Presentation on theme: "1 Chapter 22 Chemistry of The NonMetals lDescriptive chemistry of the elements is consistent with the various principles discussed earlier. lWe will focus."— Presentation transcript:
1 Chapter 22 Chemistry of The NonMetals lDescriptive chemistry of the elements is consistent with the various principles discussed earlier. lWe will focus on trends in and explanations for the observed behavior of the elements.
General Concepts: Periodic Trends lMain-group elements -- the valence electrons are filling s and p orbitals lThree types of main-group elements: metal, metalloid, nonmetal lIncreasing metallic character going down a group and from right to left across a period
3 Metallic Character 1:18
4 Group Work lList the relative properties of the metals: volatility melting and boiling points density thermal conductance electrical conductance appearance as solids brittleness typical structure bonding tendency to lose or gain electrons acidity/basicity of oxides non-volatilehighhighhighhighshiny/lustrousmalleable atomic crystals metalliclosebasic
5 lList the relative properties of the non-metals: volatility melting and boiling points density thermal conductance electrical conductance appearance as solids brittleness typical structure bonding tendency to lose or gain electrons acidity/basicity of oxides volatilelowlowlowlowdull soft or brittle discrete molecules covalentgainacidic Group Work
6 Metalloids lMetalloids: physical properties more like those of metals, but chemical reactivity is more like that of nonmetals; many atomic properties are intermediate between those of metals and of nonmetals Si As
7 Period 2 Elements are Unique lCompounds formed between nonmetals tend to be molecular. lAs we move down the periodic table bonding changes. lThe third period onwards has accessible d- orbitals that can participate in bonding. lTherefore, the octet rule can be broken for the third period onwards.
8 Period 2 Elements lThe first member of a group can form bonds more readily than subsequent members. lSi is much larger than C and the 3p orbital is much larger than the 2p orbital, so the overlap between 3p orbitals to form a 3p bond is significantly poorer than for a 2p bond.
9 Period 2 Elements lSince the Si-Si bond is much weaker than the C-C bond, Si tends to form bonds. lSiO 2 is a network solid with Si-O bonds. lCO 2 is a gas with O=C=O bonds.
10 Lightest Elements are Unique lProperties of the first element in each group are usually more distinctive, while the rest of the elements in a group have similar properties lThe unusual properties of the first element in a group can be explained on the basis of its unusually small size, which arises because the valence electrons are not shielded from the nucleus and the electrons are held relatively tightly in the atom
11 Trends in Properties of the Elements
Hydrogen Isotopes of Hydrogen lThere are three isotopes of hydrogen: Protium 1 1 H, deuterium 2 1 H, and tritium 3 1 H. lDeuterium (D) is about % of naturally occurring H. lTritium (T) is radioactive with a half-life of 12.3 yr.
13 Properties of Hydrogen lHydrogen is unique. lHydrogen has a 1s 1 electron configuration so it is placed above Li in the periodic table. lHowever, H is significantly less reactive than the alkali metals. Hydrogen can gain an electron to form H -, which has a He electron configuration. Therefore, H could be placed above the halogens. Hydrogen can gain an electron to form H -, which has a He electron configuration. Therefore, H could be placed above the halogens. lHowever, the electron affinity of H is lower than any halogen.
14 Properties of Hydrogen lElemental hydrogen is a colorless, odorless gas at room temperature. lSince H 2 is nonpolar and only has two electrons, the intermolecular forces are weak (boiling point C, melting point -259 C). What kind of intermolecular forces? lThe H-H bond enthalpy is high (436 kJ/mol). Therefore, reactions with hydrogen are slow and a catalyst needs to be used.
15 Properties of Hydrogen lWhen hydrogen reacts with air, explosions result (Hindenburg exploded in 1937): 2H 2 (g) + O 2 (g) 2H 2 O(l) H = kJ
16 Preparation of Hydrogen lIn the laboratory hydrogen is usually prepared by reduction of an acid. lZn is added to an acidic solution and hydrogen bubbles are formed. lThe hydrogen bubbles out of solution and is collected in a flask. lThe collection flask is usually filled with water so the volume of hydrogen collected is the volume of water displaced.
