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Properties of Carbon Element

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Properties of Carbon Element

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  1. Properties of Carbon Element

  2. Properties of Carbon Element • We have learnt carbon element as the basis of Organic Chemistry. • Why does carbon can form so many different compounds? • Carbon is found the second period group IVA of the periodic table.

  3. EXCEPTIONS OF CARBON COMPUDS WHICH ARE NOT ORGANIC • oxides of carbon (CO2, CO) • carbonates,bicarbonates(NaHCO3,CaCO3) • cyanides (NaCN, etc)

  4. Properties of Carbon Element • The Lewis structure for carbon shows 4 unpaired valence electrons. • To fulfill the octet rule, a carbon atom needs 4 more electrons. • A carbon atom may form 4 covalent bonds and is capable of forming long chains with single, double or triple bonds between carbon atoms. • These chains may be continuous (straight) or branched. • The 2 ends of a chain can bond together to form a ring. • Carbon compounds are divided into classes based on their chemical similarity.

  5. Hydrocarbons • Hydrocarbons are compounds containing hydrogen and carbon. Hydrocarbons may have different numbers of bonds between carbon atoms. • The four hydrocarbon classes are: alkane (single bond), alkene, (double bond), alkyne (triple bond), aromatic (benzene ring). • Alkanes contain only single C-C bonds. They contain as many hydrogen atoms as possible, and are said to be saturated. • Hydrocarbons containing double or triple bonds are unsaturated. • A homologous series is series of compounds that differ by a constant increment. Aromatic hydrocarbons include a benzene ring- 6 carbon atoms with all the bonds alternating between a single and a double bond.

  6. Properties of Carbon Element Carbon is unique • It has 6 electrons in its outer shell arranges 1s22s2sp2 • It has room for 4 bonds to 4 other atoms. • Carbon-to-carbon bonds can be single (A), • double (B), or • triple (C). • Note that in each example, each carbon atom has four dashes, which represent four bonding pairs of electrons, satisfying the octet rule.


  8. C C C C C C Hydrocarbons Alkanes Alkenes Alkynes Aromatics

  9. Properties of Carbon Element • A)The carbon atom forms bonds in a tetrahedral structure with a bond angle of 109.5O. • (B) Carbon-to-carbon bond angles are 109.5O, so a chain of carbon atoms makes a zigzag pattern. • (C) The unbranched chain of carbon atoms is usually simplified in a way that looks like a straight chain, but it is actually a zigzag, as shown in (B).

  10. Properties of Carbon Element Carbon-to-carbon chains can be • (A) straight, • (B) branched, or • (C) in a closed ring. • (Some carbon bonds are drawn longer, but are actually the same length.)

  11. Why does carbon can form so many different compounds? • There are now more than ten million organic compounds known by chemists. • Many more undoubtedly exist in nature, and organic chemists are continually creating (synthesizing) new ones. • Carbon is the only element that can form so many different compounds because each carbon atom can form four chemical bonds to other atoms, and because the carbon atom is just the right, small size to fit in comfortably as parts of very large molecules. • Having the atomic number 6, every carbon atom has a total of six electrons. • Two are in a completed inner shell, while the other four are valence electrons—outer electrons that are available for forming bonds with other atoms.

  12. Why does carbon can form so many different compounds? • The carbon atom's four valence electrons can be shared by other atoms that have electrons to share, thus forming covalent (shared-electron) bonds. • They can even be shared by other carbon atoms, which in turn can share electrons with other carbon atoms and so on, forming long strings of carbon atoms, bonded to each other like links in a chain. • Silicon (Si), another element in group 4A of the periodic table, also has four valence electrons and can make large molecules called silicones, but its atoms are too large to fit together into as great a variety of molecules as carbon atoms can.

  13. Why does carbon can form so many different compounds? • Carbon's ability to form long carbon-to-carbon chains is the first of five reasons that there can be so many different carbon compounds; a molecule that differs by even one atom is, of course, a molecule of a different compound. • The second reason for carbon's astounding compound-forming ability is that carbon atoms can bind to each other not only in straight chains, but in complex branchings, like the branches of a tree. • They can even join "head-to-tail" to make ringsof carbon atoms. • There is practically no limit to the number or complexity of the branches or the number of rings that can be attached to them, and hence no limit to the number of different molecules that can be formed. • The third reason is that carbon atoms can sharenot only a single electron with another atom to form a single bond, but it can alsosharetwo or three electrons, forming a double or triple bond. • This makes for a huge number of possible bond combinations at different places, making a huge number of different possible molecules. • And a molecule that differs by even one atom or one bond position is a molecule of a different compound.

