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Chapters 18 – The Periodic Table See Fig. 18.1 Oxygen 0.96% Neon 0.67% Carbon 0.27% Others 1.14% Helium 36.56% Hydrogen 60.4% Universe Relative Abundances Calcium 1% Sulfur 2% Aluminum 1% Nickel 2% Others 1% Iron 35% Magnesium 13% Silicon 15% Oxygen 30% Relative Abundances Whole Earth Sodium 2.83% Calcium 3.63% Iron Potassium Others 2.59% 0.83% Magnesium Titanium 2.09% 0.44% Hydrogen 0.14% 5.00% Aluminum 8.13% Oxygen 46.6% Silicon 27.72% Relative Abundances Earth's Crust Phosphorous 1.00% Calcium 1.40% Nitrogen 3.00% Hydrogen 10.00% Chlorine Potassium Sulfur Sodium 0.14% 0.34% 0.26% Iron 0.14% Zinc Magnesium 0.004% 0.003% 0.50% Other trace elements 0.21% Carbon 18.00% Oxygen 65.00% Human Body “Representative” Elements • The properties of the elements are determined by their electron configurations. • Electron configurations (and orbital properties) determine periodic trends: – Electron affinity – Ionization energy – Atomic size – Metal vs. non-metal characteristics – Bonding abilities • Each group (vertical column) shares common properties with each other and distinct differences from other groups based on electron configurations. Periodic Trends in Main Groups Periodic Trends in Main Groups: Atomic Radii Fig . 18.2: Atomic radii increase down the columns Periodic Trends in Main Groups: Ionization Energy Ionization energy decreases down the columns (electrons are easier to remove) Group 1A Elements H, Li, Na, K, Rb, Cs, Fr Alkali metals Common Features: All but H are very active metals ns1 electron configuration ( n = 1 to 7) Metals readily oxidized to X+ state 2 M+ 2 H2O 2M+ + 2OH- + H2 Vigorous reaction with water Demo: Reactivity of Li, Na and K with Water What reaction is happening here? 2Li + 2H2O 2LiOH + H2 2Na + 2H2O 2NaOH + H2 2K + 2H2O 2KOH + H2 Reaction of Li Metal with Water Reaction Equation: Li(s) + H2O(l) → Li+(aq) + OH−(aq) + ½ H2(g) The reaction of lithium with water is the slowest among the alkali metals, even though it is the most thermodynamically favorable. Ea is controlled by the lattice energy (highest for Li); ΔG o by the combination of lattice, ionization, and hydration enthalpies and entropies (lowest for Li, largely due to the favorable hydration enthalpy of the small Li+ ion). Periodic Trends in Main Groups: Group 1A Survey of Main Groups: Group 1A Hydrogen – The Simplest Element • Most abundant element in the universe (but not on earth, why?) • Colorless, odorless, highly flammable gas • Low molar mass and non-polar…low b.p. and m.p. • Sources: − Combustion of methane − Byproduct of gasoline production − Electrolysis of water • Major uses: − Production of ammonia (Haber process) − Hydrogenation of vegetable oils to produce solid shortenings • Typical nonmetal: covalent with other NMs, salts with very active metals • Hydrides: binary compounds containing H The Chemistry of Molecular Hydrogen Hydrogenation H Unsaturated fat O O H O Fatty acid O Fatty acid H2 (Ni catalyst) H H O H H O O Fatty acid O Fatty acid Saturated fat Hydrides Hydrides: binary compounds containing H Ionic hydrides: hydrogen (as H-) combines with G1A or G2A metal; contains hydride ion and metal cation; H- is strong reducing agent H + e 2Li + H2 H (Hrxn = - 73 kJ/mol) 2LiH Covalent hydrides: hydrogen combines with other nonmetals…we already know about many of these (HCl, CH4, NH3, H2O) Metallic (interstitial) hydrides: transition metal catalysts treated with H2 (g); H2 molecules dissociate at the metal surface and H atoms occupy holes in the crystal structure (potential use as a portable fuel) Reaction of Lithium Hydride with Water Reactions of ionic hydrides with water are violent. LiH(s) + H2O(l) → Li+(aq) + OH–(aq) + H2(g) The H – ion in LiH (and other column-I and -II hydrides) is a strong base, reacting