17 Preparation of Hydrogen lWhat other metals could be used to prepare hydrogen by reaction with acid? What property would we examine? lIn larger quantities, hydrogen can be prepared by the reduction of methane in the presence of steam at 1100 C: CH 4 (g) + H 2 O(g) CO(g) + 3H 2 (g) CO(g) + H 2 O(g) CO 2 (g) + H 2 (g)
18 Uses of Hydrogen lHydrogen is used for ammonia production and to hydrogenate vegetable oils to make margarine and shortening. lHydrogen is used to manufacture methanol: CO(g) + 2H 2 (g) CH 3 OH(g) lHydrogen is being considered for direct use as a fuel in automobiles. Would we want to carry around tanks of hydrogen? How else could we store it?
19 Binary Hydrogen Compounds lThree types of binary hydrogen compounds are formed: lionic hydrides (e.g. LiH, made between metals and H); lmetallic hydrides (e.g. TiH 2, made between transition metals and H); and lmolecular hydrides (e.g. CH 4, made between nonmetals and metalloids and H). lThermal stability (measured by G f ) decreases as we go down a group and increases across a period.
20 Binary Hydrogen Compounds lMost stable is HF. lMetal hydrides, such as CaH 2, react with water to give H 2 and metal hydroxide. GfGfGfGf
Oxygen Properties of Oxygen lOxygen has two allotropes: O 2 and O 3. lO 2 is a colorless, odorless gas at room temperature. lThe electron configuration is [He]2s 2 2p 4, which means the dominant oxidation state is 2-. lThe O=O bond is strong (bond enthalpy 495 kJ/mol).
22 Preparation of Oxygen lCommercially: obtained by fractional distillation of air. (Normal boiling point of O 2 is -183 C and N C.) lLaboratory preparation of oxygen is the catalytic decomposition of KClO 3 in the presence of MnO 2 : 2KClO 3 (s) 2KCl(s) + 3O 2 (g). lAtmospheric oxygen is replenished by photosynthesis (process in plants where CO 2 is converted to O 2 in the presence of sunlight).
23 Uses of Oxygen lMost widely used as an oxidizing agent. (e.g. in the steel industry to remove impurities.) lOxygen is used in medicine. lIt is used with acetylene, C 2 H 2 for oxyacetylene welding: 2C 2 H 2 (g) + 5O 2 (g) 4CO 2 (g) + 2H 2 O(g) 1:31
24 Ozone lPale blue, poisonous gas. lOzone dissociates to form oxygen: O 3 (g) O 2 (g) + O(g) H = 107 kJ lOzone is a stronger oxidizing agent than oxygen: O 3 (g) + 2H + (aq) + 2e - O 2 (g) + H 2 O(l) E = 2.07 V O 2 (g) + 4H + (aq) + 4e - 2H 2 O(l) E = 1.23 V
25 Ozone lOzone can be made by passing an electric current through dry O 2 : 3O 2 (g) 2O 3 (g)
26 Peroxides and Superoxides lPeroxides: have an O-O bond and O in the -1 oxidation state. lHydrogen peroxide is unstable and decomposes into water and oxygen: 2H 2 O 2 (l) 2H 2 O(l) + O 2 (g), H = kJ lThe oxygen produced will kill bacteria. lPeroxides are important in biochemistry: it is produced when O 2 is metabolized.
27 Peroxides lDisproportionation occurs when an element is simultaneous oxidized and reduced: 2H + (aq) + H 2 O 2 (aq) + 2e - 2H 2 O(l) E = 1.78 V E = 1.78 V O 2 (g) + 2H + (aq) + 2e - H 2 O 2 (aq) E = 0.68 V E = 0.68 V lDisproportionation: 2H 2 O 2 (aq) 2H 2 O(l) + O 2 (g) E = 1.10 V
28 Superoxides lSuperoxides: have an O-O bond and O in an oxidation state of -½ (superoxide ion is O 2 - ). lUsually form with active metals (KO 2, RbO 2 and CsO 2 ). lPotassium superoxide reacts with water vapor from the breath to form oxygen gas. 4KO 2 + 2H 2 O 3O 2 + 4KOH
The Noble Gases Noble Gas Compounds lNoble gases are very unreactive. lAll elements have high ionization energies. lHe is the most important noble gas as liquid helium is used as a coolant. lThe heavier noble gases react more readily than the lighter ones. lThe most common compounds of noble gases are xenon fluorides. lXenon fluorides have Xe in the +2 to +8 oxidation states.
