CHAPTER 6

1. CHAPTER 6 Chemical Periodicity

2. Development of the Periodic Table • Scientists needed a way to organize and refer to the growing number of known elements • Initially, elements were categorized according to similarities in their physical and chemical properties (Dobereiner’s Triads, Mendeleev’s and Newland’s Periodic Tables) • Ultimately, the way in which the elements were organized reflected the relationship between their atomic structures and properties (Moseley’s Periodic Table)

3. Mendeleev’s Periodic table • Mendeleev listed the elements in horizontal rows in order of increasing atomic mass • Regular patterns of physical and chemical properties occurred in vertical columns • Blank spaces left where unknown elements belonged, but able to predict their properties

4. Moseley’s Periodic Table • Moseley determined the atomic numbers of the atoms of the elements • He arranged the elements by order of atomic number, not by atomic mass

5. Modern Periodic Table • The horizontal rows of the periodic table are called periods (organized by increasing atomic number) • The vertical columns are called groups or families (organized by chemical properties) • Periodic Law: when the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties

6. Organization of the periodic table Group Period

7. Metals, Metalloids, and Nonmetals Metalloids

8. Main group orRepresentative elements IA VIIIA IIA IIIA IVA VA VIA VIIA p-block s-block

9. Transition elements IIIB IVB VB VIB IB IIB VIIB VIIIB d-block f-block (inner transition elements)

10. Special groups within the periodic table

11. Electron Configurations and Periodicity • Electron plays the greatest part in determining the physical and chemical properties of an element • Elements can be classified into four different categories based on their electron configurations • Notice that the period number corresponds to the principal energy level for s and p block elements only.

12. Electron Configurations and Periodicity • The representative elements(s and p block elements) Alkali metals: group 1 elements Alkaline earth metals: groups 2 elements Halogens: nonmetallic elements of group 17 Noble gases: group 18 elements The outermost s and p sublevels of these elements are completely filled. They are also called inert gasesbecause they do not participate in chemical bonding. * Other metals, non-metals and metalloids are also classified as representative elements

13. Electron Configurations and Periodicity • Transition Elements (d and f block elements) The transition metalsare elements whose outermost s sublevel and the nearby d sublevel contain electrons The inner transition metalsare elements whose outermost s sublevel and the nearby f sublevel contain electrons

14. Periodic Trends in Atomic Size • The radius of atoms: atomic radii • From one group to the next, atomic radius decreases because the electrostatic attraction between increasingly larger nuclei increases, pulling the electrons at the same energy level closer to the nucleus • From one row to the next, atomic radius increases because increasing numbers of electrons shield the electrostatic force and valence electrons are located further and further away from the nucleus

15. Formation of Ions and ionic radii • Ionic radiusis the size of an ion (an atom that has gained or lost electrons. • Cations (positive charge) are formed when metal atoms lose electrons • Cations are smaller than their corresponding neutral atom because electrons are lost from the highest energy level, leaving behind electrons that are closer to the nucleus which experience a greater electrostatic force. • Anions (negative charge) are formed when non-metal atoms gain electrons • Anions are larger than their corresponding neutral atom because addition of electrons expands the electron cloud and the nucleus exerts the same electrostatic force.

16. Periodic Trends in Ionization Energy • The energy that is required to overcome the attraction of the nuclear charge and remove an electron is the ionization energy • Increased nuclear charge that occurs from one group to the next accounts for electrons being held more tightly within the atom. The more tightly they are held, the greater the amount of energy that is required to remove and electron from an atom • From one row to the next, ionization energy decreases because increasing numbers of electrons shield the electrostatic force and valence electrons are located further and further away from the nucleus

17. Periodic Trends in Electronegativity • The electronegativity(electron affinity) of an element is the ability of an atom to attract electrons within a bond • Electronegativity increases as you go across a period form left to right and decreases as you move down a group • Electronegativity values for elements are established relative to one another.

18. PERIODIC TRENDS Ionic radius Anion radius Cation radius

19. Valence Electrons • Valence electrons are the highest energy level electrons of any atom and are always found in the s and p sublevels. • The number of valence electrons in an atom largely determines its chemical properties. • The group number of the representative elements directly corresponds to the number of valence electrons of atoms in that group (American PT)

20. Lewis Dot Structures • Lewis Dot structures represent atoms with their valence electrons. • Octet Rule: Atoms will gain, lose or share electrons when forming chemical bonds to resemble the stable electron configuration of the noble gases which have a full octet (8) of valence electrons.