CHEM 120CHAPTER 14 Main Group Chemistry: CHEM 120 CHAPTER 14 Main Group Chemistry Dr. Floyd Beckford Lyon College


Slide2: Recall that the main group elements are those in group 1A – 8A Group 1A, 2A :- s-block Group 3A – 8A :- p-block The elements may be subdivided into metals, nonmetals and the semimetals The semimetals are B, Si, Ge, As, Te, At and Sb


Slide3: Periodic Trends Left right: Zeff increases because each additional electron is not completely shielded Electrons are attracted more strongly to the nucleus :- ionization energy generally increases :- atomic radius decreases :- electronegativity increases REVIEW


Slide4: Metallic character decreases left to right From top to bottom in a group, additional shells are being filled :- atomic radius increases :- metallic character increases :- electronegativity decreases :- ionization energy decreases Metallic character increases down a group


Slide5: HYDROGEN Elemental hydrogen is a colorless, odorless, tasteless diatomic gas, H2 It has the lowest atomic weight and density of any known substance Has a very low melting and boiling point but a high dissociation energy D = 436 kJ mol-1


Slide6: There are three isotopes of hydrogen: Protium: 11H, ordinary hydrogen Deuterium: 21H or D, heavy hydrogen Tritium: 31H or T Nearly all the naturally occurring hydrogen is ordinary hydrogen The three isotopes have similar chemical properties; similar electron configuration, 1s1


Slide7: Their chemistries however show the isotope effect – the quantitative differences which results from their different masses - e.g. D2O has a higher melting and boiling point than H2O Reacts like an alkali metal but can also accept an electron to react like the halogens So it bonds in both ionic and covalent manners


Slide8: Reactivity Hydrogen reacts with both metals and nonmetals to form hydrides Hydrogen is 1s1. It can form 1. Ionic hydride, H-, by gaining an electron 2. Molecular hydride, by sharing electrons ( to form covalent bonds) Ionic hydrides form only when hydrogen reacts with an active metal


Slide9: 2M(l) + H2(g)  2[M+,H-](s) High T, press. M = Li, Na, K, Rb, Cs 2Li(l) + H2(g)  2LiH(s) (m. pt) 680 °C M(l) + H2(g)  [M2+,2H-](s) M = Ca, Sr, Ba Ionic hydrides are salt-like, crystalline, high- melting point solids All basic – form alkaline solutions in water


Slide10: COVALENT HYDRIDES Hydrogen reacts with nonmetals to form covalent or molecular hydrides H2(g) + X2(g)  2HX(g) hydrogen halides X = F, Cl, Br, I The covalent hydrides are usually gases at ordinary temperature


Slide11: THE SECOND-ROW ELEMENTS Within a group these elements have properties which differ from the heavier elements Small size and high electronegativity dictates their chemistry :- BeO is amphoteric; the others are basic :- B forms mainly covalent compounds; the others form primarily ionic compounds


Slide12: 2nd-row elements lack d-orbitals – they form a maximum of 4 bonds e.g. NH3 vs. PCl5 Their small size  form multiple bonds


Slide13: GROUP 1: The Alkali Metals


Slide14: Notice Low physical constants Very reactive due to their valence electron configuration, 1s1 :- powerful reducing agents :- low heat of atomization :- low ionization energy; high charge density :- high lattice energy


Slide15: Lithium behaves somewhat differently than the others It alone reacts with nitrogen (to form Li3N) Its small size and high charge density allows it to polarize nearby molecules - allows a large degree of covalency in its bonds So its compounds are less ionic than the others - they are less soluble in water and more soluble in polar organic solvents


Slide16: GROUP 2: The Alkaline Earth Metals


Slide17: Quite similar to the alkali metals but less reactive :- harder and have higher physical constants :- strong reducing agents :- salts are less soluble in water - higher charge densities - lattice energies Beryllium behaves differently


Slide18: All Be compounds exhibit covalent bonding In most of its compounds beryllium is electron deficient