with any molecule containing acidic H to form H2(g). These reactions can also be regarded as oxidationreduction. However, the reaction is much milder than most redox reactions, suggesting that it is best considered as acid-base from a practical standpoint. Small amounts of undissolved material are probably Li2CO3(s), formed by reaction of LiH(s) with CO2(g) and H2O(g) in the air. Group 2A Elements Be, Mg, Ca, Sr, Ba, Ra Alkaline earth metals Common Features: All are reactive metals (Be and Mg less reactive with water at 25°C) ns2 electron configuration (n = 2 to 7) Readily oxidized to X2+ state M + 2H2O M2+ + 2OH- + H2 Form “insoluble” salts (carbonates, phosphates, sulfates, fluorides) Some metal oxides have limited solubility in water (MgO, CaO, Al2O3) Periodic Trends in Main Groups: Group 2A Table 18.7 Be2+ Mg2+ Ca2+ Sr2+ Ba2+ Ra2+ Biological Properties Be Highly toxic Mg Ca Biogenic…metabolism and muscles Biogenic…bones and teeth Sr Ba Highly toxic Highly toxic; sulfate insolubility – used for x-ray enhancement Highly toxic; radioactive, t1/2 = 1,622 years Ra Chlorophyll – The most important Mg compound CO2 + H2O catalyst (CH2O) + O2 Sugars, cellulose Chromophore Respiration Group 2A Elements: Water Hardness Ca2+(aq) + CO32¯(aq) CaCO3(s) Precipitates as calc at elevated temperatures Ca2+, Mg2+ Ion Removal from Water by Ion Exchange Mg + O2 • Magnesium is a highly flammable metal • Once ignited, it is difficult to extinguish – it can burn in nitrogen, carbon dioxide and water • Magnesium powder is used in fireworks and marine flares • Mg2+ is the second most common cation in seawater Periodic Table of Elements Group 3A Elements B Non-metallic Al, Ga, In, Tl Metals Common Features: ns2p electron configuration (n = 2 to 6) Cannot achieve valence electron octets in neutral molecules B and Al form strong bonds to F and O H F B F O 3 H2 O F B O H O H Boron trifluoride Boric acid (+ 3 HF) Periodic Trends in Main Groups: Group 3A Table 18.8 Increase in metallic character going down the group. How Can Group 3A Elements Achieve Electron Octets? Bridging bonds H H B H H B + H H H H H B H B H H Hrxn = -165 kJ/mol Covalent boron hydrides: “boranes” (e- deficient, highly reactive) Fig. 18.7 How Can Group 3A Elements Achieve Electron Octets? : : F: : F B F F F B F F F Reactions with anion electron donors Tetrafluoroborate anion : : : O H B O H H B H H H Reactions with neutral electron donors Ether-BH3 complex How Can Group 3A Elements Achieve Electron Octets? Cl 2 Cl Al Cl Cl Al Al Cl Cl Cl Cl Cl nH2O O HO Al Aluminum hydroxide polymers O O O Al O Al OH Aluminum chloride dimer, Al2Cl6 Periodic Trends in Main Groups: Group 3A Aluminum Most abundant metal in Earth’s crust Metallic properties, but covalent bonds to NMs Covalency: amphoteric Al2O3, acidic Al(H2O)63+ Amphoteric Nature of Aluminum Hydroxide Substances like Al(OH)3(H2O)3(s) that can act as either acids or bases are called amphoteric. Aluminum hydroxide is formed by the following neutralization reaction: Al(H2O)63+(aq) + 3 OH–(aq) → Al(OH)3(H2O)3(s) + 3 H2O(l) Cl–, NH4+ spectator ions. The subsequent neutralization reactions are as follows: Al(OH)3(H2O)3(s) + 3 OH−(aq) → Al(OH)63−(aq) + 3 H2O(l) Na+ spectator ion. Al(OH)3(H2O)3(s) + 3 H3O+(aq) → Al(H2O)63+(aq) + 3 H2O(l) Cl− spectator ion. Thermite Reaction Reaction Equation: Fe2O3(s) + 2 Al(s) → 2 Fe(s) + Al2O3(s) ΔHo = –849 kJ/mol ΔSo = –37.48 J/mol-K ΔGo = –838 kJ/mol Uses: welding, purifying an ore (U in the Manhattan Project) In this reaction, an unfavorable ΔS is offset by a very large negative ΔH. The heat is sufficient to raise the temperature of the products past the melting point of Fe (1530 °C). Can go horribly wrong…the heat has to go somewhere and a sand pit is the best “Mythbusters”: thermite and ice…ice chunks flew 150 feet, or half a football field