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George Mason University General Chemistry 212 Chapter 14 Main Group Element Patterns

George Mason University General Chemistry 212 Chapter 14 Main Group Element Patterns Acknowledgements Course Text: Chemistry: the Molecular Nature of Matter and Change, 6 th ed , 2011, Martin S. Silberberg, McGraw-Hill

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George Mason University General Chemistry 212 Chapter 14 Main Group Element Patterns

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  1. George Mason University General Chemistry 212 Chapter 14 Main Group Element Patterns Acknowledgements Course Text: Chemistry: the Molecular Nature of Matter and Change, 6thed, 2011, Martin S. Silberberg, McGraw-Hill The Chemistry 211/212 General Chemistry courses taught at George Mason are intended for those students enrolled in a science /engineering oriented curricula, with particular emphasis on chemistry, biochemistry, and biology The material on these slides is taken primarily from the course text but the instructor has modified, condensed, or otherwise reorganized selected material.Additional material from other sources may also be included. Interpretation of course material to clarify concepts and solutions to problems is the sole responsibility of this instructor.

  2. Main-Group Elements • Chapter Overview • Application of bonding, structure, and reactivity to Main-Group Elements • Hydrogen • Period 2 Elements – Trends across Periodic Table • Group 1A – The Alkali Metals • Group 2A – The Alkaline Earth Elements • Group 3A – The Boron Family • Group 4A – The Carbon Family • Group 5A – The Nitrogen Family • Group 6A – The Oxygen Family • Group 7A – The Halogens • Group 8A – The Noble Gases

  3. Main-Group Elements • In chemistry and atomic physics the periodic table divides the elements into 4 groups • Main Group Elements are elements in groups (periodic columns) whose lightest members are represented by helium, lithium, beryllium, boron, carbon, nitrogen, oxygen, and fluorine as arranged in the periodic table of the elements. • Main group elements include elements (except Hydrogen) in groups: 1 (IA) 2 (IIA) (s-block) and 13 (IIIA) 14 (IV(A) 15 (VA) 16 (VIA) 17 (VIIA) 18 (VIIIA) (p-block)

  4. Main-Group Elements • Transition elements • The Transition elements occupy the following groups 3 4 5 6 7 8 9 10 11 12 IIIB IVB VB VIB VIIB VIIIB VIIIBVIIIB IB IIB • Zinc (Zn), Cadmium (Cd), and Mercury (Hg) in Group IIB share some properties of both the Main Group, and Transition Elements and some scientists believe they should be included as Main Group elements • Lanthanides • Elements in Period 6, group IIIB whose f-subshells are being filled • Actinides • Elements in Period 7, group IIIB whose f-subshellls are being filled

  5. Main-Group Elements In older nomenclature the main group elements are groups IA and IIA, and groups IIIA to groups IIIA) Main group elements (with some of the lighter transition metals) are the most abundant elements on the earth, in the solar system, and in the universeThey are sometimes called the representative elements Group 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 Period Main Group Elements 1 2 3 4 5 6 7 p block Transition (d block) Elements s block Lanthanides Actinides

  6. Main-Group Elements • Hydrogen (1s1) • 90% of all atoms in Universe are Hydrogen atoms • Single Electron; Small Size • No perfectly suitable position in the periodic table • Depending on the property, Hydrogen fits better in 1A, 4A, 7A • +1 oxidation State (grp 1A? • Relatively High Ionization Potential (grp 7?) • Forms diatomic molecule (H2 - grp 7?) • Shares electrons (grp 4?) • Half-filled valence shell; ionization energy; electron affinity, electronegativity, and bond energies most similar to group 4

  7. Hydrogen Chemistry • Hydrogen Bonding Dipole-Dipole force between Hydrogen (H) and small, highly electronegative atoms with lone electron pair: Nitrogen (N); Oxygen (O); Fluorine (F) • Highly reactive, combining with nearly every element • Ionic (salt like) hydrides • Group 1A & 2A metals 2Li(s) + H2(g)  2LiH(s) Lithium Hydride Ca(s) + H2(g)  CaH2(s) Calcium Hydride • In H2O, H- is a strong base that pulls H+ from water Na+H-(s) + H2O  Na+(aq) + OH-(aq) + H2(g) • Hydride ion is also a strong reducing agent Ti4+Cl4(l) + 4LiH(s) = Tio(s) + 4LiCl(s) + 2H2(g)