30 Noble Gas Compounds lNoble gas compounds violate the octet rule.
31 Noble Gas Compounds lIn the presence of water, xenon fluorides form oxyfluorides: XeF 6 (s) + H 2 O(l) XeOF 4 (l) + 2HF XeF 6 (s) + 3H 2 O(l) XeO 3 (aq) + 6HF lThe only other noble gas compound known is KrF 2, which decomposes at -10 C. lXenon fluorides are more stable than the oxides and oxyfluorides.
Group 7A: The Halogens lF, Cl, Br, I, At lMost common are chlorine, bromine, and iodine lFluorine has properties atypical of the group lAstatine is radioactive and exists naturally only in very small amounts
33 Halogens lThe halogens exist as diatomic molecules lAt room temperature, fluorine is a yellow gas, chlorine is a pale green gas, bromine is a red liquid, and iodine is a purple solid lThe elements have very high ionization energies, typical of nonmetals
34 Properties of the Halogens lOuter electron configurations: ns 2 np 5. lAll halogens have large electron affinities. lMost common oxidation state is -1, but oxidation states of +1, +3, +5 and +7 are possible. lHalogens are good oxidizing agents. lEach halogen is the most electronegative element in its row. 2:35 1:07
35 Properties of the Halogens lThe properties of the halogens vary regularly with their atomic number.
36 Fluorine lThe bond enthalpy of F 2 is low. Hence fluorine is very reactive. lWater is oxidized more readily than fluoride, so F 2 cannot be prepared by electrolysis of a salt solution. F 2 is an extremely reactive gas, which reacts with all the elements, except oxygen and the lighter noble gases, to form stable fluorides, often explosively. lF 2 is such a strong oxidizing agent that it can convert oxides, including water, to molecular oxygen
37 Group Work lWrite a balanced equation for the reaction of F 2 with H 2 O (to form O 2 and ??).
38 Chlorine lChlorine exists as chlorides in sea water, salt lakes, and brine deposits lCl 2 gas prepared industrially by the electrolysis of sodium chloride solutions lAlso a by-product of the preparation of metals by electrolysis of molten salts such as NaCl, MgCl 2, and CaCl 2
39 Chlorine lMost Cl 2 is used as a raw material in the production of other chemicals, in the synthesis of herbicides and insecticides, in the bleaching of textiles and paper, in purifying drinking water, and in the production of plastics such as polyvinyl chloride (PVC)
40 Bromine lBromine exists in small quantities in the form of bromides co-existing with chlorides lPrepared by reacting a solution containing bromide ion with chlorine
41 Bromine lUses of bromine: las a bleach lin the manufacture of bromide compounds lethylene bromide, C 2 H 4 Br 2, used as an antiknock agent in gasoline
42 Iodine lIodine exists as iodides in brines and seaweed, and as iodates in deposits of sodium nitrate (NaNO 3, also called Chile saltpeter) Recovered by oxidation of I - with Cl 2 or by reduction of IO 3 - with HSO 3 - Recovered by oxidation of I - with Cl 2 or by reduction of IO 3 - with HSO 3 - lUsed as an antiseptic and disinfectant and as a reagent for chemical analysis
43 Astatine lNot much is known about its chemistry lIt is highly radioactive; all chemical studies have used small quantities added to iodine solutions and measure behavior by determining where the radioactivity ends up. lThe total amount in the Earth’s crust is estimated to be < 30 g at any one time.
44 Halogens lIn addition to the common -1 and 0 oxidation numbers, the halogens (except for fluorine) exist with each positive oxidation number through +7 lHalogen oxides are known with oxidation numbers as high as +7; most are very strong oxidizing agents 06m02vd1.mov Gummy bears + KClO 3
45 Oxyacids and Oxyanions lFluorine only forms one oxyacid: HOF. Oxygen is in the zero oxidation state. lAll are strong oxidizing agents. lAll are unstable and decompose readily. lOxyanions are more stable than oxyacids. lAcid strength increases as the oxidation state of the halogen increases.