Group 4A Elements C Non-metallic Si, Ge Metalloid Sn, Pb Metals Common Features: ns2p2 electron configuration (n = 2 to 6) Can achieve valence electron octets in neutral molecules All form covalent bonds with other nonmetals C: s and p bonds, as in CH4, O=C=O, H-CN, etc. Si: prefers s bonds, as in SiCl4 , SiH4, or network solids (SiO2) Ge: prefers s bonds, as in GeCl4 Sn: forms covalent s bonds, as in SnH4, but also ionic bonds, as in SnCl2. Pb: forms covalent s bonds, as in Pb(C2H5)4 ,(tetraethyl lead), but also ionic bonds, as in PbSO4 (paint filler). Toxic!! Group 4A Elements C: graphite, diamond, fullerene second most abundant element in human body (more to come in Ch 21) Si: semimetal (semiconductor for electronics) second most abundant element in Earth’s crust Ge: semiconductor for electronics (transistors) Sn: tin foil, bronze, solder, pewter, protective coating for steel Pb: anti-knocking agent in fuels (problematic for catalytic converters), lead crystal (extra sparkly, but toxic!), car batteries Periodic Trends in Main Groups: Group 4A Note: Si-Si compounds do exist, but are more reactive than the corresponding C-C bonds. This is why we see more Si-O bonds than Si-Si bonds. Reactions of Main Group Elements: Group 4A Group 5A Elements N, P Non-metallic As Metalloid Sb, Bi Metals Common Features: ns2p3 electron configuration (n = 2 to 6) Electronegativity decreases down the group Metallic character increases down the group N and P: form 3- anions in salts with active metals Bi and Sb: form mostly 3+ rather than 5+ ions Group 5A elements can form molecules involving 3, 5, or 6 covalent bonds to the Group 5A atom MX3: NH3, PH3, NF3, AsCl3, AsF3 contain a lone pair of electrons (Lewis base, pyramidal shape) MX5: PF5, PCl5, SbCl5, AsF5 the small atomic size of nitrogen makes it difficult to accommodate 5 covalent bonds; prediction of VSEPR model – trigonal bipyramidal MX6: MX5 + X- MX6e.g. PF5 + F- PF6AsF5 + F- AsF6- (phosphorous hexafluoride anion) (arsenic hexafluoride anion) The Chemistry of Nitrogen • Elemental nitrogen exists as N2 molecules N N, 2 p bonds 941 kJ/mol bond energy • Other elements in Group 5A do not form p bonds, but exist as aggregates with single bonds (e.g. P4, As4, etc.) • The strong N N bond makes N2 very unreactive toward most other substances (inert atmosphere) • Many N-containing compounds decompose exothermically to N2, releasing large amounts of energy Examples: N2O(g) N2(g) + ½ O2(g) ∆H = -82 kJ/mol N2H4(g) N2(g) + 2 H2(g) (hydrazine) ∆H = -95 kJ/mol 4 C3H5N3O9(l) 6 N2(g) + 12 CO2(g) (Nitroglycerine) + 10 H2O(g) +O2(g) + Energy! 4 moles of liquid 29 moles of gas with large bond energies! 2 C7H5N3O6(s) 12CO(g) + 5H2(g) + 3N2(g) + 2C(s) + Energy! (trinitrotoluene, TNT) 2 moles of solid 20 moles of gas + energy! Reacts spontaneously with strong oxidants (H2O2, NO2) – used as self-igniting fuel in torpedoes, rockets, space shuttle Some Important Nitrogen Reactions 1. Nitrogen is “fixed” industrially in the Haber process under high temperature and heat, and with a catalyst: N2 (g) + 3 H2 (g) 2 NH3 (g) 2. Further industrial reactions convert NH3 to NO, NO2, and HNO3, (e.g. Ostwald process): 4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g) nitric oxide 3. NO can also be made directly from N2 at high temperatures (like in an internal combustion engine): N2(g) + O2(g) 2 NO(g) A Few Important Compounds of Nitrogen 1. Ammonia, NH3. First substance formed when atmospheric N2 is used to make N-containing compounds. Annual multimillion-ton production for use in fertilizers, explosives, rayon, and polymers such as nylon, urea-formaldehyde resins, and acrylics. 2. Hydrazine, N2H4. Used in rockets as a propellant, and as a starting point for anti-tuberculin drugs. 