  8. Covalent Hydrides • Hydrogen reacts with nonmetals to form covalent hydrides CH4 NH3 H2O HF • Conditions for forming Covalent Hydrides depend on the reactivity of the nonmetal – the more stable, the more temperature & pressure required for formation Ex: Ammonia – 400oC & 250 atm N2(g) + 3H2(g)  2NH3(g) Horxn = -91.8 kJ At low temperatures (-196oC) Hydrogen combines readily with reactive Fluorine (F2) F2(g) + H2(g)  2HF Horxn = -546 kJ Catalyst

  9. Metallic (Interstitial) Hydrides • Many Transition Elements form metallic (interstitial) hydrides, where Hydrogen molecules (H2) and Hydrogen atoms (H) occupy the holes in the metal’s crystal structure. • These are not compounds, but rather gas-solid solutions • They lack a Stoichiometric formula because metal can incorporate a variable amount of hydrogen, depending upon temperature and pressure

  10. Trends Across Periodic Table • Electrons fill 1 ns – 3 np orbitals according to Pauli Exclusion Principle and Hund’s Rule • d orbitals in lower Periods can be used to accommodate additional oxidation states • Atomic size generally decreases • 1st ionization potential increases • Electronegativity increases • Metallic character decreases with increasing nuclear charge • Reactivity highest at right & left sides, less in middle • Bonding - metallic  covalent  none (noble gas) Continued on next Slide

  11. Trends Across Periodic Table • Bonding between each element and an active nonmetal changes from ionic to polar covalent • Bonding between each element and an active metal changes from metallic to polar covalent to ionic • Acid-Base behavior of common element oxide in water changes from basic to amphoteric (acts as acid or base (H2O) to acidic as bond between element and oxygen becomes more covalent • Reducing strength decreases through the metals • Oxidizing strength increases through the nonmetals

  12. Group 1A - Alkali Metals (ns1) • Lithium (Li), Sodium (Na); Potassium (K); Rubidium (Rb); Cesium (Cs); Francium (Fr) • Single electron relatively far from nucleus • weak metallic bonding - attraction between delocalized electrons and metal-ion cores in crystalline structure • Low melting points, soft consistency • Reactive Metals • Powerful reducing agents – lose 1 electron becoming 1+ cations, donating the electron to other elements • ns1 configuration forms salts readily (+1 cations) • Low Heat of Atomization (oHatom ) – Recall Lattice Energy) • Energy to convert solid into individual gaseous atoms oHatom (Li>Na>K>Rb>Cs)

  13. Group 1A - Alkali Metals (ns1) • Low Ionization Energy (IE) – Each alkali element has the largest size and the lowest IE in its Period • Size of atom decreases considerably when valence electron is lost • Lattice Energy – The atomic radius increases as you move down a group.  Since the square of the distance is inversely proportional to the force of attraction, lattice energy decreases as the atomic radius increases • For a given anion, the Lattice Energy become smaller as the cation becomes larger

  14. Group 1A - Alkali Metals (ns1) • Solubility – Despite strong ionic attractions, the Group 1A salts are water soluble – attraction between the ions and the polar Water molecule creates highly Exothermic Heat of Hydration (Hhydr) • Entropy – Entropy increases as ions disperse going into solution overcoming the high lattice energy • Magnitude of Hydration Energy decreases as ionic size increases H = -Hhydr (Li+ > Na+ > K+ > Rb+ > Cs+

  15. Group 1A - Alkali Metals (ns1) • Anomalous Behavior of Lithium • Lithium ion (Li+) is small and highly positive • Dissociation of Lithium salts, such as LiF, Li2CO3, LiOH, and Li3PO4, in water is much more difficult than similar salts of sodium (Na) and Potassium (K) • Only member of Alkali group that forms simple Oxide and Nitride, Li2O & Li3N, on reaction with O2 & N2 in air • Only Lithium forms organo-metalic molecular compounds with hydrocarbon groups from organic Halides 2Li(s) + CH3CH2Cl(g)  CH3CH2Li(s) + LiCl(s)

