1. TOPIC 3: PERIODICITY

2. History  During the 19 th century, it was becoming apparent that some elements displayed similar properties, and some patterns were beginning to emerge.  In 1869, Dmitri Mendeleev of Russia and Lothar Meyer in Germany published nearly identical tables.  Mendeleev is given credit for the modern periodic table since he advanced his ideas more vigorously.

3. History (cont.)  He even proposed that certain elements were missing, and even predicted what the properties of these missing elements would be! After the discovery of germanium and gallium, it was found that Mendeleev was correct!

4. IB Core Objective  3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number.  Describe: Give a detailed account.

5. 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number.  Elements are arranged from left to right in order of increasing atomic number.

6. IB Core Objective  3.1.2 Distinguish between the terms group and period.  Distinguish: Give the differences between two or more different items.

7. Periods Groups or Families 3.1.2 Distinguish between the terms group and period.

8.  Group is a vertical column consisting of elements with the same number of valence electrons.  Period is a horizontal row where atomic number increases and chemical properties gradually change from left to right: From metals (except hydrogen), to metalloids, then non metals, and on the far right are the noble gases.  Review: Valence electrons are electrons in the outermost energy level (orbital) of an atom.

9. IB Core Objective  3.1.3 Apply the relationship between the electron configuration of elements and their position in the periodic table up to Z = 20.  Apply: Use an idea, equation, principle, theory or law in a new situation.

10. 3.1.3 Apply the relationship between the electron configuration of elements and their position in the periodic table up to Z = 20.  What do you remember from Topic 2 on electron configuration as you move across a period on the periodic table?  Each period begins with the first element having one electron in a new main energy level.  Period 1 fills the n=1 level, period 2 fills the n=2 level, etc.  What about the group number?

11. 3.1.3 Apply the relationship between the electron configuration of elements and their position in the periodic table up to Z = 20.  The group number gives the number of electrons in the valence shell. For example, electron configuration in group 1:  H: 1 Li: 2,1 Na: 2,8,1 K: 2,8,8,1  They all have one valence electron!  Group 1=1 valence electron  Group 2?  Group 2=2 valence electrons

12. 3.1.3 Apply the relationship between the electron configuration of elements and their position in the periodic table up to Z = 20.  Why is knowing the electron configuration (and valence electrons) important?  Because I said it’s important.  Also, it is the valence electrons that take part in chemical reactions and determine the physical and chemical properties for an element.

13. IB Core Objective  3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.  Apply: Use an idea, equation, principle, theory or law in a new situation.

14. 3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.  Look at the periodic table for trends regarding number of electrons in the highest occupied energy level.

15. 3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table. What energy level is sulfur (S) in? A: n=3 How many valence electrons does it have? A: 6 Speed Game Valence Electrons and Naming

16. IB Core Objective  3.2.1 Define the terms first ionization energy and electronegativity.  Define: Give the precise meaning of a word, phrase or physical quantity.

17. 3.2.1 Define the terms first ionization energy and electronegativity.  Ionization energy: Taught to you by the friendly folks from HL! Review Ionization Energy Energy required to rip off one electron from an element in its gaseous state. X (g)  X + (g) + e -

18. 3.2.1 Define the terms first ionization energy and electronegativity. Electronegativity How strongly the atom attracts electrons in a covalent bond. Mr. F can really attract the electrons.Mr. K, not so much

19. IB Core Objective  3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3.  Describe: Give a detailed account.

20. 3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3.  Use the Excel spreadsheets you have created on atomic radii and ionization energy to help determine these trends.  Break into small groups. HL students—using the graphs, explain the trends in ionization energy to the SL students.  Why is sodium the lowest in period 3?  Why is aluminum lower than magnesium in period 3?  Why is the ionization energy so high for argon in period 3?

21. 3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3.  Look at the atomic radii graphs you have made. What trends do you notice?  Does the atomic radii go up as the atomic number goes up?  Why do you think these trends exist?

22. 3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3.  When going across a period on the periodic table, the number of protons in the nucleus increases.  What happens to the charge on the nucleus when protons increase (not taking into account the electrons)?  A: It increases  When the nucleus is more positively charged, the electrons become more attracted.  Attraction moves them closer, decreasing the atomic radii. + - When you’re positive, I feel so much closer to you.

23. 3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3.

24. Electronegativity Going across period 3, electronegativity increases. The elements further to the right have a higher affinity for electrons—the ones on the left tend to lose electrons.

