Periodic Relationships Among the Elements Chapter 8

Periodic Relationships Chapter 8 Topics History of periodic table and periodic trends Periodic Classification of element Periodic trend of Physical properties (Atomic size, Ionization energy, Electron Affinity, and Electronegativity ) Ions electron configuration Ions Sizes (isoelectronic ions) Group elements and its oxides Periodic Relationships Dr. Ali Bumajdad

History of periodic table and periodic trends Lother Meyer (German, 1830-1895) and Ivanonich Mendeleev (Russian, 1834-1907) put the form of periodic table Mendeleev given most credit because he predict the existence and properties of unknown elements and he corrected several values for atomic masses

Mendeleev’s periodic Table Atomic mass Not atomic number Corrected (they thought its atomic mass is 76)

When the Elements Were Discovered

The current periodic table is different than Mendeleev by: 1) Element order by atomic number not by atomic mass 2) Contain many more elements discovered after Mendeleev.

Classification of the Elements

Atomic radius Ionization energy Electron affinity Electronegativity Atomic radius Ionization energy Electron affinity Electronegativity

-Electron is added on same principle quantum number Effective nuclear +ve charge increase due to no shielding effect  Size decrease -Electron is added on different principle quantum number Effective nuclear +ve charge not much change due to the shielding effect  Size increase

(1) Atomic size (2) Ionization energy (3) Electron affinity Effective nuclear charge (Zeff): the “positive charge” felt by an electron (1) Atomic size (2) Ionization energy (3) Electron affinity Atomic radius is half the distance between two neighboring atoms in their solid state Energy required to remove e from an isolated gaseous atom, ion or molecules (kJ/mol) Energy released or absorbed when an electron is added to an isolated gaseous atom, ion (kJ/mol) Na(g) + E Na+ + 1e Cl(g) + 1e Cl- + E As you go down a group the atom size increase (Zeff is approx. constant or even decrease and e placed in a new energy level) Release =exothermic Endothermic process (+ve E) because e is attracted to nucleus Cl-(g) + 1e + E Cl2- As size increase it is easier to remove e Absorbed = endothermic

(1) Atomic size (2) Ionization energy (3) Electron affinity As you go left to right across a period atom get smaller (Zeff increase and e placed in the same energy Level. All element (except H) have more than 1 ionization energy All element have more than 1 Electron Afffinity 1st Eionization < 2ed Eionization <… (removing e from neutral atom is easier than from +ve ion) 1st Eaffinity usually exothermic because e placed into environment where it experience attraction to nucleus For transition elements and inner transition elements, the last e,s are added onto an inner shell (3d) not the valence shell. This cause the Zeff to increase but not much (there is a slight shielding effect) atomic size decrease but slowly. Q) Why the 2nd Eionization for alkali metal is much more than the first one? It has a noble gas electron configuration 2st Eaffinity endothermic because the second e added to a negatively charge ion

(1) Atomic size (2) Ionization energy (3) Electron affinity 1 Why? 2 Q) Why the 3ed Eionization for earth alkali metal is the highest change? It has a noble gas electron configuration Q) Why Zr have similar size as Hf? A raw of Lanthanide elements Some Irregularities a cross a period 1 P Eaffinity < As affinity For ions: +ve ion < -ve ion O Eaffinity < S Eaffinity Why? Because e are crowding in shell 2 more than in shell 3 more repulsion in shell2 less energy evolve less Eaffinity 2 C has -ve Eaffinity while N has +ve Eaffinity

(1) Atomic size (2) Ionization energy (3) Electron affinity Why? Some Irregularities a cross a period Because the e in c inter a vacant 2p orbital while for N e must be placed in an already occupied orbital Be Eionization > B Eionization N Eionization > O Eionization P Eionization > S Eionization Why? For Be and B, it is easier to remove e from 2p1 than from 2s2 because it is further apart from nucleus For N and O, it is easier to remove e when there is e-e repulsion in an orbital

Effective Nuclear Charge (Zeff) increasing Zeff increasing Zeff

Adding electron to the same principal quantum number Adding electron on different principal quantum number

Sa Ex. : Eionization of P = 1060kJ/mol and for S = 1005 kJ/mol, explain this irregularity. The valence electron configuration for S contains one filled orbital and hence there is e-e repulsion and hence easier to be removed Sa Ex. : Among the following elements (Ne, Na, Mg) (a) Which atom has the largest 1st ionization (b) Which atom has the smallest 2nd ionization energy Ne Mg (the first and second e’s are valence electrons) Be2+ < Mg2+ < Ca2+ < Sr2+

The Alkali Metals Li Cs More reactive (why) Easier to loss electron (less effective nuclear charge) In aqueous system less reducing ability (why) Difficult to be solubilized in water due to its large size (less effective nuclear charge), less polar

Ions electron configuration When two nonmetals react they share electrons, so that the valence electron configuration completes for both atoms When nonmetal react with metal, the ions form so that the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom and the valence orbital of the metal are emptied Na [Ne]3s1 Na+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. Ca [Ar]4s2 Ca2+ [Ar] Al [Ne]3s23p1 Al3+ [Ne] H 1s1 H- 1s2 or [He] Atoms gain electrons so that anion has a noble-gas outer electron configuration. F 1s22s22p5 F- 1s22s22p6 or [Ne] O 1s22s22p4 O2- 1s22s22p6 or [Ne] N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Cations and Anions Of Representative Elements +1 +2 +3 -3 -2 -1

Electron Configurations of Cations of Transition Metals 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]4s23d6 Mn: [Ar]4s23d5 Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5 Fe3+: [Ar]4s03d5 or [Ar]3d5

Ions Sizes (isoelectronic ions) Isoelectronic ions: ions contain the same number of electrons Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne No. proton 11 13 9 8 7 10 The ion with less protons is bigger N3- > O2- > F- > Ne > Na+ > Al3+ Dr. Ali Bumajdad

Se2- > Br- > Rb+ > Sr2+ Q) What neutral atom is isoelectronic with H- ? H-: 1s2 same electron configuration as He Sa. Ex. : Arrange the ions in order of decreasing size ? Se2-, Br-, Rb+, Sr2+ Se2- > Br- > Rb+ > Sr2+ Sa. Ex. : Choose the largest ion in each of the following Groups: Li+, Na+, K+, Rb+, Cs+ Ba2+, Cs+, I-, Te2-

Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.

Dr. Ali Bumajdad

The Radii (in pm) of Ions of Familiar Elements

Sa Ex. : Predict the trend in radius for : Be2+ ,Mg2+ ,Ca2+ ,Sr2+

Group 1A Elements (ns1, n  2)

Group 2A Elements (ns2, n  2)

Group 3A Elements (ns2np1, n  2)

Group 4A Elements (ns2np2, n  2)

Group 5A Elements (ns2np3, n  2)

Group 6A Elements (ns2np4, n  2)

Group 7A Elements (ns2np5, n  2)

Group 8A Elements (ns2np6, n  2) Completely filled ns and np subshells. Highest ionization energy of all elements. No tendency to accept extra electrons. 8.6

Properties of Oxides Across a Period basic acidic