TRIGONAL PLANAR MOLECULES

2For the three redcircles to befarthest apart, theyspread out so thateach is 120degrees from theothers!TRIGONALPLANARMOLECULES

3notactually90degreesThese hydrogen atoms might appear at first glance tobe 90 degrees apart, but remember that moleculesexist in THREE DIMENSIONS, not two!Each hydrogen atom is actually 109.5 degreesapart, forming a TETRAHEDRON.These atoms are in the plane of the paper!This atom is behind the paper!This atom is pointing out at you!To see the tetrahedron in three dimensions WITHOUT buying amolecular model kit, just take four balloons, blow them up, and thentie them together. The knot will be the central atom, and the balloonswill line themselves up to be 109.5 degrees apart.

4VSEPR shapes"Groups" can be either BONDS or LONE PAIRS!VSEPR shapes:Groups aroundcentral atom23Shapelineartrigonal planar4tetrahedral / pyramidal / bent5trigonal bipyramidal (andderivatives)6octahedral (and derivatives)Bond angle(s)in degrees180120109.590 and 12090

5More on "4 things around a central atom":- A compound that obeys the octet rule can have a maximum of four groups around its centralatom. But we describe the molecular shape based on how ATOMS are arrnaged aroundthe center. What if some of those groups aren't atoms, but pairs of UNSHARED electrons?These atoms are in the plane of the paper!This atom is behind the paper!With four ATOMSaround the center,we call the shape"TETRAHEDRAL"This atom is pointing out at you!With two ATOMSand two LONE PAIRS,we call the shape"BENT"With three ATOMSand one LONE PAIR,we call the shape"PYRAMIDAL"

6SHAPES OF EXPANDED VALENCE MOLECULESThere are five atoms bonded to the central phosphorus atom, and they willattempt to get as far apart as possible from one another!The top and bottom atoms are90 degrees apart from the atomsaround the center.The atoms around the center are120 degrees apart from each other.There are acually two DIFFERENT bond angles in this structure. It's calledTRIGONAL BIPYRAMIDAL.There are several derivatives of the trigonal bipyramidal shape (like the tetrahedralshape) - depending on how many things around the central atom are atoms!

7There are six atoms bonded to the central sulfur atom, and they willattempt to get as far apart as possible from one another!All bond angles in this arrangement are90 degrees!This shape is called OCTAHEDRAL, since it haseight sides.Like the tetrahedral and trigonal bipyramidal arrangements, there are several derivativesof the octahedron - depending on how many of the six things around the center areatoms!

8Examples:Shape? The central atom is surrounded by fourchlorine atoms (and no lone pairs), so this isa TETRAHEDRAL molecule.skeletalstructurefinal structureShape? Two atomsbonded to the centralcarbon atom (and nolone pairs) give thismolecule aLINEAR shape.Shape? This molecule has four groups around thecentral atom, but only three of them are atoms(one lone pair). So, this molecule isPYRAMIDAL (with the same bond angles astetrahedral - about 109.5 degrees)

9lineartetrahedralpyramidal

10Shape? This molecule has TWO "central"carbon atoms, so we'll just describe theshape of the molecule around each one.skeletalstructureEach carbon has THREE other atomsattached to it and NO lone pairs, so thefinalstructure shape of the molecule around eachcarbon is TRIGONAL PLANAR.Structure tip: Multiplecarbon atoms meanmultiple "central atoms"skeletalstructurefinalstructureShape? Three groups around the centralcarbon atom. All three are atoms, so thismolecule is TRIGONAL PLANAR. It doesn'tmatter that one atom is oxygen and theother two are hydrogen atoms when you'redescribing the shape.

11"ethene"Each carbon isTRIGONAL PLANAR"formaldehyde"TRIGONAL PLANAR

12VSEPR and large molecules- Large molecules have more than one "center" atom- Describe the molecule by describing the shape around each "center".Each of the three carbon centers is TETRAHEDRAL, sinceeach are surrounded by four groups.The shape around this oxygen atom is BENT.These carbon atoms have TETRAHEDRAL geometry.