46 Periodic Acid lPeriodic (HIO 4 ) and paraperiodic (H 5 IO 6 ) acid have iodine in the +7 oxidation state. lPeriodic acid is a strong acid, paraperiodic acid is a weak acid (K a1 = 2.8 10 -2, K a2 = 4.9 ). lThe large iodine atom allows 6 oxygen atoms around it. lSmaller halogens cannot form this type of compound.
47 Strength of Oxoacids
48 Halogens lMany aqueous halogen species are susceptible to disproportionation: l ClO 2 l HClO 2 l HOCl l Cl 2 (in base) l HOBr l Br 2 (in base) l HOI l IO - l I 2 (in base) lHow do we decide which species will disproportionate?
The Other Group 6A Elements: S, Se, Te, and Po (Chalcogens) lTrend that affects other properties is the increase in metallic character down the group, indicated by the decreases in ionization energy and electronegativity O2O2O2O2 S8S8S8S8 Se Te
50 Group 6A lNonmetallic character dominates in this group lNonmetallic O exists as diatomic molecules lNonmetallic S exists as various covalently bonded polyatomic forms lMetalloids Se and Te are more metallic than S, but bear some resemblance to S lPo is even more metallic, but its behavior is not well known since it is a rare, radioactive element
51 Group 6A lOxygen dominated by oxidation number -2 lThe other elements have oxidation numbers from -2 through +6, the higher ones being quite common, especially in combination with oxygen
52 Sulfur lS found in earth's crust as sulfide and sulfate minerals and as the free element lAlso a small but critical constituent of plant and animal tissue lOccurs as sulfur dioxide and sulfur trioxide in the atmosphere Frasch process
53 Sulfur lElemental sulfur is a tasteless, odorless, combustible yellow solid existing in a variety of allotropes with different molecular structures lRhombic and monoclinic forms consist of S 8 rings
54 Sulfur lMajor use of sulfur is the preparation of sulfuric acid, which is used primarily to make phosphate fertilizers and impure phosphoric acid from phosphate rock lSulfur forms binary compounds with all the elements except iodine and the noble gases
55 Sulfur lHydrogen sulfide (H 2 S): lusually prepared by reaction of a metal sulfide with an acid lgas well known for its "rotten-egg" odor lextremely poisonous llargest source of sulfur in the atmosphere lSulfur reacts with oxygen to form two oxides, sulfur dioxide and sulfur trioxide, which form oxoanions (SO 3 2- and SO 4 2- ) and oxoacids (H 2 SO 3 and H 2 SO 4 ) by reaction with metal oxides or water
56 Sulfur lSulfuric acid is a powerful dehydrating agent, strong acid and moderate oxidizer. lSulfuric acid removes H 2 O from the sugar leaving a black mass of C. Steam is produced because the reaction is very exothermic.
57 Group Work lWhat are the structures of: lSO 2 lSO 3 SO SO SO SO lH 2 SO 3 lH 2 SO 4
Nitrogen lElectronic configuration of the Group VA (15) elements is ns 2 np 3 lLittle resemblance between the chemistry of nitrogen and the other elements in this group
59 Nitrogen lNitrogen as an element is the colorless, odorless, diatomic molecule N 2, the major constituent of air lAn essential component of all living matter in protein and amino acids lNitrogen compounds are important components of chemical fertilizers lMost uses of nitrogen involve its compounds, such as ammonia and nitrogen oxides.
60 Group Work lList the formulas for all the oxides that you expect nitrogen to form.
61 Preparation of Nitrogen lN 2 is produced by fractional distillation of air. lNitrogen is fixed by forming NH 3 (Haber Process). lNH 3 is converted into other useful chemicals (NO, NO 2, nitrites and nitrates).
62 Nitrogen lPositive oxidation numbers of nitrogen occur in the oxides, including N 2 O, NO, N 2 O 3, NO 2, N 2 O 4, and N 2 O 5. lAqueous N 2 O 3 is converted to nitrous acid (HNO 2 ), and N 2 O 5 to nitric acid (HNO 3 ). lNitrous oxide, N 2 O, occurs naturally in the atmosphere, as a result of the natural degradation of proteins.