Slide19: GROUP 3A The elements: B, Al, Ga, In, Tl; all metals except boron :- valence electron configuration, ns2 np1 – so the +3 oxidation state is expected to be most stable Gallium:- unusually low melting point, 298 °C :- long liquid range; 30 – 2403 °C :- used in transistors and high T thermometers


Slide20: The lower oxidation states become more important down the group :- ionic Tl exist primarily in the +1 oxidation state Related to the inert pair effect – the fact that the heavier elements lose their np electrons but not their ns electrons Lower oxidation state oxides are more basic


Slide21: BORON Boron shows nonmetallic properties – all boron compounds are covalent Compounds Halides: the boron halides are highly reactive They are volatile, covalent compounds 2B + 3X2  2BX3 X = F, Cl, Br, I


Slide22: BF3 and BCl3 (gases); BBr3 (liquid); BI3 (low melting point solid) BX3 acts as Lewis acids Hydrides: boron compounds of hydrogen are called boranes Molecular, volatile compounds; BnHm Simplest borane – diborane, B2H6


Slide23: Boranes have unusual structures and bonding Notice that there are two types of bonds in boranes :- the regular two-center 2e- bond :- a three-center 2e- bond


Slide24: GROUP 4A Carbon, silicon, germanium, tin and lead Silicon and germanium are important for use in electronics Carbon is nonmetal; Si and Ge are metalloids; tin and lead are metals Have valence electronic configuration, ns2 np2 :- most common oxidation state, +4


Slide25: CARBON Carbon exist in 3 allotropic forms: diamond, graphite, and fullerene Diamond has a covalent network structure – the 3D structure makes diamond very hard :- has a very high melting point Graphite has a 2D sheet-like structure Unlike diamond the carbon is sp2 hybridized


Slide26: Remaining p orbital is used to form a -bond -bonds are delocalized – graphite is an electrical conductor in the plane The sheets are parallel to each other – held together by London forces :- graphite is soft Fullerenes: contains C60 molecules – shape like soccer balls


Slide27: Compounds Catenation – the ability to bond to itself - is a unique feature of carbon Forms organic compounds Oxides: Two major oxides; carbon monoxide, CO and carbon dioxide, CO2 2C(s) + O2(g)  2CO(g) 2CO(g) + O2(g)  2CO2(g)


Slide28: CO form strong bonds to hemoglobin – which makes it very toxic CO2 is an odorless, colorless, nontoxic gas It is implicated as one of the greenhouse gases which may cause global warming It is a by-product of carbohydrate metabolism


Slide29: SILICON A hard semiconducting element A diamond-like structure and a fairly high melting point Found in nature mainly as silica, SiO2 SiO2(l) + 2C(s)  Si(l) + 2CO(g) Si for semiconductors must be ultra-pure Si(s) + 2Cl2(g)  SiCl4(l) SiCl4(g) + 2H2(g)  Si(s) + 4HCl(g)


Slide30: Unlike carbon, silicon rarely catenates As is obvious from the previous slide, silicon halides are reactive (more so than those of carbon) Silicon forms a variety of compounds with oxygen :- silicates , (SiO)n :- silicones


Slide31: GROUP 5A Nitrogen, phosphorus, arsenic, antimony and bismuth Valence electron configuration, ns2 np3 :- exhibit maximum oxidation state +5 :- minimum oxidation state –3 N and P show all oxidation states between +5 and -3


Slide32: NITROGEN Elemental nitrogen, N2 is a colorless gas – makes up 78 % of the earth’s atmosphere Obtained form the fractional distillation of air Nitrogen unreactive because of the strong triple bond – D = 945 kJ Compounds Ammonia, NH3: produced by the Haber process


Slide33: Ammonia is the starting material for many other nitrogen compounds - fertilizer Ammonia – Lewis base (lone pair) :- soluble in water to form basic solutions


Slide34: Nitrogen also forms hydrazine, N2H4 2NH3(aq) + OCl- (aq)  N2H4(aq) + H2O(l) + Cl-(aq) Oxides N2 forms a large number of oxides The major three: N2O, nitrous oxide; NO, nitric oxide and NO2, nitrogen dioxide


