• valence orbitals –outer shell orbitals

• 1869 Dmitri Medeleev in Russian, Lothar Meyer in Germany
• similar chemical and physical properties are periodic when elements arranged with increasing atomic weight
• Henry Moseley (1887-1915)
• Developed concept of atomic numbers
• Determined frequencies of X-rays emitted from elements after being bombarded with high-energy electrons
• Frequency increased with increasing atomic mass

7.2.1 Electron Shells in Atoms

• radial electron density – probability of finding the electron at a particular distance from the nucleus
• as nuclear charge increases, 1s shells pulled closer to nucleus

• atomic radius – estimation of radius of atoms
• estimate by assuming that atoms are spheres that touch each other
• atomic radius increases down a group, decrease down row
• atomic radius affected by principal quantum number and effective nuclear charge
• increase principal quantum number, increases size of orbital
• increase effective nuclear charge, reduces size or orbital

• ionization energy – minimum energy required to remove an electron from the ground state
• first ionization energy – energy needed to remove the first electron
• second ionization energy – energy needed to remove second electron
• the greater ionization energy, harder it is to remove electrons
• I_1< I_{2 <} I₃
• Positive nuclear charge remains same, number electrons decreases à effective nuclear charge increases
• Greater effective nuclear charge, greater energy required to remove electron
• Sharp increase in ionization energy when inner-shell electron removed
• Only outer most electrons involved in sharing and transfer of electrons in bonding and reactions

• I₁ increases with increasing atomic number
• Alkali metals lowest ionization energy, noble gases greatest ionization energy
• Ionization energy decreases down each group
• Ionization energy of transition metals increase slowly across row
• Ionization energy depends of effective nuclear charge and average distance of the electrons from the nucleus
• Increase attraction of electrons to nucleus – increase effective nuclear charge or decrease distance from nucleus
• Across period – increase in effective nuclear charge and decrease in atomic radius, increase in ionization energy
• Down group – atomic radius increases, effective nuclear charge remains constant, ionization energy decreases

• electron affinity – energy change when electron added to atom
• greater attraction between atom and electron, the more negative electron affinity
• across row – becomes more negative
• electron affinity of noble gases is positive
• down group – does not change much

• low ionization energy
• transition metal – form more than one positive ion
• compounds of metals and nonmetals are ionic
• metal oxides dissolve in water to form metal hydroxides
• metal oxides react with acids to form salts and water

• seven nonmetals exist under ordinary conditions as diatomic molecules
• H₂, O₂, N₂, F₂, Cl₂, Br_2(l), I_2(s)
• Nonmetals gain electrons
• Compounds composed of all nonmetals are molecular substances
• Nonmetal oxides – acidic oxides

• properties of both metals and nonmetals

• soft metallic solids
• characteristics of metals
• low densities and melting points
• down group – increasing atomic radius, decreasing first ionization energy
• lowest first ionization energy
• most active metals
• combine directly with nonmetals
• react with hydrogen to form hydrides, with sulfur to form sulfides, with chlorine to form chlorides
• hydride ion – H^-1
• react vigorously with water – produce hydrogen gas and solutions of alkali metal hydroxides

• solids, harder, more dense, higher melting point that alkali metals
• tend to lose 2 electrons to form 2+ ions

• ionization energy higher than active metals
• reacts with nonmetals to form molecular compounds
• reacts with active metals for form hydrides

• oxygen has two forms – O₂ and O₃
• allotropes – different forms of the same element in the same state
• 21 percent dry air is O₂
• oxygen attracts electrons from other elements
• oxygen with other metals from oxides
• peroxide O₂^2- , superoxide O₂^-
• compounds with these ions react with themselves to produce an oxide and O₂
• sulfur – several allotropic forms
• most common and most stable – S₈, usually written as S(s)
• sulfur and its compounds burned in oxygen to form sulfur dioxide

• salt formers
• halogens are typical nonmetals
• melting and boiling points increase with increasing atomic number
• exist as diatomic molecules
• have highly negative electron affinities
• gain electrons to form halide ions
• Fluorine removes electrons from almost any substance

• nonmetals, gas at room temperature
• monoatomic
• completely filled s and p subshells
• 1962 Neil Bartlett
• first noble-gas compound by reacting Xe with F₂(g) to form XeF₂ and XeF₄ and XeF₆
• KrF₂, only single stable compound