25. IB Core Objective  3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li → Cs) and the halogens (F → I).  Describe: Give a detailed account.  Explain: Give a detailed account of causes, reasons or mechanisms.

26. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li → Cs) and the halogens (F → I).  What do you think would happen if an electron was lost, like in one of the alkali metals?  A: It would lose a shell, so the ionic radius would decrease.  What would happen if an electron was gained in one of the halogens?  A: The extra electron would cause repulsion, making the ionic radius bigger.

27. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li → Cs) and the halogens (F → I). Questions What happens to the atomic radii as you move down the alkali metals or halogens? A: It increases.

28. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li → Cs) and the halogens (F → I). Questions What happens to the ionic radii when the alkali metals become ions? A: They become smaller. What happens to the ionic radii when the halogens become ions? A: They become bigger.

29. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li → Cs) and the halogens (F → I). Questions What happens to the first ionization energy as you go down the alkali metal or halogen group? A: It goes down

30. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li → Cs) and the halogens (F → I). Electronegativity When going down the periodic table, electronegativity decreases.

31. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li → Cs) and the halogens (F → I). Melting Points What causes melting? The attractive forces holding the particles in the structure of a solid are overcome, and the particles are then free to move around. The melting point depends on the strength of these attractive forces and the way in which particles are arranged in the solid state.

32. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li → Cs) and the halogens (F → I). Melting Point Melting point decreases going down the alkali metals on the periodic table. Melting point decreases because as the atoms get larger, the forces of attraction between them decrease. Melting point increases going down the halogens on the periodic table. Why is this? The solids of halogens are non-polar diatomic molecules, and are thus weakly attracted to each other by van der Waals’ forces. The forces of attraction in van der Waals’ increase as the mass of the molecules increase.

33. IB Core Objective  3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.  Compare: Give an account of similarities and difference between two (or more) items, referring to both (all) of them throughout.

34. 3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.  Which is more electronegative? Sodium or potassium?  A: Sodium  Which is more electronegative? Chlorine or bromine?  A: Chlorine  Which is more electronegative? Bromine or nitrogen?  A: Nitrogen  This is a little harder to figure out using just the trends, but there’s a key word you can learn….

35. 3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table. FONClBrISCHP !! A fun way to remember the order of electronegativity! Which is more electronegative? Bromine or sulfur? A: Bromine!

36. 3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.

37. IB Core Objective  3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group.  Discuss: Give an account including, where possible, a range of arguments for and against the relative importance of various factors, or comparisons of alternative hypotheses.

38. 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. Alkali metals Low melting points. Soft and malleable. Low densities (why are the densities low?) A: They are the largest atoms in their period of the periodic table. Alkali metals are chemically very reactive. Why? A: The outer electron is very easily lost, and combine easily with reactive non-metals. So what do you think might react easily with alkali metals?

39. 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group.  All of the metals react with water to form a metal hydroxide and hydrogen gas.  2M (s) + 2H 2 O (l) → 2M + (aq) + 2OH - (aq) + H 2(g)  Lithium: The reaction occurs slowly and steadily.  Sodium: The reaction is vigorous, producing enough heat to melt the sodium (98 ° C)  Potassium: Reaction is violent and the heat produced is enough to ignite the hydrogen gas involved.  The solution becomes strongly alkaline (basic) after these reactions.

40. Metal Sodium (Na) in Water

41. Metal Potassium (K) in Water

42. Large Quantity of Na Large Amount of Na Larger amount of Na

43. 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. Halogens Very reactive non-metals. How many electrons do they require to fill their valence shell? A: one They exist as diatomic molecules. What are the diatomic molecules in the halogen group? A: F 2, Cl 2, Br 2, I 2

44. 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group.  What do you remember about fluorine (or Mr. F) when it comes to attracting electrons?  It has a strong electron affinity, or it pulls in electrons strongly. (High electronegativity)  Fluorine is the strongest oxidizing agent known (another way to say it easily takes electrons).  It is so reactive that its use is not allowed in most school laboratories.

45. 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group.  The ability to oxidize decreases as you go down Group 7.  For example, if one reacted aqueous chlorine with sodium bromide, the chlorine would take the electron:  Cl 2(aq) + 2Br - (aq) → 2Cl - (aq) + Br 2(aq)  Chlorine is also very reactive, and was used as a poisonous gas in World War I.

