1. Topic 3
Periodicity

2. 4.1 The Periodic Table: Learning Goals
Understand how the elements in the periodic table are arranged
Understand the terns ‘group’ and ‘period’
Understand how the electronic configuration of an element relates to its position in the periodic table

3. The Periodic Table Arranged in order of increasing atomic number
Vertical columns and horizontal periods
Most of the elements are metals
The elements bordering between metals and non-metals are sometimes referred to as semi- metals or metalloids (Si, Ge, Sb)
Some of the groups have specific names

6. History of Chemistry
In 1869 Russian scientist Dmitri Mendeleev published the original periodic table
He arranged the elements based on increasing atomic weight
He also left gaps for elements that had not yet been discovered and predicted their properties
The modern periodic table is arranged by increasing atomic number

7. The Periodic Table and Electron Configurations
The group number of an element indicates the number of electrons in the element’s highest main energy level (outer shell)
Group 1= 1 outer electron
Group 2= 2 outer electrons, etc.
The period number indicates the number of main energy levels (shells) in an atom
Period 1= 1 main energy level
Period 2= 2 main energy levels, etc.

8. 4.2 Physical Properties: Learning Goals
Define first ionization energy and electronegativity of an element
Understand trends in atomic radius, first ionization energy and electronegativity across a period
Understand trends in atomic radius, first ionization energy and electronegativity down group 1 and group 7
Explain the variation of melting points for elements across period 3 and down group 1 and group 7

9. Electronegativity
In covalent bonds between 2 different types of atoms, the atoms do not attract the electron pair equally
How strongly electrons are attracted depends on the size of individual atoms and their nuclear charge
Electronegativity decreases down a group
Because, the size of atoms increase down a group
The larger an atom is, the less pull from the nucleus thus, lower electronegativity

10. Example: HF vs. HCL

11. Electronegativity Across a Period
Increases across a period
Because, increases in nuclear charge with no significant change in shielding
Shielding remains approximately constant because atoms in the same period have the same number of inner shells
Also, the number of protons attracting the electrons increases

12. First Ionization Energy
The energy required to remove the outermost electron from a gaseous atom
Energy for: M(g) M+(g) + e-
Down a group:
I.E. decreases
Because, size of the atom increases so the outer electron is further from the attraction of the nucleus
Across a period:
I.E. increases
Because, the nuclear charge increases

13. Atomic Radius Size of an atom; mostly due to electron orbits Trends:
Increases down a group
Because, there are more electrons when going down a group
Decreases across a period
Because increase in nuclear charge with no increase in shielding (similar as electronegativity and ionization energy)
Across period 3, Na is the larges and Cl is the smallest
Ar is not counted because it does not bond, thus the radius cannot be measured

14. Ionic Radius Measure of the size of an ion
Positive ions are smaller than their atomic radii
Negative ions are greater than their atomic radii
Quiz

15. Ionic Radii Trends Not as simple as neutral atomic radii
Due to the type of ions changing from side to side
Ex. Na+ and Mg2+
They have the same number of electrons, but Mg2+ has the highest nuclear charge (pull of nucleus on electrons) and is smallest
Ex. P3- and S2-
Equal electrons, but S2- has the stronger nuclear charge, thus is smaller
Trend is generally similar to atomic radii

16. Melting Point Depends on the type of bonding in an element
Does not have a consistent trend
The alkali metals are metallic m.p. decrease
Because, the delocalized electron sea is further from the pull of the nucleus
The halogens increases down the family
Because the relative atomic masses of X2 increases and van der Waal’s forces increases
van der Waal’s Intermolecular forces resulting from temporary dipole-dipole interactions

17. Variation in M.P. across period 3
Must look at bonding to understand changes
Increases from Na to Mg to Al all metallically bonded
Silicon increases sharply has a giant covalent structure which are very difficult to break
There is a large decrease from Si to P (covalent) van der Waal’s forces are weak, thus not a lot of energy is required to break them
Slight increase from P to S covalent compound, but has a higher molecular mass and requires more E to break the bonds
Decrease from S to Cl to Ar also covalently bound, but much smaller masses thus, easier to break

