1. IB CHEMISTRY
Topic 3 Periodicity
Higher level

2. 3.1 The periodic table
OBJECTIVES
• The periodic table is arranged into four blocks associated with the four sub-levels—s, p, d, and f.
• The periodic table consists of groups (vertical columns) and periods (horizontal rows).
• The period number (n) is the outer energy level that is occupied by electrons.
• The number of the principal energy level and the number of the valence electrons in an atom can be deduced from its position on the periodic table.
• The periodic table shows the positions of metals, non-metals and metalloids.
• Deduction of the electron configuration of an atom from the element’s position on the periodic table, and vice versa.

3. Periodic Table and orbitals
Atomic number (Z) – the number of protons in the nucleus of an atom of that element
The atomic number of each element increases left to right across each period
The s,p,d,f atomic orbitals are arranged in blocks of the periodic table.

4. Group vs. Period
Group – vertical columns of the periodic table which contain elements having similar chemical and physical properties
The groups to be known are 1 alkali metals, 17 halogens, 18 noble gases, transition metals, lanthanoids and actinoids
Period – horizontal rows of the periodic table

5. PT and groups

6. PT and electron shells

7. Metals vs nonmetals Metals Conductors of heat and electricity
Malleable (bent into shapes)
Ductile (drawn into wires)
Lustre (shiny)
Oxidized (lose electrons)
Nonmetals
Insulators of heat and electricity
Brittle
Dull
Reduced (gain electrons)

8. PT metals vs nonmetals

9. 3.2 Periodic trends
OBJECTIVES
• Vertical and horizontal trends in the periodic table exist for atomic radius, ionic radius, ionization energy, electron affinity and electronegativity.
• Trends in metallic and non-metallic behaviour are due to the trends above.
• Oxides change from basic through amphoteric to acidic across a period.
• Prediction and explanation of the metallic and non-metallic behaviour of an element based on its position in the periodic table.
• Discussion of the similarities and differences in the properties of elements in the same group, with reference to alkali metals (group 1) and halogens (group 17).
• Construction of equations to explain the pH changes for reactions of Na2O, MgO, P4O10, and the oxides of nitrogen and sulfur with water.

10. Atomic Radii
Decrease across a period, increase down a group.

11. Ionic Radii
cations get smaller
anions get larger

12. First Ionization Energy
First Ionization Energy – The energy required to remove one mole of electrons from a mole of atoms or ions in the gaseous phase
X(g) X+(g) + e-

13. Explaining IE
IE increases across a period because the nuclear charge increases, attracting the electrons and the number of electrons in the shell are increasing.
IE decreases down a group because the electrons further away from the nucleus and electrons in lower shells are blocking the attraction causing electron shielding.
Anomaly type 1:
B has a lower IE than Be because the 2p electrons are slightly higher in energy than the 2s electrons, and so the ionization for B is lower than for Be
Anomaly type 2:
O has a lower IE than N because the px, py, and pz only contain one electron. The extra electron in O causes a pair and repulsion making it easier to remove, hence giving it a lower IE than N.

14. Electron affinity
Electron affinity – is the energy released when 1 mol of electrons is attached to 1 mole of neutral atoms or molecules in the gas phase X(g) + e-  X-(g)
Increasing Eea

15. Electronegativity
Electronegativity – is a measure of the ability of an atom to attract bonded electron pairs to itself when in a covalent bond

16. Electronegativity and metallic nonmetallic character
Metals have small electronegativity values, non-metals have high values. Differences greater than 1.8 will form an ionic bond rather than a covalent bond.

17. Comparing electronegativies

18. Summary of trends in the Periodic Table

19. Alkali metals low melting and boiling points
melting and boiling points decrease due to increase shielding and less nuclear attraction
very reactive due to need to lose just one electron to have full electron shell
large atoms so metals are soft and not dense

20. Alkali metals and water
Alkali metals react vigourously with water to create hydrogen and a base
Na(s) + H2O(l) → NaOH(aq) + H2(g)
Alkali metals with halogens
Alkali metals react vigourously with halogens to form salts
2Na(s) + Cl2(g) → 2NaCl(s)

21. Halogens
very reactive due to need to gain just one electron to have full electron shell
very electronegative as just have to gain one electron
melting and boiling points increase due to increased London dispersion forces (IMF) between the simple covalent molecules, and increased molecular weight

22. Halogens become darker are move from gas to solid down the group
A halogen higher on up the group will displace (is more reactive) than one lower down

23. Halogens with halide ions
The more reactive halogen (further up the group) will will take an electron from a halide ion to itself become a halide ion. Cl2(aq) + 2Br-(aq)  2Cl-(aq) + Br2(aq)