46. Poem

47. 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group.  Halogens are only slightly soluble in water since they are non-polar.  In aqueous solutions halogens dissociate slightly to form an acidic solution: Cl 2(aq) + H 2 O (l) → H + (aq) + Cl - (aq) + HOCl (aq)

48. 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group.  Halogens will bond with metals to give ionically bonded salts.  2K (s) + Br 2(l) → 2KBr (s)  2Na (s) + Cl 2(g) → 2NaCl (s)  These salts are usually white and soluble in water (exceptions are lead and silver halide salts, which are insoluble).

49. IB Core Objective  3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.  Discuss: Give an account including, where possible, a range of arguments for and against the relative importance of various factors, or comparisons of alternative hypotheses.

50. 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.  Look at the first two elements (Na & Mg). What is the ionization energy compared to elements further to the right?  A: They have relatively low ionization energies.  What kinds of bonds would these elements typically form?  A: Ionic  What is an oxide?  A compound with oxygen bonded with at least one other element.

51. 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.  So oxides with metals are ionic bonds.  Since the oxide ion can form a bond with hydrogen ions, they act as bases in water.  Na 2 O (s) + H 2 O (l) → 2Na + (aq) + 2OH - (aq)  The aluminum oxide is amphoteric. What do you think amphoteric means?  It can act as either an acid or a base.  Al 2 O 3(s) + 6H + (aq) → 2Al 3+ (aq) + 3H 2 O (l)  Al 2 O 3(s) + 2OH - (aq) + 3H 2 O (l) → 2Al(OH) 4 - (aq)

52. 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.  Moving to the right, we reach silicon. Silicon is a metalloid. What is a metalloid?  A: Possesses properties of a metal and a non-metal.  The ionization energy for silicon becomes too great for it to effectively form an ion. So what kind of bond would it form?  A: Covalent  It is slightly acidic in solution.

53. 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.  The remaining elements form covalent bonds (except argon), and form acidic solutions in water.  For example, sulfur trioxide reacts to form sulfuric acid, a strong acid: SO 3(s) + H 2 O (l) → H + (aq) + HSO 4 - (aq)

54. 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3. Two other equations you need to know…  Magnesium oxide…what do you think the equation would be if it reacts with water? MgO(s) + H 2 O(l) → Mg(OH) 2 (aq)/(s) Acidic or basic in solution? A: Basic What about phosphorous pentoxide in water? P 4 O 10 (s) + 6H 2 O(l) → 4H 3 PO 4 (aq) Acidic in solution

55. 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.  To summarize the trend across period 3:  Ionic  Covalent  Very Basic  weaker basic  amphoteric  weakly acidic  Very acidic  Na 2 O + H 2 O  2NaOH (react with water)  Na 2 O + HCl  Na + + Cl - + H 2 O (react with acid)  Al 2 O 3 is amphoteric

56. IB HL Objective  13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure.  Explain: Give a detailed account of causes, reasons or mechanisms.

57. 13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure.  Can you deduce what the oxides would be for the first four elements across period 3?  A: Na 2 O, MgO, Al 2 O 3, SiO 2  What about phosphorous?  Phosphorous is….different. You will learn down the road that phosphorous can form five covalent bonds, and can have different oxidation states (Topic 9).  The two oxides for phosphorous you will need to know is P 4 O 6 and P 4 O 10.  The rest are SO 2, SO 3, Cl 2 O and Cl 2 O 7.  What is the general trend for number of oxides bonding as you move from left to right on period 3?

58. 13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure.  As number of valence electrons increases, there is an increase in the number of electrons available for bond formation.  Thus the number of oxides goes up.  Think about the electronegativity differences across period 3 compared to oxygen. What is the trend?

59. 13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure.

60.  The metals on the left (Na, Mg, and Al) all have a larger electronegativity difference with oxygen.  Oxygen likes to attract their electrons.  So oxygen becomes negatively charged, and the metals become positively charged.  So they form an ionic bond.  This causes them to have high melting points, and they can conduct electricity when molten.

61. 13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure.  SiO 2 (silicon dioxide) also has a high melting point, because it forms a complex lattice with strong covalent bonds. Silicon dioxide is also commonly called silica. Do you know what this is used for?

62. 13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure.  Oxides of phosphorous, sulfur and chlorine form covalent bonds because the electronegativity difference is small.  So they need to share.  They have a low melting point.  Low boiling point.  And don’t conduct electricity.