3. Nitric oxide (NO), nitrogen dioxide (NO2), and nitric acid (HNO3). Used in fertilizer manufacture, nylon production, metal etching, explosives industry, and biological signaling, among other things. 4. Amino acids, H3N+-CH(R)-COO- (R = one of 20 different organic groups). Occur in every organism, both free and linked together into proteins. Essential to growth and function of all cells. Phosphorus No double bonds P P P P White phosphorus - soluble in non-polar solvents - toxic - highly reactive to oxidation P P P P P P P P P Red phosphorus - network/polymeric - insoluble - less reactive 3 P4(s) + 10 KClO3(s) 3 P4O10(s) + 10 KCl(s) Heat of reaction = - 9,425 kJ/mol The head of "strike anywhere" matches contain an oxidizing agent such as potassium chlorate together with tetraphosphorus trisulfide (P4S3), glass and binder. The phosphorus sulfide is easily ignited, the potassium chlorate decomposes to give oxygen, which in turn causes the phosphorus sulfide to burn more vigorously. The head of safety matches are made of an oxidizing agent such as potassium chlorate, mixed with sulfur, fillers and glass powder. The side of the box contains red phosphorus, binder and powdered glass. The heat generated by friction when the match is struck causes a minute amount of red phosphorus to be converted to white phosphorus, which ignites spontaneously in air. This sets off the decomposition of potassium chlorate to give oxygen and potassium chloride. The sulfur catches fire and ignites the wood. Group 6A Elements O, S, Se Non-metallic Common Features: Te, Po Metalloid ns2p4 electron configuration (n = 2 to 6) Metallic properties increase down the column, but no G6A element behaves like a true metal Common behavior: reaction with metal to become 2- ion in ionic compound; for most metals, most common minerals are oxides or sulfides Covalent bonds with other NMs; series of covalent hydrides (H2X) All but O have d orbitals available, so more than an octet is common Te and Po can be 4+ cations, but with limited chemistry Se interest growing in recent years (medical research) Some Important Simple Oxygen Compounds 1. Water, H2O. Perhaps the single most important compound on earth! 2. Oxygen, O2. 21% of the atmospheric gas composition. Necessary for all "aerobic" life forms as a source of energy. Good oxidant thermodynamically, poor oxidant kinetically. Most of energy we need on earth: exothermic reaction of O2 with carbon-containing compounds 3. Hydrogen peroxide, H2O2. Used as an oxidizing agent, disinfectant, bleach, and in the production of peroxy compounds for polymerization. 4. Ozone, O3. Exists naturally in upper atmosphere. Absorbs UV solar radiation, protecting the Earth's surface. Created as emission from combustion engines- In the lower atmosphere, it is a deleterious oxidant. Ozonolysis is used to kill bacteria in water. 5. Superoxide, HOO. (or OO.-). Incompletely reduced form of O2. Good oxidant, bad for biological systems- superoxide reductases From Lecture #1 this quarter… (slide 13) 142 Review POLYATOMIC (COVALENT) IONS Some covalent molecules are stable with a net charge. Most common examples: Name (“… Ion”) Formula Example Related Compound Acid Ammonium Hydronium Metal Aqua Hydroxide Acetate Carbonate Nitrate Phosphate Sulfate Perchlorate NH4+ H3O + M(H2O)62,3+ OH CH3CO2 CO32 NO3 PO43 SO42 ClO4 NH4Cl HCl(aq) Cr(H2O)6Cl3 KOH CH3CO2K K2CO3 KNO3 K3PO4 K2SO4 KClO4 Name Related (“… Acid”) Oxide Hydrochloric H2O CH3CO2H H2CO3 HNO3 H3PO4 H2SO4 HClO4 Acetic Carbonic Nitric Phosphoric Sulfuric Perchloric CO2 N2O5 P4O10 SO3 Cl2O7 Periodic Table of Elements Sulfur Found in deposits of free element and widely distributed in ores (sulfides, sulfates) 60% in free element, 40% from purification of fossil fuels or when sulfur-containing fuels are burned S2 molecules only exist in gas phase at high T; stronger s bonds than p bonds, so typically found in aggregates Sulfur The S8 molecule (ring) Figure 18.22 Sulfur Chains, Sn (n up to 10,000) S Odoriferous Compounds CH3CH2CH2CH2-S-H Skunk CH2=CH-CH2-N=C=S Garlic CH3SH Gas odorizer H2 S Rotten eggs and body odors Important Reactions of the Oxygen Family (E) 1. Halides can be formed with many Group 6A elements: E(s) + X2 (g) various halides (E = S, Se, Te ; X = F, Cl) 2. The other elements in the group are oxidized by O2: E(s) + O2 (g) EO2 (g) (E = S, Se, Te, Po) SO2 is oxidized further, and the product is used in the final step of H2SO4 manufacture. 