63 Nitrogen lN 2 O can be formed by thermal decomposition of NH 4 NO 3 lN 2 O is used as an anesthetic (in laughing gas) lNitric oxide, NO, is formed by reaction of Cu metal with dilute aqueous nitric acid or in high temperature combustion processes and in the oxidation of ammonia gas commercially
64 Nitrogen lNO reacts rapidly with O 2 to form reddish- brown NO 2. lDinitrogen trioxide results from reaction between NO and NO 2. lN 2 O 3 is blue as a liquid. lGaseous N 2 O 3 reacts with water to form the weak acid, nitrous acid, HNO 2.
65 Group Work lWrite an equation for the preparation of nitrogen dioxide from common laboratory chemicals.
66 Nitrogen lNitrogen dioxide, NO 2, is a poisonous, reddish-brown gas with an irritating odor, which exists in equilibrium with colorless N 2 O 4 lDinitrogen pentoxide, N 2 O 5, is a volatile low-melting white solid; dissolved in water, it forms HNO 3, nitric acid
The Other Group 5A Elements: P, As, Sb, and Bi (Pnictogens) lAlthough nitrogen is found in nature primarily as unreactive N 2, the other elements are found only in compounds lMetallic nature increases down the group P As Bi Sb
68 Group Work lPredict the following properties for phosphorus: lmetallic character ltype of bonding lacid/base character of oxides lformulas of oxides
69 Phosphorus lPhosphorus is essentially nonmetallic, forms covalent bonds and has acidic oxides
70 Group Work lPredict the acid/base character of the oxides of As, Sb, and Bi, relative to those of P. How can we explain the predicted trend?
71 Group 5A lArsenic has properties between those of a nonmetal and a metalloid, with amphoteric (though more acidic than basic) oxides lAntimony is mostly metallic, but with some properties of a metalloid, and amphoteric (more basic than acidic) oxides lBismuth, is metallic, with basic oxides lThese trends are consistent with dramatic decreases in ionization energy down the group
72 Group 5A lTrends in oxidation number: lN and P range from -3 to +5 lAs primarily +3 and +5 lSb mostly +3, occasionally +5 lBi almost exclusively +3, rarely as +5 lThe lower, more metallic elements in the group have fewer stable oxidation numbers than N
73 Phosphorus lPhosphorus found as phosphates in 190 different minerals, most importantly apatite, Ca 5 (PO 4 ) 3 OH lPhosphates are an important constituent of all bone tissue
74 Phosphorus lMainly occurs in phosphorus minerals (e.g. phosphate rock, Ca 3 (PO 4 ) 2 ). lElemental P 4 produced by reduction 2 Ca 3 (PO 4 ) 2 (s) + 6SiO 2 (s) + 10C(s) P 4 (g) + 6CaSiO 3 (l) + 10CO(g) 2 Ca 3 (PO 4 ) 2 (s) + 6SiO 2 (s) + 10C(s) P 4 (g) + 6CaSiO 3 (l) + 10CO(g) lPhosphoric acid, made from phosphate rock, is one of the ingredients in Coca-Cola.
75 Phosphorus lPhosphorus occurs as 19 allotropes, the principal ones being white (tetrahedral P 4 ), red, and black phosphorus
76 Phosphorus lWhite phosphorus is a soft solid, with molecules held together by weak intermolecular forces. lWhite phosphorus is poisonous and causes painful skin burns. lWhite phosphorus is quite reactive and ignites spontaneously in air Barking Dogs
77 Phosphorus lRed phosphorus is an amorphous solid formed by heating white phosphorus lRed phosphorus involves P 4 tetrahedral molecules bonded to one another in long chains lBlack phosphorus is still less reactive lBlack phosphorus may be amorphous or have a graphite-like structure; it is metallic in appearance and an electrical conductor
78 Group Work lWhat acids are formed when the oxides, P 4 O 6 and P 4 O 10, are reacted with excess water?
79 Phosphorus lTwo oxides of phosphorus, P 4 O 6 and P 4 O 10, formed by burning phosphorus lReact with water to form phosphorous acid, H 3 PO 3 and phosphoric acid, H 3 PO 4 lP 4 O 10 is used as a drying agent because of its affinity for water.