13All bond angles in the propane moleculeare 109.5 degreesLike propane, the bondangles in ethanol are alsoclose to 109.5 degrees.

14POLARITY and shape:- A polar molecule has an uneven distribution of electron density, making it have ends (poles) thatare slightly charged.POLARITY influences several easily observable properties.- Melting point. (Polar substances have higher melting points than nonpolarsubstances of similar molecular weight.)- Boiling point. (Polar substances have higher boiling points than nonpolarsubstances of similar molecular weight.)- Solubility. (Polar substances tend to dissolve in other polar substances, whilebeing insoluble in nonpolar substances. Nonpolar substances dissove othernonpolar substances, and generally have poor solubility in polar solvents.)- Polar molecules contain POLAR BONDS arranged in such a way that they do not canceleach other out. but how can we tell whether or not a bond will be POLAR? Use experimental data onELECTRONEGATIVITY!ELECTRONEGATIVITY:-A measure of how closely to itself an atom willhold shared electrons- A bond where there is a LARGE electronegativity differencebetween atoms will be either POLAR or (for very large differences)IONIC!- A bond with little or no electronegativity difference between atomswill be NONPOLAR

15ELECTRONEGATIVITY TRENDS- You may look up elecronegativity data in tables, but it helps to know trends!IAIIA2LiBe3Na MgIIIB IVB VB VIB VIIBINCREASINGELECTRONEGATIVITYIIIA IVA VA VIA VIIA4K5Rb Sr6Cs Ba La Hf Ta W7FrCa Sc TiYZrRa Ac RfVIB IIBVIIIBCr Mn Fe Co NiBCNOFAlSiPSClCu Zn Ga Ge As Se BrNb Mo Tc Ru Rh Pd Ag Cd InRe Os IrDb Sg Bh Hs MtPt Au Hg TlSnSb TePb BiIPo At"inner" transition metals go here- FLUORINE is the most elecronegative element, while FRANCIUM is the least!- All the METALS have low electronegativity, and metal/nonmetal combinations form IONICbonds- HYDROGEN is similar in electronegativity to CARBON, so C-H bonds areconsidered NONPOLAR

16Examples:Polar molecule?POLAR BONDS? Yes. C-F bonds are polar. Large electronegativitydifference between C and FSHAPE? Tetrahedral, with a fluorine at each point. Sincethe molecule is symmetric, there's no negative "side" or positive"side". This is a NONPOLAR molecule.Polar molecule?POLAR BONDS? Yes. C-F bond is polar. C-H bonds arenonpolar.SHAPE? Tetrahedral, but since the molecule isn't perfectlysymmetrical due to the 3 H atoms and only a single F, thisis a POLAR molecule.Polar molecule?POLAR BONDS? Yes (C-F).SHAPE? Tetrahedral, with F atoms at two points, and Hatoms at the other two. This gives the molecule a hydrogen"side" and a fluorine "side" . making it POLAR.Polar molecule?POLAR BONDS? C O is polar.SHAPE? LINEAR. The C O bonds are arrangedsymmetrically, so this molecule is NONPOLAR.

17"fluoromethane"Fluorine is able to pull electron density throughthe molecule, as it is being opposed by much lesselectronegative hydrogen atoms."difluoromethane"In 2D, the fluorine atomsappear to be on theopposite sides of themolecule, butin 3D they are on thesame side.

18VALENCE BOND THEORY- an attempt to explain why molecules behave in the way that the VSEPR model predicts.- Describes the formation of bonds in terms of the OVERLAP of ORBITALS from the bondingatoms.Bonds are formed when two atoms are close enough together so thattheir ORBITALS OVERLAP (share the same space).Each SET of overlapping orbitals can contain at most a total of TWOelectrons. So, two orbitals with one electron each may bond. An orbitalwith two electrons can only bond with an EMPTY orbital (This is called aCOORDINATE COVALENT BOND.)These 1s orbitals overlap to formwhat we call a "sigma bond" withoverlap BETWEEN the twoatomic nuclei.