63. 13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure. Chlorides How would you write the first three elements bonded with chlorine? A: NaCl, MgCl 2, Al 2 Cl 6 NaCl and MgCl 2 are ionically bonded, solids at room temperature and have high melting points. When molten or added to water, they conduct electricity. Al 2 Cl 6 is a poor conductor of electricity and has a much lower melting point than NaCl and MgCl 2.

64. 13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure.  How would you write the chlorides for the next three elements across period 3?  A: SiCl 4, PCl 5, PCl 3, and Cl 2  Chlorine usually forms only one bond, so the same extended lattice with Si is not the same as with oxides.  All of these molecules are covalently bonded and have low melting and boiling points.

65. IB HL Objective  13.1.2 Describe the reactions of chlorine and the chlorides referred to in 13.1.1 with water.  Describe: Give a detailed account.

66. 13.1.2 Describe the reactions of chlorine and the chlorides referred to in 13.1.1 with water.  NaCl added to water gives a neutral solution with a pH of 7.  MgCl gives a slightly acidic solution with water.  When AlCl 3 is added to water, a very exothermic reaction takes place, and hydrochloric acid is formed.  2AlCl 3(s) + 3H 2 O (l) → Al 2 O 3(s) + 6HCl (aq)

67. 13.1.2 Describe the reactions of chlorine and the chlorides referred to in 13.1.1 with water.  The other chlorides also react vigorously with water to produce acidic solutions of hydrochloric acid.  SiCl 4 (l) + 4H 2 O(l) → Si(OH) 4 (aq) + 4HCl(aq)

68. IB HL Objective  13.2.1 List the characteristic properties of transition elements.  List: Give a sequence of names or other brief answers with no explanation.

69. 13.2.1 List the characteristic properties of transition elements.  Which subshell block are the transition elements located?  A: d-block  Transition elements are those in which the element has a partially filled d-sublevel.

70. 13.2.1 List the characteristic properties of transition elements.  d-block contain all dense, hard metallic elements.  The oxidation state can be variable.  Because of their small size, transition metals can attract species that are rich in electrons, and form complex ions.  Compounds of transition metals are colored.  Transition elements have catalytic properties.

71. IB HL Objective  13.2.2 Explain why Sc and Zn are not considered to be transition elements.  Explain: Give a detailed account of causes, reasons or mechanisms.

72. 13.2.2 Explain why Sc and Zn are not considered to be transition elements. To be a transition metal, it must have a partially filled d suborbital in one of the common oxidation states (ion)  Sc: [Ar] 3d 4s 2  Sc 3+ [Ar] (no d-suborbital)  Zn: [Ar] 3d 10 4s 2  Zn +2 [Ar] 3d 10 (d-suborbital is not partially full)

73. IB HL Objective  13.2.3 Explain the existence of variable oxidation number in ions of transition elements.  Explain: Give a detailed account of causes, reasons or mechanisms.

74. 13.2.3 Explain the existence of variable oxidation number in ions of transition elements.  The difference in energy levels between 3d and 4s is small  Typically d-block will form +2 ions due to removal of e - from the 4s instead of 3d.  Oxidation numbers can also be predicted based on losing all the electrons in 4s and 3d.  Sc: [Ar] 4s 2 3d 1  Sc 3+  Ti: [Ar] 4s 2 3d 2  Ti 4+  Fe: [Ar] 4s 2 3d 6  Fe 3+ and Fe 2+ Stable due to half filled d-subshell

75. 13.2.3 Explain the existence of variable oxidation number in ions of transition elements. Question If iron (III) formed an oxide, what would the formula be? A: Fe 2 O 3

76. 13.2.3 Explain the existence of variable oxidation number in ions of transition elements.  Question: What are the 2 exceptions to the electron configuration rules for non-ions in the D-block?  A: Chromium and copper.  Chromium: [Ar] 4s 1 3d 5  Copper: [Ar]4s 1 3d 10  It is more favorable for them to half-fill and completely fill the 3d sub-level than spin-pair two in the 4s orbital.