18. Graph the boiling points
Na 98
Mg 649
Al 660
Si 1410
P 44
S 119
Cl -101
Ar -189

19. 4.3 Chemical properties of elements in group 1 and group 7
Chemical properties of elements in the same group
Chemical reactions depend on the amount of outer electrons an atom contains
Thus, families have basically the same properties

20. Reactions of elements in group 1
Very reactive metals that react readily with oxygen, water, halogens, and other things
Nearly all of the reactions involve the single outer electron being removed
Reactions are more vigorous going down the column due to lesser ionization energy
Reaction with oxygen:
React vigorously with oxygen
4M(s) + O2(g) 2M2O(s)

21. Con’t Reaction with water Rapid reactions with water
2M(s) + 2H2O(l) 2MOH(aq) + H2(g)
The alkali metal hydroxides are strong bases
Reactions get more vigorous descending down the column

22. Reactions of Group 7 Elements
Reactions generally are due to the elements gaining an electron to have an octet
Reactivity decreases down the family
Fluorine is the most reactive element known
Basically reacts with every element on the periodic table
Variations in reactivity are not as easily explained as group 1
Several factors must be considered when explaining

23. React with alkali metals to form salts
Halogen Reactions
React with alkali metals to form salts
2M(s) + X2(g) 2MX(s)
Salts formed are white or colorless
Displacement Reactions
Reactions between a solution of a halogen and a solution containing halide ions
Ex: Cl2(aq) + 2KBr(aq)  2KCl(aq) + Br2(aq)

24. 4.4 Properties of the Oxides of Period 3 Elements
Look at table 4.3 p.156
Sodium oxide, magnesium oxide, aluminum oxide
Have giant ionic structures
Ions held by strong electrostatic forces causing high melting points and boiling points

25. Covalent Oxides
Look at 4.4 p. 158

26. Reactions of Period 3 Oxides with Water
In general, metallic oxides are basic and non-metallic oxides are acidic
Ex. Na2O(s) + H2O(l)  2NaOH(aq)
P4O10(s)+ 6H2O(l)  4H3PO4(aq)

27. HL first years

28. 4.5 Properties of the Chlorides of Period 3 Elements
Table 4.6 p. 161
Bonding of aluminum chloride is complex and undergoes changes of structure and bonding as it changes states
We assume bonding is covalent molecular in all states
In the Lewis structure Al only has 6 electrons and will bond with another AlCl3 to form a dimer
Overall structure is non-planar
The bonds are weak and easily broken, thus low melting/ boiling points

29. Con’t SiCl4 and PCl3 covalent molecular liquids PCl5 complex bonding
SiCl4 is tetrahedral and PCl3 is trigonal pyramidal
van der Waals’ forces are weak, so they have low melting/ boiling points
Neither conduct electricity
PCl5 complex bonding
Solid at room temperature
Forms a covalent molecular liquid when melted
van der Waals’ are stronger than PCl3, has a slightly higher melting/ boiling point
Does not conduct electricity

30. Con’t Cl2 covalent molecular gas Diatomic molecule
Weak van der Waals’ forces in liquid chlorine
Low molecular mass
Bonds can be broken at temperatures below room temperature
Does not conduct electricity

31. Reactions of Period 3 Chlorides with Water
Table 4.7 p. 162
NaCl(s) Na+(aq) + Cl-(aq)
[Mg(H2O)6]2+(aq)  [Mg(H2O)5(OH)]+(aq) + H+(aq)
AlCl3(s) + 3H2O(l)  Al(OH)3(s) +3HCl(g)
AlCl3(s) + 6H2O(l)  [Al(H2O)6]3+(aq) + 3Cl-(aq) (with excess water)
SiCl4(l) + 4H2O(l)  Si(OH)4(s) + 4HCl(aq)
PCl3(l) + 3H2O(l)  H3PO3(aq) +3HCl(aq)
PCl5(l) +4H2O(l)  H3PO4(aq) + 5HCl(aq)
Cl2(aq) + H2O(l) HCl(a) + HOCL(aq)

