Generally hard, non-malleable solids In pure state they are non-conductors but with controlled impurities they form semi-conductors (computer chips are made of Si)

17. Properties of Oxides Across a Period basic acidic

18. 18

19. Na : [Ne]3s 1 Na + : [Ne] Ca : [ Ar ]4s 2 Ca 2 + :[ Ar ] Al : [Ne]3s 2 3p 1 Al 3+ : [ Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration H : 1s 1 H - : 1s 2 or [He] F : 1s 2 2s 2 2p 5 F - : 1s 2 2s 2 2p 6 or [Ne] O: 1s 2 2s 2 2p 4 O 2- : 1s 2 2s 2 2p 6 or [Ne] N: 1s 2 2s 2 2p 3 N 3 - : 1s 2 2s 2 2p 6 or [Ne] Atoms gain electrons so that anion has a noble-gas outer electron configuration ELECTRON CONFIGURATIONS OF CATIONS AND ANION OF REPRESENTATIVE ELEMEN TS

20. Electronic Configuration

21. Na + : 1s 2 2s 2 2p 6 or [Ne] Al 3 + : 1s 2 2s 2 2p 6 or [Ne] O 2- : 1s 2 2s 2 2p 6 or [Ne] Na + , Al 3+ , F - , O 2- , and N 3- are all isoelectronic with Ne What neutral atom is isoelectronic with H - ? S ame electron configuration as He ISOELECTRONIC - same number of electrons, and hence the same ground-state electron configuration H - : 1s 2 or [He] F - : 1s 2 2s 2 2p 6 or [Ne] N 3 - : 1s 2 2s 2 2p 6 or [Ne] He: 1s 2 Ne: 1s 2 2s 2 2p 6

22. When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the ( n – 1) d orbitals . Fe: [ Ar ]4s 2 3d 6 Fe 2+ : [ Ar ]4s 3d 6 or [ Ar ]3d 6 Fe 3+ : [ Ar ]4s 3d 5 or [ Ar ]3d 5 Mn : [ Ar ]4s 2 3d 5 Mn 2+ : [ Ar ]4s 3d 5 or [ Ar ]3d 5 ELECTRON CONFIGURATIONS OF CATIONS OF TRANSITION METALS

23. Isoelectronic elements

24. PERIODIC TRENDS Across Period Down Group

25. Atomic and Ionic Radii Electronegativity Metallic Character Ionization Energy Electron Affinity PERIODIC TRENDS

26. Effective nuclear charge ( Z eff ) is the “positive charge” felt by an electron. Na Mg Al Si 11 12 13 14 10 10 10 10 1 2 3 4 186 160 143 132 Z eff Core Z Radius (pm) Z eff = Z - s 0 < s < Z ( s = shielding constant) Z eff  Z – number of inner or core electrons

27. Effective Nuclear Charge ( Z eff ) increasing Z eff increasing Z eff

28. Atomic radius Across a period () - P roton is added to the nucleus of each element - I ncreases the nuclear charge - Increases the attraction between the nucleus and the electrons - E lectrons are pulled closer to the nucleus, thereby decreasing the size of the atomic radius

29. Down a group () - Extra electron shells are added (outer electron enter new energy levels) - Outer electrons are further away & increasingly screened from the positive nucleus - Not held so tightly, thus atomic radius increase Atomic radius

30. 30

31. b) Ionic radius Positive ions (metal ions) - smaller ionic radius than parent atom - loss of e - , no of e -  - E lectrons are pulled much closer to the nucleus,  ionic radius Negative ions (non-metal ions) - larger ionic radius than parent atom - addition of e - , no of e -  - n ucleus is less able to pull the electrons towards it,  ionic radius

32. Ionic Radii Cation - smaller than atom from which it is formed. Anion - larger than atom from which it is formed.

33. 34 COMPARISON OF ATOMIC RADII WITH IONIC RADII

34. c) Electronegativities of the elements

35. Metallic character of an element : Defined as ability of an atom to lose an electron. Non-metallic relates to the ability to gain electrons. d) Metallic Character

36. Across Period Down a Group Atomic size Metallic Character Atomic size Metallic Character Alkali Metal Alkali Earth Metal Non Metals Metal

37. A minimum energy (kJ/mol) required to remove an electron from an atom in its ground state . e) Ionization Energy, IE I 1 < I 2 < I 3 Ionization Energy of Sodium (Na) (kJ mol -1 ) First Second Third Fourth 495.8 4562.4 6912 9543 Proton no for Na : 11 Electronic Configuration: Na: 1s 2 2s 2 2p 6 3s 1