24. Oxide reactions and pH
The oxides of elements have increasing acidity across a period (Al is amphoteric being both acidic and basic).
Metal oxides are basic, non-metal oxides are acidic.
Na2O(s) + H2O(l)  2NaOH(aq)
MgO(s) + H2O(l)  Mg(OH)2(aq)
P4O10(s) + 6H2O(l)  4H3PO4(aq)
SO2(g) + H2O(l)  H2SO3(aq)
SO3(l) + H2O(l)  H2SO4(aq)
(All these and previous equations must be learnt.)
Increasing acidity across a period

25. 13.1 First-row d-block elements
Higher level
13.1 First-row d-block elements
OBJECTIVES
• Transition elements have variable oxidation states, form complex ions with ligands, have coloured compounds, and display catalytic and magnetic properties.
• Zn is not considered to be a transition element as it does not form ions with incomplete d-orbitals.
• Transition elements show an oxidation state of +2 when the s-electrons are removed.
• Explanation of the ability of transition metals to form variable oxidation states from successive ionization energies.
• Explanation of the nature of the coordinate bond within a complex ion.
• Deduction of the total charge given the formula of the ion and ligands present.
• Explanation of the magnetic properties in transition metals in terms of unpaired electrons.

26. Transition metals
Higher level
A transition element is defined as an element that possesses an incomplete d-sublevel in one or more oxidation states (ie. as an ion)
All the elements in group 12, Zn, Cd, Hg, and Cn are not transition metals as they contain full d-sublevels with 10 d-electrons as ions (lose s electrons).

27. Transition metals
Higher level
Transition metals have an empty d orbital. The d orbital splits into two energy sublevels and electrons moving between these gives them their properties
Note: for Cr and Cu it is more energetically favourable to half-fill and completely fill the d sub-level respectively so they contain only one 4s electron

28. Properties Produces colours and allows for complex ion formation
Higher level
Produces colours and allows for complex ion formation
It gives them variable oxidation numbers and makes them good catalysts
Most materials are diamagnetic (repelled by a magnet), some transition metals are paramagnetic (ferromagnetic – attracted by a magnet) due to unpaired electrons allowing spin in one direction to form poles.
When transition metals lose electrons they lose the 4s electrons first
All transition metals can show an oxidation state of +2 and occurs when they lose both s orbital electrons

29. Oxidation numbers
Higher level
Variable oxidation states come from the fact that they have relatively small differences in their successive ionization energies. Cf. 1st and 2nd IE of Na vs one of the transition metals.
Common oxidation states from the data booklet:

30. Levels in transition metals not so pronounced

31. Electron configurations
Higher level
Element Z 3d 4s Sc 21
[Ar]   Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30

32. Higher level
Complex ions
A complex consists of a central atom, which is usually a metal atom or ion, and attached groups called ligands
The coordination number is the total number of points at which a central atom or ion attaches ligands

33. Higher level
The region surrounding the central atom or ion and containing the ligands is called the coordination sphere
A substance consisting of one or more complexes is called a coordination compound

34. Higher level
The number of lone pairs bonded to the metals ion is known as the coordination number
Coordination
number
Shape 6
octahedral 4
tetrahedral
or square planar 2
linear
Square planar compounds are rare, but usually d8 configurations with strong field ligands.

35. Co-ordination number examples
Higher level
Co-ordination number examples
Co-ordination number 6 4 2
Examples
[Fe(CN)6]3-
[CuCl4]2-
[Ag(NH3)2]+
[Fe(OH)3(H2O)3]
[Cu(NH3)4]2+
All of these complex ions are bonded to monodentate ligands which means they all consist of one type of ligand.
4 lobes is normally tetrahedral, but with full d8 (Cu) and strong ligands it is square planar.

36. Higher level
Ligands
A ligand is a neutral molecule or anion which contains a non-bonding pair of electrons, these electron pairs, from the ligand, form coordinate bonds with the metal ion to form complex ions
Complex ions form with transition metals because of their small size d-block ions attract species that are rich in electrons

37. Coordinate bonds and numbers
Higher level
Coordinate bonds and numbers
A common ligand is water and most (but not all) transition metal ions exist as hexahydrated complex ions in aqueous solution, e.g. [Fe(H2O)6]3+
(1) The lone pair of electrons from the water molecules (ligands) form the coordinate bonds.
(2) Iron is forming 6 bonds, so the coordination number of the iron is 6

38. Higher level
EDTA
Ethylenediaminetetraacetate (EDTA) is a chelate or polydentate ligand as it grabs onto the metal with more than one donor atom.
It has several important uses including the removal of heavy metal ions such as treatment for lead poisoning, as well food preservation in preventing transition metals catalyzing food rancidity (going off).

