History and Trends

1. The PERIODIC TABLEofELEMENTS

2. The development

3. J.W. Dobereiner Classified elements into several sets of triads TRIADS Li, Na, K Ca, Sr, Ba Have similar chemical properties Properties of the middle element are approximate averages of the first and third elements

4. J.A.R Newlands 62 elements are already known “law of octaves” (increasing atomic mass) Similar properties: 8th and 1st , 9th and 2nd , 10th and 3rd

5. Dmitri Mendeleev Together with Lothar Meyer, they published nearly identical schemes for classifying the elements Wrote the names of the elements nd its properties in cards and arranged the cards in various ways.

6. He noticed that a periodic repetition of the properties of the elements could be observed when the elements were arranged in increasing atomic masses. He eventually produced the first periodic table of elements.

7. H.G.J. Moseley Correctly hypothesized the fundamental property of each elements – the amount of positive charge in the nucleus = Atomic no. Proposed that the correct way to arrange elements was with the increasing atomic number

8. The PERIODIC LAW Basis for the Periodic Table When elements are arranged in increasing atomic number, heir physical and chemical properties show a periodic pattern.

9. Reading the periodic table

10. Groups or Families The arrangement of the elements in vertical columns Each family has the similar properties

11. Period Horizontal rows in the periodic table

12. Labeling and naming groups

14. Periodic Trends

15. Metals and Non - Metals

16. Effective Nuclear Charge The measure of the attraction between the nucleus and the electron Also defined by the equation Zeff = Z – S Wherein, Z = number of proton in the nucleus S = average number of electrons that are between the nucleus and the electron in question

18. Greater Zeff The greater the attraction between the nucleus and the electron The electron are drawn closer to the nucleus The atomic size is reduced

19. Greater Shielding constant The lower the Zeff The lesser the attraction between electrons and the nucleus

20. Electrons in the inner shell (lower n values), effectively shields the electrons in the outer shell (higher n values) However the electrons in the same shell do not effectively shield one another Example: electrons in the fourth energy level

22. Trend? Effective nuclear charge increases across any row in the periodic table. Effective nuclear charge increases slightly moving down a family/group

23. Atomic Size Half of the internuclear distance between adjacent atoms (atomic radii) Trend? Across a period (L-R) = decreases This is due to the fact that moving from left to right across a period, the atomic number increases while the shielding factor, does not significantly increase; therefore, the Zeff increases thus pulling the electrons towards the nucleus

24. Down a group = increases Down a group the principal quantum number of the outermost electron increases

26. Check-up Arrange the following atoms in order of increasing size: P, S, As, Se 2. Arrange the following atoms in order of decreasing atomic radius: Na, Be, Mg

27. Ionic Size An estimate of the size of an ion in a crystalline ionic compound From the relationship between the nuclear size and the atomic size, the size of the ion relative to its parent atom can be predicted

29. Isoelectronic series isoelectronic – ions having the same number of electrons Size decreases as the nuclear charge (atomic number) increases.

30. Check-up Arrange the atoms and ions in order of decreasing size: Mg2+, Ca2+, Ca Which of the following atoms and ions is the largest ? S2-, S, O2- Arrange the ions in order of decreasing size: S2-, Cl-, K+, Ca2+

31. Ionization Energy The minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. First Ionization energy Energy needed to remove the first electron from a neutral atom Second Ionization energy Energy needed to remove the second electron The greater the ionization energy, the more difficult it is to remove an electron

33. Check-up Referring to a periodic table, arrange the following atoms in order of increasing first ionization energy. Ne, Na, P, Ar, K

34. Electron Configuration of ions When electrons are removed from an atom to form cation, they are always removed first from the orbitals with largest available principal quantum number, n. Example : Li  Li +

35. When electrons are added to an atom to form anion, they are added to an empty or partially filled orbital Example: F  F-

36. Check-up Write the electron configurations for the Ca2+ ion Co3+ ion S2- ion

37. SW

38. Electron Affinity Energy change that occurs when an electron is added to an atom Measures the attraction, or affinity, of the atom for the added electron

39. Ionization energy vs Electron Affinity Ionization energy Measures the ease with which an atom loses an electron Electron affinity Measures the ease with which an atom gains an electron

43. Electronegativity Ability of an atom to attract electrons o itself

47. Worksheet