2 SO2 (g) + O2 (g) 2 SO3 (g) SO3 (g) + H2O(l) H2SO4 (aq) Important Reactions of the Oxygen Family 3. Sulfur is recovered when hydrogen sulfide is oxidized: 8 H2S(g) + 4 O2 (g) S8 (s) + 8 H2O(g) This reaction is used to obtain sulfur when natural deposits are not available. 4. The thiosulfate ion is formed when an alkali sulfite reacts with sulfur: S8 (s) + 8 Na2SO3 (aq) 8 Na2S2O3 (aq) Some Important Compounds of the Oxygen Family 1. Hydrogen sulfide, H2S. Vile toxic gas formed during anaerobic decomposition of plant and animal matter, in volcanoes, and in deep sea thermal vents. Used in the manufacture of paper. Atmospheric traces cause silver to tarnish through formation of black Ag2S. 2. Sulfur dioxide, SO2. Colorless, choking gas formed in volcanoes or whenever a S-containing compound (coal, oil, metal sulfide ores, etc.) is burned. More than 90% of SO2 produced is used to make sulfuric acid. Also used as a fumigant and preservative of fruit, syrups, and wine. As a reducing agent, removes excess Cl2 from industrial waste water, removes O2 from petroleum handling tanks, and prepares ClO2 for bleaching paper. Major atmospheric pollutant in acid rain. More Important Compounds of the Oxygen Family 3. Sulfur trioxide (SO3) and sulfuric acid (H2SO4). SO3, formed from SO2 over V2O5 catalysts, is then converted to sulfuric acid. Sulfuric acid is the cheapest strong acid and is so widely used in industry that its production level is an indicator of a nation’s economic strength. Strong dehydrating agent that removes water from any organic source. 4. Sulfur hexafluoride, SF6. Extremely inert gas used as an electrical insulator. Also used as an atmospheric tracer of air movement over great distances. O - S - Se - Te - Po Figure 18.24: Dehydration of Sucrose by H2SO4 C12H22O11 (s) + 11 H2SO4 12 C (s) + 11 H2SO4·H2O (l) Demo: Acid Rain During combustion sulfur combines with oxygen to form sulfur dioxide and sulfur trioxide: S(s) + O2(g) → SO2(g) Predominant reaction. 2 S(s) + 3 O2(g) → 2 SO3(g) Insignificant indoors, (Catalyzed by sunlight.) The sulfur dioxide and sulfur trioxide then combine with water to give acids: SO2(g) + 2 H2O(l) ↔ H2SO3(aq) SO3(g) + 2 H2O(l) → H2SO4(aq) sulfurous acid sulfuric acid The burning of fossil fuels in industry, power plants and in homes accounts for most of the SO2 emitted to the atmosphere. The SO2 is oxidized by several pathways to SO3. Both SO2 and SO3 react with rain water to form acids. The resulting acids are hazards to aquatic life and corrode buildings made of marble. In the northeastern U. S., some precipitation has a pH of 4. Periodic Table of Elements Group 7A Elements F, Cl, Br, I, At Non-metallic Common Features: ns2p5 electron configuration (n = 2 to 6) All are non-metals Properties vary smoothly down the group, except for unexpectedly low E.A. of fluorine and small bond energy of F2 Highly reactive: not found as free elements in nature, but rather as halide ions (X-) in minerals and seawater Highly electronegative…form polar covalent bonds with other NMs and ionic bonds with lower O.S. metals Compounds with metals in higher O.S. are polar covalent (TiCl4, SnCl4) Astatine: radioactive, t1/2= ~8 hours Melting Points and Boiling Points of the Halogens Larger halogens