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The Periodic Table. History Arrangement of Elements Electron Configuration Trends Periodic Trends Reactivity. A. Johann Dobreiner’s Law of Triads in 1817. B. John Newlands – Law of Octaves. Lothar Meyer (1835-1895 - German) P roperties of elements show a
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The Periodic Table History Arrangement of Elements Electron Configuration Trends Periodic Trends Reactivity
Lothar Meyer (1835-1895 - German) Properties of elements show a repetitive pattern when they are arranged by atomic mass D.Dimitri Mendeleev (1834-1907- Russian) (father of modern periodic table) Published system used today (1869) 2. Elements arranged by increasing mass 3. Left spaces for elements not yet discovered - predicted properties (scandium, gallium, germanium)
E. Henry Mosley (1887-1915) English 1.Arrange elements by increasing atomicnumber – this led to the periodic law 2. Periodic Law- properties of elements are periodic functions of their atomic # periodic repetition of physical and chemical properties
II. Arrangement of Elements • Periodic Table – arrangement of elements in order of increasing atomic number so that elements with similar properties are in the same column period – horizontal row (7) group(family)- vertical columns (1-18) periodicity – reoccurrence of similar properties of elements in groups
E. Metals – Metalloids - Nonmetals 1. Metals are on the left side – all are solids except mercury (Hg) a. elements near the left of a period are more metallic than those near the right b. elements near the top of a group are more metallic than those near the bottom 2. Metalloids – group of elements between metals and nonmetals(B,Si,Ge,As,Sb,Te) 3. Nonmetals are on the right side – all are solids or gases except bromine(Br) liquid
PROPERTY METALNONMETAL Lusterhighlow Deformabilitymalleablebrittle and ductile Conductivitygoodpoor Electron gain/loselosegain Ion formedcation (+)anion(-) Ionization energylowhigh Electronegativitylowhigh
IV.Periodic Trends(Main Group Elements) • Atomic Radii 1. atomic radius is ½ the distance between nuclei of identical atoms joined in a molecule 2. decreases across periods (left-right) a. caused by increasing attraction between protons and electrons 3. increases from top to bottom a. caused by adding electrons to new shells
What is the atomic radius? Atomic radii include the region in which electrons are found 90% of the time
Atomic Size } • Atomic Radius = half the distance between two nuclei of a diatomic molecule. Radius
B. Ionization Energy 1. Energy required to remove an electron from an atom of an element (KJ/mol) 2. Increases across periods (left to right) a. result of increased nuclear attraction 3. Decreases down groups (families) a. electrons added to higher energy levels b. shielding effect of inner shell electrons c. repulsion of inner shell electrons 4. Energy to remove second and third electron is greater
Symbol First Second Third 11810 14840 3569 4619 4577 5301 6045 6276 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne
C. Electronegativity • Measures how strongly one atom attracts the electrons of another atom when they form a compound • Increases across periods (left to right) a. Fluorine has greatest value of 4 3. Decreases down groups a. electrons far from the nucleus in larger atoms have less attraction b. Cesium and Francium with large radii have the smallest electronegativity
D. Ionic Radii • Ion – atom that acquires a charge by gaining or losing electrons a. cation (+) ion anion (-) ion 2. Period trends a. cation radii decrease across periods b. anion radii increase across periods 3. Group trends a. increase in cation and anion radii down groups
E. Electron Affinity 1. Energy change that occurs when an electron is added to a neutral atom 2. If it is easy to add an electron to an atom the energy value is negative a. halogens have large negative values 3. If it is difficult to add an electron to an atom the energy value is positive a. atoms in groups 2 and 18 have high positive values (due to filled subshells) b. usually higher values in larger atoms
V. Reactivity • Reactivity – measure of the tendency of an element to engage in chemical reactions by losing, gaining or sharing electrons 1. atoms of reactive elements are very likely to gain, lose or share electrons 2. atoms of reactive elements are likely to form chemical bonds with other elements
B. Reactivity and the Periodic Table 1. alkali metals (group 1) most reactive metals 2. alkaline earth metals (group 2) second most reactive group of metals 3. halogens (group 17) most reactive nonmetals 4. noble gases (group 18) least reactive C.Ionization Energy and Electronegativity 1. elements with very high and very low values are very reactive
Electron Configuration • S block [groups 1 and 2] • P block [groups 13,14,15,16,17,18] • D block [groups 3,4,5,6,7,8,9,10,11,12] • F block (lanthanide and actinide series)
H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 1s1 group 1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1
S- block s1 s2 • Alkali metals all end in s1 • Alkaline earth metals all end in s2 • Should include He but helium has the properties of the noble gases. - its outer shell is filled with the maximum number of electrons allowed for the first shell (2)