Carbon lCarbon constitutes about % of the earth’s crust. lCarbon is the main constituent of living matter. lStudy of carbon compounds is called organic chemistry.
81 Carbon lCarbon forms more compounds than all other elements except hydrogen; typical compounds are the hydrocarbons and their derivatives
82 Carbon lExists as diamond, graphite, and an amorphous form, as well as the recently discovered allotrope C 60, called buckminsterfullerene
83 Carbon lDiamond: clear crystalline form of carbon, one of the hardest substances known, can be synthesized from graphite with high temperature and pressure and a metal catalyst
84 Group Work lWhy is diamond so hard, while graphite is soft and slippery, even though both are pure carbon?
85 Carbon lGraphite: lslippery gray-black solid lstrong covalent bonds hold atoms together in each layer, but the layers are bonded only by weak van der Waals forces, so the layers slide across one another readily lfound widely distributed in the earth's crust and synthesized from amorphous carbon
86 Carbon lUses of graphite: lcrucibles llubricants lpencils lnuclear reactors (to slow down fast neutrons) lelectrodes for electrolysis reactions
87 Carbon lAmorphous carbon: lcarbon blacks lcharcoal lactivated carbon lsoot lcoke lEssentially microcrystalline forms of graphite with no layering lFormed by thermal decomposition or partial decomposition of coal, petroleum, natural gas, and wood with an insufficient supply of oxygen
88 Carbon lAmorphous carbon: lburning oil gives lampblack lheating coal in the absence of air produces coke lheating wood in the absence of air gives charcoal lCarbon black used as a filler in rubber tires to increase toughness and prevent brittleness lLampblack used in inks, paints, coating on carbon paper
89 Carbon lCharcoal used lin filters to adsorb odors lin gas masks to adsorb poisonous gases lin the decolorizing of sugar lin water treatment lin the reclamation of dry-cleaning solvents lCoke used in the extraction of metals from their oxide ores
90 Carbon lBuckminsterfullerene consists of molecules of C 60 formed by laser or high-temperature carbon arc vaporization of graphite lOne member of a class of new forms of carbon called fullerenes, which consist of clusters of carbon atoms containing even numbers from 44 to 84
91 Fullerenes in Solution
92 Carbon lC 60 exists as a truncated icosahedron, which contains 12 pentagonal faces and 20 hexagonal faces lRemarkable physical stability, but chemically reactive lNow being prepared as tubes, into which metals can be inserted. These are the thinnest capillary tubes known.
93 Carbon lElemental carbon is relatively unreactive at room temperature lInsoluble in water, dilute acids and bases, and organic solvents lAt high temperatures, carbon becomes highly reactive and combines directly with many elements, including oxygen lCarbon at high temperatures also reduces water, metal oxides, oxoanions (e.g., phosphate in phosphate rock), hydrogen
94 Oxides of Carbon lCarbon forms CO and CO 2. lCO is very toxic (binds irreversibly to Fe in hemoglobin, causing respiratory arrest). lCO also has a lone pair on C, which is unusual. lCO is a good Lewis base Ni(CO) 4 forms when Ni is warmed in CO lCO can be used as a fuel 2CO(g) + O 2 (g) 2CO 2 (g) H = -566 kJ lCO is a good reducing agent Fe 3 O 4 (s) + 4CO(g) 3Fe(s) + 4CO 2 (g)
95 Oxides of Carbon lCO 2 is produced when organic compounds are burned in oxygen: C(s) + O 2 (g) CO 2 (g) CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) C 2 H 5 OH(l) + 3O 2 (g) 2CO 2 (g) + 3H 2 O(l) lCO 2 is produced by treating carbonates with acid.
96 Oxides of Carbon lFermentation of sugar to produce alcohol also produces CO 2 : C 6 H 12 O 6 (aq) 2C 2 H 5 OH(aq) + 2CO 2 (g) lAt atmospheric pressure, CO 2 condenses to form CO 2 (s) or dry ice. lCO 2 is used as dry ice (refrigeration), carbonation of beverages, washing soda (Na 2 CO 3.10H 2 O) and baking soda (NaHCO 3.10H 2 O).