39. Higher level
13.2 Coloured complexes
OBJECTIVES
• The d sub-level splits into two sets of orbitals of different energy in a complex ion.
• Complexes of d-block elements are coloured, as light is absorbed when an electron is excited between the d-orbitals.
• The colour absorbed is complementary to the colour observed.
• Explanation of the effect of the identity of the metal ion, the oxidation number of the metal and the identity of the ligand on the colour of transition metal ion complexes.
• Explanation of the effect of different ligands on the splitting of the d-orbitals in transition metal complexes and colour observed using the spectrochemical series.

40. Crystal field theory (CFT)
Higher level
Crystal field theory (CFT)
CFT suggest that the dxy, dyz, dxz orbitals are of a lower energy state (more stable) than the 𝑑 𝑥 2 −𝑦 2 and 𝑑 𝑧 2 orbitals creating a split d-sublevel.
The x, y, z axis is
where the atomic axis
where ligands join.

41. Ligands and energy levels
Higher level
Ligands and energy levels
Strong field ligands cause spin paired splitting which has higher energy levels as the electrons in the ligands are repelling the electrons in the metal which are on the same axis.
These orbitals are not directly in line with the ligands.

42. Crystal field splitting energy (∆0)
Higher level
Crystal field splitting energy (∆0)
The free electron pair orbitals of the ligands are attracted to the nucleus of the metal cation. The overlapping orbitals with the 𝑑 𝑥 2 −𝑦 2 and 𝑑 𝑧 2 can cause electrons to move to the split lower energy orbitals of dxy, dyz, dxz which do not have direct overlap with the ligands. This difference in energy is the ∆0 and varies in size, hence movement of electrons between these split d orbitals will produce different wavelengths (seen as different colours).

43. Higher level
1. ∆0 and ligand strength
More negative charge density in ammonia makes ∆0 greater than than water. Similarly, as fluorine has more electron density than chlorine, so ∆0 increases.

44. Higher level
2. ∆0 and metal ions
Descending down a group the ∆0 increases due to more orbital overlap.

45. Higher level
3. ∆0 and oxidation state
The higher the oxidation state of the cation the more ∆0 increases due to greater attraction of the nucleus to the ligand and hence more orbital overlap.

46. Crystal field splitting energy (∆0)
Higher level
Crystal field splitting energy (∆0)
The crystal field splitting energy is the difference in energy between these two split sublevels.

47. Higher level
Ligands and ∆0
The strength of the ligand determines the amount of splitting. Splitting can be spin free or spin paired.
HIGH energy split, low spin
LOW energy split, high spin

48. Diamagnetism and paramagnetism
Higher level
Diamagnetism and paramagnetism
Eg. CN is a strong-field ligand creating diamagnetism, H2O is a weak field ligand creating paramagnetism.
diamagnetic
paramagnetic

49. Electron configurations and ∆0
Higher level
Electron configurations and ∆0
The lower energy state is written as 𝑡 2𝑔 and the higher energy state is written as 𝑒 𝑔 . The configurations are written as follows: 𝑡 2𝑔 5 𝑒 𝑔 0 𝑡 2𝑔 4 𝑒 𝑔 2

50. Spectrochemical series
Higher level
Spectrochemical series
The spectrochemical series (from data booklet) shows the splitting strength of the ligands:
Weak field ligands
Strong field ligands
M3+ metal ion
M2+ metal ion
Weak field ligands cause spin free splitting, <∆0.
Strong field ligands cause spin paired splitting, >∆0.

51. Incomplete square planar explanation
Higher level
Incomplete square planar explanation
d8 configurations with strong field ligands (low spin) will look like this  Hence with the greater repulsion on the lower energy orbitals, the ligands will best line up on the x and y axis of the dx2-y2(square planar) to be as far apart from each other as possible (rather than dz2) or in a tetrahedral shape.
This is to answer a textbook Q

52. Other configurations Higher level 4 lobes 2 lobes 6 lobes
Linear as z2 is slightly closer to nucleus
Square planar – d8 with strong ligand (Cu)
Tetrahedral as visually looks like the easiest acess to the nucleus with a 3D model
Octahedral – only symmetrical place for 6 ligands

53. Colours
Higher level
Many complex ions are colored because the energy differences between d orbitals match the energies of components of visible light
The colour absorbed can be determined by taking the colour transmitted (observed) and finding the opposite wavelength on the colour wheel (from data booklet).
Yellow light transmitted
Violet light absorbed

54. Why?
Higher level
Corundum Al2O3 
Ruby Al2O3: 1% Cr3+

55. Higher level
TOK Question here: Is it red?

56. Why?
Higher level
Beryl Be3Al2Si6O18 
Emerald Be3Al2Si6O18: 1% Cr3+

57. Higher level
Short medium long wavelengths

58. Different colours of chromium – induce the strength of the ligand…
Higher level
Different colours of chromium – induce the strength of the ligand…
Induce from the colours absorbed, whether Cl or water is a stronger field ligand. ANSWER: red is lower energy so splitting must be less, so Cl is a weaker ligand than water.