  14. Why does carbon can form so many different compounds • The fourth reason is that the same collection of atoms and bonds, but in a different geometrical arrangementwithin the molecule, makes a molecule with a different shape and hence different properties. • These different molecules are called isomers. • The fifth reason is that all of the electrons that are not being used to bond carbon atoms together into chains and rings can be used to form bonds with atoms of several other elements. • The most common other element is hydrogen, which makes the family of compounds known as hydrocarbons. • But nitrogen, oxygen, phosphorus, sulfur, halogens, and several other kinds of atoms can also be attached as part of an organic molecule. • There is a huge number of ways in which they can be attached to the carbon-atom branches, and each variation makes a molecule of a different compound.

  15. The Greater Stability of C-C Bonds • Since the average bond dissociation energy of C-C is greater than the average bond energies between different atoms. • Thus the energy released when carbon atom bonds to another carbon atom is greater than the energy released when the other atoms like B,N,O,Si,P and S bonds to each other. • Thus C-C bond is more stable than the others like B-B,N-N, O-O,Si-Si,P-P and S-S.

  16. The atoms closer to C in the periodic table are B,N,O,Si,P and S. The ability of these atoms to bond each other to form chains is lower than C. For examle Si can produce chains made of at most 11 atoms of it and N at most three atoms it. Although the ability to form chains between their atoms for P and S is greater than Si and N but it is very much smaller compared to C. Ability to Form Chains Between Their Atoms

  17. Ability to Form Chains Between Their Atoms The greater ability of carbon to form chains compared to atoms closer to it in the periodic table can be explained by two reasons: • The average bond dissociation energies of them is lower than that of carbon. • The electronegativity values B,Si and P lower than that of C.atoms.Thus the attraction forces between these atoms are smaller than that of carbon.This is also true when these atoms are bonded to the other atoms like hydrogen or halogens.

  18. Electronegativity • Electronegativity: • a measure of an atom’s attraction for the electrons it shares with another atom in a chemical bond • Pauling scale • generally increases left to right in a row • generally increases bottom to top in a column

  19. Greater Bonding Capacity of C compared to N and O The electronegativity values of N and O are greater than that of C. But their bonding capacities are smaller than that of C since they have lower number of unpaired electrons. Lewis Dot Diagrams of Selected Elements

  20. Summary… • Compared to C atom B,Si,P,N and O atoms can not be expected to form greater number of compounds and unbrached and branched chains and cyclic compounds. • Carbon compounds are more stable than Si4,P4,O3,S8 and B4 molecules.

  21. Electron Configuration of Elements

  22. Lewis Dot Structures… • Gilbert N. Lewis • Valence shell: • the outermost occupied electron shell of an atom • Valence electrons: • electrons in the valence shell of an atom; these electrons are used to form chemical bonds and in chemical reactions • Lewis dot structure: • the symbol of an element represents the nucleus and all inner shell electrons • dots represent valence electrons

  23. Lewis Dot Structures • Table 1.4 Lewis Dot Structures for Elements 1-18

  24. Lewis Model of Bonding… • Atoms bond together so that each atom acquires an electron configuration the same as that of the noble gas nearest it in atomic number • an atom that gains electrons becomes an anion • an atom that loses electrons becomes a cation • the attraction of anions and cations leads to the formation of ionic solids • an atom may share electrons with one or more atoms to complete its valence shell; a chemical bond formed by sharing electrons is called a covalent bond • bonds may be partially ionic or partially covalent; these bonds are called polar covalent bonds

  25. Covalent Bonds! • The simplest covalent bond is that in H2 • the single electrons from each atom combine to form an electron pair • the shared pair functions in two ways • simultaneously; it is shared by the two atoms and fills the valence shell of each atom • The number of shared pairs • one shared pair forms a single bond • two shared pairs form a double bond • three shared pairs form a triple bond

  26. Energy (KJ/mol) 0.74 A H – H distance Hydrogen Molecule Formation Potential Energy Diagram - Attraction vs. Repulsion 0 balanced attraction & repulsion no interaction increased attraction increased repulsion - 436 (internuclear distance) Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 318

  27. Lewis Structures! • To write a Lewis structure • determine the number of valence electrons • determine the arrangement of atoms • connect the atoms by single bonds • arrange the remaining electrons so that each atom has a complete valence shell • show a bonding pair of electrons as a single line • show a nonbonding pair of electrons as a pair of dots • in a single bond atoms share one pair of electrons, in a double bond they share two pairs of electrons, and in a triple bond they share three pairs of electrons

  28. Table of Lewis Structures! • In neutral molecules • hydrogen has one bond • carbon has 4 bonds and no lone pairs • nitrogen has 3 bonds and 1 lone pair • oxygen has 2 bonds and 2 lone pairs • halogens have 1 bond and 3 lone pairs

  29. Resonance! In chemistry, resonance is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis formula. A molecule or ion with such delocalized electrons is represented by several structures called resonance structures.