19Hybridization- Look at carbon's electron configuration:You would expect that carbon would form several differentkinds of bonds in a molecule like methane. But, methane'sbonds are experimentally all identical. How doescarbon form the four equivalent C-H bonds we see inmethane?We observe that thesebonds are IDENTICAL!Same bond energy,distance, and angle.

20- In valence bond theory, atomic orbitals can COMBINE to make new orbitals that can then goon to bond with other molecules.- When orbitals combine to make HYBRID ORBITALS, .The overall NUMBER OF ORBITALS does not change.The overall NUMBER OF ELECTRONS around the atom doesnot changeThe energy of the orbitals is between the energies of the orbitalsthat combine.These sp3 orbitals were formed fromthe combination of carbon's original2s and 2p orbitals. These orbitals areall identical, and are spread 109.5degrees apart from one another.Hybrid orbitals are named from theorbitals that go into making thehybrid. 2s 3 2p orbitals "sp3"!

21Types of hybrid orbitals:Hybrid typeNumber oforbitalsMolecular shapesp2linearsp23trigonal planarsp34tetrahedral (or derivatives)sp3d5trigonal bipyramidal(or derivatives)sp3d26octahedral (or derivatives)

22MULTIPLE BONDS and VALENCE BOND THEORY- Valence bond theory provides an explanation of multiple (double and triple) bonding thatexplains some interesting observations about these kinds of bonds.Each carbon has a TRIGONAL PLANARgeometry. This suggests that the carbonsare "sp2 hybridized".ethyleneOne unchanged 2porbitalThree sp2 hybrids thatare 120 degrees apart"Original"unbondedcarbon atomCarbon atom withsp2 hybrid orbitals

23sp2 hybrid orbitals in BLUE2p orbital in REDThe 2p orbitals overlap above and below the axis betweenthe two carbon atoms. This OFF-AXIS overlap is called aPI BOND.The sp2 hybrid orbitals overlap ON THE AXIS betweenthe two carbon atoms. This bond is called a SIGMABOND.As you can see, the carbon-carbon double bond in ethylene is made up of TWO DIFFERENTKINDS OF BONDS!

24Some notes on sigma and pi bonds:SIGMA bonds are formed when orbitals overlap along the axis between two atoms. Thesebonds have good overlap between the bonding orbitals, meaning that they arestrong. Single bonds are always sigma bonds. Double and triple bonds contain onesigma bond each.PI bonds are formed when off-axis orbitals (usually p orbitals) overlap. Since theoverlapping orbitals do not face each other as in the sigma bond, the overlap inpi bonds tends to be poorer than in sigma bonds, As a result, pi bonds tend tobe weaker than sigma bonds. Double bonds contain a single pi bond, and triple bondscontain two pi bonds.Experimentally, we observe that the bond energy of the C C bond is less thanthe bond energy of two C-C bonds. This suggests that the second bondin a double bond is different from the first!Molecules may rotate around SIGMA bonds, since rotation around the axis betweentwo atoms will not affect the overlap and break the bond. Off-axis PI BONDS preventrotation because rotation would break the pi bond.

25ROTATION, ISOMERS, and VALENCE BOND THEORY- Consider this molecule:"1,2-dichloroethane". are these two structures different?No! The molecule is free to rotate around the C-C single (sigma) bond, and we do not observetwo different versions of 1,2-dichloroethane. Both of the forms drawn above are equivalent.

26The molecule isfree to rotateabout thecarbon-carbonbond!

27. now consider "1,2-dichloroethene":. are these two structures different?YES! The two carbon atoms in these structures are held together by a DOUBLE BOND,which contains a pi bond. The molecule cannot rotate around the C C double bondwithout breaking the pi bond, so the form with the two chlorine atoms on opposite sidescannot freely flip over to the form with the chlorine atoms on the same side.These two Lewis structures actually represent DIFFERENT MOLECULES. They are calledISOMERS, since they have the same chemical formula but different arrangements ofatoms.

28For this rotation to take place, the PI BOND must break and then re-form!