Topic 3 Periodicity

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

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

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

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.

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

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

Example: HF vs. HCL

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

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

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

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

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

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

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

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

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

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)

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

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

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)

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

Covalent Oxides Look at 4.4 p. 158

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)

HL first years

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

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

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

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)

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

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

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

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+

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

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

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

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+

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

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

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

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

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