77. 13.2.3 Explain the existence of variable oxidation number in ions of transition elements. Question If copper lost the 4s electron, what would the oxidation state be? A: Cu +1 What would the formula be if copper(I) formed an oxide? A: Cu 2 O

78. 13.2.3 Explain the existence of variable oxidation number in ions of transition elements. Metals with 4+ oxidation state Metals with a 4+ oxidation state are rare. For example, Mn 4+ is small and highly charged…it excites the oxygen ligands into a vacant d orbital into the manganese. Therefore, the bonding is closer to covalent! MnO 2 This complex ion will then ionically bond with a metal, such as potassium, creating…. KMnO 4

79. 13.2.3 Explain the existence of variable oxidation number in ions of transition elements.  Why is the oxidation state of Mn 4+ rare?  Look at the electronic configuration! [Ar]3d 3  Not a very stable formation!  What do you think another oxidation state for Mn would be?  Mn 7+ ….lose all of the electrons! High oxidation number equals not very stable state.  Bonds with 4 oxygen to make MnO 4 -. This complex ion then bonds with potassium to make KMnO 4

80. 13.2.3 Explain the existence of variable oxidation number in ions of transition elements.  What if chromium lost all of it’s electrons?  Then it would be Cr 6+.  Another oxidation state for chromium is Cr 3+.  If chromium (III) bonded with chloride, what would the formula be?  A: CrCl 3  If chromium (VI) bonded with oxygen, what would the formula be?  A: Cr 2 O 7 2-

81. IB HL Objective  13.2.4 Define the term ligand.  Define: Give the precise meaning of a word, phrase or physical quantity.

82. 13.2.4 Define the term ligand.  Ligands: neutral molecules or negative ions that contain a non-bonding pair of electrons.  In an ordinary covalent bond, each atom will share an electron.  Because of their small size, transition metal ions attract species that are rich in electrons.  Therefore, a ligand and the metal form a dative covalent bond where the ligand provides both of the electrons.

83. H H N H Lone pair of e - 2+ Cu H H N H H H N H H H N H 2+ 13.2.4 Define the term ligand.

84. IB HL Objective  13.2.5 Describe and explain the formation of complexes of d-block elements.  Describe: Give a detailed account.  Explain: Give a detailed account of causes, reasons or mechanisms.

85. 13.2.5 Describe and explain the formation of complexes of d-block elements. Coordination Numbers  # of ligands formed around a central atom  6 ligands: Form octahedral structure  4 ligands: Form tetrahedral  2 ligands: Form linear structure  Complex ions can form  (+) Cations (NH 4 ) +  (-) Anions (OH) -

86. 13.2.5 Describe and explain the formation of complexes of d-block elements.  Water is a common ligand, and will form hexahydrates with most transition metals: [Fe(H 2 O) 6 ] 3+  Cyanide (CN - ), also forms octahedral structures with transition metals: [Fe(CN) 6 ] 3- What is the oxidation state for the iron in both of these complex ions?

87. 13.2.5 Describe and explain the formation of complexes of d-block elements.  Chloride is another example of a ligand.  It can form a tetrahedral structure with copper(II). So what is the formula for this complex ion? A: [CuCl 4 ] 2-  Ammonia is a neutral ligand that can form linear structures with silver(I). What would the formula be for this complex ion? A: [Ag(NH 3 ) 2 ] +

88. 13.2.5 Describe and explain the formation of complexes of d-block elements.  Ligands can often replace other ligands:  Ex: [CuCl 4 ] 2- + 4H 2 O ↔ [Cu(H 2 O) 4 ] 2+ + 4Cl-  The color will then go from yellow/green to pale blue.  [Cu(H 2 O) 4 ] 2+ + 4NH 3 ↔ [Cu(NH 3 ) 4 ] 2+ + 4H 2 O  The color will go from pale blue to deep blue.

89. IB HL Objective  13.2.6 Explain why some complexes of d-block elements are coloured.  Explain: Give a detailed account of causes, reasons or mechanisms.

90. 13.2.6 Explain why some complexes of d-block elements are coloured.  If an atom is all alone, all orbitals have equal energy….BUT  If the atom is surrounded by charged or polar molecules (ligands), the effect of the electric field from these ligands have a different effect on the d- orbitals.  For example, if six ligands bond to a transition metal ion to form an octahedral complex, the 3d sub-level will be split into two.

91. 13.2.6 Explain why some complexes of d-block elements are coloured. Cu 2+ H H O H H O H H O H H O H H O H H O Equal energy Energy Gap corresponds to visible light Un-equal energy

92. 13.2.6 Explain why some complexes of d-block elements are coloured.  What causes color?  Wavelengths of light. Some of the wavelengths are being absorbed from white light (promoting the electron to a higher energy level).  The colors that you see are the wavelengths which are not being absorbed (they are transmitted back).  If there are no electrons in the d-orbitals, then it will appear colorless (i.e. Sc 3+ and Ti 4+ )