38. 42

39. 1 st IE the energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of singly charged positive ions: 2 nd IE the energy required to remove one mole of electrons from one mole of singly-charged positive ions in the gas phase to form one mole of doubly-charged positive ions: X(g)  X + (g) + e - X + (g)  X 2+ (g) + e - Ionization energy (IE)

40. Ionization energy (IE) 1 st IE & 2 nd IE are always positive number (endothermic process) More energy is required to remove the 2 nd electron than to remove 1 st electron This is because 2 nd electron is being removed from a positive ion, which is smaller than the original atom and experiences a greater force of attraction Even more energy would be required to remove the 3 rd electron, and so on

41. Attraction between nucleus & outermost e -   IE ,  distance Size of positive nuclear charge   IE ,  positive nucleus – greater attraction for outer e - Inner shells of e - repel outer e - , screening or shielding it from the nucleus   IE ,  e - shells – between outer e - and nucleus,  firmly held outer e - Factors that affect the IE

42. First IE increased First IE decreased

43. Ionization energy (IE) As the size of the nuclear charge increases the force of the attraction between the negatively charged electrons and the positively charged nucleus increases . + + Small nuclear charge  Large nuclear charge    Small force of attraction  Smaller ionisation energy Large force of attraction  Greater ionisation energy

44. Cont … As the energy of the electron increases , the electron is farther away from the nucleus. As a result the force of attraction between the nucleus and the electron decreases . + Electrons closer to positive nucleus  Large force of attraction  Greater ionisation energy Electrons further away from positive nucleus  Small force of attraction  Smaller ionisation energy +

45. Trends across a period () + + + + Going across a Period Going across a period, the size of the 1 st ionisation energy shows a general increase . This is because the electron comes from the same energy level, but the size of the nuclear charge increases . Ionization energy (IE)

46. 1 st ionization energy The end of each period is marked by ↑IE of a noble gas stable electronic structure unreactive

47. 1 st ionization energy (Period 2 & 3) Cont.. From the graph, 1 st IE do not increase smoothly across a period due to the presence of subshells within each shell . IE Be > IE B IE Mg > IE Al IE N > IE O IE P > IE S Removing 1 electron from an atom of B & Al removes the single electron in the p subshell However, for Be & Mg, an electron must be removed from a full s subshell. Full subshells particularly stable N & P contain half-full outer p subshell (greater stability) , therefore removing 1 electron from N & P require more energy Removing an electron from O & S removes the 4 th electron in the p subshell , leaving a half-full p subshell

48. Trends down a group () + + + + Down the Group The electron removed during the first ionisation is from a higher energy level and hence it is further from the nucleus. The nuclear charge also increases, but the effect of the increased nuclear charge is reduced by the inner electrons which shield the outer electrons. Ionization energy (IE)

49. Ionization energy (IE) Trends down a group () Trends across a period () Easier to lose e -  distance between nucleus and outer e -  nuclear charge,  no of inner e -  energy to bring about ionisation Harder to remove e - I ncreases positive nuclear charge without addition of any extra e - shells to screen outer e -  atomic radius, e - are held more firmly  energy to bring about ionisation Summary

50. Ionization Energy

51. 1 st EA negative energy change accompanying the gain of one mole of electrons by one mole of gaseous atoms to form one mole of singly charged negative ions 2 nd EA energy change accompanying the gain of one mole of electrons by one mole of singly charged negative ions in the gas phase to form one mole of double negatively charged ions X(g) + e -  X - (g) X - (g) + e -  X 2- (g) f) electron affinity(EA)

52. EA increased EA decreased

53. Facer, G., As Chemistry 2 nd Edition, Philip Allan Updates , pp 38, 2008. Cont… 1 st EA = negative value (exothermic) negatively charged e - being added is brought towards the positively charged nucleus. Force of attraction occur Energy is released when the two are brought closer together Eg : O(g) + e - → O - (g) EA = -142 kJ mol -1 2 nd & 3 rd EA = positive value (endothermic) energy is required to add e - to an already negative ion Incoming e - is repelled by the negative ion Energy has to be supplied to bring the ion and the electron together Eg : O - (g) + e - → O 2- (g) EA = +844 kJ mol -1

54. 1 st EA Period 1 & 2 Facer, G., As Chemistry 2 nd Edition, Philip Allan Updates , pp 38, 2008. http://cnx.org/content/m31451/latest/

55. Cont… e - is added to the outer orbit of the atom As e - is negatively charged and is being brought towards a positive nucleus, energy is released The general trend is upwards because atomic radii decrease from Li to F, which causes the force of attraction increase The dip at Be occurs because the added electron goes into an already singly occupied 2s-orbital . Therefore, it experiences considerable repulsion and less energy is released The dip at nitrogen is caused by the extra repulsion of putting a 2 nd electron the singly occupied 2p x -orbital