97 Carbonic Acid and Carbonates lWhen CO 2 dissolves in water (moderately soluble) carbonic acid forms: CO 2 (aq) + H 2 O(l) ⇌ H 2 CO 3 (aq) lCarbonic acid is responsible for giving carbonated beverages a sharp acidic taste. lPartial neutralization of H 2 CO 3 gives hydrogen carbonates (bicarbonates) and full neutralization gives carbonates. lMany minerals contain CO 3 2-.
98 Carbonic Acid and Carbonates lAt elevated temperatures CaCO 3 decomposes: CaCO 3 (s) CaO(s) + 2CO 2 (g) lThis reaction is the commercial source of lime, CaO. lCaO reacts with water and CO 2 to form CaCO 3 which binds the sand in mortar: CaO(s) + H 2 O(l) Ca 2+ (aq) + 2OH - (aq) Ca 2+ (aq) + 2OH - (aq) + CO 2 (aq) CaCO 3 (s) + H 2 O(l)
99 Carbides lCarbon combines with elements with a lower electronegativity to form carbides, which exist in three classes ¶Salt-like carbides form with the most electropositive metals and are ionic, so they are hydrolyzed by water or dilute acid to give hydrocarbons
100 Carbides ·Interstitial carbides are formed with transition metals, are very hard and have very high melting points, high metallic conductivity, and metallic luster; they consist of a metal with carbon atoms located in some of the interstitial sites (or holes) in the metal structure
101 Carbides ¸Covalent carbides include those of silicon and boron, which are close in size and electronegativity to carbon; they are completely covalent and form infinite network structures, are exceptionally hard materials widely used as abrasives
The Other Group 4A Elements: Si, Ge, Sn, and Pb Si Sn
103 General Characteristics of Group 4A Elements lThe electronegativities are low. lThe dominant oxidation state for Ge, Sn and Pb is +2. lCarbon has a coordination number of 4, the other members have higher coordination numbers. lC-C bonds are very strong, so C tends to form long chains. lBecause the Si-O bond is stronger than the Si-Si bond, Si tends to form oxides (silicates).
104 Group Work lCompare the following properties on going down Group 4A: lMetallic character lIonization energies lMelting points lAcid/base nature of oxides
105 Group 4A lSi is nonmetallic/metalloid with only some of its chemistry similar to carbon lGe is a metalloid lSn and Pb are metallic lIonization energies and melting points decrease down the group, reflecting the change from nonmetal to metal
106 Group Quiz lWrite a formula for the acid, if any, that would be formed by reaction of these oxides with water: lCO 2 lCO lCaO lSO 3 lP 4 O 10 lHint: You should have three formulas.
107 Group 4A lCarbon bonds readily to itself lThis tendency diminishes on going down the group because the bond strength decreases considerably as the elements get larger lOxidation number +4 dominates near the top of the group, +2 becomes more stable down the group
108 Silicon lSilicon exists in the earth's crust as silicon dioxide and over 800 silicate minerals lElemental silicon obtained by reduction of SiO 2 with C or CaC 2 at high temperature, purified by zone refining SiO 2 (l) + 2C(s) Si(l) + 2CO(g)
109 Silicon lSilicon: lbrittle gray-black metallic-looking solid lquite hard lhigh melting point ldiamond-like tetrahedral network structure linert at room temperature but reactive at high temperatures lUsed in semiconductor devices
110 Silicon lSilicon hydrides or silanes arise from reaction of Mg 2 Si with acids, giving a mixture of SiH 4, Si 2 H 6, Si 3 H 8, Si 4 H 10, Si 5 H 12, and Si 6 H 14
111 Group Work lWhat is the structure and bonding for Si 2 H 6 ?
112 Silicon lSilanes are quite reactive: lThe first two are stable. lThe others decompose to give SiH 4, Si 2 H 6, and H 2. lMuch more reactive than the corresponding alkanes because of the availability of empty 3d orbitals that can be used to form bonds with another reactant lThey are spontaneously flammable in air.