have greater polarizabilities larger dispersion forces higher temperatures for phase transitions See Table 18.18 Important Reactions of the Halogens - I 1. The halogens (X2) oxidize many metals and non-metals. The reaction with hydrogen, although not used commercially for HX production (except for high-purity HCl), is characteristic of these strong oxidizing agents. X2 (g) + H2 (g) 2 HX(g) 2. The halogens disproportionate in water: X2 (g) + H2O(l) X = Cl, Br, I HX(aq) + HXO (aq) In aqueous base, the reaction goes to completion to form hypohalites and, at higher temperatures, halates; for example: 3 Cl2 (g) + 6 OH-(aq) ClO3-(aq) + 5 Cl-(aq) + 3 H2O(l) Important Reactions of the Halogens - II 3. Molecular fluorine, F2 is produced electrolytically at moderate temperature: 2HF (as KHF2, a solution of KF in HF) H2 (g) + F2 (g) A major use of F2 is in the preparation of UF6 for nuclear fuel. 4. Glass (amorphous silica) is etched with HF: SiO2 (s) + 6 HF(g) H2SiF6 (aq) + 2 H2O(l) F - Cl - Br - I - At Some Important Compounds of the Halogens 1. Fluorspar (fluorite), CaF2. Widely distributed mineral used as a flux (chemical cleaning agent) in steel making and in the production of HF. 2. Hydrogen fluoride, HF. Colorless, extremely toxic gas used to make F2, organic fluorine compounds, and polymers. Also used in aluminum manufacture and in glass etching. 3. Hydrogen chloride, HCl. Extremely water-soluble gas that forms hydrochloric acid, which occurs naturally in stomach juice of mammals (humans produce 1.5 L of 0.1 M HCl daily!) and in volcanic gases (from reaction of H2O on sea salt). Made by reaction of NaCl and H2SO4 and as a byproduct of plastics (PVC) production. Used in the “pickling” of steel (removal of adhering oxides) and in the production of syrups, rayon, and plastics. More Important Compounds of the Halogens 4. Sodium hypochlorite, NaClO, and calcium hypochlorite, Ca(ClO)2 Oxidizing agents used to bleach wood pulp and textiles, and disinfect swimming pools, foods, and sewage (also used to disinfect the Apollo 11 on return from the moon). Household bleach is 5.25% NaClO by mass in water. 5. Ammonium perchlorate, NH4ClO4. Strong oxidizing agent. 6. Potassium iodide, KI. Most common soluble iodide. Table salt additive to prevent thyroid disease (goiter). Used in chemical analysis because it is easily oxidized to I2, which forms a colored end point. 7. Polychlorinated biphenyls, PCBs. Mixture of chlorinated organic compounds used as nonflammable insulating liquids in electrical transformers. Production discontinued due to persistence in the environment, where it becomes concentrated in fish, birds, and mammals, and causes reproductive disturbances and possibly cancer. Group 8A Elements He, Ne, Ar, Kr, Xe, Rn Non-metallic Common Features: ns2p6 electron configuration (n = 2 to 6) – FULL OCTET All are non-metals All are very non-reactive, but some compounds can be formed (XeF4, XeF2, KrF2, etc.) The Periodic Table How can we organize these? The Periodic Table The Periodic Table Can we predict what is missing? The Periodic Table The Periodic Table • Is a convenient way to represent the elements • It displays the trends in – electron configuration – metallic character – group commonalities • It is not the only way to do this… …there are alternative views Internet search: “alternative periodic tables” MayanPeriodic.com The shells are shown as concentric circles. Each row in the tabular form is shown as a ring. This spiral periodic table was presented by Professor Theodore Benfey in 1964. The elements form a two-dimensional spiral, starting from hydrogen, and folding their way around two islands, the transition metals, and lanthanides and actinides. Mendeleev flower Timothy Stowe's physicists periodic table – each grouping is for a different principal quantum number.