GENERAL CHEMISTRY I CHEM-1030 INSTRUCTOR S LECTURE NOTES CHANG, CHEMISTRY CHAPTER 8

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1 GENERAL CHEMISTRY I CHEM-1030 INSTRUCTOR S LECTURE NOTES CHANG, CHEMISTRY CHAPTER 8 Periodic Relationships Among the Elements Once an intriguing mystery. Similarities now know to be due to similar outer electron configurations. Early work by Dmitri Mendeleev: When the known elements were arranged in order of atomic mass, certain properties of the elements recurred at regular intervals (periods). Mendeleev s prediction of the existence of Ga (eka-aluminum) Mendeleev knew nothing of noble gases. Problems with out of sequence atomic masses of Ar & K and Te & I. Discrepancy solved by Henry Mosely (1913) discovery of Atomic Number Now could arrange elements unambiguously according to position Could predict missing elements with certainty. Recurring properties in families no longer a mystery Elements have similar properties when their outermost (valence) electron configurations are similar. Names of various periodic table regions: Metals Left and bottom of table Nonmetals Upper right of table Metalloids Along zigzag line joining metals and nonmetals (open to interpretation) Representative Elements: filling of s and p shells Transition Metals filling 3d, 4d, or 5d shells (colors, magnetism, and variable valences) Noble Gases filled 2p, 3p, 4p, 4p, 5p, 6p shells (p 6 configuration) Lanthanides filling 4f shell Actinides filling 5f shell (Group Numbers used are open to debate) Electron configurations of Anions and Cations Most representative elements form cations or anions isoelectronic with the nearest noble gas atom. All metals form cations, many nonmetals form anions. Why? Special Stability of np 6 Configurations (also 1s 2 for He) Electron Configurations of the following (using noble gas core notation) Group 1A, Group 2A, Group 3A (Al only) H F, Cl O, N, P, (C) Why are transition metal cations not isoelectronic with preceding noble gas atoms? Complexity of electron interactions in d orbitals gives unpredictable and variable cation charges. Also true in 4f and 5f orbitals of lanthanides and actinides

2 2 Periodic Variation in Physical Properties Effective Nuclear Charge Effective shielding by inner shell electrons Less effective shielding by electrons in valence orbitals Measurement of First Ionization Potential Atomic radius is a periodic property Different ways of measuring atomic radius (Atoms in space can be thought of to have infinite size.) Atomic Radius Slide Atomic radius decreases across a period from left to right because: Valence electrons do not shield one another as effectively as core electrons shield valence electrons. Atomic radius increases going down a family because of increasing orbital size and # of electrons. Higher energy orbitals are larger than lower orbitals. Ionic Radius Influenced by # of occupied orbitals (total # of electrons) and especially by ionic charge. Remember mutual repulsion of electrons Three useful categories for comparing ionic sizes: Decreasing size of isoelectronic ions. Cations generally much smaller than anions N -3, O -2, F -1, Ne, and Na +1. Ionic radii increase down a family because of increasing # of electrons. Cations of higher positive charge smaller than those of lesser positive charge Fe +2 and Fe +3 Ionization Energy: For the general reaction: X(g) + energy X +1 (g) + e (Why is the gaseous state specified?) Ionization of an atom is always endothermic (ΔH is positive) (Why?) Subsequent ionization energies are always greater than the first one. (Why?) The low first ionization energies of the Group IA atoms explains their high chemical reactivity Both the first and second ionization energies of group IIA metals are relatively low. The first ionization energies of the elements increase from left to right as metallic character decreases. First ionization energies of elements increase from top to bottom as distance from nucleus increases. Note the extremely high first ionization energies of the noble gases, especially He. (Consequences for their chemical unreactivity?) Electron Affinity: Electron Affinity is the negative of the first ionization energy. Applicable primarily to the affinity of nonmetal atoms for electrons. The Electron Affinities of Group IA atoms are very low.