Slide35: Nitric oxide, NO: a very important biological species 3Cu(s) + 2NO3-(aq) + 8H+(aq)  3Cu2+(aq) + 2NO(g) + 4H2O(l) NO is also useful in killing microorganisms NO2 is a toxic red-brown gas Has an odd # of electrons – so paramagnetic


Slide36: NO2 reacts with water to form acids: 2NO2(g) + H2O(l)  HNO2(aq) + HNO3(aq) HNO3, nitric acid is made commercially by the Ostwald Process 2NO(g) + O2(g)  2NO2(g) 2NO2(g) + H2O(l)  2HNO3(aq) + NO(g)


Slide37: PHOSPHORUS Found in nature mostly as calcium phosphate, Ca3(PO4)2 Phosphate groups: an important component of the nuclei acids DNA and RNA Phosphorus exist in two common allotropic forms: white and red phosphorus White phosphorus consists of discrete P4 units


Slide38: White phosphorus :- molecular substance with a low melting point and is nonpolar :- very reactive Red phosphorus has a polymeric structure :- the more stable form :- has a higher melting point


Slide39: Compounds Forms compounds in oxidation states +5 to -3 Hydrides: forms phosphine, PH3, an extremely toxic gas Halides: reacts with all halogens to form P(III) or P(V) halides Limited X2: P4 + 6X2  4PX3 Excess X2: P4 + 10X2  4PX5


Slide40: The halides react with water to form acids PCl3(l) + 3H2O(l)  H3PO3(aq) + 3HCl(aq) PCl5(l) + 4H2O(l)  H3PO4(aq) + 5HCl(aq) Oxides: Acidic Limited O2: P4(s) + 3O2(g)  P4O6(s) Excess O2: P4(s) + 5O2(g)  P4O10(s) P4O6(s) + 6H2O(l)  H3PO3(aq) P4O10(s) + 6H2O(l)  H3PO4(aq)


Slide41: OXYGEN Properties Oxygen is paramagnetic in all three phases Has electron configuration 1s2 2s2 2p4 It is obvious that oxygen can form an octet by accepting or sharing 2e- This is how oxygen reacts With active metals, oxides are formed 4Li(s) + O2(g)  2Li2O(s)


Slide42: With nonmetals, covalent oxides are formed C(s) + O2(g)  CO2(g) Oxygen reacts directly with all elements in the periodic table (except the noble gases) Some of these reactions are very slow which is sometimes critical e.g. the rusting of iron


Slide43: Oxides Oxygen compounds are generally referred to as oxides With metals the compounds may be (a) oxides; O2- - oxidation number (–2) (b) peroxides; O22- - oxidation number (–1) (c) superoxides; O2- - oxidation number (- ½ ) Oxides may be acidic or basic


Slide44: Basic oxides are formed with the metallic elements on the left of the periodic table – they Dissolve in water to form bases K2O(s) + H2O(l)  2KOH(aq) Acidic oxides are the covalent oxides formed from the reaction with nonmetals - they dissolve in water to form acids


Slide45: SULFUR In nature sulfur exists in the elemental form as well as in minerals; FeS, pyrite and PbS, galena Sulfur exists in many allotropic forms notably rhombic sulfur – contains S8 rings In monoclinic sulfur the S8 rings are packed differently


Slide46: Compounds Oxides: Sulfur dioxide, SO2 and sulfur trioxide, SO3 are the two most important oxides SO2 is toxic, colorless gas S(s) + O2(g)  SO2 The combustion of sulfur-containing petroleum fuels is a major cause of acid rain


Slide47: H2SO4, sulfuric acid, is made by the Contact Process S(s) + O2(g)  SO2(g) H2SO4 is a strong acid :- also a good oxidizing agent – strength depends on the conditions


Slide48: THE HALOGENS Electron configuration, ns2 np5 Achieve an octet by gaining or sharing an electron Form an important set of oxoacids :- HXO, HXO2, HXO3, HXO4 :- oxidation states of X; +1, +3, +5, +7