  16. Group 1A - Alkali Metals (ns1) • Reactions & Compounds of Alkali Metals • Alkali metals reduce Hydrogen in Water to form Hydrogen gas 2E(s) + 2H2O  2E+ + 2OH-(aq) + H2(g) Where E = any alkali metal (Li, Na, K, Rb, Cs) Reaction becomes more vigorous down group • Alkali metals reduce oxygen, but product depends on the metal 4Li(s) + O2(g)  2Li2O(s) oxide K(s) + O2(g)  KO2(s) superoxide • Alkali metals reduce Hydrogen to form ionic hydrides 2E(s) + H2(g) 2EH(s)

  17. Group 1A - Alkali Metals (ns1) • Reactions & Compounds of Alkali Metals • Alkali metals (E) reduce Halogens (X) to form Halides 2E(s) + X2 2EX(s) X = F, Cl, Br, I) • Sodium Metal (Na) can be produced from Molten NaCL and electricity 2NaCl(l)  2Na(l) + Cl2(g) • Sodium Hydroxide (Lye) can be produced from Salt (NaCl), water (H2O) and electrolysis 2NaCl(s) + H2O(l)  2NaOH(aq) + H2(g) + Cl2(g) • In an ion-exchange process, water can be “softened” by removal of dissolved hard-water cations to displace Na+ ions from a “resin” M2+(aq) + Na2Z(s)  MZ(s) + 2Na+(aq) (M = Mg, Ca: Z = resin)

  18. Group 1A - Alkali Metals (ns1) atomic properties physical properties

  19. Group 2A - Alkaline Earth Metals (ns2) • Be, Mg, Ca, Sr, Ba, Ra (E2+ ions) • Oxides (except Be) give basic (alkaline) solutions: Ca(OH)2, Mg(OH)2 • High melting points (higher lattice energy than 1A) • Atomic & Ionic sizes • Smaller radii and higher ionization energy • Increase in size down the group • Combination of size, extra electron, and metallic bonding result in stronger attractions between delocalized electrons and the atom cores • Thus, Melting Points and Boiling Points are much higher than 1A alkali metals • Harder & more dense than Alkali metals, but soft and lightweight compared to transition metals (Fe, Cr, etc)

  20. Group 2A - Alkaline Earth Metals (ns2) • Even though the Alkaline Earth metals have higher ionization potential, they still form ionic compounds (E2+), but Beryllium (Be) is an exception forming covalent bonds • Like Alkali metals, Alkaline Earth metals are strong reducing agents • Group 2A (Alkaline Earth) elements are reactive because the higher lattice energy of their compounds more than compensates for the large total Ionization Energy (IE) needed to form the 2+ cations • The higher Lattice Energy (from the smaller cation size) and higher Charge Density results in lower solubility • Ion-Dipole attraction is so strong that many slightly soluble 2A salts crystallize as “Hydrates” Epsom salt – MgSO47H2O Gypsum – CaSO42H2O

  21. Group 2A - Alkaline Earth Metals (ns2) • The anomalous behavior of Beryllium • Beryllium has smallest size; highest Ionization energy, and highest Electronegativity of the Alkaline Earth elements • Combined with the high charge density of the ion (Be2+) it polarizes the nearby electron clouds very strongly and causes extensive orbital overlap; this results in covalent bonding • BeF2 is the most ionic of the Beryllium compounds, but its melting point and electrical conductivity are relatively low compared to the other alkaline earth Fluorides • Unlike the other Alkaline Earth Metals, whose oxides are basic, BeO is amphoteric and does not react with water to form OH- ions

  22. Group 2A - Alkaline Earth Metals (ns2) • Diagonal relationships: Lithium and Magnesium • Certain Period 2 elements exhibit behaviors that are very similar to those of the Period 3 elements immediately below and to the right • 3 relationships • Li, Mg • Be, Al • B, Si • Lithium and Magnesium reflect similar atomic and ionic size • Both elements form: • Nitrides, • Hydroxides and Carbonates (CO3) that decompose with heat, • Organic compounds with polar covalent metal-carbon bonds • Salts with similar solubilities

