UEQ What else does the Periodic Table tell us about our atoms?

LEQ What did the early Periodic Tables look like?

Chapter 6 The Periodic Table and Periodic Law A. Development of the Periodic Table (Read details in the text) 1. Antonie Lavoisier: Compiled a list of 23 known elements. 2. John Newland: Arranged known elements by mass resulted in a repeat of properties every eight element. He called this the Law of Octave.

3. Lothar Meyer and Dmitri Mendeleev: Like Newland, arranged known elements in increasing atomic mass. Mendeleev then published and is given credit for the first periodic table. 4. Henry Moseley: Arranged elements in increasing atomic numbers (the # of protons in the nucleus). The arrangement still repeated in a regular pattern and is called the periodic law.

LEQ What are the different parts of the Periodic Table?

Periodic Table s, p, d, f s and p are representative elements # of electrons in valance Periods and Groups (Oxidation #,s) # of electrons in transitional elements Dot configuration Hybrid formation Metals (loss e-, cation) Non-metal (gain e-, anion)

The Modern Periodic Table Metals(ion forms) Non-metals (ion forms) Periodic Box Representative elements Transitional elements Alkali metals Alkaline earth metals Lanthanide series Actinide series Noble gases (inert gases) Metalloids Halogens

Periodic Trends 1. s, p, d, and f block 2. Valance electrons (fig 6.7) a. electron configurations (shorthand) b. ionic configuration 3. Dot configurations 4. Chemical and physical properties based on repeatable trends.

Review Sampler 1. Who gave us the Law of Octave? ans: John Newland 2. Who is the ‘Father of the Modern Periodic Table’? ans: Henry Mosley 3. Who gave us the ‘Periodic Law’ ans: Henry Mosley 4. What does the Periodic Law state? ans: properties of elements repeat in a regular pattern of eight based on the atomic number of the element. 5. Who proposed a periodic table based on the atomic mass? ans: Dmitri Mendeleev

Review Sampler 6. What part of the periodic table will you find the representative elements? ans: In the s and p block, Group A elements 7. Group 3A has how many valance electrons? ans: 3, s 2 p 1 8. What is the oxidation number for Cl ? ans: 1-

Review Sampler 9. What is the oxidation number for Mg and Sr ? ans: both are Write the dot configuration for AlMgC 11. What is a cation? Give an example. ans: + ion or metal. Any metal like Fe or K. 12. What is an anion? Give and example. ans: - ion or non-metal. Any non-metal like Cl or O

Review Sampler 13. What Group # will we find the Halogens? ans: Group 7A 14. Given an example of an alkali metal? ans any element in Group 1A 15. Given an example of a metalloid ? ans: Si, Ge, As, Sb, Te, Po, At 16. In which ‘series’ will you find Uranium? ans: Actinium series 17. Why are the Noble gases also called inert? ans: each has a filled valance shell with 8 electrons 18. Which has a 2+ oxidation number? Na Cr Ca N Cl

19. Write the oxidation number for the following: OKMg NiFeBr BaLiS BPXe

LEQ How are the different trends in the Periodic Table interpreted?

5. Atomic Radius: One-half the distance between two adjacent nuclei. (fig 6.11 and 6-12) a. Within the A Groups: Atomic radius decrease left to right within a period. Reason: the number of valance electrons and the atomic number (the # of positive charged particles) increase left to right and the number inner electrons (shielding effect) remains the same.

Atomic radius increase from top to bottom within a group. Reason: the number of valance electrons remains the same while the number of inner shielding electrons increases.

Atomic Radii

Ionization energy Ionization energy: the amount of energy required to remove an electron from a gaseous atom. Electrons removed are from the outer most portion of the electron cloud.

Trends: (fig 6-16 and 6-17) From left to right across within a single period, the amount of ionization energy increases (takes more and more energy to remove an electron). How does this relate to the atomic radius for the same atoms? Top to bottom within a group, the ionization energy decreases (takes less energy to remove an electron). How does this relate to the atomic radius for the same atoms?

Note: Table 6-5, page 192. What are the trend in the values for the 1st to the 2nd to the 3rd ionization energy relative to the position of the atom on the periodic table?

Ionization Energy

Ionic Radius A. The size of the atom as a result of the gain or loss of electron(s). Or when an ion is formed. (fig 6-14 and 6-15) Within the A Group: For positive ions, the ionic radius is always smaller than the atomic radius for the same atoms. For negative ions, the ionic radius is always larger than the atomic radius for the same atom.

Trend: - From left to right within a single period, the ionic radius decreases according to the charge. -From top to bottom within a single group, the ionic radius increased according to the charge. Isoelectroic ions: same configuration w/ different atoms

Electronegativity Electronegativity: the ability of an atom to attract the electrons of another atom. - The values range from less than one to 4.0. The larger the value, the more likely the atom will attract electrons. Fluorine is the most electronegative with a value of 4.0 and Francium is the least electronegative with a value of 0.7.

Trends: The trends are more general than other trends. (Fig 6-18) From left to right within a single period, the electronegativity tends to increase. Top to bottom within a single group, the electronegativity tends to decrease. Calculate EN value: Used to define type of bonds.

Electron Affinity The amount of energy absorbed when an electron is added to an atom to form an ion with a 1- charge. * Elements with very negative electron affinities gain electrons easily to form a negative ion (anion). - non-metal >> metals * A negative value for electron affinity will indicate that energy is released.