Periodic Table Chapter 4.1, 4.2, 4.3

4-1: Development of the Periodic Table 1) History of the Periodic Table – By the end of the 1700’s scientists had identified only 30 elements (ex. Cu, Ag, Au, H2, N2, O2, C). By the mid 1800’s by using spectroscopy additional elements were identified by using their line spectra and about 65 elements had been identified.

A. J.W. Dobereiner: 1829 Organized the elements into groups with similar properties. He called these groups triads. The middle element is often the average of the other two. Ex) Cl – 35.5 Br – 79.9 I – 126.9 Ca Avg Sr Ba

Triads on the Periodic Table

B. J.A.R. Newlands: 1867 Law of Octaves. He said properties repeated every 8th element. There were 62 known elements at the time. He was also a musician.

C. Dimitri Mendeleev: 1869 1. Organized the 1st periodic table according to increasing atomic mass and put elements with similar properties in the same column.

2. He arranged some elements out of atomic mass order to keep them together with other elements with similar properties. He also left three blanks in his table and correctly predicted the properties of these 3 unidentified elements that were later identified and matched his predictions.

D. Moseley: 1915 Each element has a certain amount of positive charge in the nucleus which are called protons. 1. Moseley reorganized the periodic table by Atomic Number. 2. The Periodic Law: When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern.

Glenn Seaborg “Seaborgium” Sg #106 Born in 1912 in Michigan, Seaborg proposed reorganizing the Periodic Table one last time as a young chemist working on the Manhattan Atomic Bomb Project during WWII by pulling the “f-block” elements out to the bottom of the table. He was the principle or co-discoverer of 10 transuranium elements. He was awarded the Noble prize in 1951 and died in 1999.

4-2: Reading the Periodic Table Information in each square: B. Organizing the Squares 1. Vertical Columns – groups & families 2. Horizontal Rows - periods Atomic # (protons) 13 Al Aluminum 26.9815 Mass # - protons & neutrons particular isotope of aluminum like Al-26 or Al-27 Symbol Name Atomic Mass Weighted average of all an element’s isotopes

Parts of the Periodic Table

C. Electron Configuration & Families Carbon has 4 valence electrons Valence electrons – outermost electrons responsible for bonding. Elements in the same group have the same number of valence electrons.

E. Metal, Nonmetals and Metalloids (Semimetals): 3. Metalloids - Properties of both metals & nonmetals - good conductors of heat & electricity - most solids at room temperature (except Hg) - high luster (shiny) - ductile (can be drawn into thin wire) - malleable (bends without breaking) - high melting points - high densities - react with acids - brittle (shatters when struck) - low luster (dull) - neither ductile nor malleable - nonreactive with acids - nonconductors (Semimetals)

Alkali Metals have 1 valence electron: Brainiac Alkali Metal Video English Video – Noble Gases Clip

4-3: Trends in the Periodic Table Atomic Radius 1. The distance from the center of the nucleus to the outermost electron. 2. Atoms get larger going down a group and smaller going across a period. Ex) Na is larger than Mg Na is smaller than K

Atomic Radii of the Representative Elements

Atomic Radii vs Atomic Number

Ionic Size When atoms gain electrons, they become (-) and get larger.

Positive Ion Size When atoms lose electrons, they become (+) and get smaller. Ions get larger as you go down a group.

Relative Sizes of Positive & Negative Ions The sodium ion lost an electron, and therefore the positive protons in the nucleus exert a stronger pull on the remaining negative electrons, shrinking the orbitals. Thus positive ions are smaller than their atoms. The chloride ion gained an electron, and therefore the fewer positive protons in the nucleus exert a weaker pull on the extra negative electrons, increasing the size of the orbitals. Thus negative ions are larger than their atoms.

Electron Attraction in a Bond & Ion Size

C. Ionization Energy: The energy needed to remove one electron from an atom. Elements that do not want to lose their electrons have high ionization energies. Elements that easily lose electrons have low ionization energies. I.E. decreases down a group (opposite of atomic radius) I.E. increases across a period. (opposite of atomic radius) Ex) Na IE smaller than Mg Na IE larger than K

Ionization Energy of the 1st 20 Elements

Ionization Energy vs. Atomic Number

Sublevels by the Periodic Table

D. Successive Ionization Energies: Energy required to remove electrons beyond the 1st electron. Ionization energies will increase for every electron removed. 3. Na [Ne]3s1 Na• 1st = ____ kJ 2nd = ____ kJ 4. Mg [Ne]3s2 Mg: 1st = ____ kJ 2nd = ____ kJ 3rd = ____kJ 5. Al [Ne]3s23p1 Al: 1st = ____kJ 2nd = ____kJ 3rd = ____kJ 4th = ___kJ 496 4560 738 1450 7730 11,600 577 1816 2744

E. Electron Affinity: Energy change that occurs when an atom gains an electron. A highly negative electron affinity attracts electrons. (nonmetals) A positive electron affinity does not attract electrons. (metals)

Electron Affinity decreases

F. Electronegativity: 1. Reflects an atoms ability to attract electrons in a chemical bond. 2. E.N. decreases going down a group 3. E.N. increases going across a period. 4. Examples: NaCl and H2

G. Octet Rule: H. Oxidation Number: An atom will tend to lose, gain or share electrons in order to acquire a full set of valence electrons. “Octet” = 8 = s2p6 configuration H. Oxidation Number: The charge on an ion when it gains or loses electrons to acquire a stable octet.

Which of the following would have the largest? Atomic Radii? Ionization Energy? Electron Affinity? Element D Element C Element A

Which of the following would have the smallest? Atomic Radii? Ionization Energy? Electron Affinity? Element B Element D Element D

Electron Shielding