  23. Group 2A - Alkaline Earth Metals (ns2) • Reactions & Compounds (E = Mg, Ca, Sr, Ba) • Metals reduce Oxygen (O2) to form Oxides 2E(s) + O2(g)  2EO(s) Ba + O2  BaO2 (Barium Peroxide) • Larger metals reduce water to form hydrogen gas E(s) + 2H2O(l)  E2+aq) + 2OH- (aq) + H2(g) • Metals reduce Halogens to form ionic halides E(s) + X2  EX2(s) (X = F, Cl, Br, I) • Most metals (Be exception) reduce Hydrogen to form ionic hydrides E(s) + H2(g)  EH2 (s) (except Be)

  24. Group 2A - Alkaline Earth Metals (ns2) • Reactions & Compounds (E = Ca, Mg, Sr, Ba) • Most elements reduce Nitrogen to form ionic Nitrides 3E(s) + N2(g)  E3N2(s) (except Be) • Element Oxides are Basic (except for amphoteric BeO) EO(s) + H2O(l)  E2+(aq) + 2OH-(aq) • All Carbonates undergo thermal decomposition to the oxide ECO3(s)  EO(s) + CO2(g) (CaO – Lime) • Beryl (Be3Al2Si6O18) - Gemstone, source of Be • Magnesium oxide (MgO) – Refractory material for furnace bricks • Alkyl Magnesium Halides – RMgX (R=Hydrocarbon) Grignard Reagents – organic compound synthesis heat

  25. Group 2A - Alkaline Earth Metals (ns2) atomic properties physical properties

  26. Group 3A – Boron Family (ns2np1) B Al Ga In Tl • Boron heads family, but other elements in group 3A exhibit diverse properties • Boron & Aluminum, especially Aluminum, are much more abundant than the others, but still quite rare • Group 3A elements include “p” orbitals for first time • In Period 4 (transition elements) the “d” orbitals are present • Physical Properties are influenced by type of bonding

  27. Group 3A – Boron Family (ns2np1) • Boron is a network covalent metalloid - Black, hard, very high melting point • A network solid or covalent network solidis a chemical compound in which the atoms are bonded by covalent bonds in a continuous network • In a network solid there are no individual molecules and the entire crystal may be considered a macromolecule • Boron (metalloid) is much less reactive than the others members of the 3A group because it forms covalent bonds

  28. Group 3A – Boron Family (ns2np1) • Other group members are metals – shiny, relatively soft with low melting points • Aluminum is more ionic; its low density and 3 valence electrons make it a good electrical conductor • Although Aluminum is a metal, its halides exist in the gaseous state as covalent dimers - AL2Cl6 (contrast salts of group 1 & 2 metals) • Aluminum Oxide, Al2O3,is amphoteric (can act as an acid or base) rather than basic like the Group 1A & 2A metals • Although the other Group 3A elements are basically ionic they exhibit more Covalent character than similar 2A compounds. • 3A cations are smaller with more charge density than 2A cations and they polarize an anion’s electron cloud more effectively

  29. Group 3A – Boron Family (ns2np1) • Oxidation-Reduction (REDOX) behavior in Group 3A • Presence of Multiple Oxidation States • In Groups 3A – 6A many of the larger elements (down the group) exhibit an oxidation state “two lower” than the A-Group number • This lower state occurs when the atoms lose their np electrons, not the ns electrons. • The lower oxidation state is the result of lower bond energies • Bond energies decrease as the size of the atom and the bond length increase for elements lower in the group

  30. Group 3A – Boron Family (ns2np1) • Increasing prominence of the low oxidation state • When a group exhibits more than one oxidation state, the lower state becomes more prominent going down the Group • All members of the 3A group exhibit the +3 state, but the +1 state appears first with some compounds of Gallium (Period 4) • The +1 state becomes the most important state of Thallium (Period 6)

  31. Group 3A – Boron Family (ns2np1) • Relative Basicity of Group 3 oxides • Recall: A1 oxides (ionic charge +1 and more metallic) are more basic than A2 oxides (ionic charge +2 and less metallic) • In general, oxides with the element in a lower oxidation state (less positive) are more basic than oxides with the element in a higher oxidation state • For Indium oxides in Group A3, In+12O acts more like a metal and is more basic than In+32O3 • The lower charge density of In+1 does not polarize the O-2 ion as much as the In+3 ion • Thus, in E2O compounds, the E-O bonding is more ionic than in E2O3 compounds, thus; the O-2 ion is more available to act as a base – donate electron pair or accept a proton