113 Group Work lSiO 2, found as quartz, is quite hard. What feature of its structure gives rise to this hardness? RoseQuartz
114 Silicon lSiO 2 exists as polymeric (SiO 2 ) n, which has silicon covalently bonded to four bridging oxygen atoms. lExtended covalent bonding network lHard, high-melting solid Amethyst Smoky Quartz
115 Silicon lGlass is formed by heating together silicon dioxide, alkali metal and alkaline earth metal oxides, and sometimes other oxides, the particular mixture controlling the properties of the glass. lSilicate structure is somewhat random in contrast to crystalline silicates
116 Silicon lSome bonds are under more strain than others, so the glass melts over a range of temperatures and can be softened without melting
117 Silicon l90 % of the earth’s crust is composed of compounds of Si and O. lSilicates are compounds where Si has four O atoms surrounding it in a tetrahedral arrangement. lThe oxidation state of Si is +4. lThe silicate tetrahedra are building blocks for more complicated structures.
118 Silicon lMany silicates naturally with the basic structural unit being the SiO 4 tetrahedron occurring in several varieties: lsingly lsmall groups sharing oxygen atoms lsmall cyclic groups linfinite chains ldouble-stranded chains (or bands) linfinite sheets lThese few structures form many hundreds of minerals by combination of silicate anions with metal cations.
119 Silicate Minerals
123 Silicon lThe silicate structure is reflected in the physical structure of minerals such as mica and asbestos.
Boron lBlack crystalline element, extremely hard and brittle, low density, high melting point and boiling point, low electrical conductivity (so classified as a semiconductor), used in semiconductor electronics and added to steel to increase strength and to copper to increase electrical conductivity
125 Boron lSimilarities of some properties to those of carbon, of others to those of silicon (due to the diagonal relationship leading to similar size and electronegativity) lCarbon, silicon, and boron all form covalently-bonded extended network solids and covalent halides
126 Boron lOccurs in deposits of borax, Na 2 B 4 O 7. 10H 2 O lBorax used as an additive to laundry detergents to soften the water. lWhich U.S. president was associated with Twenty Mule Team Borax? lBoron recovered by acidifying, heating, and reduction of the resulting oxide by heating with Mg, purified by zone refining
127 Boron lResembles metals in its physical properties but is more like nonmetals chemically lChemical behavior is complex and unusual: ionization energy is unusually high, so formation of a cation is difficult, high electronegativity (comparable to nonmetals) lAll its compounds are covalent
128 Group Work lWhat is the valence electron configuration of boron? How many covalent bonds will it normally form?
129 Boron lValence electron configuration 2s 2 2p 1, so forms only three normal covalent bonds, but electron deficiency makes it a good Lewis acid lOxidation number +3 common, but others found in boranes B 3 + does not exist in aqueous solution. Why not? B 3 + does not exist in aqueous solution. Why not?
130 Boron lReacts with F 2 and Cl 2 to give trihalides lGreat affinity for O 2, which is used to remove oxygen from metal oxides to purify molten metals
131 Boron lB reacts with N 2 at high temperature to give solid BN. lBN is very stable due to its graphite-like structure, which arises from the presence of only 3 valence electrons and the tendency to use sp 2 hybrid orbitals.
132 Boron lBN is also known in a diamond-like structure, which is formed by application of high temperature and pressure, and is extremely hard and is used as an abrasive.
133 Boron lLarge number of hydrides have been prepared lBH 3 is known but very unstable (reactive) lSimplest stable one is diborane, B 2 H 6, which decomposes to other boranes, e.g. B 5 H 9, when heated
134 Boron lUnusual structure and bonding in the boranes, in which hydrogens act as bridges between boron and the B-H-B arrangement shares two electrons between three atoms (called three center bonding) lBecause of their electron deficiencies, boranes are highly reactive lDiborane is very reactive: B 2 H 6 (g) + 3O 2 (g) B 2 O 3 (s) + 3H 2 O(g) H = kJ
135 Boron lSome boranes are reactive (B 5 H 9 ) while some are stable in air at room temperature (B 10 H 14 ). lB 2 O 3 is the only important boron oxide. lBoric acid, H 3 BO 3 or B(OH) 3 is a weak acid (K a = 5.8 ). lBoric acid is used as an eye wash.