32. 4.6 The Transition Elements: Learning Goals
Describe the characteristic properties of transition metals
Explain why transition metals have variable oxidation states
Explain why formation and describe with shape of complex ions
Explain why transition metal complex ions are colored
Describe some uses of transition metals and their compounds and catalysts

33. The Transition Metals The ‘d-block’ elements 3d to 7d elements
Defined as: an element that forms at least one stable oxidation state (other than 0) with a partially filled d subshell
Scandium and zinc excluded from transition metals

34. Properties of Transition Metals
All typical metals; high m.p., b.p. and densities
Ionization energies increase from Ti to Cu, but only slightly
Radii decrease from Ti to Cu, but only slightly
Can exhibit more than one oxidation number in compounds/ complexes
Form complex ions
Usually form colored complexes/ compounds
In compound/ complex form can act as catalysts in many reactions

35. Ionization of Transition Elements
The transition elements form positive ions
Common transition metal ions
Element
Ion(s) Cr
Cr2+ and Cr3+ Mn
Mn2+ Fe
Fe2+ and Fe3+ Co
Co2+ Cu
Cu+ and Cu2+

36. Variable Oxidation Numbers
All transition metals show an oxidation state of 2+
Because they have a full 4s subshell and removal of these electrons would result in a 2+ charge
Table 4.25 Pg. 167

37. 4s and 3d subshells are close in energy
Cause of Multiple O.N.
4s and 3d subshells are close in energy
Removing electrons from either require nearly the same amount of energy
Thus, lost electrons depend on several factors
Lattice enthalpy
Ionization energy
Hydration of enthalpy

38. Complex Ions
Consist of a central transition metal ion surrounded by a ligand
Ligand negative ions or neutral molecules that use lone pairs of electrons to bond to a transition metal ion to form a complex ion
Dative covalent bonds are formed between the ligand and the transition metal ion
Look and Fe complex ion
Except for Ti, al transition metals form an octahedral complex ion to form [M(H2O)6]2+ in solution
Draw form pg 168

39. Oxidation number of a Transition Metal in a Complex Ion
May be determined based on charge of ligand
Ligands are either negative or neutral
Deduce the charge of Ni in the complex ion [Ni(CN)4]2-
Neutral Ligands
1- Ligands
H2O
Cl-
NH3
CN- CO
Br- 2+

40. Transition metal complexes do not obey the VSEPR theory rules
Shapes of Complex Ions
Transition metal complexes do not obey the VSEPR theory rules
Six coordinate (dative covalent bonds) complexes are nearly always octahedral
Four coordinate complexes are either tetrahedral or square planar
Table 4.9 P. 169

41. Formation of Complex Ions
May undergo substitution reactions
Ex. Water replaced with another ligand
[Cu(H2O)6]2+(aq) + 4Cl-(aq)  [CuCl4]2-(aq) + 6H2O(l)
The ligands are considered Lewis bases and according to Le Chatelier’ principle, a strongly acidic solution means that the concentration of H+ ions is high and the position of equilibrium is shifted to the left-hand side

42. Formation of Colored Complexes
Complex ions form colored solutions as a result of their 3d orbitals being split into 2 groups by ligands
The electrons are promoted from the group of lesser energy to the group with higher energy, causing an emission of colored light
Ions with no electrons in the 3d subshell are colorless
Similarly, ions with 10 electrons in the 3d subshell are colorless because their electrons will not move between the orbitals

43. Catalytic Ablility
The elements and their compounds/ complexes are able to act as catalysts
Ex.
Iron is the catalyst in the Haber reaction to produce ammonia
Vanadium (V) oxide can be used in the Contact process of convert sulfur(IV) oxide to sulfuric(VI) acid
Manganese(IV) oxide is used in the decomposition of hydrogen peroxide
Nickel is used in the hydrogenation of alkenes to form alkanes

44. Catalysts of Haber and Contact Processes
Catalysts are used to speed up a reaction without themselves being used up
With a catalyst, less energy and time are required to gain product
The use of the catalysts in both the Haber and Contact processes is to run the reaction at a lower temperature and increase the yield of product
Improves the overall cost effectiveness of these processes in industry