  32. Group 3A – Boron Family (ns2np1) • Boron Chemistry • Boron compounds are covalent (unique within group) • Forms network covalent compounds or large molecules with metals, H, O, N • Electron deficient; uses two approaches to complete octet • Accepting a Bonding Pair from Electron-Rich atom BF3(g) + NH3(g)  F3B-NH3(g) (BF3 acts as acts as Lewis acid in accepting the electron pair from the Nitrogen in NH3) B(OH)3 + H2O(l)  B(OH)4-(aq) + H+(aq) (Acts as acid by accepting electron pair from H2O) Note: Water is acting as the base

  33. Group 3A – Boron Family (ns2np1) • Boron Chemistry • Two approaches to filling octet (con’t) • Accepting electron pair from Electron-Rich atom (con’t) • Boron-Nitrogen compounds are similar in structure to elemental Carbon and some of its organic compounds • Size, Ionization Energy, Electronegativity of Carbon is between Boron & Nitrogen • Ethane & Amine – Borane have the same number & electron configuration       C – C B – N      

  34. Group 3A – Boron Family (ns2np1) • Boron Chemistry • Two approaches to filling octet (con’t) • Forming Bridge Bonds with Electron-Poor Atoms • Boron Hydrides - Boranes • 2 types of B – H bonds • Normal electron-pair bond • sp3 orbital of B overlaps 1s orbital of H in each of the four terminal B-H bonds

  35. Group 3A – Boron Family (ns2np1) • Hydride Bridge Bond (3-center, 2 electron bond) • Each B – H – B grouping is held together by only two electrons • Two sp3 orbitals, one from each B, overlap an H 1s orbital between them • Two electrons move through this extended bonding orbital – one from one of the B atoms and the other form the H atom – and join the 2 B atoms via the H atom bridge

  36. Group 3A – Boron Family (ns2np1) atomic properties physical properties

  37. Group 3A – Boron Family (ns2np1) • Reactions & Compounds • Elements react sluggishly, if at all, with water (H2O) 2 Ga(s) + 6H2O(hot)  2Ga3+(aq) 6OH-(aq) + 3H2(g) 2Tl(s) + 2H2O(steam)  2Tl+(aq) +2OH-(ag) + H2(g) Note different oxidation numbers for Ga3+ & Tl+ • All members form oxides when heated in pure O2 4E(s) + 3O2(g)  2E2O3(s) (E = B, Al, Ga, In) 4Tl(s) + O2 2Tl2O3(s) • Oxide acidity decreases down the group: B2O3 > Al2O3 > Ga2O3 > In2O3 > TlO2 (weakly acidic) (strongly basic) The +1 oxide (TlO2) is more basic than the +3 oxide

  38. Group 3A – Boron Family (ns2np1) • Reactions & Compounds • All members reduce Halogens 2E(s) + 3X2 2EX3 (E = B, Al, Ga, In) 2Tl(s) + X2 2TlX(s) • Trihalides of AL, Ga, In are mostly ionic butexist as dimers in the gas phase • Acid (H2SO4) treatment of Al2O3 produces Al2SO4, a colloid (coagulant) used in water purification • Al2O3 + 3H2SO4  Al2SO4(s) + 3H2O(l)

  39. Group 4A – Carbon Family (ns2np2) • The whole range of elemental behavior occurs within the 4A group • Non metalic Carbon (C) • Metalloids (Silicon (Si) & Germanium (Ge) • Metallic (Tin (Sn) & Lead (Pb) • Newly synthesized element at bottom of group • Carbon forms the basis of “Organic Chemistry”  20,000,000 compounds • Polymer Chemistry • Biochemistry based on Carbon • Geochemistry • Electronic technologies bases on Si

  40. Group 4A – Carbon Family (ns2np2) • Bonding effects on Physical Properties • Silicon has a much lower melting point than Carbon because of the longer, weaker bonds. • The melting point difference between Germanium (Ge) and Tin (sn) is due to the change from network covalent to metallic • Going from Group 3 to group 4 there are large increases in melting point and the Hfusbecause of the change from metallic to network covalent bonding