  30. Tautomerization Tautomerization usually involves the movement of a hydrogen atom between a different location on the molecule, resulting in two or more molecular structures. These structures are called tautomers, which exist in dynamic equilibrium with each other. Enol form Keto form

  31. Molecular Geometry and Bonding Theories

  32. H H H C H H C H H H tetrahedral shape of methane H 109.5o C H H H CH4 molecular shape molecular formula structural formula ball-and-stick model tetrahedron

  33. Cl H Cl H Cl H C C Cl H H 109.5o C H H H space-filling model Methane & Carbon Tetrachloride molecular formula structural formula molecular shape ball-and-stick model CH4 CCl4

  34. H 109.5o C H H H 104.5o 107.3o Molecular Geometry 180o Trigonal planar Linear Tetrahedral Bent Trigonal pyramidal H2O CH4 AsCl3 AsF5 BeH2 BF3 CO2

  35. .. H N C H H H H 107o 109.5o H H .. O H .. 104.5o H CH4, methane NH3, ammonia H2O, water

  36. B A B B A B B B A B B B Molecular Shapes Three atoms (AB2) Four atoms (AB3) • Linear (180o) • Bent • Trigonal planar (120o) • Trigonal pyramidal linear trigonal planar Five atoms (AB4) • Tetrahedral (109.47o) tetrahedral Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 313.

  37. B : N : : O Bonding and Shape of Molecules Number of Bonds Number of Unshared Pairs Covalent Structure Shape Examples -Be- 0 0 0 1 2 2 3 4 3 2 Linear Trigonal planar Tetrahedral Pyramidal Bent BeCl2 BF3 CH4, SiCl4 NH3, PCl3 H2O, H2S, SCl2 C

  38. Molecular Shapes AB2 Linear AB3 Trigonal planar AB2E Angular or Bent AB2E2 Angular or Bent AB4 Tetrahedral AB3E Trigonal pyramidal

  39. Molecular Polarity Molecular Structure Courtesy Christy Johannesson

  40. + - H Cl Dipole Moment • Direction of the polar bond in a molecule. • Arrow points toward the more electronegative atom.

  41. Determining Molecular Polarity • Depends on: • dipole moments • molecular shape Courtesy Christy Johannesson

  42. F BF3 B F F Determining Molecular Polarity • Nonpolar Molecules • Dipole moments are symmetrical and cancel out. Courtesy Christy Johannesson

  43. O net dipole moment H2O H H Determining Molecular Polarity • Polar Molecules • Dipole moments are asymmetrical and don’t cancel . Courtesy Christy Johannesson

  44. H net dipole moment CHCl3 Cl Cl Cl Determining Molecular Polarity • Therefore, polar molecules have... • asymmetrical shape (lone pairs) or • asymmetrical atoms Courtesy Christy Johannesson

  45. Bond dipoles .. .. Bond dipoles m = Q r Dipole moment, m Coulomb’s law Dipole Moment In H2O the bond dipoles are also equal in magnitude but do not exactly oppose each other. The molecule has a nonzero overall dipole moment. C O O Overall dipole moment = 0 O Nonpolar H H The overall dipole moment of a molecule is the sum of its bond dipoles. In CO2 the bond dipoles are equal in magnitude but exactly opposite each other. The overall dipole moment is zero. Overall dipole moment Polar Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 315

  46. .. .. Polar and Nonpolar Molecules .. F N O H Cl H H H B H H F F Polar Polar Nonpolar Polar Cl Cl H Cl C C Cl H Cl H Nonpolar Polar A molecule has a zero dipole moment because their dipoles cancel one another.

  47. H H atomic orbitals H Be Be H H s p H Be s p Formation of BeH2 using pures and p orbitals Be = 1s22s2 BeH2 Be p s No overlap = no bond! atomic orbitals atomic orbitals The formation of BeH2 using hybridized orbitals hybrid orbitals BeH2 Be sp p All hybridized bonds have equal strength and have orbitals with identical energies.

  48. Hybridization - The Blending of Orbitals + = Poodle + Cocker Spaniel = Cockapoo + = sp orbital s orbital + p orbital =

  49. 1s 2s 2p Be atom with one electron “promoted” 1s 2s 2p hybrid orbitals 1s sp 2p n = 2 sp Be atom of BeH2 orbital diagram n = 1 hybridize Be H H s orbital p orbital sp hybrid orbitals shown together (large lobes only) two sp hybrid orbitals sp Hybrid Orbitals Ground-state Be atom Energy px py pz s