3 3 Adding an electron to a neutral Na atom for instance, puts a second electron into the 3s orbital already shielded by the filled 12, 2s and 2p subshells. In Group VII A, Electron Affinity decreases from top to bottom as distance form nucleus increases. Trends in Chemical Properties of Representative Elements Hydrogen (1s 1 ) does not belong with Group IA or any other elements. H is unique because if its small size, high ionization energy and lack of underlying core electrons. 1s 1 forms covalent bonds (sharing of electrons) with most nonmetal atoms, in H2O 1s 1 1s 0 to form common H +1 ion (acids) 1s 1 1s 2 to form uncommon H -1 hydride ion (rare, unstable NaH, CaH2, NaAlH4 Group IA Alkali Metals (ns 1 ) : Low first ionization energy leads to explosive reactivity, especially with water. 2 Na(s) + 2 H2O(l) 2 NaOH(aq) + H2(g) All metals are soft, reactive Large 1s 1 cloud (large atom) leads to low density, below that of water for Li, Na, K Na, K and Li in human physiology Group IIA Alkaline Earth Metals (ns 2 ) : Be, Mg, Ca, Sr, Ba, Ra always lose two electrons, never one. Less reactive than Group IA metals. Ca(s) + H2O(l) Ca(OH)2(s) + H2(g) slowly Be is atypical (small size) and very toxic. Physiological similarity of Mg, Ca, Sr, Ba and Ra Chemical toxicity of Ba versus dangers from radioactive isotopes of Sr and Ra. Group IIIA Elements (ns 2 np 1 ): Boron considered a nonmetal (or metalloid). Al is a metal (somewhat atypical) and always forms + 3 cation, never +1. Gallium, indium, thallium form +1 and +3 cations (toxicity of thallium) Now trend changes to more nonmetallic behavior from left to right. First ionization energy increases making cation formation more difficult. Electron affinity increases, making sharing of electrons or anion formation more likely. Group IVA, Carbon Group. (ns 2 np 2 ) : Carbon is a nonmetal. Carbon will form neither anions or cations. Si and Ge are metalloids. Tin and Lead are metals, forming the +2 and +4 cations (note 5s 2 5p 2 and 6s 2 6p 2 configurations) Why is lead (II) so toxic?

4 4 Group VA, Nitrogen Group. (ns 2 np 3 ) : N is a nonmetal, a diatomic gas. N forms mostly covalent (molecular) compounds. Six oxides of N are possible: N2O, NO, NO2, N2O4, N2O3, N2O5. Sometimes lumped as NOx Importance of N oxides in medicine, pollution studies. The nitride ion N -3 forms rarely and is unstable to water. Nitride usually forms only with Na, Mg, etc., as Na3N and Mg3N2. Important compound is nitric acid, HNO3. P is a nonmetal, existing as "white" P4 and red or black P. P -3 phosphide ion very rare and unstable to water. H3PO4 is very important in industry and in biochemistry. Also As, Sb and Bi (a metal). Group VIA, Oxygen Group. (ns 2 np 4 ) Oxygen, sulfur and selenium are nonmetals (or Se considered a metalloid). Te and Po are metals. Oxygen, O2 is a diatomic gas. Oxygen forms many covalent compounds and also many ionic oxides (O -2 ). Elemental sulfur occurs mostly as S8. S forms many ionic sulfides (S -2 ) and covalent compounds. Important compounds are SO2, SO3, H2S and H2SO4. Group VIIA, Halogens, (ns 2 np 5 ) Reactivities and physical properties: Never found free in nature, but as halides (salts). All halogens readily form X -1 (halide ion) isoelectronic with adjacent noble gas. F2; Reactivity with glass, water and asbestos Cl2; Used to disinfect drinking water and sewage Br2; Used to disinfect swimming pools I2; Used as an antiseptic. F -1 : Fluoridated water and toothpaste, rat poison. HF weak acid but eats through glass alt C -1 : HCl (strong acid) in industry and stomach acid. Cl -1 in NaCl and bloodstream. Br -1 in strong (rare)acid HBr. Old bromide sedatives. I -1 in iodized salt to prevent goiter. A few interhalogen compounds. Many other uses and occurrences. Group VIIIA, Noble Gases, (ns 2 np 6 ) Once called inert, now called noble gases. He (1s 2 ), Ne, Ar, Kr, Xe, Rn. Special stability of closed 1s 2 shell and np 6 shells. First ionization energy of Xe is low enough that a few Xe compounds form with F and O.

5 5 Explosiveness of XeF4 (anecdote) Use of He in balloons, welding. Ne in neon signs. Ar to protect tungsten filament in light bulbs. Xe in white strobe lights. Hazards of breathing (radioactive) radon. Comparison of Oxides Across a Period Oxygen will combine with most elements Oxides of most elements can be classed as acidic or basic. Acidic oxides will produce H +1 in aqueous solution. Basic oxides will produce OH -1 in aqueous solution. Interesting trend in oxide behavior on going across a period form left to right. Na2O + H2O 2 NaOH NaOH is a strong, soluble base (lye, caustic soda) MgO + H2O Mg(OH)2 Mg(OH)2 is not very soluble but a strong base (milk of magnesia) Al2O3 + H2O Al(OH)3 Al(OH)3 is an insoluble base, very weak with some amphoteric properties. SiO2 + 2 H2O Si(OH)4 But SiO2 shows some acidic behavior by reacting with concentrated base. (Hence, do not store concentrated bases in Pyrex glass, or any glass.) P4O10 + H2O 4 H3PO4 a borderline weak/strong acid SO3 + H2O H2SO4, the most-used strong acid. Cl2O7 + 2H2O 2 HClO4, the strongest acid.