  41. Group 4A – Carbon Family (ns2np2) • Allotropism: Different Forms of an Element • Elemental Carbon – Graphite & Diamond • Different crystalline & molecular forms with different physical properties • Carbon Allotropes • Graphite – Black, “greasy”, soft, more stable than diamond • Diamond – Colorless, electrical insulator, extremely hard • Bucky-Balls (Buckminsterfullerene) – soccer ball-shaped with the formula C60 • Tin Allotropes • -tin – stable at room temperature & above • -tin – stable below 13oC

  42. Group 4A – Carbon Family (ns2np2) • Bonding Changes in Group 4A • Carbon – Covalent (intermediate EN) • Si & Ge – strong polar bonds (silicate minerals) • Tin (Sb) & Lead(Pb) – Metallic (Ionic) • Multiple Oxidation States • Carbon (+4) • Silicon (+4 more stable than +2) • Lead (+2 more stable than +4) • Elements with lower oxidation states act more like metals (more basic)

  43. Group 4A – Carbon Family (ns2np2) • Highlights of Carbon Chemistry • Carbon, like other elements in “Period 2, is the anomalous element in the group • Carbon forms bonds with: • Smaller Group 1A & 2A metals • Many transition metals • Halogens • Neighbors B, Si, N, O, P, S • Exhibits all possible oxidation states from +4 in CO2, and Halides to -4 in CH4

  44. Group 4A – Carbon Family (ns2np2) • Highlights of Carbon Chemistry • Two main features of Carbon Chemistry • Catenation and the ability of Carbon to form multiple bonds • Carbon can form chains, branches, and rings (aromatic & aliphatic) • Multiple bonds – sigma (), Pi (), Triple () • The C-C bond is short enough for side-to-side overlap of two half-filled 2p orbitals to form  bonds that give rise to many diverse structures and reactivities of organic compounds

  45. Group 4A – Carbon Family (ns2np2) • The Other 4A elements • E-E bonds become longer going down the group, with decreasing bond strength C – C > Si – Si > Ge – Ge • The empty d shell orbitals make these chains susceptible to chemical attack – they are reactive • The long bonds are not suitable for overlap of p orbitals; thus, no  bonds

  46. Group 4A – Carbon Family (ns2np2) • Carbonates • Metal Carbonates are the main mineral Marble, Limestone, Chalk, Coral, others • Antacids – Calcium Carbonate & Stomach Acid CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2(g) + H2O(l) • Limestone (CaCO3) deposits help moderate the effects of acid rain (H2SO4 & HNO3) • Carbon Dioxide (CO2) • Essential to all life as primary source of carbon in plants & animals through photosynthsis • Atmospheric buildup from motor vehicles and fossil fuel powerplant severely affect global climate

  47. Group 4A – Carbon Family (ns2np2) Trisilylamine (SiH3)3N trigonal planar • Silicon Chemistry • Silicon Halides are more reactive than Carbon Halides because Si (3s, 3p, 3dorbitals) has empty 3d orbitals available for bond formation • The Si – X bond is long but stronger than corresponding C – X bond • Si – X bond has some double bond character because of the presence of a  bond and a different type of  bond called a p,d- (side-to-side overlap of the Si d orbital and a Halogen p orbital Trimethylamine (CH3)3N The impact of p,d-p bonding on the structure of trisilylamine trigonal pyramidal

  48. Group 4A – Carbon Family (ns2np2) • Silicon chemistry is dominated by the Silicon-Oxygen – (Si–O) bond • C – C bonds can repeat endlessly; similarly, the Si–O bonds can also repeat forming large chains in Silicate minerals in the earths crust and Silicones, which are synthetic polymers with a large number of industrial applications • Silicate Minerals • From common sand (SiO2) and clay to semiprecious amethyst, Silicate minerals are the dominant form of matter on the earth • Oxygen is the most common element on earth and Silicon is the next most abundant

  49. Group 4A – Carbon Family (ns2np2) • The Orthosilicate (–SiO4 –) grouping is the building unit for Silicate minerals • Zircon ZrSiO4 1 Unit • Hemimorphite 2 Units [Zn4(OH2Si2O7H2O] • Beryl 6 Units [Be3Al2Si6O18] • Silicon Polymers • Manufactured Substances • Alternating Si & O atoms with two Organic groups bonded to each Silicon atom

  50. Group 4A – Carbon Family (ns2np2) atomic properties physical properties

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