1. PERIODICITY " THERE ARE ONLY TWO WAYS TO LIVE YOUR LIFE. ONE IS AS THOUGH NOTHING IS A MIRACLE. THE OTHER IS AS THOUGH EVERYTHING IS." -- ALBERT EINSTEIN Ch.6 J.C. Rowe Windsor University School of Medicine

2. Modern Periodic Table

3. What is the Periodic Table?  An arrangement of the chemical elements in which the elements are arranged by order of atomic number in such a way that the periodic properties (chemical periodicity) of the elements are made clear.  The standard form of the table includes periods (usually horizontal in the periodic table) and groups (usually vertical). Elements in groups have some similar properties to each other.  The periodic table is a masterpiece of organized chemical information. The evolution of chemistry's periodic table into the current form is an astonishing achievement with major contributions from many famous chemists and other eminent scientists

4. What is the Periodic Table? Cont’d  there are a number of ways that the elements have been arranged--we will use the "Modern Periodic Table" which is derived from Mendeleev's version of the organization of the elements.  Modern Periodic Table is based upon increasing atomic number of the elements and allows for organization of the elements as well as grouping of the elements according to chemical and physical properties Groups – are the vertical groupings or columns of elements in the table Periods – are the horizontal groupings or rows of elements in the table

6. The Mendeleev Periodic Table

8. Groups & Periods  The vertical columns in the Periodic Table are called GROUPS.  Groups are numbered with Roman numbers (I, II, III….)  The horizontal rows are called PERIODS.  Periods are numbered with ordinary numbers. groupsgroups periods

9. Three major groupings of types of elements  Metals -- generally hard, shiny, dense, conductive with a shiny luster examples would be copper, lead, sodium know which elements on the periodic table are metals -- in the periodic table in the front of your eyes, metals are in yellow  Non-metals -- do not exhibit the properties of metals, they are generally insulators, brittle, and often gaseous examples would be chlorine, neon, sulfur know which elements on the periodic table are non-metals -- in front of your eyes, non-metals are green  Metalloids – have properties in between those of metals and non-metals -- in front…..., metalloids are purple they are boron, silicon, germanium, arsenic, antimony, tellurium, astatine, know these

10. What is the Periodic Table?

11. Periodic Trends  Elements that border on the amphoteric line are Metalloids (dark gray). They have characteristics of both.  Metals (blue)  Nonmetals (light gray)

12. There are assorted named sub-groupings  There are assorted named sub-groupings which you should be familiar with:  Representative elements -- all Group A elements  Transition elements – all Group B elements Group # Name IA alkali metals IIA alkaline earth metals VA nitrogen group VIA oxygen group or chalcogens VIIA halogens VIIIA nobel gases

13. Group A vs. Group B

14. Atomic Number(Z) vs. Mass number (A)  two numbers on the periodic table which you must become very familiar with are:  the atomic number--it is the smaller number and is always a whole number  the mass number--it is the larger number and is a decimal number Atomic # Mass #

15. Characteristics 1. Have lusters 2. Are malleables & ductille 3. Conduct heat & electricity 4. Tend to lose electrons 1. Are dull 2. Are brittle 3. Do not conduct heat or electricity very well 4. Tend to gain electrons MetalsNonmetals

16. Stable electronic configurations  Metals tend to lose their outershell electrons in order to be stable.  Stable positively charged ions (cations)  The ability to lose electrons makes them good reducing agents.  Non-metals tend to easily gain electrons in order to be stable.  Stable negatively charged ions (anions)  The ability to gain electrons makes them good oxidising agents Metals / cations Non- metals/ anions

17. Ionic states :Cation or Anion  LOSS of electron leads to CATION (+)  GAIN of electron leads to ANION (-)

18. How Elements React  When an element reacts with another element, the outershell of elements of each atom is what first comes into contact with the other atoms.  When elements react, they are usually either : 1. Transferring electrons from their outershells to another element 2. Transferring electrons to their outershells from another element 3. Sharing electrons in their outershells with another element.

19. How Elements React Cont’d.  Often atoms react with other atoms in order to achieve an octet ( 8 electrons in their outer shell), so the number of electron in their outer shell determines how many electrons they need to gain or lose to achieve an octet.  So it is the number of electrons in the outer shell that is most important in determining how an element will react with another elements.

20. Valence Electrons

21. Valence electrons Cont’d.  Within a group, elements have very similar valence electron configurations--same number of electrons in the same type of subshells but the only difference is in the major shell.  The electron configurations of the valence electrons are very important.  the most important electrons in an atom are those on the outer-most edge of the atom (the outer major shell) – they come into contact with other atoms first and participate fully in reactions – they are the valence electrons  valence shells contain the valence electrons

22. s-p-d-f long form periodic table  s-block Group IA and IIA valence electrons are in the ns subshell where 'n' denotes the major shell or row  p-block Group IIIA - VIIIA valence electrons are in the ns and np subshell where 'n' denotes the major shell or row  d-block Group IB - VIIIB (Transition Metals) valence electrons are in the ns and (n-1)d subshell  f-block Inner Transition Metals sometimes called the lanthanides & the actinides valence electrons are in the ns and (n-2)d subshell

23. The electron configurations of the valence electrons are very important  s-block Group IA and IIA valence electrons are in the ns subshell where 'n' denotes the major shell or row  p-block Group IIIA - VIIIA valence electrons are in the ns and np subshell where 'n' denotes the major shell or row  d-block Group IB - VIIIB (Transition Metals) valence electrons are in the ns and (n-1)d subshell  f-block Inner Transition Metals sometimes called the lanthanides & the actinides valence electrons are in the ns and (n-2)d subshell

24. Family trends

25. Family

26. Atomic Size  atoms increase in size as the number of shells increase—in other words, as the valence electrons occupy larger shells the atoms become larger going down a group  atoms decrease in size as electrons are added to the same shell—in other words, the atoms become smaller across the periodic table from left to right, due to the contraction effect as more electrons and protons are added to the same shell (as you move from atoms on the left to atoms on the right), the attraction between positive and negative increases and the size of the shell is contracted

28. Atomic Radius

29. Ionic Size  positive ions are smaller than the atoms from which they are derived  the greater the size of the positive charge, the smaller size of the ion relative to the atom  negative ions are larger than the atoms from which they are derived  the greater the size of the negative charge, the larger the size of the ion relative to the atom

30. isoelectronic ions a series of isoelectronic ions (different ions with the same number of electrons) such as those above, are not the same size

31. Trends of Electron Affinity

32. Ionization Energy  ionization energy is the energy required to remove an electron in the formation of an ion—the more difficult it is to remove an electron, the greater the ionization energy  smaller atoms have greater ionization energy since the valence electrons are closer to the nucleus and more strongly attracted and, therefore, more difficult to remove  ionization energy increases going up a group and across a row in the periodic table

33. Ease of Ionisation

34. Summary

35. The primary duty of the University to a student is to provide him with such instructors as will make him realize that the responsibility for progress is his own and no one else's. S.E. Whitnall, 1933